SCIENCE BOOKS

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Models and structure of the atom

1. Standard notation
2. The atom: definitions
3. Multiple choice: Standard nuclear notation
4. Standard notation
5. Structure of the atom
6. Isotopes
7. Standard notation
8. Isotopes: relative atomic mass
9. Models of the atom

Key unit Competence

To relate Bohr’s atomic model with atomic spectrum of Hydrogen, write electronic configuration of atoms and ions using s, p, d and f atomic orbitals and interpret graphical information related to ionization energy of elements.

Learning objectives

By the end of this unit, I will be able to:

•Explain the stability of atoms using the concept of quantization of energy.

•Explain the achievements and limitations of Bohr’s atomic model.

•Explain the existence of energy levels using the data from emission spectra.

•Describe hydrogen spectral lines and spectral line series

•Explain the types of spectra in relation with the nature of light

•Explain the quantum theory of the atom using the quantum numbers.

•Determine the number and shapes of orbitals in given level or principal quantum number

•Explain the rules governing the electronic configuration: Aufbau principle and Hund’s rule

•Explain the relationship between the electronic configuration and the stability of the atoms

•Interpret the graphs of first ionisation energy against the atomic number.

•Describe the factors which influence the first ionisation energy.

2.1. Bohr’s atomic model and concept of energy levels

The potential energy of a person walking up ramp increases in uniform and continuous manner whereas potential energy of person walking up steps increases in stepwise and quantized manner. This can be explained by the values of energy which are continuous for the person walking up ramp while they are discrete (discontinued) for the person walking up steps (Figure 2.1(a) and Figure 2.1(b).

Niels Bohr (1885-1962) a young Danish physicist working in Rutherford’s laboratory, suggested a model for the hydrogen atom and predicted the existence of line spectra. In his model, based on Planck’s and Einstein’s ideas about quantized energy, Bohr proposed three postulates:

•An electron can rotate around the nucleus in certain fixed orbits of definite energy without emission of any radiant energy. Such orbits are called stationary orbits.

•An atom can make a transition from its stationary state of higher energy E2 to a state of lower energy E1 and emit a single photon of frequency ν. Conversely, an atom can absorb an energy at the lower level E1 and transit to the higher energy level E2. That is, the change in energy for a system which can be represented by the equation:

where n is an integer (1,2,3,...) and h is Planck’s constant

determined from experiment and has a value of 6.626x10-34J.s; ⱱ is the frequency of the electromagnetic radiation absorbed or emitted. Each of these small “packets” of energy is called photon also called quantum of energy. Energy can be gained or lost only in whole-number multiples of the quantityνh,That is, the change in energy for a system can be represented by the equation:, where n is an integer (1,2,3,...).

An atom does not release energy when it is in one of its stationary states. That is, the atom does not change energy while the electron moves on a given orbit. When an electron on a given orbit absorbs an appropriate quantum of energy, it jumps, i.e. is promoted to a higher energy orbit; this process is called “excitation”of electron. On the contrary, if the electron loses an appropriate quantum of energy, it falls on the lower energy orbit by emission of a light corresponding to the lost quantum of energy and the process is called “de-excitation” of electron. As there are many energy levels on which electrons can be excited and de-excited, an atom will have many lines of absorption, each corresponding to a quantum of energy absorbed: this appears as a series of lines called absorption spectrum. In the same way the series of emission lines will produce an emission spectrum (see Fig. 2.4).

2.1.1. Achievements of Bohr’s Atomic Model

• Explanation of the stability of an atom

Based on Rutherford’s atomic model, the electrons move around the nucleus in circular paths called orbits. According to the classical theory of electromagnetism, a charged particle revolving around a charged nucleus would release energy and end up by spiraling into the nucleus; thus the atoms would be unstable. The Bohr’s atomic model makes an assumption of discreet orbit (allowed orbit) to explain why an atom is stable; by doing so, Bohr introduces the concept of quantization of energy. Bohr’s atomic model explains the origin of atomic absorption and emission spectra.

• Explanation of the production of the absorption and emission spectra

The Bohr’s atomic model explains the origin of atomic absorption and emission spectra.

2.1.2. Limitations of Bohr Model

1.Bohr’s theory fails to explain the origin of the spectral lines of multi-electron atoms.

It only explains the origin of the spectrum of hydrogen-like species having only one electron such as H, He+, Li2+, Be2+, ........

The model fails to explain the spectral lines of atoms or species with more than one electron.

2. According to Bohr, the circular orbits in which electrons revolve are planar. However, modern research has shown that an electron moves around the nucleus in the three dimensional space.

3. Bohr’s theory fails to account for Zeeman Effect and Stark Effect. Zeeman Effect is the splitting of the spectral lines into thinner and closely- spaced lines when an excited atom is placed in a magnetic field. Stark Effect consists of the splitting of the spectral lines into thinner and closely-spaced lines in presence of electric field.

4. Bohr’s theory is in contradiction with Heisenberg’s uncertainty principle. Bohr assumes that the electron revolves around the nucleus in circular orbits at fixed distance from the nucleus and with a fixed velocity. However, according to W. Heisenberg, it is not possible to know simultaneously the accurate position and the velocity of a very small moving particle such as an electron.

Checking up 2.1

1. Find out two more examples that you can use to illustrating the concept of quantization.

2. Discuss the main weakness of Rutherford’s nuclear atom.

2.2. Hydrogen spectrum and spectral lines

Activity 2.2

Look at the picture of neon tube light below and do research about how this neon tube light works to produce light and present your findings

Bohr’s atomic model allows to explain the emission spectra of atoms. This happens when excited electrons lose energy in form of electromagnetic radiation and fall to lower energy levels.

The wave-particle nature of the light

Light as a waveThe light is a wave-like phenomenon as shown in Figure 2.2.

It is characterized by its wave length, generally symbolized by the Greek letter lambda, λ, and its frequency, represented by the Greek letter nu1, ν.

As shown in the Figure below, the wavelength represents the distance between two successive summits/peaks (or two successive troughs).The frequency represents the number of complete wavelengths made by the light per second, also called cycles per second.

Visible light is composed by different visible lights with different λ and ν.

But all those lights have the same speed, the speed of light, which, in a vacuum, is equal to: 3.00x108 m/s; although different types of light have different λ and ν, they move at the same speed c. This results in the relation between the speed of light and its wavelength and frequency: c = νλFrom this relation, and since c is constant, we can conclude that:

•Light with long wavelength has low frequency, whereas

•Lights with short wavelength has high frequency.

Let’s take an example to illustrate: light1 has λ1 equal to 105m whereas light2 has λ2 equal to 10-5 m. After 1 second, both would have travelled 3.00x108m, the speed of light, but their frequencies will be different:

ν1 = c/λ1 = 3.00x108ms-/105m = 3.00x103s-= 3,000 cycles/s

v2 = c/λ2= 3.00x108ms-/10-5m= 3.00x1013s-= 30,000,000,000,000 cycles/sLight is energy, represented by: E = hν where h is Plank’s constant (h=6.626x10-34Js)

Hence energy in the light is proportional to its frequency; higher the frequency of the light, higher is its energy and vice-versa.

The different colors of the visible light differ by their wavelength as shown in the

As illustrated in Figure 2.2 below, the right side of the spectrum consists of high-energy, high-frequency and short wavelength radiations. Conversely, the left side consists of low-energy, low-frequency and long wavelength radiations.

The letter gamma, γ, may also be used.

When an electron is excited or de-excited, the energy absorbed or emitted corresponds to the difference of energy, ΔE, between the final energy level of the electron, E2, and the starting energy level of the electron, E1: E2 – E1 = ΔE = hν. ΔE is positive when E2>E1, this is the case of absorption and excitation of electron; on the other hand ΔE may be negative when E2<E1, in case of emission and de-excitation of electron.

Figure 2.4 below shows the different series of emission spectra of hydrogen. As you can see, the difference between those series is the final energy level where the electron fall after de-excitation.

The series have been named according to the scientists who discovered them.

Ionization of an atom or loss of an electron corresponds to excitation of an electron to the level n=∞.

Checking2.2

1. What is the meaning of infinity level in the hydrogen spectral lines?

2. Given a transition of an electron from n=5 to n=2. Calculate energy

ii) Frequency

iii) Wavelength

2.3. Atomic spectra

Activity 2.3

Observe the picture above, discuss in groups and answer the following questions.

a. What do you see on the above photo?

b. State the physical phenomenon which is related to the above photo.

c. Think of any other means of producing the same pattern. List two of them.

d. What property can you attribute to light with reference to the above process?

Checking up 2.3

1. Different metals, when exposed to a flame, emit different flame colors. Explain the origin of that difference.

2. Would you expect to see the emission rays and the absorption rays of a given element to appear at the same place of a photographic plate or not. Explain your answer.

3. How do you explain the many spectral rays for the same element?

2.4 Orbitals and Quantum Numbers

Activity 2.4

1. a)Write the electronic configuration of aluminium atom(Z=13)

b) Indicate the number of electrons in each energy level/quantum shell

c) The shells are numbered from inside-outward starting from 1, 2, 3, 4 ... which other name is given to these shells?

d) How did you obtain the exact number of electrons in each energy level/quantum shell in (c) above?

We have seen the weakness and critics against the atomic Bohr’s model. In order to answer to the questions not answered by that model, other atomic models were proposed. One of those models is the Quantum model that has been developed by the Austrarian physicist Erwin Schrödinger (1887-1961). The model is based on a mathematical equation called Schrödinger equation. This model is based on the following assumptions or hypotheses:

•An electron is in continuous movement around the nucleus but cannot be localized with precision; only the high probability of finding it in a cartain region around the nucleus can be known.

•The region where the probability of finding electron is high, at more than 95%, is called “orbital”; in other words, the orbital is the volume or the space (tridimensional) around the nucleus where there is a high probability of finding the electron.

Without going into the mathematical development of the Schrödinger equation, we can say that the energy of the electron depends on the orbital where it is located. And an atomic orbital is described by a certain number of “quantum numbers”

according to the solution of Schrodinger equation, i.e. 3 whole numbers:

1) The principal quantum number is a positive integer which varies from 1 to ∞.

The principal quantum number indicates the energy level in an atom where electrons can be located: the higher the n value, the higher the energy level. An electron in energy level n=1 has lowest energy in an atom. The principal quantum number, n, has been traditionally given names by the letters: K(n=1), L(n=2), M(n=3), N(n=4), O(n=5), P(n=6).

In the Bohr’s atomic model, K, L, M, ... were used to represent different orbits or shells of electrons. Later on, the term shell sometimes is used to describe a group of orbitals with the same principal quantum number. The term subshell describes a group of orbitals with the same principal and second quantum number. The maximum number of orbitals and electrons that can be found in an energy level n are n2and 2n2, respectively (Table 2.1). The maximum number of sub shells in an energy level n equals n.

Table 2.1 Relation between the principal quantum number, the number of orbitals and the maximum number of electrons.

2)The angular momentum quantum number (l)

The second quantum number is the angular quantum number represented by the letter, l: it is an integer which can take any value from zero or higher but less than n-1, i.e. equal to: 0,1, 2, 3,....up to n-1. For example if n= 1, l is equal to 0, if n= 2, l can be 0, 1. It is also called secondary or azimuthal quantum number. It indicates the shape of the orbital and is sometimes called the orbital shape quantum number. By tradition, those different shapes of orbitals have been given names or letter symbols: l = 0 = s, l =1 = p, l = 2 = d, l=3 = f

3) Magnetic quantum number (ml)

The magnetic quantum number describes the orientation of the orbital. It is an integer that varies from -l to +l. For example if: l = 0, ml can only be 0; if l = 1, ml = -1, 0, +1; if l=2, ml = -2, -1, 0, 1, 2. As you can see for each value of l there are (2l+ 1) values of ml corresponding to (2l + 1) orientations under the influence of magnetic field. The s orbital where l is zero and ml has no orientation; it has the shape of a sphere as shown in

Table 2.2: Relationship between the n, l and ml

The table 2.2 shows that, apart s sub-level that has only one orbital, other sub-levels have a certain number of different orbitals; those orbitals have the same energy but differ in their specific orientations. Example p orbitals are 3 with different orientations: pxpypz.

Different sub-levels belonging to the same principal quantum number have different energies as follows: s < p < d < f

4) The spin quantum number (S)

The fourth quantum number is the spin quantum number, represented by the symbol S (or ms in some books). The electron behaves as a spinning magnet.The spin quantum number is the property of the electron, not the orbital.This number describes the spinning direction of the electron in a magnetic field. The direction could be either clockwise or counterclockwise. The electron behaves as if it were spinning about its axis, thereby generating a magnetic field whose direction depends on the direction of the spin. The two directions for the magnetic field correspond to the two possible values for the spin quantum number, S (ms). Only two values are possible: s = +1/2 and -1/2 as shown in the Figure 2.7 below.

ms = +1/2 and ms = -1/2 are commonly represented by ↑ and ↓ respectively.In conclusion an electron in any given atom is decribed by 4 quantum numbers: (i) three quantum numbers which describe the orbital where the electron is located: n, l and ml and (ii) one quantum number describing the spin of the electron, S or ms.

2.5 Electronic configuration of atoms and ions

Activity 2.5

1. Write the electronic structure of the following chemical speciesK (Z=19), Ne (Z=10), Al3+ (Z=13), Cl (Z=17), O2- (Z=16)

2. Using information in question 1,

a. Determine the group and period of K and Cl.

b.Which species have a stable electronic configuration? Explain.

The electron configuration is the distribution of electrons of an atom in its atomic orbitals. The electronic configuration of an atom is governed by three main rules.

1) Aufbau Principle

The Aufbau principle explains how to build the electronic configuration of an atom.The Aufbau principle states that atomic orbitals of lower energy must be filled before the higher energy orbitals. The Aufbau principle is referred to as the “building-up” principle.

According to that principle, if an atom has only one electron, this electron will occupy the lowest principal quantum energy level, n=1, and the lowest energy orbital in that principal quantum energy, i.e. l = 0 = s. This is represented as: 1s1, meaning one electron in s orbital of the 1st energy level. If the atom has 2 electrons, the second electron will be filled in the same s orbital to give the structure 1s2.

Note the way of notation: the first number indicates the principal energy level or principal quantum number, followed by the letter indicating the orbital, followed by the number of electrons present in the orbital, as super-script.

2) Pauli Exclusion Principle

What happens if we have 3 electrons in an atom? Can we squeeze them in 1s orbital?

There is a principle called Pauli Exclusion Principle states: in an atom, two electrons cannot have the four quantum numbers n, l, ml, and S(ms) identical. This explains why the two electrons in the same orbital must have their spin opposite:

The Pauli exclusion principle doesn’t allow us to put a 3rd electron in the same orbital, since if we put 3 electrons in the same orbital, 2 electrons will have 4 identical quantum numbers and this is not allowed

In other words, the Pauli Exclusion Principle is telling us that you cannot put more than 2 electrons in an orbital. i.e. the maximum number of electrons in an orbital is 2.

Hence, the 3rd electron must go in energy level n = 2, since energy level n=1 if full; it has only one orbital s. Then for that atom with 3 electrons, the electronic structure is: 1s22s1

3) Hund’s rule

Atom with 4 electrons: 1s22s2

Atom with 5 electrons: 1s22s22p1

What about an atom with 6 electrons? Are we putting the 6th electron in the same orbital as the 5th electrons? Remember that there are 3 p orbitals pxpypz of the same energy!

The Hund’s rule answers to that question.

It states that orbitals of equal energy are each occupied by one electron before electrons begin to pair up into the same orbital. By convention, all the unpaired electrons are given the same orientation spin. A slight preference for keeping electrons in separate orbitals helps to minimize the natural repulsive forces that exist between two electrons.

Therefore, atom with 6 electrons: 1s22s22px1py1pz0

When building the electronic configuration of elements, you must be guided by the principles and rules seen above and: writing the principal quantum number in Arabic number, followed by the orbitals immediately followed by the number of electrons in the orbital as superscript.

An atom X: 1s2: has only two electrons in s orbital at the 1st energy level

An atom Y: 1s22s22p3: has electrons in 2 levels of energy: level n=1, and level n= 2. In level 1, it has 2 electrons in s orbital. In level 2, it has 2 electrons in s orbital and 3 electrons in p orbitals.

Figure 2.8 is a useful and simple aid for keeping track of the order in which electrons are first filled for each atomic orbital. The different orbitals are filled in the order 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.Notice that as energy levels increase starting from n=3, 4s orbital is filled before 3d, 5s before 4d, etc... as shown in the diagram below. But when ionized, 4s electrons are ionized before 3d, and 5s before 4d.

Checkup 2.5 (a)

1. Build the electronic configuration of the following atoms: 1H, 3Li, 5B, 11Na, 18Ar,19K, 21Sc, 24Cr, 26Fe, 29Cu

2. Write the electronic configuration for each of the following pairs of ions. State the more stable ion in gaseous state and explain your choice.

a. cu+ and cu2+

b.Fe2+ and Fe3+

Expanded notation

Expanded notation is another method of writing the s, p, d and f notation. The method uses the same concept as s, p, d and f notation except that each individual orbital of a sub-level having many orbitals is represented with a subscript letter indicating the orientation of the orbital. This applies for p, d, and f orbitals. Considering that p-orbital has three componentsxp,ypand pz, the expanded electronic configuration of some elements is given hereafter.

Checkup 2.5 (b)

Write the expanded electronic configuration for each of the following atom/ions. S(z=16), P3-(z=15), Mg2+(z=12)

Orbital box representation

An orbital box representation consists of a box for each orbital in a given energy level, grouped by sublevel, with an arrow indicating an electron and its spin.

Note that two electrons in the same orbital have necessarily opposite spins as indicated in the examples below.

The table 2.4 shows the electronic configuration of some elements using orbital box representation and applying Hund’s rule.

Table 2.4: Electronic configuration using orbital box representation

Checkup 2.5(c)

Using boxes to represent orbitals, draw the electronic configuration ofN3- (z=7), Ti4+(z=22), Mg2+(z=12), Ar(z=18)Identify the isoelectronic species that are present.

Noble Gas Notation

All noble gases have completely filled subshells and can be used as a shorthand way of writing electron configurations for subsequent atoms.When using this method, the following steps are respected.

a. Identify the noble gas whose electronic configuration is included in that of the concerned element.

b.Write the chemical symbol of the identified noble gas within square brackets. We call this the noble gas core.

c. Add electrons beyond the noble gas core. Note that electrons that are add-ed to the electronic level of the highest principal quantum number (the outermost level or valence shell) are called valence electrons.

Example: Given the electronic configurations of the noble gases Ne and Ar, one can write the electronic configuration of some elements in noble gas notation of some elements as:

Checking up 2.5 (d)

Using the noble gas notation, write the electronic configuration of the following atoms/ions.

a. Ge (Z=32)

b. S (Z=16)

c. Co2+ (Z=27)

d. Br- (Z=35)

e. Sr (Z=38)

2.6. Relationship between ionization energy, energy levels and factors influencing ionization energy

Activity 2.6

Write the electronic configuration of the following elements/ions, use s, p, d, ...)

Sodium, magnesium, magnesium ion (Mg2+), aluminium, aluminium ion (Al3+), oxygen ion (O2-)

Identify the common feature of ions in (1) and why do they have such feature

Suggest what happened to aluminium atom when it changed to aluminium ion (Al3+)

Identify the group and the period of aluminium, sodium and oxygen atom

2.6.1. Concept of Ionization energy

The ionization energy is a measure of the energy needed for an atom, in gaseous state, to lose an electron and become positive ion.

The first ionisation energy is the energy required to remove one electron from an atom in its gaseous state. The example below shows how to represent the successive ionization energies of an atom M.

Second ionisation energy and nth ionisation energy: Two or more electrons can be removed and we have successive ionization energies.

The ionization energy is usually expressed in kiloJoules per mole (kJ.mol-1). This energy is required to overcome the attractive force between the nucleus and the electron and then remove the electron. Theoretically there are as many successive ionisation energies as there are electrons in the original atom. In figure 2.9, someone can make an interpretation of successive ionization energies of an atom

2.6.2. Interpretation of a graph of successive ionization energies of an atom

The graph shows that the energy to remove electron increases as more electrons are successively removed.

a) The energy required to remove the first electron is relatively low. This corresponds to the loss of one 3s electron.

b) To remove the second electron needs greater energy because this electron is closer to the nucleus in a 2p orbital. There is a steady increase in energy required as elec-trons are removed from 2p and then 2s-orbitals.

c) The removal of the tenth and eleventh electrons requires much greater amounts of energy, because these electrons are closer to the nucleus in 1s orbital.

2.6.3. Factors influencing the extent of ionization energy

The ionization energy is a physical property of elements that can be influenced by some factors:

1) Size of atom

The atomic size is the distance between the nucleus and valence shell. As the number of energy levels (shells) increases, the force of attraction between nucleus and valence electron decreases. Therefore, the valence electrons are loosely held to the nucleus and lower energy is required to remove them, i.e. ionization energy decreases with increase in atomic size and vice versa.This is what happens when you go down a Group.

2) Nuclear charge

The nuclear charge is the total charge of all the protons in the nucleus. As the nuclear charge increases, the force of attraction between nucleus and valence electrons on the same valence energy level increases and hence makes it difficult to remove an electron from the valence shell. The higher the nuclear charge, the higher the ionization energy. This is what happens when you cross a period from left to right.

3) Screening effect or Shielding effect

The Screening effect or Shielding effect is due to the presence of inner electrons which have a screening or shielding effect against the attraction of the nucleus towards the outermost electrons. The electrons present in inner shells between the nucleus and the valence shell reduce the attraction between nucleus and the outermost electrons. This shielding effect increases with the increasing number of inner electrons. A strong Shielding effect makes it easier to remove an external electron and hence lowers the ionisation energy.

2.6.4. Importance of ionization energy in the determination of the chemistry of an element

Ionization energy provides a basis to understand the chemistry of an element. The following information is provided.

Determination of metallic or non- metallic character.

The I.E informs us how the atom will behave chemically: a low I.E indicates that the element behaves as metal whereas a high I.E indicates that the element behaves as non-metal.

The first ionization energies of metals are all nearly below 800kJ mol-1 while those of non- metals are all generally above 800 kJ mol-1.

Down the group ionization energies decrease so that the elements became more metallic. In groups 14 and 15 there is change from non metallic to metallic character. Across a period from left to right 1st I.E. increases. The elements become less metallic to non- metallic.

Example: The first three ionization energies for elements A, B, C, and D are given in the table below

This table shows that for a given element: 1st IE < 2nd IE < 3rd IE.

From the 1st I.E. of the elements it can be predicted that elements B and C have typical metallic properties since their 1st ionization energies are low.

D is expected to be non-metal because of its high 1st IE.

Checking up 2.6

1. Given elements: Cl, Ca, and Na and the following 1st IE: 456, 578.8 and 1251 KJ/mol. Respectively.Match those 1st IE with the three element and justify.

2. Given elements: bromine and iodine and 1st IE: 1008 and 1140 KJ/mol. Which 1st IE correspond to which element? Explain.

3. Explain why 2nd IE is always greater than the 1st IE?

2.7. End unit assessment

1. Which of the following is the correct representation of the ground-state electron configuration of molybdenum? Explain what is wrong with each of the others.

2. Which of the following electron configurations are correct and which ones are wrong? Explain.

3. Photosynthesis uses 660 nm light to convert CO2 and H2O into glucose and O2. Calculate the frequency of this light.

4. Which of the following orbital designations are incorrect: 1s, 1p, 7d, 9s, 3f, 4f, 2d?

5. The data encoded on CDs, DVDs, and Blu-ray discs is read by lasers. What is the wavelength in nanometers and the energy in joules of the following lasers?

6. Concerning the concept of energy levels and orbitals,

a. How many subshells are found in 3=n?

b. What are the names of the orbitals in 3=n?

c. How many orbitals have the values 4=nand3=l?

d. How many orbitals have the values 23==,lnand2−=lm?(e) What is the total number of orbitals in the level4=n?

7. A hypothetical electromagnetic wave is pictured here. What is the wavelength of this radiation?

8. Consider the following waves representing electromagnetic radiation:

a. Which wave has the longer wavelength?

b. Calculate the wavelengths of the two radiations

c. Which wave has the higher frequency and larger photon energy?

d. Calculate these values.

9. Order the orbitals for a multielectron atom in each of the following lists according to increasing energy:a. p5 ,d5b. s4, p3c. s6,d4

10. According to the Aufbau principle, which orbital is filled immediately after each of the following in a multielectron atom?

a. 4s

b. 3d

c. 5f

d. 5p

11. According to the Aufbau principle, which orbital is filled immediately be-fore each of the following?

a. 3p

b. 4p

c. 4f

d. 5d

12. Four possible electron configurations for a nitrogen atom are shown be-low, but only one represents the correct configuration for a nitrogen atom in its ground state. Which one is the correct electron configuration? Which configurations violate the Pauli Exclusion Principle? Which configurations violate Hund’s rule?

13. Explain the variation in the ionization energies of carbon, as displayed in this graph.

14. The first seven ionization energies of an element W are shown below

What factors determine the magnitude of the first ionization energy

Key unit competence

Describe how properties of ionic compounds and metals are related to the nature of their bonding

Learning objectives

By the end of this unit, I will be able to:

•Explain why atoms bond together;

•Explain the mechanisms by which atoms of different elements attain stability;

•Explain the formation of ionic bonds using different examples;

•Represent ionic bonding by dot-and-cross diagrams;

•Describe the properties of ionic compounds based on observations;

•Perform experiments to show properties of ionic compounds;

•Assemble experimental set up appropriately and carefully;

•State the factors that influence the magnitude of lattice energy ;

•Relate the lattice structure of metals to their physical properties;

•Describe the formation of metallic bonds;

•State the physical properties of metals and forces of attraction that hold atoms of metal.

People like to bond with each other for many reason such as: to unite their forces and be stronger, to exchange idea and produce big things, to found a family, etc. we cannot live in isolation. People can have strong connection .Similarly some atoms can also have strong bonds between them. Some atoms have weak connections, just like two people can have connections.

Some atoms may not need to bond with others; they are self-sufficient as some people, a small number, may be self-sufficient.

Connections between atoms are called chemical bonds. Solids are one of the three fundamental states of matter. In molecules, atoms or ions are held together by forces called chemical bonds.There are 3 types of chemical bonds: Ionic, Covalent and Metallic bonds.

The type of a bond in molecules is determined by the nature and properties of the bonding atoms. However, in this unit we will only emphasize on ionic and metallic bonding.

3.1. Stability of atoms and why they bind together

Activity 3.1

1. In pairs discuss and write electronic configuration of sodium , neon, argon, magnesium, aluminium, oxygen and chlorine

2. What happens when oxygen and chlorine gain electrons?

3. What happens when sodium, magnesium and aluminium lose electrons?

4. Discuss on how atoms of elements can gain their stabilities by either loosing or gaining electron(s) on the valence shells and show with evidence that an atom is stable?

5. How does the formation of an ionic bond between sodium and chlorine reflect the octet rule?

Like people always relate and connect to others depending on their values, interests and goals so does unstable atoms. They are combine together to achieve stability. We know that noble gases are the most stable elements in the periodic table. The noble gases are extremely unreative: they do not tend to form compounds or combine to themselves

What do the noble gases have in common? They have a filled outer electron energy level. When an atom loses, gains, or shares electrons through bonding to achieve a filled outer electron energy level, the resulting compound is often more stable than individual separate atoms. Neutral sodium has one valence electron. When it loses this electron to chlorine, the resulting Na+ cation has an outermost electron energy level that contains eight electrons.

It is isoelectronic (same electronic configuration) with the noble gas neon. On the other hand, chlorine has an outer electron energy level that contains seven electrons. When chlorine gains sodium’s electron, it becomes an anion that is isoelectronic with the noble gas argon.

Below are examples of how magnesium bonds with oxygen and calcium with chlorine:

Example 1

Example 2

As you can see, the noble gases have the most stable electronic configuration, characterized by the presence of 8 electrons in the outermost shell: this structure is called the “Octet Structure” or “Octet Rule”; octet meaning eight in Latin.Any atom that does not have that structure, will combine with another atom (of the same element or different element) to achieve the octet structure.

Checking Up 3.1

Write the shorthand electronic configuration for copper.Predict the ions that will be formed from atoms in the table, name the isoelectronic noble gas and establish the electronic configuration (E.C) of the ion formed.

3.2. Ionic bonding

Atoms have many ways of combining together to achieve the octet structure, and one of them is the formation of an ionic bond.

In an ionic bond, electrons are transferred from one atom to another so that they form oppositely charged ions; in other words, one atom loses electron(s), the other gains electrons(s).The resulting strong force of attraction between the oppositely charged ions is what holds them together. Ionic bonding is the electrostatic attraction between positive and negative ions in an ionic crystal lattice.

3.2.1. Formation of ionic bond

Activity 3.2

Draw diagrams to illustrate the formation of ionic compounds in magnesium oxide, magnesium chloride, sodium peroxide, and sodium sulphide.

The transfer of electrons from one atom to another followed by attraction between positive and negative ions is called ionic bonding. This type of bonding occurs between metals and non-metals. The compounds formed are called ionic compounds. As stated previously, metals try to lose their outer electrons while non metals look to gain electrons to obtain a full outer shell. When metals lose their outer electrons they form positively charged ions called cations. When non-metals gain electrons they form negatively charged ions called anions. An example is shown below:

The curved arrow between sodium and fluorine represents the transfer of an electron from a sodium atom to a fluorine atom to form opposite ions. These 2 ions are strongly attracted to each other because of their opposite charges. A bond is now formed and the resulting compound is called Sodium Fluoride

Another example of ionic bonding is the bonding of Beryllium and Fluorine to form beryllium fluoride.

The transfer of electrons from beryllium results in the formation of an ionic bond. Beryllium now has a positive (+2) charge and Fluorine now has a negative charge. The resulting compound is called Beryllium Fluoride (BeF2).

Other examples showing the formation of ionic compounds using dot and cross diagrams.

a. Formation of calcium chloride

b. Formation of magnesium oxide

c. Formation of bonds in sodium fluoride

Checking Up 3.2

1. For each of the following ionic bonds: Sodium + Chlorine, Magnesium + Iodine, Sodium + Oxygen, Calcium + Chlorine and Aluminium + Chlorine

a. Write the symbols for each element.

b. Draw a Lewis dot structure for the valence shell of each element.

c. Draw an arrow (or more if needed) to show the transfer of electrons to the new element.

d. Write the resulting chemical formula.

e. Write the electron configurations for each ion that is formed. Ex. H1+ = 1s0

2. Solid sodium chloride and solid magnesium oxide are both held together by ionic (electrovalent) bonds. a. Using s,p and d notation write down the symbol for and the electronic configuration of (i) a sodium ion; (ii) a chloride ion; (iii) a magnesium ion; (iv) an oxide ion.

b. Explain what holds sodium and chloride ions together in the solid crystal

c. Sodium chloride melts at 1074 K; magnesium oxide melts at 3125 K. Both have identical structures. Why is there such a difference in their melting points?

3.2.2. Physical properties of ionic compounds

Activity 3.3(a)

Determination of Relative Melting Point of different substances

Procedure:

1. Cut a square of aluminum foil that is about5 by 5 cm

2. Set up a ring stand with an iron ring attached.

3. Place the aluminum square on the iron ring, as shown at right in Figure 3.6

4. Obtain a small pea-sized sample of NaCl. Place the sample on the aluminum foil, about 5cm from the center of the square.

5. Obtain a small pea-sized sample of table sugar. Place the sample on the aluminum foil, about 1 cm from the center of the square, but in the opposite direction from the salt.

6. Your square of aluminum foil should look like in Figure 3.7.

7. Light the Bunsen burner and adjust the flame height so that the tip of the flame is just a cm or so below the height of the aluminum foil.

8. Observe as the two compounds heat up.

9. Set up another sheet of aluminum foil and determine the relative melting points (low vs. high) of the four unknowns.

10. Record your results in the table 3.1 belowCaution: if the compounds burn with sparks do not panic.

Study question:

1. Which compound melts first?

2. Giving reasons compare the melting points of the two compounds.

Table: 3.1 relative melting points of different substances

Conclusions:

The melting points of ionic compounds are higher than those of covalent compounds; this is due to strong electrostatic forces between opposite charges in the ionic substances compared to the week forces of attraction between molecules in covalent substances . This also explains why all ionic compounds are solid at room temperature.

Activity 3.3(b)

Conductivity in Solution

Procedure:

1. Dissolve a spoonful of NaCl in water.

2. Connect the apparatus as shown in figure 3.8

3. Make an observation and record your results as in table 3.2 below

4. Repeat the procedure 1 to 3 above using sugar solution, ethanol and copper(II) sulfate solution

5. Record your results in the table below.

Table 3.2: relative conductivity of different substances

Study questions:

1. Give reasons for your observations above.

2. Solid sodium chloride does not conduct electricity whereas an aqueous solutionof sodium chloride does. Explain

Conclusion:

Based on our tests with salt and sugar, the ability to conduct electricity in solution of ionic compounds is much higher than in covalent compounds

Activity 3.3(c) Solubility test

Procedure:

1. Using forceps, place 5-8 crystals of each of sodium chloride, magnesium chloride, copper sulphate, calcium carbonate, copper carbonate, sodium sulphate (a small pinch) of the compound into one of the test tubes in test tube rack.

2. Half-fill the test tube with distilled water and stir with a clean stirring rod.

3. Observe if the crystals dissolve in water.

4. Record your findings in a suitable table.

Table 3.3: Solubility of different substances

Study questions:

1. Classify substances as soluble and slightly soluble in water.

2. Giving reasons, suggest an explanation for your observation.

Conclusion:

Water is a good solvent for many ionic compounds but not a solvent for covalent compounds, apart few exceptions (you will learn about later on).

Shattering: Why are Ionic compounds generally hard, but brittle?

It takes a large amount of mechanical force, such as striking a crystal with a hammer, to force one layer of ions to shift relative to its neighbour. However, when that happens, it brings ions of the same charge next to each other (Figure3.9). The repulsive forces between like-charged ions cause the crystal to shatter. When an ionic crystal breaks, it tends to do so along smooth planes because of the regular arrangement of the ions.

Checking Up 3.3

1. The diagrams below show the electric conductivity of distilled water, solid sodium chloride and a solution of sodium chloride respectively. Use the diagrams to explain the observations from the set up.

i) no light is given out by bulb in A

ii) no light is given out by bulb in B

iii) light is given out in C

2.Why are ionic compounds brittle?

3.Why do ionic compounds have high melting points?

4.What happens when an electric current is passed through a solution of an ionic compound?

3.2.3. Lattice energy

Activity 3.4

By using information in this student’s chemistry book and other books from the school library, attempt to answer the following questions.

1. Define lattice energy

2. Explain how the lattice energy is used to describe high melting points of ionic compounds.

3. What is the bonding force present in ionic compounds?

4. Why is the melting temperature of magnesium oxide higher than that of magnesium chloride, even though both are almost 100% ionic?

5. How is lattice energy of ionic compounds related to their high melting points?

It is a type of potential energy that may be defined in two ways. In one definition, the lattice energy is the energy required to break apart an ionic solid and convert its component ions into gaseous ions (Endothermic process). On the other hand lattice energy is the energy released when gaseous ions bind to form an ionic solid (Exothermic process). Its values are usually expressed with the units’ kJ/mol.

Lattice Energy is used to explain the stability of ionic solids. Some might expect such an ordered structure to be less stable because the entropy of the system would be low. However, the crystalline structure allows each ion to interact with multiple oppositely charge ions, which causes a highly favourable change in the enthalpy of the system. A lot of energy is released as the oppositely charged ions interact. It is this that causes ionic solids to have such high melting and boiling points. Some require such high temperatures that they decompose before they can reach a melting and/or boiling point.

Checking Up 3.4

Define lattice energy

Which one of the following has the greatest lattice energy?

MgO or NaCl

LiCl or MgCl2

Which one of the following has the greatest Lattice Energy?

NaCl or CaCl2

AlCl3 or KCl

How do the ionic radius and ion charge affect the lattice energy of an ionic substance?

Factors affecting lattice enthalpy

Activity 3.5

1. Basing on the information obtained from the definition of lattice energy, suggest the factors that affect the lattice energy.

2. Which has the larger lattice energy: NaCl or CsI?

Basing on the information obtained from the definition of lattice energy, suggest the factors that affect the lattice energy.

Which has the larger lattice energy: NaCl or CsI?

There are two main factors that affect lattice enthalpy.

a) The charges on the ions

Sodium chloride and magnesium oxide have exactly the same arrangements of ions in the crystal lattice, but the lattice enthalpies are very different.

From the above diagram the lattice enthalpy of magnesium oxide is much greater than that of sodium chloride. This is because in magnesium oxide, +2 ions are attracting -2 ions; in sodium chloride, the attraction is only between +1 and - 1ions.

b. The radius or the size (volume) of the ions

The lattice enthalpy of magnesium oxide is also increased relative to sodium chloride because magnesium ions are smaller than sodium ions, and oxide ions are smaller than chloride ions.

It means that the ions are closer together in the lattice, and that increases the strength of the attractions.

For example, as you go down Group 17 of the Periodic Table from fluorine to iodine, you would expect the lattice enthalpies of their sodium salts to fall as the negative ions get bigger - and that is the case:

As the negative ion gets bigger in size, the distance between the centres of oppositely charged ions increase and attractions so decrease. For instance as you go down group 1 halides, the enthalpy decreases.

3.3. Formation of metallic bonds and physical properties of metals

Activity 3.6

The figure above shows materials commonly used at home. If you reflect back around your house/home you will see hundreds of objects made from different kinds of materials.

1. Observe the objects (in picture) and classify them according to the materials they are made of.

2. Have you ever wondered why the manufacturers choose the material they did for each item?

3. Why are frying saucepans made of metals and dishes, cups and plates often made of glass and ceramic?

4. Could dishes be made of metal? And saucepans made of ceramic and glass

3.3.1. Formation of metallic bond

Another way of combining is the combination between metal atoms to form metallic bond.

When metal atoms combine together, there is no transfer of electrons since the combining atoms are of the same nature, i.e. all are metals and no one is ready to give up or to capture electrons.In metallic bonding, all metal atoms put together their valence electrons in a kind of pool of electrons where positive metallic cations seem to bathe. This model is called “Elecron Sea Model” (Fig. 3.15)

Metals have a sea of delocalized electrons within their structure. These electrons have become detached and the remaining atoms have a positive charge. This positive charged is attracted to the delocalized sea of electrons due to electrostatic forces of attraction (forces which result from unlike charges), and as a result has a strong interaction. It is this interaction which makes the metals so hard and rigid. Figures 3.13 and 3.14 are representation of metallic bond.

3.3.2. Physical properties of metals

Activity 3.7: Looking at metals

1. Collect a number of metal items from your home or school. Some exam-ples are listed below: hammer, electrical wiring, cooking pots, jewellery, burglar bars and coins, nails,

2. What is the function of each of these objects?

3. Discuss why you think metal was used to make each object. You should consider the properties of metals when you are answering this question.

a. Electrical conductivity

Activity 3.8

Procedure

1. Take a dry cell/battery, a torch bulb/ bulb, connecting wires, crocodile clips and connect them. As in the figure 3.17

2. Repeat the experiment above using different metals

3. Record your results in a suitable table.

Study questions:

1. Compare the relative conductivity of the metals used in the above experi-ment.

2. Suggest the purpose of the resistor in the experimental set up.

Due to the mobile valence electrons of metals, electricity can pass through the metals easily. So they are conductors of electricity. Silver and copper are the best conductors of electricity.

Note: mercury is a poor conductor of electricity.Thermal conductivity

Activity 3.9

Experiment to demonstrate the ability of different substances to conduct heat

Apparatus/apparatus: two cups (made from the same material e.g. plastic); a metal

Procedure:

1. Pour boiling water into the two cups so that they are about half full.

2. At the same time, place a metal spoon into one cup and a plastic spoon in the other.

3. Note which spoon heats up more quickly.

4. Record your observations.

Study questions:

1. Which one heats faster plastic spoon or metallic spoon and why?

2. Why do we use plastic cups?

3. Why are cooking pots made of metallic materials not plastics?

Results: The metal spoon heats up more quickly than the plastic spoon. In other words, the metal conducts heat well, but the plastic does not.

Conclusion: Metals are good thermal conductors, while plastic is a poor thermal conductor. The reason is due to the mobility of electrons with transfer of kinetic energy between electrons. This explains why cooking pots are metallic, but their handles are often plastic or wooden. The pot itself must be metal so that heat from the cooking surface can heat up the pot to cook the food inside it, but the handle is made from a poor thermal conductor so that the heat does not burn the hand of the person who is cooking.

c. Malleability and ductility

Activity 3.10

Experiments to demonstrate the malleability and ductility of metals

Materials: wires, nails, hammer, piece of cloth.

Procedure:

1. Wrap the material to be test in a heavy plastic or cloth to avoid pieces flying from the material.

2. Place the material on a flat hard surface

3. Use a harmer to pound the material flat

4. Record your observations as malleable or non-malleable.

Metals can have their shapes changed relatively easily in two different ways i.e.

Malleable: can be hammered into sheets or

Ductile: can be drawn into rods and wires

As the metal is beaten into another shape the delocalised electron cloud continues to bind the “ions” together.

The positive ions in the lattice are all identical. So the planes of ions can slide easily over one another attractive forces in the lattice are the same whichever ions are adjacent and remain the same throughout the lattice

Some metals, such as gold, can be hammered into sheets thin enough to be translucent.

Note: If a material is pounded into a flat shape it is called malleable but if it breaks or does not change it is called non-malleable.

d) Metal appears hiny/lustrous

Activity 3.11

Demonstration of shininess in metals

Procedure:

1. Hold a small piece of sodium metal using forceps.

2. Place it on a hard surface and cut it into two parts

3. Observe the cut surface. What do you observe?

4. Look at the surface of aluminium sheets, how does it appear?

Study question:

Explain what makes metal surfaces appear shiny/luster.

Light is composed of very small packages of electromagnetic energy called photons. We are able to see objects because light photons from the sun (or other light source) reflect off of the atoms within the object and some of these reflected photons reach the light sensors in our eyes and we can see the objects.

When photons of light hit the atoms within an object three things can happen:

The photons can bounce back from the atoms in the object, can pass through an object such as glass or can be stopped by the atoms within the object.Objects that reflect many photons into our eyes make the objects appear shiny. Objects that absorb photons and reflect less photons appear dull or even dark black to our eyes.

Did you know? Of all of the metals, aluminium and silver are the shiniest to our eyes. Gold is also one of the more shiny metals. However, gold is not as shiny as silver and aluminium. Mercury, a liquid metal, is also shiny and special telescope mirrors have been made of mercury.

(e)Melting and boiling points

Activity 3.12:

1. Why do metals have variable melting points?

2. Why do metals have high melting points compared to non-metals?

Melting point is a measure of how easy it is to separate individual particles. In metals it is a measure of how strong the electron cloud holds the positive ions. The ease of separation of ions depends on the electron density and Ionic / Atomic size.

Melting point increases across the period (from Na---Al), the electron cloud density increases due to the greater number of valence electrons contributed per atom. As a result the ions are held more strongly. Hence more energy will be needed to separate the ions as shown in the table 3.4.

Table 3.4: Valence electron density as a factor that influence melting and boil-ing points of metal

Conversely, as the size of the atom/ion increases down the group the melting and boiling points decrease as shown in the table 3.5

Table 3.5. Atomic/ionic size as a factor that influences melting and boiling points of metals

(*) Sevenair & Burkett(1997), Introductory Chemistry, Investigating the Molecular Nature of Matter; WCB.

N.B: You may find, in another book, different values of atomic/ionic sizes, due to different methods used for the determination the sizes.

3.3.3. Factors affecting the strength of metallic bonds.

Activity 3.13

1. Explain different strengths of metallic bonds in different metals?

2. Compare the metallic strength of the following metals:

(i) Sodium and magnesium

(ii) Sodium and potassium

The three main factors that affect the metallic bond are:

•Number of protons/ Strength of nuclear attraction: The more protons the stronger the force of attraction between the positive ions and the delocalized electrons

•Number of delocalized electrons per atom: The more delocalized electrons the stronger the force of attraction between the positive ions and the delocalized electrons

Size of atom: The smaller the atom, the stronger the force of attraction between the positive ions and the delocalized electrons and vice-versa, the larger the atom, the weaker the force of attraction between the positive ions and the delocalised electrons.

The strength increases across a period from left to right because:

The atoms have more protons. There are more delocalized electrons per atom.Electrons are added to the same energy level. Group1 elements have 1 electron in their outer shells and so contribute 1 electron to the sea of electrons, Group 2 elements contribute 2 electrons per atom, and Group 3 elements contribute 3 electrons per atom

If the atoms/ions are smaller; there is therefore a greater force of attraction between the positive ions and the delocalized electrons.

In group 1 elements, the melting and boiling points decrease as the size increases hence attraction between the delocalized electrons and metal cations decreases down the group as shown table 3.6

Table 3.6: Variation of melting and boiling points of group 1 elements

Checking Up 3.6

1. Look at the table below, which shows the thermal conductivity of a number of different materials, and then answer the questions that follow:

The higher the number in the second column, the better the material is at conducting heat (i.e. it is a good thermal conductor). Remember that a material that conducts heat efficiently will also lose heat more quickly than an insulating material. Use this information to answer the following questions:

1. Name two materials that are good thermal conductors.

2. Name two materials that are good insulators.

3. Explain why:

a. cooler boxes are often made of polystyrene

b. Homes that are made from wood need less internal heating during the cold months.

c. Igloos (homes made from snow) are so good at maintaining warm temperatures, even in freezing conditions.

d. Houses covered by iron sheets and houses covered by tiles can be compared in their capacity of keeping the interior of the house hot of fresh during a sunny and hot day.

4. Magnesium has a higher melting and boiling point than sodium. This can be explained in terms of the electronic structures, the packing, and the atomic radii of the two elements.

a. Explain why each of these three things causes the magnesium melting and boiling points to be higher. b. Explain why metals are good conductors of electricity.

c. Explain why metals are also good conductors of heat.

5. Pure metals are usually malleable and ductile.

a. Explain what those two words mean. If a metal is subjected to a small stress, it will return to its original shape when the stress is removed. However, when it is subjected to a larger stress, it may change shape permanently. Explain, with the help of simple diagrams why there is a different result depending on the size of the stress.When a piece of metal is worked by a blacksmith, it is heated to a high temperature in a furnace to make it easier to shape. After working it with a hammer, it needs to be re-heated because it becomes too difficult to work. Explain what is going on in terms of the structure of the metal. Why is brass harder than either of its component metals, copper and zinc?

3.4. End Unit Assessment

1. Choose from a list of words and fill in the missing words in the text below:

List of words:

Conduct electricity, electrodes, electrolysis, electrostatic attraction, free electrons, good conductivity, great malleability, high density, high melting points, ionic bond,metal, negative ion, non-metal, positive ion,regular crystal shape and attractive forces.

Te x t :

Metals have a layered structure of ................... in fixed positions but between them are oppositely charged ............... that can move around at random between the metal atoms. There is a strong ...................................... between these oppositely charged particles which gives them ..........................The strong forces also give a ....................... making the average ........... heavier than an average ................ The presence of ........................... in the structure keeps the bonding intact when metals are bent or hammered giving them.......................................... Also, these ....................give metals ..............as regards heat and electricity.When electrons are transferred from (usually) ............ atom (e.g. sodium) to................. atom (e.g. chlorine) an ionic bond is formed. Sodium loses an electron to form a singly charged ........................... and chlorine gains an electron to form a singly charged negative ion. In an ionic compound, the ionic bond is the electrostatic attraction between the neighbouring positive ions and negative ions.

The strong forces holding this giant ionic lattice together give these ionic compounds............................ and...................................................................When ionic compounds are melted they are found to ................ in a process called ...........using electrical contacts called............ In this process, move to the negative electrode (cathode) and metalsare released. At the same time, ....................move to the positive electrode (anode) and ................ are formed. Research from the internet or text books to find out other physical properties of metals and ionic compounds that are not mentioned above.

Answer these questions by choosing the best alternative repre-sented by letters from A, B, C and D.

1. Metals lose electrons from their lattice to become

a. positive ions

b. negative ions

c. alkalis

d. non- metals

2. Neither ions nor electrons are free to move in

a. liquids

b. metals

c. ionic solids

d. All of the above

3. Attractive forces between metal ions and delocalized electrons can be weakened or overcome by

a. hammer

b. high temperature

c. water

d. All of the above

4. Metals are good conductors due to

a. ionic lattices

b. crystalline lumps

c. mostly solids

d. delocalized electrons

5. Most atoms adopt one of three simple strategies to achieve a filled shell. Which of the following is NOT one of these strategies?

a. They accept electrons

b. They share electrons

c. They give away electrons

d. They keep their own electrons

6. Which of the following is NOT a type of chemical bond?

a. Covalent

b. Metallic

c. Valence

d. Ionic

7. In metallic bonding...

a . One atom takes the outer shell electrons from another atom.

b. A couple of atoms share their electrons with each other.

c. Some electrons are shared by all the atoms in the material.

d. Bonding takes place between positively charged areas of one atom with a negatively charged area of another atom.

8. Which of the following is NOT a characteristic of metals?

A. Shiny /lustre

B. Brittle/Shatters easily

C. Conducts electricity

D. Malleable

9. When two or more metal elements are combined they form an...

a. bronze

b. alloy

c. Covalent bond

d. Brass

10. Sulphur is a solid non-metallic element at room tempera-ture, so it is?

a. A good conductor of heat

b. A substance with a low melting point

c. Easily bent into shape

d. A good conductor of electricity

11. Copper is a metallic element so it is likely to be a?

a. substance with a low boiling point

b. poor conductor of electricity

c. good conductor of heat

d. substance with a low melting point

12. Sodium chloride is a typical ionic compound formed by combining a metal with a non-metal. Sodium chloride will?

a. have a low melting point

b. consist of small NaCl molecules

c. conduct electricity when dissolved in water

d. not conduct electricity when molten

13. Copper is a metallic element so it is likely to be a?

a. Substance with a shiny surface

b. Poor conductor of electricity

c. Poor conductor of heat

d. Substance with a low melting point

14. When an ionic bond is formed between atoms of different elements?

a. Protons are transferred

b. Electrons are transferred

c. Protons are shared

d. Electrons are shared15. Sodium chloride has a high melting point because it has:

a. Many ions strongly attracted together

b. Strong covalent double bonds

c. A giant covalent 3-dimentional structure

d. Molecules packed tightly together

16. Which substance is likely to have a giant ionic structure:

a. Melts at 1400oC, insoluble in water, good conductor of electricity either when solid or molten

b. Melts at 2800oC, insoluble in water, non-conductor of electricity when molten or solid

c. Melts at 17oC, insoluble in water, non-conductor of electricity either when solid or molten

d. Melts at 2600oC, dissolves in water, non-conductor of electricity when solid,undergoes electrolysis in aqueous solution

17. Sodium chloride conducts electricity when:

a. Solid or molten

b. Solid or in solution

c. Molten or in solution

d. Non of the above

18. The structure of magnesium oxide is a

a. Giant covalent lattice

b. Giant ionic lattice

c. Simple ionic lattice

d. All the above19. What is the formula for magnesium chloride (contains Mg2+ and Cl? ions)?a. MgClb. Mg22Clc. MgCl2d. MgCl320. Why does sodium chloride have a lower melting point than magnesium chloride?

a. Its positive ions are smaller and have a smaller charge

b. Its positive ions are larger but have a smaller charge

c. Its positive ions are smaller but have a larger charge

d. All the above

21. Explain the conductivity of sodium chloride

a. It conducts electricity when molten because it contains free electrons

b. It conducts electricity when molten because sodium has metallic bonding

c. It conducts electricity when molten because its ions are free to move.

d. None of the above

Short and long answer questions

22.(a) Explain why the lattice dissociation enthalpy of NaBr is a bit less than that of NaCl.

(b) Explain why the lattice dissociation enthalpy of MgO is about 5 times greater than that of

NaCl

23.a) The table (using figures for lattice energies from gives experimental and theoretical values for the silver halides.(The values are listed as lattice dissociation energies.) compare the values and give a detailed explanation.

(b) For AgF, the experimental and theoretical values are very close. What does that show?

Key Unit Competence

Demonstrate how the nature of the bonding is related to the properties of covalent compounds and molecular structures.

Learning objectives

By the end of this unit, I will be able to:

•Define octet rule as applied to covalent compounds.

•Explain the formation of covalent bonds and describe the properties of covalent compounds. •Describe how the properties of covalent compounds depend on their bonding.

•Explain the rules of writing proper Lewis structures

•Draw different Lewis structures

•State the difference between Lewis structures from other structures.

•Apply octet rule to draw Lewis structures of different compounds.

•Make the structures of molecules using models.

•Write the structures of some compounds that do not obey octet rule.

•Explain the formation of dative covalent bonds in different molecules.

•Compare the formation of dative covalent to normal covalent bonding.

•Describe the concept of valence bond theory.

•Relate the shapes of molecules to the type of hybridization.

•Differentiate sigma from pi bonds in terms of orbital overlap and formation.

•Explain the VSEPR theory.•Apply the VSEPR theory to predict the shapes of different molecules/ions.

•Predict whether the bonding between specified elements will be primarily covalent or ionic.

•Relate the structure of simple and giant molecular covalent compounds to their properties.•Describe simple and giant covalent molecular structures.

•Describe the origin of inter-molecular forces.

•Describe the effect of inter and intra molecular forces on the physical properties of certain molecules. •Describe the effect of hydrogen bonding in the biological molecules.

•Relate the physical properties to type of inter and intra molecular forces in molecules.

•Compare inter and intra molecular forces of attraction in different molecules.

4.1. Overlap of atomic orbitals to form covalent bonds

Activity 4.1

1. Using dot and cross diagrams, show how a covalent bond is formed in:

(a) Hydrogen molecule (H2)

(b) Hydrogenchloride molecule (HCl)

(c) Chlorine molecule (Cl2).

In UNIT 3, you have learnt that atoms have different ways of combination to achieve the stable octet electronic structure; two of those ways of combination led to the formation of ionic bond and metallic bond. But what happens where the two combining atoms need electrons to complete the octet structure and no one is willing to donate electrons? For example the combination of 2 hydrogen atoms or the combination of 2 chlorine atoms?

When this happens, the combining atoms share a pair of electrons where each atom brings or contributes one electron. In other words there is an overlapping of two orbitals, one orbital from one atom, each orbital containing one electron (see Fig.4.1): this bond is called “Covalent bond”. The attraction between the bonding pair of electrons and the two nuclei holds the two atoms together.

The covalent bond is a bond formed when atoms share a pair of electrons to complete the octet. Similarly, people need each other irrespective of their race, economic, political and social status for the success of human race. Some compounds that exist in nature such as hemoglobin in our blood, chlorophyll in plants, paracetamol,

esponsible for transport of oxygen, green color in plants and as pain killer respectively are made of the covalent bond. The covalent bonds mostly occur between non-metals or between two of the same (or similar) elements.Two atoms with similar electronegativity do not exchange an electron from their outermost shell; the atoms instead share electrons so that their valence electron shell is filled.

In general, covalent bonding occurs when atoms share electrons (Lewis model), concentrating electron density between nuclei. The build-up of electron density between two nuclei occurs when a valence atomic orbital of one atom combines with that of another atom (Valence bond theory).In Valence bond theory, the bonds are considered to form from the overlap of two atomic orbitals on different atoms, each orbital containing a single electron.

The orbitals share a region of space, i.e. they overlap. The overlap of orbitals allows two electrons of opposite spin to share the common space between the nuclei, forming a covalent bond.

As example, we can consider the case of hydrogen molecule. The molecule of hydrogen is made of two hydrogen atoms bonded together, this is made by the two spherical 1s orbitals overlap, and contains two electrons with opposite spins (Figure 4.1)

These two electrons are attracted to the positive charge of both the hydrogen nuclei, with the result that they serve as a sort of ‘chemical glue’ holding the two nuclei together.

The figure (Figure 4.1) shows the distance between the two nuclei. If the two nuclei are far apart, their respective 1s-orbitals cannot overlap and no covalent bond is formed. As they move closer each other, the orbital overlapping begins to occur, and a bond starts to form.

The examples below represent different atoms overlapping in order to form covalent bonds.

Example 1: Water formation

Water molecule is formed by two atoms of hydrogen and one atom of oxygen, its formula is H2O.

The steps of water formation are indicated in the Figure 4.2

The oxygen atom has 6 electrons in the valence shell and needs two electrons to complete its outer shell. Similarly, each hydrogen atom has one electron valenceand needs one electron to complete its outer shell. Therefore, oxygen can only share 2 of its 6 electron valence otherwise it will exceed the maximum of 8; the two lone pairs remain.

Example 2: Formation of hydrogen chloride (Figure 4.3)

The molecule of hydrogen chloride is made by covalent bond between one chloride atom and a one hydrogen atom as shown below.

The chlorine atom has 7 electrons in the valence shell and needs one electron to complete its outer shell. Similarly, the hydrogen atom has one electron valence and needs one electron to complete its outer shell. The two atoms share one single electron in their outer shell to make a covalent bond in order to become stable.

Example 3: Formation of ammonia (NH3) (Figure 4.4)

The ammonia is a compound formed by one atom of nitrogen and three atoms of hydrogen as indicated below.

•Nitrogen atom needs three electrons to complete its outer shell and each hydrogen atom needs one electron to complete its outer shell. Then, nitrogen can only share 3 of its 5 electrons in order not to exceed the maximum of 8. A lone pair remains.

4.1.1 Properties of covalent molecules

Covalent molecules are chemical compounds in which atoms are all bonded together through covalent bonds. The covalent compounds possess different properties and some are emphasized below.•Covalent compounds exist as individual molecules, held together by weak van der Waals forces.

•Due to the weak van der Waals forces that hold molecules together, covalent compounds have low melting and boiling points; because the weak forces between molecules can be broken easily to separate the molecules. That is why covalent compounds can be solid, liquid and gaseous at room temperature.

•Covalent compounds do not display the electrical conductivity either in pure form or when dissolved in water. This can be explained by the fact that the covalent compounds do not dissociate into ions when dissolves in water.

•Generally non-polar covalent compounds do not dissolve in water; but

many polar covalent compounds are soluble in water( a polar solvent)

Non-polar covalent compounds are soluble in organic solvents (themselves non-polar covalent).

The two statements above are at the origin of the say by chemists: “Like dissolves like”.

Checking up 4.1

Using dot and cross diagram draw the diagrams to show the formation of bonds in the following molecules:

1 a) Ammonia (NH3) (b) carbon dioxide gas ( CO2) (c) Nitrogen molecule (N2) (d) tetra chloromethane (CCl4)

2. Explain the properties of covalent molecules compared to ionic compounds.

4.2. Theories on the formation of covalent bond

4.2.1. Lewis theory and structures

Activity 4.2(a)

1. Draw the diagrams indicating only the valence electrons of the following :Chlorine molecule (Cl2), Carbon atom (C), Phosphorus atom (P)

2. Draw the diagram to show how all electrons are shared in a molecule of

(i) NH3 indicating all unshared electrons.

(ii) HCl

(iii) N2

3. Identify the common feature possessed by the diagrams drawn above in 2.

4. Suggest the name given by those diagrams drawn above.

The Lewis structure is a representation of covalent molecules or polyatomic ions where all valence electrons are distributed to the bonded atoms as shared electron pair (bond pairs) or unshared electron (lone pair).We then combine electrons to form covalent bonds until we come up with a Lewis structure in which all of the elements (with the exception of the hydrogen atoms) have an octet of valence electrons.

The representation of Lewis structures follows different steps. Let us use the nitrate ion (NO3-) as a typical example. Determine the total number of valence electrons in a molecule or ion.

1) Determine the total number of valence electrons in a molecule or ion.

2) Draw a skeleton for the molecule or ion which connects all atoms using only single bonds. In simple molecules, the atom with the most available sites for bonding is usually placed at the center. The number of bonding sites is determined by considering the number of valence electrons and the ability of an atom to expand its octet. As you will progress in your study of chemistry, you will be able to recognise that certain groups of atoms prefer to bond together in a certain way!

(3) Of the 24 valence electrons in NO3-, 6 are required to make the skeleton. Consider the remaining 18 electrons and place them so as to fill the octets of as many atoms as possible (start with the most electronegative atoms first then proceed to the more electropositive atoms).

(4) Are the octets of all the atoms fulfilled? If not then fill the remaining octets by making multiple bonds (make a lone pair of electrons, located on a more electronegative atom, into a bonding pair of electrons that is shared with the atom that is electron deficient).

Check that you have the lowest formal charge (F.C.) possible for all the atoms, without violating the octet rule; F.C. = (valence e-) - (1/2 bonding e-) - (lone electrons).

Thus the Lewis structure of NO3- ion can be written in the following ways:

The 4 steps are used to find Lewis structures. If the octets are incomplete, and more electrons remain to be shared, move one electron per bond per atom to make another bond. Note that in some structures there will be empty orbitals (example: the B of BF3), or atoms which have ten electrons (example: P in PF5) or twelve electrons (example: S in SF6).

The first step in this process involves the calculation of the number of valence electrons in the molecule or ion. In the case of a neutral molecule, this is nothing more than the sum of the valence electrons on each atom. If the molecule carries an electric charge, we add one electron for each negative charge or subtract an electron for each positive charge.

In Lewis structure, the least electronegative element is usually the central element, except H that is never the central element, because it forms only one bond.

Another way of finding Lewis structure

1. Calculate n (the number of valence (outer) shell electrons needed by all atoms in the molecule or ion to achieve noble gas configurations for instance,NO3-, n=1× 8(for N atom) + 3×8 (for O atom) = 32 electrons.

2. Calculate A, number of electrons available in the valence (outer) shells of all the atoms. For negatively charged ions, add to this number the number of electrons equal to the charge of the anions. For cations you subtract the number of electrons equal to the charge on the cation.

For instance: NO3-,

A= 1×5(for N) +3×6 (for O atom) +1(for -1 charge) = 5+18+1=24 electrons.

3. Calculate S, total number of electrons shared in the molecule or ion, using the relationship

S = n-A

S= n-A= 32-24 =8 electrons shared (4pairs of electron shared)

4. Place S electrons into the skeleton as shared pair’s. Use double and triple bonds only when necessary. Lewis formulas may be shown as either dot formula or dash formulas.

5. Place the additional electrons into the skeleton as unshared (lone) pairs to fill the octet of every A group element (except H, which can share only 2 electrons) check that the total number of valence electrons is equal to A from step 2.

Check that you have the lowest formal charges (f.c.) possible for all the atoms, without violating the octet rule; F.C. = (valence e-) - (1/2 bonding e-) - (lone electrons).

Thus the Lewis structure of NO3- ion can be written in the following ways:

Checking up 4.2(a)

1 .What do you consider when drawing Lewis structures of different molecules?

2. Draw the Lewis diagram for the following compounds(a)PCl3 (b) H2S (c) CH3Cl (d) C2H2

3.What is meant by the term Lewis structure?

4. Using water differentiates between Lewis structures from other structures.

5. Deduce the importance of Lewis structures in covalent molecules.What critics can you make to those strcutures?

Exceptions to the octet rule

Activity 4.2(b)

1. Draw the Lewis structure of aluminium chloride (AlCl3)?

2. How many electrons surround the aluminium atom in the centre?

3. Does it obey the octet rule? And why?

There are three general ways in which the octet rule doesn’t work:

•Molecules with an odd number of electrons

•Molecules in which an atom has less than an octet

•Molecules in which an atom has more than an octet

a. Odd number of electrons

Consider the example of the Lewis structure for the molecule nitrous oxide (NO):Total electrons: 6+5=11

Bonding structure:

There are currently 5 valence electrons around the nitrogen. A double bond would place 7 around the nitrogen, and a triple bond would place 9 around the nitrogen. We appear unable to get an octet around each atom.

1. Draw the diagram to show how the single and double bonds are formed ini)Oxygen molecule (O2) (ii) nitrogen molecule (N2) (iii) ethene molecule(C2H4)

Key unit competency

Use atomic structure and electronic configuration to explain the trends in the physical properties of elements.

Learning objectives

By the end of this unit, I will be able to:

•Outline the historical back ground of the Periodic Table.

•Explain the trends in the physical properties of the elements across a period and down a group.

•Classify the elements into respective groups and periods using electronic configuration.

•Relate trends in physical properties of the elements to their electronic configuration.

•Classify the elements into blocks (s, p, d, f-block).

5.1. Historical Background of the Periodic Table

Activity 5.1

Who is the father of the periodic table? Explain your answers.

Differentiate the laws of triads and octaves

149ChemistrySenior Four Student's Book5.1. Historical Background of the Periodic TableActivity 5.1Who is the father of the periodic table? Explain your answers.Differentiate the laws of triads and octavesDuring the nineteenth century, many scientists contributed to the development of the periodic table. In the beginning, a necessary prerequisite to the construction of the periodic table was the discovery of the individual elements. Although elements such as gold, silver, tin, copper, lead and mercury have been known since antiquity, the first scientific discovery of an element occurred in 1649 when Hennig Brand discovered phosphorous. The periodic table of elements is a chart created in order to help to organize the elements that had been discovered at that time. By 1869, a total of 63 elements had been discovered. As the number of known elements grew, scientists began to recognize patterns in properties and began to develop classification schemes.

Some important dates help us to understand more about how the periodic table has been created.

5.1.1. Law of Triads

In 1817 Johann Dobereiner noticed that the atomic weight of strontium fell midway between the weights of calcium and barium, elements possessing similar chemical properties.

n 1829, after discovering the halogen triad (three) composed of chlorine, bromine, and iodine and the alkali metal triad of lithium, sodium and potassium he proposed that nature contained triads of elements the middle element had properties that were an average of the other two members when ordered by the atomic weight (the Law of Triads).

Between 1829 and 1858 a number of scientists (Jean Baptiste Dumas, Leopold Gmelin, Ernst Lenssen, Max von Pettenkofer, and J.P. Cooke) found that these types of chemical relationships extended beyond the triad. During this time fluorine was added to the halogen group; oxygen, sulfur, selenium and tellurium were grouped into a family while nitrogen, phosphorus, arsenic, antimony, and bismuth were classified as another. Unfortunately, research in this area was hindered by the fact that accurate values were not always available.

5.1.2. Law of Octaves

In 1863, John Newlands, an English chemist suggested that elements be arranged in “octaves”. He wrote a paper in which he classified the 56 established elements into 11 groups based on similar physical properties, noting that many pairs of similar elements existed which differed by some multiple of eight in atomic weight. This law stated that any given element will exhibit analogous behavior to the eighth element following it in the table. However, his law of octaves failed beyond the element calcium.

In 1669, Hennig Brand a German merchant and amateur alchemist invented the Philosopher’s Stone; an object that supposedly could turn metals into pure gold. He heated residues from boiled urine, and a liquid dropped out and burst into flames. He also discovered phosphorus.

•In 1680Robert Boyle also discovered phosphorus without knowing about Henning Brand’ discovery.

•In 1809, curiously 47 elements were discovered and named, and scientists began to design their atomic structures based on their characteristics.

•In 1869, Dimitri Mendeleev based on John Newlands’ ideas started the development of elements organized into the periodic table. The arrangement of chemical elements were done by using atomic mass as the key characteristic to decide where each element belonged in his table. The elements were arranged in rows and columns. He predicted the discovery of other elements, and left spaces open in his periodic table for them. At the same time, Lothar Meyer published his own periodic table with elements organized by increasing atomic mass.

•In 1886, French physicist Antoine Becquerel first discovered radioactivity.During the same period of 1886, Ernest Rutherford named three types of radiation; alpha, beta and gamma rays.

•In 1886, Marie and Pierre Curie started working on the radioactivity and they discovered radium and polonium. They discovered that beta particles were negatively charged.

•In 1895, Lord Rayleigh discovered a new gaseous element named argon which proved to be chemically inert. This element did not fit any of the known periodic groups.

•In 1898, William Ramsay suggested that argon be placed into the periodic table between chlorine and potassium in a family with helium, despite the fact that argon’s atomic weight was greater than that of potassium. This group was termed the “zero” group due to the zero valency of the elements. Ramsey accurately predicted the future discovery and properties neon.

•In 1913, Henry Moseley worked on X-rays and determined the actual nuclear charge (atomic number) of the elements. He has rearranged the elements in order of increasing atomic number.

•In 1897 English physicist J. J. Thomson discovered small negatively charged particles in an atom and named them as electrons;John Sealy Townsend and Robert A. Millikan investigated the electrons and determined their exact charge and mass.

•In 1900, Antoine Becquerel discovered that electrons and beta particles as identified by the Curies are the same thing.

•In 1903,Ernest Rutherford proclaimed that radioactivity is initiated by the atoms which are broken down.

•In 1911, Ernest Rutherford and Hans Geiger discovered that electrons are moving around the nucleus of an atom.

•In 1913, Niels Bohr suggested that electrons move around a nucleus in discreete energy levels called orbits. He observed also that light is emitted or absorbed when electrons transit from one orbit to another.

•In 1914,Rutherford identified protons in the atomic nucleus. He also transformed a nitrogen atom into an oxygen atom for the first time. English physicist Henry Moseley provided atomic numbers, based on the number of electrons in an atom, rather than based on atomic mass.

•In 1932James Chadwick discovered neutrons, and isotopes were identified. This was the complete basis for the periodic table. In that same year Englishman Cockroft and the Irishman Walton first split an atom by bombarding lithium in a particle accelerator, changing it to two helium nuclei.The last major changes to the periodic table give rise from Glenn Seaborg’s work in the middle of the 20th Century. In 1940, he discovered plutonium and all the transuranic elements from 94 to 102.

•In 1944, Glenn T. Seaborg discovered 10 new elements and moved out 14 elements of the main body of the periodic table to their current location below the lanthanide series. These elements were known as Actinides series.

•In 1951, Seaborg was awarded the Nobel Prize in chemistry for his work. Element 106 has been named seaborgium (Sg) in his honor.

•Presently, 118 elements are in the modern Periodic Table.

Although Dimitri Mendeleev is often considered the “father” of the periodic table, however the work of many scientists contributed to its present form. The representation of a modern Periodic Table of Elements is shown below.

Checking up 5.1

Theperiodic table is an important tool used in chemistry:

1.Explain why the elements are classified in groups and periods of the periodic table

2. Chose one elements of Group 1 and one of group 17 and make their electronic configurations using orbitals.

5.2. Comparison of Mendeleev’s Table and Modern Periodic Table

Activity 5.2

1. Discuss the similarities and differences of Mendeleev’s table and modern periodic Table.

2. How were the positions of cobalt and nickel resolved in the modern periodic table?

The periodic table is the arrangement of chemical elements according to their chemical and physical properties. The modern periodic table was created after a series of different versions of the periodic table. The Russian Chemist/Professor Dmitri Mendeleev was the first to come up with a structure for the periodic table with columns and rows. This feature is the main building block for the modern periodic table as well. The columns in the periodic table are called groups, and they group together elements with similar properties. The rows in the periodic table are called periods, and they represent sets of elements that get repeated due the possession of similar properties. The main difference between Mendeleev and Modern Periodic Table are shown in the Table below (Table 5.1).

Table 5.1. Differences between Mendeleev’s table and the modern Periodic Table

Checking up 5.2

1. The periodic table is an arrangement of elements based on their properties.Explain the gaps found in the Mendeleev periodic table compared to the modern one?

2. How many elements does the modern periodic table contain?

3. Look at the modern periodic table and write down four things it tells you.

5.3. Location of Elements in the Periodic Table Based On the Electronic Configuration

Activity 5.3

1. Based on knowledge gained in the previous years:

a. Represent the electronic configuration of the elements 25X and 11Y.

b. Discuss the information given by the number of electrons in the last orbitals of the above element about their position in the periodic table?

c. Explain the period and the group of the periodic table in which the above elements are located.

2. Is it possible to have an element with atomic number 1.5 between hydrogen and helium?

5.3.1. Major Divisions of the Periodic Table

The periodic table is a tabular of the chemical elements organized on the basis of their atomic numbers, electron configurations, and chemical properties.

In the periodic table, the elements are organized by periods and groups. The period relates to the principal energy level which is being filled by electrons. Elements with the same number of valence electrons are put in the same group, such as the halogens and the noble gases. The chemical properties of an atom relate directly to the number of valence electrons, and the periodic table is a road map among those properties such that chemical properties can be deduced by the position of an element on the table. The electrons in the outermost or valence shell are especially important because they participate in forming chemical bonds.

Elements are presented in increasing atomic number. The main body of the table is a 18 × 7 grid. There are four distinct rectangular areas or blocks such as s, p, d and f blocks. The f-block is usually not included in the main table, but rather is floated below, as an inline f-block would often make the table impractically wide. Using periodic trends, the periodic table can help predict the properties of various elements and the relations between properties. It therefore provides a useful framework for analyzing chemical behavior and is widely used in chemistry and other sciences (Petrucci et al., 2007).

5.3.2. Location of elements in modern Periodic Table using examples

In the periodic table, the elements are located based on groups and periods.

a) Finding the Period of elements

The Period of an element is equal to the highest energy level of electrons or principal quantum number which is being filled by electrons.

For a better understanding the following is an example:

Consider the element 16S with the electronic configuration: 1s22s22p63s23p4, the number 3 is the highest energy level or principal quantum number of electrons. Thus the period of S is 3.

If we consider another example, the electronic configuration of 23Cr is: 1s22s22p63s23p64s13d5 The number 4 is the highest energy level or principal quantum number of electrons. Thus the period of Cr is 4.

b) Finding the Group of elements

The Group of an element is equal to the number of outermost or valence electrons of element or number of electrons in the highest energy level of elements. Another way of finding the group of element is looking at sub shells. If the last sub shell of electron configuration is “s” or “p”, then the group becomes “1 to 18” with the number of group corresponding to the electrons occupying the last orbitals according to the given examples.

Consider the element 19K with the electronic configuration:1s22s22p63s23p64s1.Since the last sub shell is “s” with one electron, the element K is located in group 1of the periodic table.

Similarly, the element 35Br has the corresponding electronic configuration:1s22s22p63s23p64s23d104p5. Since the last sub shell is “p”, the element bromine (Br) is located in group 17of the modern periodic table.

The elements of groups 3 to 12 have the electron configuration ns and (n-1)d, total number of electrons in these orbitals. This informs us that those elements are located in group of element corresponding to transition metals.

Let us consider the following examples. The element 26Fe has the electronic configuration:1s22s22p63s23p64s23d6, 6+2 = 8 correspond to the group 8 of the transition elements.

Here are some evidences for everyone to find group number of elements.

Last Orbital Group: Considering the last orbital of an element, we are able to identify in which group of the periodic table it is located. The following tables show the location of representative elements in different groups (Table 5.2) and transition elements (Table 5.3 based on the last orbital.

Table 5.2. Group number of representative elements

For the transistion elements or d block elements, their valence electron configuration is ns1-2and (n-1)d1-10, hence the total number of valence electrons in these orbitals gives the group number.

Example: Sc (4s23d1), its group number is 3

Cu (4s13d10), its group number is 11

Table 5.3. Group number of transition elements

Example: Find the period and the group of the elements 16X and 24X using its electronic configuration:

16X: 1s22s22p63s23p4; the period of X is 3 and its group is 16

24X: 1s22s22p63s23p64s23d4; the period is 4 and the group is 4 + 2 = 6

The elements in the periodic table are also located using atomic orbitals.

Checking up 5.3

The elements X, Y, Z, T and U are given in the picture below.

Which one(s) of the following statements are correct, which one(s) are false for these elements.

a. X is alkaline metal

b. Y is in p block

c. Z is halogens

d. U is lanthanide

e. T is noble gas

2. Explain the major parts of the periodic table?

3. Differentiate the blocks included in a periodic table?

a. By using examples, locate at least two elements in each block

b. Indicate the group and the period of the chosen element.

5.4. Classification of Elements into Blocks (s, p, d, f-block)

Activity 5.4.

1. Apart from blocks, the elements in the periodic table are classified as representative elements, transition metals, lanthanides and actinides. Match these categories with s, p, d and f-blocks of the periodic table.

2. Classify the following elements into s, p, d or f-block and justify your answer: Al (z=13), K (z= 19), Ca (z =20) and Fe (z=26).

In the Periodic Table, the elements are organized into different blocks according to their electron configurations. They are classified into four blocks such as s-block, p-block, d-block and f-block depending on the type of atomic orbitals that are being filled with electrons. This division is based upon the name of the orbitals which receives the last electrons. The s-block has two groups of reactive metals: Group 1 and 2.

p-block is composed of metals and nonmetals of Group 13 to 18.

d-block is made of transition metals:Group 3 to Group 12, and f-block is made of lanthanide and actinide series or inner transitionmetals.

The division of elements into blocks is primarily based upon their electronic configuration as shown in Figure 5.1. Two exceptions to this categorization can be mentioned. Helium is placed in p-block although its valence electrons are in s orbital because it has a completely filled valence shell (1s2) and as a result, displays properties representative of other noble gases. The other exception is hydrogen. It has only one s-electron and hence can be placed in group 1 (alkali metals); but in many modern Periodic Tables, hydrogen is left hanging above the Periodic Table and doesn’t belong to any group. This is due to the particular properties of hydrogen:

•Hydrogen is the smallest chemical element

•Hydrogen is a gas while the other elements of group 1 are solids,

•Hydrogen is not a metal whereas the other elements of group1 are metals,

•In some compounds where hydrogen combine with non-metals, it behaves like a metal, e.g. in the polar molecule Hδ+Clδ-, hydrogen tends to lose an electron,

•When combined with very active metals, it behaves as a non-metal and forms a negative ion H-, hydride ion; e.g. Na+H-(sodium hydride).

Elements within the same group have the same number of electrons in their valence (outermost) shells, and they have similar valence electron configurations. They exhibit similar chemical properties. Elements within the same period have different numbers of electrons in their valence shells and the number of electrons is increasing from left to right. Therefore, elements in the same period are chemically different, changing from metals to non-metals across the period from left to right.

5.5. Characteristics of different blocks of the periodic table

a. The s-block elements

The s-block comprises elements of Group 1 (alkali metals) and Group 2 (alkaline earth metals) which have ns1 and ns2 outermost electronic configuration. The elements of group 1 and 2 are classified as reactive metals with low ionization enthalpies and highly electropositive. They lose their valence electron(s) willingly to form a positive ion with the charge +1 in the case of alkali metals or +2 in the case of alkaline earth metals. The metallic character and reactivity of s-block elements increases as we move down the group. Because of high reactivity they are never found pure in nature. Most of the metals of this block give a characteristic flame color. They are soft metals with low melting and boiling points. The compounds of s-block elements, with the exception of those of lithium and beryllium are predominantly ionic. There are 14 s-block elements in the periodic table. The Table 5.4 below shows an example of the electronic configuration of some elements of Group 1.

Table 5.4. Electronic configuration ofthe group 1 elements

b. The p-block elements

The p-Block comprises elements belonging to Group 13 up to 18. The elements of p-block and s-block are called the representative elements or main group elements.

The valence electron of p-block elements varies from ns2np1 to ns2np6 in each period. At the end of each period appears a noble gas element with a closed valence shell ns2np6 configuration. The p-block includes 6 groups:13, 14, 15, 16, 17 and 18. The atoms of the elements belonging to these groups accept the last electron in 2p, 3p, 4p, 5p and 6p orbitals. Group 18 elements are called noble gases and all these elements have closed shell ns2 np6 electronic configuration in the outermost shell. All the orbitals in the valence shell of the noble gases are completely filled by electrons and it is very difficult to alter this stable arrangement by the addition or removal of electrons. The noble gases thus exhibit very low chemical reactivity.

The p-block elements consist of both metals and non-metals but the number of non-metals is considerably higher than that of metals. The metallic character increases as we move down the group and non-metallic character increases from left to right along a period. The ionization energiesof p-block elements are relatively higher compared to those of s-block elements. Most of them form covalent compounds. Some of p-block elements exhibit more than one oxidation state in their compounds and their oxidizing character increases from left to right in a period and reducing character increases from top to bottom in a group. Most of the p-block elements are highly electronegative and form acidic oxides. The p-block contains some elements that show intermediate properties between metals and non-metals; those elements are called “metalloid” or “semi-metals”, e.g. Si and Ge; they are also called “semi-conductors”.

c) The d-block elements

Elements in which the last electron enters any one of the five d orbitals are called d- block element. These elements comprise Group 3 to 12 in the centre of the Periodic Table (Figure 5.2). These elements are characterized by the filling of inner d orbitals by electrons and are therefore referred to as d-Block Elements. Their outer electronic configuration is (n-1)d1-10ns0-2. As these elements are transition metals, mostly of them form coloured ions and they display variable valences or oxidation states. However, Zn, Cd and Hg which have the electronic configuration, (n-1) d10ns2 do not show most of the properties of transition elements.

The properties of d-block elements are midway between those of s- block and p- block elements, they are divided into four series called 1st, 2nd, 3rd and 4th transition series. The first transition series contains 10 elements from scandium (Z=21) to zinc (Z=30); for those elements, 3d orbitals are being progressively filled in. The second transition series contains the elements from Y through Ag in which the 4d orbitals are being progressively filled. The third transition series also contains the elements La and the element Hf through Au in which the 5d orbitals are progressively filled. The fourth transition series comprise elements also encloses the elements from Rf to Rg with 6d orbitals being filled.

The elements of d-block are hard, malleable and ductile metals with high melting and boiling point. Mercury is an exception, it is liquid at room temperature. Because they are metals, they are good conductors of heat and electricity.

Their ionization energy are between s and p block elements. Transition metals show variable oxidation states reason why they form both ionic and covalent compounds. In chemistry, catalysts used in different reactions are transition metals (i.e: V, Cr, Mn, Fe, Co, Ni, and Cu).

d. The f-Block elements

The f-block elements are found within the two rows of elements at the bottom of the Periodic Table. They are called the Lanthanides and include Ce (Z = 58) to Lu (Z = 71) and Actinides, Th (Z = 90) to Lr (Z = 103). All of these elements are characterized by the outer electronic configuration (n-2)f1-14 (n-1)d 0–1 ns2. The last electron added to each element is filled in f- orbital. These two series of elements are hence named the inner transition elements (f-block elements). They are all metals and display similar properties within each series. The f-block elements exhibit variable oxidation states, they have high melting and boiling point. They form complex compounds and most of the elements of the actinide series are radioactive.

Checking up 5.5

1. Among the common blocks, s, p, and d; which block has a tendency to form complex compounds?

2. Why d-block are called transition elements?

3. Why f-block are called inner transition elements?

5.6. Variation of Physical Properties down the Groups and across the Periods

Activity 5.6

1. The elements in the periodic table display many trends which can be used to predict their physical properties. Explain three of the factors that you think can influence the physical properties of elements in the periodic table.

2. Discuss the trends of the above factors across a period and down a group in the periodic table.

The elements in the periodic table are arranged in order of increasing atomic number. All of these elements display several other trends and we can use the periodic table to predict their physical properties. There are many noticeable patterns in the physical and chemical properties of elements as we descend in a group or move across a period in the Periodic Table.

Those trends can be observed in: ionization energy, electronegativity, electro positivity, electron affinity, melting and boiling point, density and metallic character and hereafter are some factors which cause those trends.

5.6.1. Atomic radius

The atomic radius of an atom is defined as half the distance between the nuclei of two atoms of the same element that are joined together by a single covalent bond.

Atomic radius of elements decreases as we move from left to right in periodic table. This is explained by the number of outer electrons and protons which increase while there is no change in the energy level. The results increase the attracting forces making the radius smaller.

Increasing nuclear charge (more protons) pulls the electrons closer to the nucleus, and the screening effect of inner electron shells will be the same for all members of a given period. The combined effect of both factors results in the electrons being pulled closer to the nucleus and a smaller radius.

On the other side, in the same group, as we go down, the atomic radius of elements increases. This is due to the energy level which increases when you move down in group of the periodic table, the attraction of external electrons by nucleus decreases and atomic radius increases.

In general, atomic radii increase down a group because a new shell is added for each successive member of a group, leading to a greater radius. Then an increased screening effect of extra electron shells i.e. the nucleus has less of a pull on the outer electrons.

5.6.2. Electronegativity

Electronegativity is a measure of the tendency an atom to attract to itself the shared pair of electrons making a bond. The charge in the nucleus increases from left to right across a period.The electronegativity of atoms is affected by both the charge of the nucleus and the size of the atom. The higher its electronegativity, the more an element attracts electrons. In general, the electronegativity of a non-metals is greater than that of metals. Trends are observed in the period (Figure 5.3) or in a group of the Periodic Table (Figure 5.4).

•In a period, the electronegativity increases from left to right This is explained by the fact that as we go from left to right, there in an increase of positive charge in the nucleus, since the number of protons increases; but the electrons are being added to the same energy level. This results in the reduction of the volume or radius of the atoms from left to right and explains why attraction of external electrons by the nucleus increases from left to right.

•In a group, the electronegativity decreases from top to bottom. This is due increase of energy levels down in a group, and thus there is an increased distance between the valence electrons and the nucleus, or a greater atomic radius.The positive charge of the nucleus is further away from the valence electrons and the nucleus cannot attract efficiently external electrons.

Note:

•Since noble gases do not react or do not form chemical bonds, their electronegativity cannot be determined.

•For the transition metals, the electronegativity does not vary significantly across the period and down a group. This is because their electronic structure affect their ability to attract electrons easily as for the other elements.

•The lanthanides and actinides possess more complicated chemistry that does not generally follow any trend. Therefore, they do not have electronegativity values.According to these two general trends, the most electronegative element is fluorineand Francium is the least (Figure 5.4 and Figure 5.5).

The charge in the nucleus increases across a period. Greater is the number of protons, greater is the attraction for bonding electrons.

5.6.3. Ionization energy (I.E)

Ionization energy: it is the amount of energy required to remove an electron from a neutral gaseous atom. The lower this energy is, the more readily the atom loses electron and becomes a cation. Therefore, the higher this energy is, the more unlikely it is the atom to become a cation. We can distinguish, first, second, and third ionization. Helium is the element with the highest ionization energy (Zumdahl and Zumdahl, 2010). The noble gases possess very high ionization energies because of their full valence shells compared to the elements of group 1 (Table 5.5).

The table shows that generally the IE decreases down the Group, as the size of the atoms increases down the Group.

Table 5.5. First ionization energies (kJ/mol) for the alkali metals and noble gases

As mentioned above, different types of ionization energy exist such as first and second ionization energy. In general, the second ionisation energy of an element is always greater than the first ionisation. This is explained as follows: every time you remove an electron from an atom, the remaining electrons are more strongly attracted by the nucleus and it require more energy to remove other electrons from the atom.

Hence: 1st IE < 2nd IE < 3rd IE

Ionisation energy of rare gases or any species with an octet electronic structure show very high IE because the electron is being removed from a very stable electronic structure.

The ionization energy varies across a period and down a group.Across a period ionisation energies increase because the nuclear charge increases (greater positive charge on the nucleus) and holds the outer electrons more strongly. More energy needs to be supplied to remove the electron.Down a group ionisation energies decrease because the outer electrons are further away from the nucleus. The screening effect of the inner electron shells reduces the nuclear attraction for the outer electrons, despite the increased (positive) nuclear charge

5.6.4. The melting points and boiling points

Melting points and boiling points show some trends in groups and periods of the Periodic Table.As you already know, the Periodic table can be subdivided into two main area or regions:

•the left region where you find only metallic elements

•the right region where you find both metallic and non-metallic elements; all non-metallic elements are in the extreme right part of that region.

The general trends of melting and boiling points depends on the regions:

•in the left region, melting and boiling points generally decrease down the groups due to the decrease of strength of the metallic bond down the groups;

•on the contrary, in the right region at the extreme right in groups 17 and 18, there is a general increase of melting and boiling points down the group due to the increase of the molecular mass;

•from left to the middle of the periodic table, there is an increasing of melting and boiling points from left to right in a periode due the the increasing of the strength of the metallic bond;

•whereas from the middle of the periodic table, there is a decrease of melting and boiling points from left to right due to the progressive increase of non-metallic character where elements exist as simple molecules.The melting and boiling points vary in a regular way or pattern depending on their position in the Periodic Table. In general the forces of attraction for elements on the left of the table are strong metallic bonds; they require higher energy to be broken, hence higher melting and boiling points.

As we cross toward the right side of the periodic table, the non-metal character of elements increases and elements, except few elements, form molecules that are held together by weak intermolecular forces; hence their melting and boiling points are generally low.

For example going down in group 1, the melting point and boiling point of the alkali metals decrease. This is due to the weakning of metallic bond down the group. However, going down in group 17 of the halogens the melting point increases meaning that there is an increase in the force of attraction between the molecules. The illustrations below show the variation of melting and boiling point for some elements of the periodic table (Figures 5.6 and 5.7).

5.6.5. The density

The density of a substance is its mass per unit volume, usually in g/cm3. The density is a basic physical property of a homogeneous substance; it is an intensive property, which means it depends only on the substance’s composition and does not vary with size or amount.

The trends in density of elements can be observed in groups and periods of the periodic table. In general in any period of the table, the density first increases from group 1 to a maximum in the centre of the period because the mass increases while the size decreases, and then the density decreases again towards group 18 because of the nature of bonds.

Going down a group gives an overall increase in density because even though the volume increases down the group, the mass increases more. The variation of density with atomic number is shown in the Figure 5.8.

5.6.6. Electrical and thermal conductivity

The electrical conductivity is the ability of a substance to conduct an electric current.

The electrical conductivity of elements increases from non-metals to metals. Metals are good conductor of electricity. This is due to the presence of free electrons in metallic lattice. The capacity of metals to conduct heat is called thermal conductivity of metals. Elctrical conductivity results from the transfer or mobility of electrons, whereas the thermal conductivity in metal is due to heat transfer by free electrons from one end of metal to another end.

As we move across the period from the left to the right, the electrical conductivity increases for the metals as the number of free electrons increases and then decreases for the non-metals because they do not have free and mobile electron.

1. Metallic character

Metallic character refers to the level of reactivity of a metal. Metals tend to lose electrons in chemical reactions, as indicated by their low ionization energies. Within a compound, metal atoms have relatively low attraction for electrons, as indicated by their low electronegativities.

Metals are located in the left and lower three-quarters of the periodic table, and tend to lose electrons to nonmetals. Nonmetals are located in the upper right quarter of the table, and tend to gain electrons from metal. Metalloids are located in the region between the other two classes and have properties properties.

•Metallic character is strongest for the elements in the leftmost part of the periodic table and tends to decrease as we move to the right of any period.

•Within any group of the representative elements, the metallic character increases progressively going down.

2. The electron affinity (E.A)

The electron affinity is the ability of an isolated gaseous atom to accept an electron. Unlike electronegativity, electron affinity is a quantitative measurement of the energy change that occurs when an electron is added to a neutral gas atom. The more negative the electron affinity value, the higher an atom’s affinity for electrons. In the periodic table, the first electron affinities of elements are negative in general except the group 18 and group 2 elements. The second electron affinities of all elements are positive. This is because the negative ion creates a negative electric field. And if now the other electrons enter the negative field, energy has to be applied to the system to overcome the repulsion between the negative electric field and incoming electron.

The more the electron affinity value is negative, the higher is the electron affinity of an atom. Electron affinity decreases down a group of elements because each atom is larger than the atom above it (refer to atomic radius trend).This means that an added electron is further away from the atom’s nucleus compared with its position in the smaller atom. With a larger distance between the negatively-charged electron and the positively-charged nucleus, the force of attraction is relatively weaker. Therefore, electron affinity decreases down the group.

Moving from left to right across a period, the electron affinity increases because the electrons added to energy levels become closer to the nucleus and there is a stronger attraction between the nucleus and electrons.

Checking up 5.6

1. The following table shows a part of a periodic table. Students have to answer the following

Fill in the blank space with the correct term based on the above table.The element with the least nuclear charge is .......and the one with the highest nuclear charge .....Nuclear charge of S is ....than the nuclear charge of Se.As you go from Na to Cl along the period nuclear charge ....Effective nuclear charge....from B to Ga while it....from Na to Ar. Shielding or screening effect ..... down the group but ......along the period from left to right.Atomic size of Li is.... than that of K. Element with the least atomic size is (10)......and the element with the highest atomic size is .......Atomic size of Ca is .... than atomic size of Be because number of .......increases down the group.Atomic size ........ from K to Kr because electrons are filled on the same shell, the .......continuously increase and attraction force increases.Analyze and complete the following concept map using: ionization energy, atomic size, electron affinity, electronegativity and metallic character.

5.7. End unit assessment

1. The following are coded groups/families of the representative elements of the periodic table (first 4 periods, s, p blocks only). The groups are in number of particular order. Use the hints below to identify the group and place of three elements of each group in their correct location in the periodic table: AOU, BVW, CKM, DLQ, ENT, FIJ, GPY, and HRS.

Hints

A has only one electron in p subshell

B is more electronegative than V

C has a larger atomic radius than both M and W

D has electronic configuration ending in p5

E is one of the most reactive metals

F has a smaller ionization energy than J

G has only 1 energy level with any electrons

H has one more proton than O and is in the same period as O

I is the largest alkaline earth metal

J has one more proton than E

K has electron configuration ending in p3

L has more filled energy levels than D

M is larger than K

N has the largest radius in its family

O is smaller than F but in the same energy level as F

P is smaller than Y

Q is the most reactive non-metal

R has the highest electronegativity in its family

T has the lowest density in its family

U more easily loses electrons (think about ionization energy) than either A or O

V has only 4 electrons in a p-subshell

W has 3 completely filled energy levels

Y has the lowest ionization energy in its family.

2. Based on the variation of ionization energy in groups and periods, how should you explain the variation of first and second ionization energy down a group and across a period?

3. Justify the following statements:

a) The first ionization energy of nitrogen is higher than that of oxygen even though nuclear charge of nitrogen is less compared to oxygen.

b) Noble gases are having high ionization energies.

4. Give reason

a. Alkali metals (group 1 elements) are not found free in nature.

b. Atomic radius of gallium is smaller than that of aluminium.(Z of Al = 13, Z of Ga = 31)The following are coded groups/families of the representative elements of the periodic table (first 4 periods, s, p blocks only). The groups are in number of particular order. Use the hints below to identify the group and place of three elements of each group in their correct location in the periodic table: AOU, BVW, CKM, DLQ, ENT, FIJ, GPY, and HRS.

Key unit competence: Compare and contrast the chemical properties of the Group 1 elements and their compounds in relation to their position in the Periodic Table.

Learning objectives

By the end of this unit, I will be able to:

•Describe and explain the physical properties of Group 1 elements in terms of metallic character and strength of metallic bond

•Describe and explain the reactivity of Group 1 elements with oxygen, water and halogens•State and explain the properties of Group 1 oxides and hydroxides

•Explain the trends in the solubility of Group 1 compounds

•State the uses of Group 1 elements and their compounds

•Compare the reactivity of Group 1 elements•Interpret the trends in the thermal decomposition of Group 1 carbonates and nitrates

•Perform experiments to test the alkalinity of Group 1 hydroxides•Carry out flame test for the presence of Group 1 metal cations in solution.

6.1. Occurrence and physical properties of group 1 elements, physical state, metallic character, physical appearance and melting point

Activity 6.1.a

1. Study the following table of data and answer the questions that follow

Data

a) What’s meant by 1st ionization energy ?

b) Explain the trend in 1st ionization energy for group 1 metals down the group.

c) What is the relationship between melting point and atomic radius in group 1 element? Explain this trend.

2. What do you know about Group 1 elements?

3. Do you know any compound of a group 1 element? If yes, what is its application in our daily life?

Group 1 elements are: lithium (3Li), sodium(11Na), potassium(19K), rubidium(37Rb), cesium(55Cs) and Francium(87Fr); they are characterised by valence electronic structure ns1 and that is why they are called s-block and Group 1 elements.

Although hydrogen has the same valence of ns1, it behaves differently from the others in many aspects; for example: hydrogen is a gas at room temperature whereas all other members are solid. Hydrogen behaves as a metal in some compounds when it combines with non-metals, and as a non-metal when it combine with very active metals. For this reason, the study of hydrogen is presented separately from the other members of group 1 elements.

Francium is exceptionally rare. It is formed by the radioactive decay of heavier elements.

Because Francium is both rare and highly radioactive, few of its properties have been determined and we are not going to talk much about it in this Unit.

Physical Properties of Alkali Metals

Activity 6.1.b

In groups, learners make research in libraries / internet and discuss on the physical properties of group 1 elements and or explain the following statements

a) Group 1 elements show weak metallic bonding

b) Atomic radius of Na is smaller than the corresponding ionic radius of Na+

c) The shining appearance of metals disappears after a certain period of time.

Group 1 elements are grey metals, soft, and can be easily cut with a knife to expose a shiny surface which turns dull on reaction with oxygen in air.

1. They have low melting and boiling points. They show relatively weak me-tallic bonding as only one valence electron is attracted by the nucleus. Also they have a big atomic radius and the attraction of nucleus toward the va-lence electron is weak.

2. They are good conductors of heat and electricity

3. They have a low ionisation energies that decreases down the group

4. They have low density compared to other metals and Li, Na and K are less dense than waterGroup 1 metals color flames: When alkali metals are put in a flame they produce characteristic colors.

Table 1: Physical properties of group 1 elements

Checking up 6.1

1. Discuss how the ionization energies vary in function of atomic radius for group 1 elements.

2. Is there any relationship between the atomic radius and the melting point in group 1 elements? Yes or No. Justify your answer

3. Why group 1 elements are said to be good conductors of electricity? Illustrate your answer

4. The following table shows 3 unknown group 1 metals X, Y, Z and some of their physical properties. Predict among the alkalis metal (Cs, Li, K), which one should correspond to X,Y,Z. Justify your answer

6.2. Reactivity of group 1 elements with oxygen, water and halogens

Activity 6.2 (a)

Analyse the case study below and answer the related question:1. In groups learners discuss about the following scenario and compare their findings to the reactivity of a chemical element in term of the variation of atomic radius

Suppose that a hen walking in the garden with its chicks. Some of the chicks are feeding themselves just near the hen (mother). Other chicks are feeding themselves far away from the mother. Which ones of the two groups of chicks will be an easy prey of a predator? Explain.

Now extend your reasoning to the behavior of group 1 element and explain the following statement

Group 1 elements react by losing their single electron in outermost shell. Arrange them in order to show which one loses easily the single electron and which one loses electron with difficult.

2. Consider a case of people who are warming around a fire; some are near, other are a little bit far; who will feel the heat of the fire more than the other?

These activities show the distance between two object affect interactions between them. The hen mother cannot protect the chicks that are far from her. In the other case, people who are far from the fire feel less the heat of the fire.

This resembles to the interactions between the nucleus and the valence electrons. The attraction between the nucleus and the valence electron decreases with increasing distance between the nucleus and the valence electrons.

When the atomic radius and volume increase, the distance between the nucleus and the valence electrons increases, the attraction between the nucleus and the valence electrons decreases, and it becomes easier to remove the valence electrons.

This explains the origin of properties of group 1 metals: they have low ionization energy; lose easily the only one valence electron to form a mono-positive cation (M+) with a rare gas electronic structure, and consequently are very active metals.

Activity6.2 (b)

a) Given the following element of group 1(Na (Z=11), Li(Z=3), Cs(Z=55), K(Z=19) , Rb (Z=37). Arrange them according to their increasing reactivity and justify why.

b)Establish the electronic structure for the following species and explain how you get it Na+,Na2+,Na-.Which one is stable and why?

6.2.1. Reactions with Oxygen

Activity 6.2.1

Experiment: burning an alkali metal in air/oxygen

Apparatuses: deflagrating spoon, Bunsen burner, glass beaker, filter paper,

Chemicals: lithium, sodium, potassium, water, and red litmus paper

Other requirements: knife, match box and petroleum gas

Procedure:

1. Cut a small piece of lithium and wrap it in piece of filter paper to remove the oil.

2. Place it on to a deflagrating spoon and heat it in a non luminous flame.

3. Observe what happens

4. When combustion is complete, dip the deflagrating spoon into a beaker of 100 ml filled up to 50 ml of water.

5. Stir the water with the spoon and then drop a piece of litmus papers into the solution in the beaker. Observe

6. Repeat the experiment with sodium and potassium

Task on the experiment:

a) Write the equations of reactions that take place when each metal is burnt in air

b) Name the product that was formed in each case.

c) What are the color changes when the aqueous solutions above are tested with litmus paper? Explain why?

d) Write the equations of reaction between each product in (b) with water

Alkali metals form oxides when they burn in air or in oxygen. The nature of the oxides varies among the elements. For example lithium forms normal oxide (Li2O), sodium forms peroxide (Na2O2) and the remaining elements mainly form superoxides (MO2) with very little amounts of peroxides:

The reason for formation of peroxides and superoxide as we move down the group is due to increased reactivity of Group 1 metals down the group.

Oxides formed by group 1 metals are highly soluble in water giving strongly alkaline solutions:

Hence Group1 metals form basic oxides.

All group 1 metals are active metals, but there are differences within the group. The reactivity increases down the group due to the increase of atomic radius which results in the decrease of attraction between the nucleus and the valence electron, which requires less energy to remove the electron down the group.

6.2.2. Reaction with water

Activity 6.2.2

Experiment to investigate the reaction of alkali metals with water

Procedure:Cut a small piece of sodium metal and put it on water in wide beaker and observe.Test the obtained solution with red and blue litmus papers and observe.

Questions

1. Why sodium is so easy to cut?

2. What are the observations when sodium is placed on water in the beaker

3. Explain the observation when tests with red and blue litmus papers are performed on the resulting solution and write the equation of the reaction that takes place to explain you answer

Alkali metals react with water and form hydroxides along with liberation of hydrogen:

Where M represents any group 1 metal, MOH represents the corresponding hydroxide such as sodium hydroxide. This reaction explains why it is said that group 1 metals reduce water (where the ion H+ from dissociation of water is reduced to H2molecule). This reaction where group 1 metals react with water to form an alkaline solution is at the origin of their name of “alkali metals”.

As said previously, the reactivity of Group 1 metals increases as we move down the group.

The reaction with water can be used to illustrate the increasing reactivity on descending the Group. Li reacts slowly with water, with effervescence; sodium reacts violently fizzing (a hissing sound is heard) and skating about on the water surface, a colorless gas is evolved; K reacts more violently and ignites on contact with water; Cs sinks in water, and the rapid generation of hydrogen gas under water produces a shock wave that can shatter a glass container.

The chemistry of Li shows some anomalies, as the cation Li+ is so small that it forms covalent compounds.

6.2.3. Reaction with halogens

Activity 6.2.3

a) In terms of s, p, d, f orbitals write the electronic configuration of chlorine (Z=17), bromine (Z=35) and iodine (Z=53)

b) Deduce the valency of each element.

c) Write molecular equations, complete and balance them, when

Sodium reacts with bromine        Potassium reacts with iodine            Lithium reacts with chlorine

Alkali metals react with halogens very easily forming halides. The reactivity of alkali metals with halogen increases from Li to Cs.

Elements of group 1 are highly reactive and form a variety of compounds. Two of the most important are Oxides and Hydroxides.

6.3.1 Oxides

Alkali metals form oxides when they burn in air or in oxygen as seen in point (6.2.1):

They are also formed by thermal decomposition of corresponding hydroxides

Checking up 6.2

1. An element J has 19 as atomic number while the element A has 35 as atomic number.

a) Write their electronic structures in term of s, p, d, f orbitals and deduce their respective valencies.

b) The element J is able to react with oxygen gas by forming two types of oxides

i) Write the formula of the 2 oxides that can be formed between the ele-ment J and oxygen

ii) Write the formula of the compound formed between J and A

iii) What type of bond does exist between J and A. Justify your answer.

c) Show how you would write the equation of reaction between J and water supposing that J stands for the real symbol of the element.

d) When the reaction stated in (c) takes place a colorless solution and a colorless gas are formed.

i) Which test would you use to identify each product of the reaction, by stating the reagent and related observations?

Oxides of Group 1 metals dissolve in water to give strong alkaline solutions; that is why they are said to form basic oxides.

6.3.2. Hydroxides

As said above hydroxides are formed when the metals or metals oxides are dissolved in water. In solid state, these hydroxides dissolve very easily in water and in alcohol. They dissociate completely in water to form alkaline solution; hence they are strong bases. The basic character of the hydroxides increases as we move down the group.

Checking up 6.3

1. Complete and balance the reactions when water is reacted with the following:

a) Potassium metal

b) Potassium oxide

2. Lithium hydroxide decomposes on heating. White powder X and a colorless gas Y is released and condenses in a colorless liquid.

i) Write the chemical formula of X.

ii) Propose a chemical test to identify Y

3. Explain why Group 1 metals form ionic compounds?

6.4. The effect of heat on Group 1 carbonates and nitrates

6.4.1. Heating the nitrates

Activity 6.4.1

Experiment: effect of heat on nitrates.

In groups learners perform the following experiment, discuss and make conclusions by explaining the observed phenomena, and write involved chemical reactions

.Apparatus: glass test tubes, pair of tongs, wooden splint/match stick, Bunsen burner/heat source and spatula.

Chemicals: Lithium nitrate, potassium nitrate Other requirements: match box

Procedure:

I.1.Take two spatula end full of lithium nitrate into a test tube and heat it strongly until there is no further change.

2. Test the gases evolved with a damp blue litmus paper and a glowing splint.

3.Observe and make conclusions on your observations.

II .Repeat the procedure but using potassium nitrate/sodium nitrate

Group 1 compounds are generally stable on heating.

Group 1 nitrates (except lithium nitrate) decompose on heating in Bunsen flame to form their corresponding nitrites (which are stable) and oxygen

All nitrates from sodium to cesium decompose in this same way.

Lithium nitrate behaves differently, producing lithium oxide, nitrogen dioxide and oxygen due tothe small size of Lithium leading to high polarizing power and formation of compounds with covalent properties.

Checking up 6.4.1

a) Differentiate between KNO3 and LiNO3 in terms of their thermal decomposition.

b) One of the two compounds decomposes releasing a colored gas that turns damp blue litmus paper to red.

i) State which one.

ii) Give the name, the formula and the color of that gas

6.4.2. Heating the carbonates

Activity 6.4.

2Experiment: effect of heat on carbonates of group 1 elements

In groups learners perform the following experiment, discuss and make conclusions by explaining the observed phenomena, and write chemical reactions.

Apparatus: glass test tubes, pair of tongs, Bunsen burner/heat source and spatula.

Chemicals: lithium carbonate, calcium carbonate, potassium carbonate and sodium carbonate, lime waterOther requirements: match box.

Procedure

I.1 Take a spatula end full of lithium carbonate into a test tube and heat it strongly until there is no further change.

2. Test the gases evolved with a damp blue litmus paper and lime water into an-other glass test tube as shown in the figure

3. Observe and make conclusions on your observations.

II .Repeat the procedure but using calcium carbonate, potassium carbonate /sodium carbonate

Most carbonates of group 1 resist to heat.

However in Group 1, lithium carbonate decomposes on heating, producing lithium oxide and carbon dioxide.

The remaining of the Group 1 carbonates does not decompose at Bunsen flame temperatures.

Checking up.6.4.2

Li2CO3 decomposes on heat releasing a colorless gas.

a) Write the equation of thermal decomposition of Li2CO3

b) What do you observe when the evolved gas is tested by lime water and damp blue litmus paper? Explain

6.5. Solubility of group 1 compounds

Activity 6.5

a) Group 1 elements form ionic compounds.

(i) Explain why.

(ii)State the properties of ionic compounds

b) Explain why lithium forms components with a covalent character contrarily to other component of the same group. State the properties of covalent com-pounds

c) Both hydroxides and carbonates of lithium are less soluble than other hydrox-ides and carbonates of group 1.Why?

All group 1 salts and hydroxides are soluble in water.However LiCO3, LiOH are less soluble of group 1 compounds due to their high covalent properties.

6.5.1. Carbonates

The carbonates of Group 1 metals are all very soluble - increasing to an astonishing 261.5 g per 100 g of water at 20oC temperature for cesium carbonate.

The least soluble Group 1 carbonate is lithium carbonate. A saturated solution of

Li2CO3 has a concentration of about 1.3 g per 100 g of water at 20°C.

Solubility of the carbonates increases as you go down the group.

6.5.2. Hydroxides

Hydroxides of Group 1 are even more soluble.

The least soluble hydroxide in Group 1 is lithium hydroxide - but it is still possible to make a solution with a concentration of 12.8 g per 100 g of water at 20°C. Solubility of the hydroxides increases as you go down Group 1.

6.5.3. Sulphates

The solubility of sulphates behaves in different way from carbonates and hydroxides. The solubility of group 1 sulphates decreases as you go down the group.

Checking up 6.5

Arrange the following compounds in ascending order of their solubility and explain why:

a) RbCl, NaCl, LiCl, KCl, CsCl

b) Cs2SO4, Li2SO4, Na2SO4, Rb2SO4, K2SO4

6.6. Flame Test for Li+, Na+ and K+

Activity 6.6:

Experiment: Flame test of alkalis metals

Materials: Mortar and pestle, beakers, Lithium carbonate, potassium sulphate, sodium sulphate

Procedure: Flame test wire /magnesia rod

NB: Wear your safety glasses.

Dip the flame test wire/magnesia rod in the salt to be tested. Some of the salt should stick to flame test wire/magnesia. Gently wave the flame test wire/magnesia rod in the flame of the Bunsen burner and note the color given of

Examples of flame test

Repeat the experiment for each of the other salts. To avoid cross contamination, use a separate splint/flame test wire for each salt. Again, note the color in each case. If you are given an unknown salt, you should be able to identify the metal in the salt from the results of your experiment. The following are the flame colors for respectively K, Na, Li

Below are examples of flame tests for Na, K, and Li.

Student Questions

What color is observed in each case when the following salts are heated in a flame: lithium carbonate, sodium sulfate, , potassium sulfate?

Alkali metals color flames

When the element or its compound isput in a flame, the heat provides sufficient energy to promote the electron of the metal (or metal ion) to a higher energy level. On returning to the ground level, energy is emitted in form of light that has a wavelength in the visible region that is seen as color.

Table 2: Group 1 metals and their corresponding flame colors

Checking up 6.6

Describe a chemical test to distinguish between 2 salt solutions below and predict expected results for:

a) Solutions of CsCl and KCl

b) Solutions of NaCl and LiCl

6.7. Uses of group 1 elements and their compounds

Basically, the use or application of any elements is determined by the chemistry and the physical properties of the particular element. Hereafter are some of the uses and applications of group 1 metals.

Activity 6.7

Make research on internet and in library about the use of group 1 elements and present your findings

•Lithium: Used for making alloys especially with aluminium to make aircraft parts, which are light and strong.

•Due to its size and electropositivity (opposite of electronegativity), lithium is used in both primary and secondary lithium batteries. Lithium is used in light weight electrical batteries of the type found in clocksand watches, heart pacemakers (a small piece of electronic equipment connected to someone’s heart to help the heart muscles move regularly.

•Lithium carbonate, Li2CO3, is used to toughen glass.

•Caustic soda, NaOH, and soda ash, Na2CO3 are the most important alkali used in industry. Both find applications in paper making, alumina, soap, and rayon.Na2CO3 (soda ash) is used in water treatment.

•NaOCl is used as bleaching agent and disinfectant

•NaCl is used in seasoning food, preparing hydrogen chloride gas, in soap production, manufacture of sodium, chlorine, sodium hydroxide and sodium carbonate.

•Molten sodium is used as a coolant in nuclear reactor. Its high thermal conductivity and low melting temperature and the fact that its boiling temperature is much higher than that of water make sodium suitable for this purpose.

•Sodium wire is used in electrical circuits for special applications. It is very flexible and has a high electrical conductivity. The wire is coated with plastics to exclude moisture.

•Sodium vapor lamps are used for street lighting; the yellow light is characteristic of sodium emission.

•Sodium amalgam (alloy with mercury) and sodium tetrahydridoborate, NaBH4, are used as reducing agents.

•Sodium cyanide is used in the extraction of silver and gold.

•Lithium: Used for making alloys especially with aluminium to make aircraft parts, which are light and strong.

•Due to its size and electropositivity (opposite of electronegativity), lithium is used in both primary and secondary lithium batteries. Lithium is used in light weight electrical batteries of the type found in clocksand watches, heart pacemakers (a small piece of electronic equipment connected to someone’s heart to help the heart muscles move regularly.

•Lithium carbonate, Li2CO3, is used to toughen glass.

•Caustic soda, NaOH, and soda ash, Na2CO3 are the most important alkali used in industry. Both find applications in paper making, alumina, soap, and rayon. Na2CO3 (soda ash) is used in water treatment.

•NaOCl is used as bleaching agent and disinfectant

•NaCl is used in seasoning food, preparing hydrogen chloride gas, in soap production, manufacture of sodium, chlorine, sodium hydroxide and sodium carbonate.

•Molten sodium is used as a coolant in nuclear reactor. Its high thermal conductivity and low melting temperature and the fact that its boiling temperature is much higher than that of water make sodium suitable for this purpose.

•Sodium wire is used in electrical circuits for special applications. It is very flexible and has a high electrical conductivity. The wire is coated with plastics to exclude moisture.

•Sodium vapor lamps are used for street lighting; the yellow light is characteristic of sodium emission.

•Sodium amalgam (alloy with mercury) and sodium tetrahydridoborate, NaBH4, are used as reducing agents

•Sodium cyanide is used in the extraction of silver and gold.

6.8 Hydrogen

Although hydrogen has the same valance electronic structure as Group 1 elements, H[1s1], it is not generally studied with the other elements of the group. Even some Periodic Tables do not put H in Group 1 column.

This is due the fact hydrogen has specific properties that distinguish it from the other elements of Group1. Hydrogen is the smallest chemical element, with only 1 electron in its electronic structure. Hydrogen is a gas, whereas other Group 1 elements are metals, all solid at room temperature. In some compounds with non-metals, hydrogen tends to lose electron and form polar compounds such as Hδ+Clδ-,

but when it reacts with active metals, it behaves as a non-metal and captures electron to form hydride such as Na+H-

.All this justifies why the study of hydrogen is separated from the study of other Group 1 elements.

6.8.1. Properties of hydrogen atom

Hydrogen is the first element in the periodic table. It is the simplest and smallest element of the chemical elements.

It is formed by only one electron and one proton for the common hydrogen(11H), with the simplest nucleus. Its electron configuration is 1s1(similar to the electron configurations of group 1 elements).Hydrogen is colorless, odorless, and tasteless gas.

Table 3: Properties of hydrogen

H2 is small and non-polar, so H2 molecules can only attract each other through weak van der Waals forces.Hydrogen is the most abundant element in the Universe and accounts for 89% of all atoms. There is little free hydrogen on Earth because H2 gas is so light that it moves very fast and can escape the Earth’s gravitational pull

6.8.2. Production of hydrogen

a) From methane gas

Example:

b) Electrolysis of water

c) Laboratory preparation

Metals are reacted with dilute acids. Group 1 metals are not used because the reaction with acid is violent and dangerous.

6.8.3. Uses of hydrogen

One third of the hydrogen produced is used for hydrometallurgical extractions of copper and other materials since hydrogen is a reducing agent. Half of the hydrogen is used in manufacturing of ammonia.

Hydrogen is also used in the manufacture of saturated oil (solid), such as margarine, from unsaturated oil (liquid) by hydrogenation reaction.

Hydrogen is used as fuel in space rocket/space shuttles.

6.8.4. Chemical properties of hydrogen

Hydrogen can form both cations (H+) and anions (H-). It has an intermediate electronegativity (2.1). Hydrogen forms covalent bonds with nonmetals where it bears a partial positive charge, Hδ+, and ionic hydrides with very active metals where it forms a hydride ion H-.

Examples: (a) Polar covalent compounds: NH3, H2O, HF, and SiH4

(b) Ionic compounds: CaH2, NaH, KH

When bonded to very electronegative elements, O, F, N, hydrogen is responsible of hydrogen bonding, a strong force between very polar covalent groups having O-H, F-H, and N-H present in a molecule.

6.9. End unit assessment

1. a)Arrange the following salts from the least soluble in water to the most soluble and justify your choice: KCl,NaCl, CsCl ,LiCl, RbCl.

b) As student in research you would like to distinguish between the above salts. Which quick chemical test will you carry out and what observations do you expect from that experiment?

2. Sodium is reacted with water:

a) What will happen?

b) Write the equation of the reaction between Na and water

c) How would you test for the presence of the products formed?

3. You are provided with the following chemical compounds KNO3, Li2CO3, LiNO3, and K2CO3.Identify which one is described as follows:

a) A soluble compound in water, but which does not decompose on heat

b) A soluble compound in water, but which decomposes on heat releasing a gas that relights a glowing splint.

c) A soluble compound in water, but which decomposes on heat releasing brown fumes and a colorless gas that supports combustion.

d) A slightly soluble compound in water, but which decomposes on heat and releases a colorless gas that turns lime water milky.

4. Group 1 consists of the elements Li, Na, K, Rb, and Cs.

a. The first member of the group often shows anomalous properties. Give two properties in which the behavior of Lithium is abnormal and explain why.

b. How does each of the following properties of the elements in Group 1 change with the increasing atomic number? Explain why.

i. Atomic radius

ii. Ionization energy

iii. Reducing properties

iv. Reactivity with water

v. Electronegativity

c. How will successive ionization energies of Na vary?

d. Why is the Na+ ion formed in normal chemical reaction rather than Na2+?

e. How are ionization energies related to the reactivity of these elements?

Key unit competence: Compare and contrast the chemical properties of the Group 13 elements and their compounds, in relation to their position in the Periodic Table.

Learning objectives

By the end of this unit, I will be able to:

•State the physical properties of Group 13 elements

•Explain the reactivity of Group 13 elements with oxygen, water, halogens, dilute acids and sodium hydroxide

•Describe the properties of oxides, hydroxides and chlorides of Group 13 elements•State the uses of Group 13 elements and their compounds

•Compare and contrast the reactivity of Group 13 elements with oxygen, water, halogens, dilute acids and sodium hydroxides

•Perform experiments to show the solubility of Group 13 compounds

•Practically illustrate the amphoteric properties of aluminium oxides and hydoxides

•Identify the anomalous properties of boron and its compounds

•Perform chemical tests for the presence of aliminium ion in the solution.

The elements of group 13 of the periodic table are boron (5B), aluminium (13Al), gallium (31Ga), indium (49In), and thallium (81Tl). Group 13 elements have a general outermost electronic configuration of ns2 np1 .This group marks the beginning of the p-block elements. It is sometimes called the Boron group.

8.1. Physical properties of group 13 elements

Activity 8.

1In groups learners make research in library or on internet, discuss and explain why

a. Group 13 elements have higher melting point than group 1and 2 elements.

b. Boron has higher ionization energy than other element of the same group.

With the exception of boron, group 13 elements are metals. Boron is a non-metal element with high melting point and low density.

Aluminium is a metal element and has a low density, it is a good conductor of heat and electricity, shiny, malleable, ductile and it has higher melting point than groups 1and 2 metals due to strong metallic bond resulting from 3 valency electrons involved in making metallic bonding in aluminium metal.

In small atoms electrons are held tightly and are difficult to remove, while in large ones they are less tightly held since they are far away from the nucleus and are easy to remove so that the ionization energy decreases down the group as the atomic radius increases.

The greater the forces of attraction and hence the boiling point and melting point decrease down the group as the atomic radius increases.

Table1: Physical properties of Group 13 elements

The 1st IE of boron is higher than the other member of the group due to its small size.

Checking up 8.1.

1a)Write the electronic structure of Al(Z=13).

b) Write the equation that shows how aluminium forms its ion.

c) Explain why aluminium forms Al3+ ion instead of Al1+2. Explain the cause of the strong bond in aluminium metal

3. Within the elements of group 13 elements down the group give and explain the relationship between atomic radius and:

i)the boiling point

ii) ionization energy

4.The following table shows 3 unkown group 13 elements A,B,C and some of their physical properties. Predict among the group 13 elements whichone should correspond to A,B,C.Justify your answer

5. Match the element in column A to the corresponding atomic radius in column B.Justify your answer

8.2. Reaction of aluminium

8.2.1. Reaction of aluminium with oxygen

Activity 8.2(a)

In groups learners make research, discuss and explain the following statement:

(a)aluminium is a metal therefore it forms compounds with the covalent character.(

b) aluminium forms amphoteric oxide.

Activity 8.3 (b)

Experiment: burning an aluminium metal in air/oxygen and amphoteric properties of Al2O3

Apparatuses: deflagrating spoon, Bunsen burner, glass beaker

.Chemicals: aluminum sheet, hydrochloric acid, sodium hydroxide Other requirements: match box and petroleum gas

Procedure:

1. Cut a small piece of aluminium

2. Place it on a deflagrating spoon and heat it in a non-luminous flame

3. When combustion is complete, divide the solid residue into two portions and place each portion in two different glass test tubes

4. Add 10ml solution of HCl To the first test tube and 10 ml of NaOH to the second test tube. Observe

Study questions

a) Write the equations of reactions that take place when aluminium metal is burnt in air

b) Name the product that was formed.

c) Write the equations of reaction between the product in (b) with HCl and with NaOH

Aluminium burns in oxygen to form aluminium oxide

•In its compounds, Aluminum occurs exclusively in the +3 oxidation state. It rapidly reacts with oxygen in air to give water in soluble coating of Al2O3. This oxide layer protects the metal beneath from further corrosion.

•All other Group 13 elements also produce compounds of the formula M2O3:

) where M represents Al, Ga, In, or Tl

Aluminum, gallium, and indium have +3 as the most stable oxidation state, whereas thallium has +1 oxidation state as the most stable.

The most common oxide form of boron, B2O3 or boron trioxide, it is produced by heating boric acid:

8.2.2. Reactions with acids

Activity 8.2 (c)

Reaction of aluminium with acids

Experiment

Learners perform experiments to investigate the reaction of aluminium with moderately concentrated HCl on warming and concentrated H2SO4

Apparatuses: Bunsen burner, gas jar, thistle, water bath, delivery tube, wooden splint

Chemicals: aluminum sheet, concentrated hydrochloric acid concentrated H2SO4, solu-tion of KMnO4. Other requirements: match box and petroleum gas

Procedure:

Put aluminium powder into two different flat bottomed flasks and connect on each a thistle, delivery tube as shown on the figure A and B.

a. Pour in the thistle of the apparatus A moderately concentrated HCl open the tap of the thistle and let HCl flow on Aluminium powder in the flat bottomedflaskCollect the gas evolved and test the gas using a burning wooden splint.

b. Pour in the thistle of the apparatus B concentrated H2SO4, open the tap of the thistle and let H2SO4 flow on Aluminium powder in the flat bottomed flaskCollect the gas evolved and test the gas using a violet solution of KMnO4.

Study Questions

a. write the equations of reaction when alumimium reacts with:

(i) moderately concentrated HCl

(ii) concentrated H2SO4

b. From your observations, explain the chemical test described for the gas evolved in each case, illustrate by chemical equations if any.

•Reaction with HCl '

Aluminium reacts when warmed with moderately concentrated hydrochloric acid forming aluminium chloride and hydrogen gas:

Hydrogen is tested using a burning splint.Apop sound is heard during the test.

•Reaction with H2SO4

Aluminium does not react with dilute sulphuric acid but reacts with concen-trated sulphuric acid forming aluminium sulphate, sulphur dioxide and water:

Sulphur dioxide gas is tested using a violet solution of potassium manganate(VII)that decolorizes to give colorless (aq) ion

Note: Aluminium does not react with HNO3 because of the insoluble layer of Al2O3 formed which prevents further reaction.

8.2.3. Reaction with alkalis

Activity 8.2 (d):

Reaction of aluminium with concentrated NaOH solution

Experiment

Learners perform experiments to investigate the reaction of aluminium with NaOH solution

Apparatuses: thermometer, pyrex beaker,stirrer

Chemicals: aluminium powder,40% sodium hydroxide solution

Procedure:

•Prepare 40% of sodium hydroxide by mixing 60cm3of water with 40g of sodium hydroxide

•Take 0.5 g of aluminium powder into a pyrex beaker

•Pour the solution of sodium hydroxide in the pyrex beaker containing aluminium powder and allow the reaction to proceed for about 5 minutes.

•Use thermometer to record the temperature during the process

Aluminium reacts vigorously with sodium hydroxide solution forming sodium aluminate and hydrogen gas.

8.2.4. Reaction with halogens

Activity 8.2.4

a. In terms of s,p,d,f orbitals give the electronic structure of aluminium (Z=13)

b. How many valency electrons does aluminium possess? Explain how you reach this conclusion

c. Give the formula of the compound formed between aluminium and chlorine(Z=17)

d. Explain why the compound AlCl is not formed

•Aluminium fluoride is made by direct combination of the metal with fluorine.

• Aluminium chloride is made by passing chlorine gas over heated aluminium metal

Aluminium can be reacted with moderately concentrated HCl

Al, Ga, In and Tl react with halogens to give binary halides of formula EX3. In this case, all 3 valence electrons are used in the reaction. All trihalides of group 13 elements are known except Tl(III) iodide ,because iodine is a weak oxidizing agent.

E:is a group 13 element, TlI3 doesn’t exist.Fluorides are ionic and have high melting points due to the small size of fluoride ion.

For the same halide ion bonded to group 13 element the ionic character increases down the group as the size of cation of group element increases. For example AlCl3is more covalent than TlCl3.

For the same element of group 13 the covalent character the covalent character increases with increasing size of the anion. For exampleGaCl3 is ionic while GaI3 is covalent

The chemistry of gallium is similar to that of aluminium.

The chemistry of indium is similar to that of aluminium and gallium except that compounds containing the 1+ ion are known, such as InCl and In2O, in addition to those with more common 3+ ions.

The chemistry of thallium is completely metallic. For example, Tl2O3 is a basic oxide. Both the +1 and +3 oxidation states are quite common for thallium; Tl2O3 and TlCl3 and TlCl are all well-known compounds.

The tendency for heavier members of group 13 to exhibit the +1 as well as the expected +3 oxidation states results from an effect called “inert pair effect”.

In this effect, since s electrons are closer and more attracted by the nucleus down the group, they are less and less available to participate in bond formation of compounds; and only p electron will participate, hence the oxidation state +1.

For lighter members such B and Al the s and p valency electrons, having almost the same energy, are always available and used at the same time to form compounds where they are in oxidation state +3.

Checking up 8.2

1. Using chemical reactions differentiate and explain the action of HCl, H2SO4 and HNO3 on aluminium metal.

2. In the groups learners discuss and explain the following statement: aluminium utensils are not washed in strong alkaline solutions.

3. Write and balance the equation of reaction between

a) Aluminium and bromine

b) Gallium and chlorine

c) Gallium and hydrochloric acid

d) If b) and c) above give different products, explain why.

e) Indium and iodine

8.3. Oxides and hydroxides of group 13 elements

Activity 8.3

Write chemical reactions to illustrate the amphoteric character of aluminium hydroxide.

Aluminium oxide is amphoteric:

•it dissolves in mineral acids to form aluminium salts: in this case it acts as a base.

•it dissolves in caustic alkali to form aluminate: in this case it acts as an acid.

Aluminium hydroxide also shows amphoteric properties.

Checking up 8.3

a) Aluminium oxide is said to be amphoteric, whereas calcium oxide is said to be basic, yet the two oxides are metal oxides? Explain the origin of that difference

b) Explain the concept of amphoterism using aluminium hydroxide

8.4. Anomalous properties of boron

Activity 8.4

In groups learners make research on internet and in library, discuss and explain the following statement:

a) Boron is a bad conductor of electricity

b) Boron has higher boiling and melting points than other member of the group

c) Boron oxide is an acidic oxide

Boron – the first member of group 13 shows anomalous behavior due to the small size and high nuclear charge/size ratio, high electronegativity. This makes boron typically non-metal whereas the other members of the group are metals. Hereafter are some other anomalies.

The melting point and boiling points of boron are much higher than those of other elements of group 13.

Boron forms only covalent compounds whereas aluminum and other elements of group 13 form even some ionic compounds. The oxide of boron is acidic in nature whereas those of others are amphoteric and basic.

The trihalides of boron (BX3) exist as monomers. The reason being that due to its small size, it cannot accommodate 4 large sized halogens atoms around it.

The hydrides of boron i.e. boranes are quite stable and are formed by many molecules such as diborane (B2H6) ,triboranes(B3H8)...while those of aluminum are unstable.

Dilute acids have no action on boron; other group 13 elements are dissolved in acids to liberate H2. Borates are more stable than aluminates.

Boron does not decompose steam water while other members do so.Boron combines with metals to give borides e.g. Mg3B2. Other members form simply alloys.

Concentrated nitric acid oxidizes boron to boric acid but no such action is noticed with other group members.

Checking up 8.4

a) What is the cause of abnormal behavior of boron

b) State any anomalous properties of boron

8.5. Identification of Al3+ ion in solution

Activity 8.5

In group of 6 learners perform experiment to test for the presence of Al3+ in solution:

Apparatuses: beaker, test tubes, droppers, test tube rack

Chemicals: aluminum salt solution, sodium hydroxide, ammonia solution

Procedure:

•Pour about 2cm3 of a solution of aluminium salt into 2 different test tubes

•To 2cm3 of solution of aluminium salt in the first test tube add drop by drop a solution of sodium hydroxide until excess. Then observe what happens.

•In a second test tube, to a 2cm3 of aluminium salt add drop by drop a solution of ammonia until excess and observe

Study question

Make a report on your findings and explain what happened

Aqueous solution of aluminium salt when reacted with a solution of sodium hydroxide, a white precipitate of Al(OH)3 is formed. Al(OH)3 precipitate dissolves in excess of sodium hydroxide as soluble complex ion:

Transparent and colorless solution

When reacted with a solution of ammonia, aluminium ion Al3+ produces a white precipitate of Al(OH)3 insoluble in excess ammonia solution

Checking up 8.5

State the reagent that should be used to distinguish between a solution containing aluminium chloride salt and that containing calcium chloride salt. Illustrate your answer by chemical equations and expected observations.

Key unit Competence

To be able to compare and contrast the chemical properties of the Group 14 elements and their compounds in relation to their position in the Periodic Table.

Learning objectives

By the end of this unit, learners should be able to:

•Compare and contrast the physical properties of Group 14 elements.

•Compare the relative stabilities of the higher and lower oxidation states in oxides.

•Distinguish between the chemical reactions of the oxides and chlorides of Group 14 elements.

•Explain the trends in thermal stability of the oxide, halides and hydrides of Group 14 elements.

•Explain the variation in stability of oxidation state of +2 and +4 down the Group 14 elements.

•Define the diagonal relationship.

9.1. Physical properties of group 14 elements.

Group 14 elements are carbon (6C), silicon (14Si), germanium (32Ge), tin (50Sn) and lead (82Pb). The elements of group 14 have one important feature in common, each of them has four valence electrons. They are characterized by the electronic configuration of 4 electrons in the outermost shell (ns2np2 ). The group is also called carbon family.

Group 14 elements show the clearest trend from non-metal to metal character down the group from carbon to lead.

In group 14 of the periodic table, there is a considerably greater change in physical properties from top to bottom in the group than there is for example, in the alkali metals of group 1 on the left and the halogens of group 17 on the right.

Carbon is a typical non-metal element, silicon and germanium are semi-metals or metalloids, tin and lead are metals. However, the elements of group 14 have one important feature in common, namely each of them has four valence electrons.

In the compounds of carbon, it almost invariably completes the valence shell by forming four covalent bonds. When it forms 4 simple covalent bonds, they have a tetrahedral arrangement around the carbon atom.

Silicon and germanium also usually form four simple covalent bonds with a tetrahedral arrangement.

Tin and lead, which are larger in volume have smaller ionization energies than carbon and silicon, and they often lose just two of their valence electrons to form the Sn2+ and Pb2+ ions.

Activity 9.1

1) Write the electronic configuration of carbon, silicon and tin in terms of s,p, d and f notation.

2) Indicate 2 uses of carbon and lead in daily use.

3) Explain how carbon forms 4 bonds in its compounds, e.g in CO2 .

4) Write the chemical formula of:

a) 2 inorganic compounds of carbon.

b) 2 organic compounds of carbon.

Table 9.1: Physical properties of Group 14 elements

•Carbon, the first element of the group has two main allotropes: graphite and diamond.

•In graphite allotrope of carbon, each carbon is bonded to 3 other carbon atoms to form a hexagonal structure. The structure of graphite is made of hexagonal layers which are attracted to each other by weak Van der Waals forces such that the layers slide over each other to make the structure soft(Fig.9.1). In graphite structure, there are delocalised double bonds with mobile electrons that allow graphite to conduct electricity.

•In diamond, each carbon is covalently bonded to 4 other carbon atoms forming a giant tetrahedral structure that makes it to be very hard (Fig.9.1). In diamond, there are no mobile electrons as in graphite, hence diamond does not conduct electricity.

.As you move down the group in the carbon family, the atomic radius and ionic radius increase while the electronegativity and ionization energy decrease.

•Atomic size increases on moving down the group due to additional electronic shells

.•Density increases as you move down the group.

•Carbon is the only element in the family that can be found in pure form in nature, as diamond and graphite.

•Lead is the only element of group 14 that does not exist in various allotropes.

•Tin occurs as white, grey and rhombic tin.

•Group 14 elements have much higher melting points and boiling points than the group 13 elements.

•Melting and boiling points tend to decrease as you move down the group mainly because•inter atomic bonding between the larger atoms reduce in strength as you move down the group.

Moving down the group, there is an increase in atomic size which results in less attraction of valence electrons by the nucleus. This change results in weaker metallic bonding down the group and therefore there is a decrease in melting point, boiling point, enthalpy change of atomization and first ionization energy.

The decrease in first ionization energy from silicon to lead is relatively little compared to that from carbon to silicon because there is a large increase in nuclear charge which counterbalances the increase in atomic radius from silicon to lead.

i) The increase in metallic character down the group causes a general increase in conductivity.

arbon is typically a solid, non-metal.Carbon graphite is a non-metal but conducts electricity due to delocalized electrons in its structure.

In its compounds, carbon almost invariably completes its valence shell by forming four covalent bonds

Silicon is solid at room temperature and pressure, it is a semi-metallic element and semi-conductor of electricity which is the second most abundant element on earth, after oxygen.

Germanium is solid at room temperature and pressure which usually forms four covalent bonds with a tetrahedral arrangement. Germanium, is a hard, brittle, greyish-white and crystalline semi-metallic element which is a semi-conductor of electricity.

Tin is a silver-white solid metal that conducts electricity at room temperature and pressure. At temperatures below 13 °C, it often changes into an allotropic (distinctly different) form known as grey tin, which is an amorphous, greyish powder.

Lead, is a dense, bluish-grey solid metallic element at room temperature and pressure which conducts electricity and has a smaller ionization energy than carbon and silicon, it often loses two of its valence electrons to form the Pb2+ ions.

Special features of carbon chemistry

i. Catenation

Catenation is the ability of an element to form bonds between atoms of the same element to form long chains or ring: -C-C-C-C-C-C-C-C-C-C-C-C-C-

Carbon forms chains and rings, with single, double and triple bonds. To be able to catenate, an element such as carbon must have a valency of 2 and above.

It should also form bonds like C-C which are similar in strength to those of C to other elements, particularly C-O bonds.

Silicon forms -Si-O-Si-bonds predominantly.

ii) Multiple bonds

Carbon forms double bonds and triple bonds between carbon atoms and that bonding is formed by one Sigma bond and one π bond for double bond, one Sigma bond and two π bonds in a triple bond.

Checking-up 9.1

1. Explain the reason why diamond has a higher melting point than silicon.

2. Discuss the increase in metallic character when moving down in group 14 elements from carbon to lead.

3. Diamond and graphite are allotropes of carbon,

a) Draw their three dimensional structures.

b) With reference to their structures, compare the hardness of diamond and graphite.

c) With reference to their structures, compare their electrical conductivity and explain.

4. Germanium has the same structure as diamond. Explain the type of bonds that exist in the two elements.

5. The first element in a group in the periodic table exhibits anomalous properties compared with other members. Use carbon to illustrate this statement.

9.2. Chemical properties of Group 14 elements

Activity 9.2 (a)

1. Get a piece of charcoal and burn it. Observe and write the chemical equation that represents the change that takes place when the charcoal burns.

2. a) Put about 1 gram of carbon charcoal in a boiling tube.

b) Add 1 ml of concentrated nitric acid.

c) Heat strongly on a Bunsen burner flame using a test tube holder.

d) Observe and note the changes during heating.

e) Deduce the chemical changes that have occurred.3. Write the molecular structure of carbon dioxide, carbonate ion and carbon monoxide.

4. Describe how CO2 gas dissolves in water and state the nature of the solution formed when it is in aqueous solution.

5. Describe 2 chemical properties of amphoteric substances.

All the elements of Group 14 form four single bonds such as in tetrahalides of formula CCl4, SiCl4, GeCl4, SnCl4 with the exception of PbBr4 and PbI4 that do not exist because bromine and iodine are not sufficiently strong oxidizing agents to convert Pb to Pb4+. Germanium, tin and lead can ionize to form Ge2+, Sn2+ and Pb2+ and also are capable of forming Ge4+, Sn4+, Pb4+ respectively.

Germanium dioxide GeO2 is more stable than germanium monoxide GeO. The oxide of tin in +4 oxidationstate is slightly more stable than that in +2 state. Silicon oxide predominantly exists as SiO2. SiO is too unstable to exist at room temperature and pressure but may exist at about 2000 0C. Carbon oxides are stable in +4 state. Thus CO which is unstable reacts exothermically to form the stable compound of CO2. In lead compounds, the +2 state is more stable than +4 state. PbO2 is unstable and is a very strong oxidizing agent.

The stability of compounds with the oxidation state of +2 in group 14 elements generally increases on moving down the group from carbon to lead.

The stability of compounds with the oxidation state of +4 in group 14 elements increases on moving upward the group from lead to carbon. The bonding in tetravalent compounds is predominantly covalent.

Inert pair effect

Inert pair effect refers to the inability of the outermost s-electrons to participate in chemical bonding in Ge, Sn and Pb elements of group 14 and hence only 2 electrons of the outermost p orbital are involved.

The outermost s sub-energy level electrons are much more tightly attracted to the nucleus than the outermost p orbital. As we move down the group, the difference in energy level between s and p orbitals becomes wider.

So if we use weak oxidizing agents, only 2 electrons in p orbitals are removed. If we use a strong oxidizing agent, 2 electrons in s orbital and 2 electrons in p orbital are all removed from the shell.

Reaction of group 14 elements with chlorine:

The group 14 elements mainly form tetrachlorides in the form MCl4 but lead doesn’t directly with Chlorine.

PbCl4 is unstable molecule and is prepared from PbO2. It decomposes to make stable PbCl2

Reaction of group 14 elements with hydrogen: Carbon graphite reacts with hydrogen gas at 30 0C to 950C to produce a mixture of methane (CH4), ethane (C2H6) and propane (C3H8). The primary product is methane and the other products are produced due to free radical reaction of the methane.

Apart from carbon, other elements of group 14 do not react with hydrogen directly.

Activity 9.2 (b)

1. Put about 0.5 g of Al2O3solid into2 test tubes then add 2ml of water.

2. Pour 3 ml of 1 mol/litre HCl solution in the third test tube and add 2 drops of universal indicator.

3. Pour 3 ml of 1 mol/litre NaOH solution in the fourth test tube and add 2 drops of universal indicator.

4. Put the first portion of Al2O3 solid in the test tube containing HCl solution and shake gently.

5. Put the first portion of Al2O3 solid in the test tube containing NaOH solution and shake gently.

6. Note the observable changes in both tests.

7. Put about 0.1 g of MgO in a test tube.

8. Add 2ml of water to the MgO solid in the test tube and shake the mixture.

9. Put 2 drops of phenolphthalein indicator in MgO solution.

10. Add 5ml of 1 mol/litre HNO3 to the solution containing MgO solution.

11. Note the observable changes.

12. Interpret the acid –base character of the solutions of CaO and Al2O3

Reaction of group 14 elements with acids and bases:

Carbon does not react with dilute acids but reacts with hot, concentrated acids:

Silicon reacts with concentrated HNO3:

Silicon reacts with concentrated sodium hydroxide (base)

Tin reacts with concentrated nitric acid to form tin dioxide and nitrogen dioxide:

Cold dilute nitric acid reacts with tin to form a mixture of tin II nitrate and ammonium nitrate solution.

Tin is sufficiently metallic to form various mixtures of Sn(SO4)2 and hydrated SnO2and sulphur dioxide when it reacts with hot, concentrated sulphuric acid.

Tin reacts with strongly non-oxidising acids in solution to form tin(II) salts and hydrogen gas:

Tin reacts with concentrated or molten sodium hydroxide base to form a stannate ion SnO22- and hydrogen gas:

Lead reacts with hot concentrated sulphuric acid according to the following equation to form lead(II) sulphate and sulphur dioxide:

The reaction of lead with concentrated nitric acid forms lead(II) nitrate and nitrogen dioxide as shown in the equation:

The reaction of lead with dilute nitric acid forms lead(II) nitrate and nitrogen monoxide as shown in the equation:

The reaction of lead with boiling, concentrated hydrochloric acid forms lead(II) chloride and hydrogen gas following the equation:

Lead which is amphoteric; reacts with hot, concentrated sodium hydroxide to form trihydroxoplumbate (II) and hydrogen gas as shown by the equation:

Checking-up 9.2

1) Explain why CO2 does not react with acids but other dioxides of group 14 react with acids.

2) State the expected observation when CCl4 and SiCl4 are treated separately with water. Write the relevant balanced equations if the reaction is possible.

3) Compare the stability of PbO and PbO2 upon heating. Write the balanced equa-tions.

4) Suggest the way that can be used to show that Pb is amphoteric.

9.3 Difference between the chemical reactions of the oxides and chlorides of Group 14 elements.

Activity 9.3

1. Measure 0.5g of lead oxide or decompose the same quantity of lead nitrate crystals by heating.

2. Divide it into 5 portions and put each portion in a test tube.

3. In the first test tube, add 2mL of dilute hydrochloric acid solution in which uni-versal indicator has been dissolved.

4. In the second test tube, add 2ml sodium hydroxide solution in which phenol-phthalein indicator has been dissolved.

5. Note the observations and deduce the acid–base nature of lead oxide.

Interpretation of results of the above activity 1

The reactions that take place are:

Conclusion

Lead (II) oxide behaves as a base when it reacts with acids and behaves as an acid when it reacts with bases, thus it is an amphoteric oxide.

Reaction of CO2 with water:

Carbon dioxide reacts with water to form carbonic acid.

he reaction is reversible, but H2CO3 dissociates to produce H+ and HCO3- ; the dissociation of carbonic acid is also reversible since carbonic acid is a weak acid.

Carbon dioxide is an acidic oxide, so H2CO3 does not react with acids.

Reaction of CO2 with sodium hydroxide:

Carbon dioxide reacts with sodium hydroxide, NaOH, to form sodium carbonate salt.

Reactions of CCl4 with water, acids and bases:

Carbon tetrachloride does not react with water, acids, and alkaline solutions.

Reaction of SiO2with water: Silicon dioxide does not react with water.

Reaction of SiO2with acids: Silicon dioxide reacts with HF acid only:

This reaction is the one used to write on the glass, for example on the windscreen of vehicle.

Reaction of SiO2 with bases: Silicon dioxide reacts with hot concentrated NaOH solution

Reaction of SiCl4with water:

Silicon dioxide undergoes hydrolysis in water, this is due to the fact that Silicon atom possesses vacant d-orbitals that can partially accommodate an oxygen atom during the reaction unlike in the case of CCl4..

Reaction of SiCl4 with acids, bases and water: SiCl4(l) does not react with acids, bases and water.

Reaction of SnO2 with water: SnO2 does not react with water.

Reaction of SnO and SnO2 with acids:SnO2does not dissolve in dilute acids but it reacts with hot concentrated sulphuric acid, H2SO4

Checking-up 9.3

1. For each of the following statements, state whether it is true or false.If you consider the statement to be false, justify your answer and if it is true, give a brief explanation or write a chemical equation.

a) All the oxides of group 14 with the formula of MO are amphoteric.

b) Some carbon allotropes can conduct electricity.

2. Write the equations to show the amphoteric nature of lead (II) oxide.

3. Discuss the reactions of tin (II) oxide and lead (II) oxide with sulphuric acid.

4. Explain why tetrachlorides of Si, Sn and Pb hydrolyse in water but carbon does not. Write the equations ofthereaction for the hydrolysis of SiCl4 and SnCl4

5. Explain why PbCl2 does not hydrolyse in water but PbCl4 does.

he thermal stability of dioxides and tetrachlorides of group 14 decreases down the group from carbon to lead due to the increase in sizeof the group 14 atoms.

The increase in size of group 14 atoms down the group results in weaker bonds. Thus PbCl4 is the least stable of the chlorides of group 14 elements.

9.4.1. Stability of oxides of group 14 elements

The thermal stability of XO increases downward the group from carbon to lead.

X represents: C, Si, Sn and Pb. The stability of XO increases in the order:

CO < SiO < SnO < PbO

Silicon monoxide is obtained by heating SiO2 and Si under special conditions.

The main valency of oxides of tin, is +4, in other words SnO2 is the most stable tin oxide; the valency of +2, SnO, is subsidiary.

On exposure to air at room temperature, SnO2 is formed.

The main valency of oxides of lead, is +2, in other words, PbO is the most stable lead oxide; the valency of +4 is subsidiary. PbO is the most stable oxide of group 14 elements. PbO2 cannot be prepared from lead monoxide because it decomposes at 3000C.

9.4.2. Thermal stability of halides of group 14 elements

CCl4, SiCl4 and GeCl4 are very stable and they do not decompose even at high temperatures.

SnCl4 decomposes only when it is heated to form SnCl2 and Cl2, so SnCl4 is more stable.

PbCl4 decomposes readily without heating to form PbCl2 and Cl2.

Thermal stability of tetrahalides of group 14 halides decreases down the group from carbon to lead because the bonds in Group 14 halides become longer and weaker down the group.

Thermal stability of tetrahalides of group 14 also decreases from fluorides to iodides.

All the tetrahalides are volatile covalent compounds except SnF4 and PbF4 which have some ionic character.

PbBr4 and PbI4 do not exist since iodine and bromine are not strong oxidizing agents enough to remove all 4 valence electrons.

9.4.3. Thermal stability of hydroxides of group 14 elements

Lead (II) hydroxide Pb(OH)2 decomposes at temperatures ranging between 1000C and 1450C to produce PbO and H2O .

Tin (II) hydroxide Sn(OH)2 decomposes easily between 600C and 1200C to produce SnO and H2O.

Carbon and silicon are non-metals, so they don’t form hydroxides.

9.4.4. Chemical tests of Sn2+, Pb2+, HCO3- :

Table 9.2: Chemical tests of ions, Pb2+,Sn2+, HCO3- (aq) and CO32-(aq) :

•Mercury chloride, HgCl2 solution is reduced to white precipitate of dimercury dichloride, Hg2Cl2 which is further reduced to a grey-black deposit of Hg by Sn (II) but not with Sn(IV)

•Excess NaOH solution causes the white precipitate of SnO and SnO2 oxides to re-dissolve.

•Sodium stannate (II) reduces the colorless bismuth nitrate, Bi(NO3)3 solution to elemental bismuth (black deposit).

•Sodium stannate (IV) has no reducing properties on Bi(NO3)3

Table 9.3: Identification of ions

When heat is applied to a solid, to test the presence of carbonate or hydrogencarbonate, use calcium hydroxide solution (lime water) which turns milky in the presence of

CO2 formed by the decomposition of the sample.

CO2 can evolve from compounds containing: CO32-, HCO32-, C2O42- .

Checking-up 9.4

1.Give an explanation why carbon tetrachloride molecule is non-polar despite the fact that the electronegativity difference between carbon and chlorine is so big.

2. State what is observed when lead dioxide is warmed with concentrated HCl acid solution and write the equation for the reaction.

3. Lead hydroxide Pb(OH)2 can react with an acid and with a base. Write the ionic equation for the reaction of Pb(OH)2 with: a) An acidb) A base.

4. Explain the reason why lead tetrachloride is unstable such that it decomposes easily compared to other tetrachlorides of group 14 tetrachlorides.

5. You are given a mixture of HCO3- and SO42- ions, explain how you can show that the mixture contains HCO3- ions.

6. Give a chemical reagent test(s) you can use to confirm that a solution contains Sn2+ ions.

7. a) Give a chemical reagent test you can use to differentiate between HCO3- and CO32- ions and deduce the observable changes.

b) Write the chemical equations of reactions to illustrate the reactions that have taken place.

8.Explain the trend of thermal stability of group 14 tetrachlorides and dichlorides on descending down the group.

9.5 Variation in stability of oxidation state +2 and +4 down the Group 14 elements.

Activity 9.5

1. Write the electronic configuration of carbon and silicon.

2. Predict the stability of dichloride compounds in group 14 elements from car-bon to lead and explain

3. Discuss the influence of variation in electronegativity to the bond strength in compounds of group 12 from beryllium to barium.

Use the bond strength between MgCl2 and CaCl2 to explain the answer.

The 2 oxidation states exhibited by elements of group 14 are +2 and +4 in their compounds.

The oxidation state+4 involves loss of 4 valence electrons, s2 and p2, while the oxidation state +2 involves loss of only 2 electrons in the p2.

The stability of the +2 oxidation state increases from carbon to lead while +4 oxidation state increases from lead to carbon.

Inert pair effect refers to the inability of the outermost s-electrons to participate in chemical bonding in Ge, Sn and Pb elements of group 14 and hence only the outermost p-electrons are involved.

The electrons in s orbital are much more tightly bound to the nucleus than p-electrons. As we move down the group, the difference in energy level between s sub-shell and p sub-shell becomes wider.

So if we use weak oxidizing agents, only 2-p electrons are removed. If we use a strong oxidizing agent 2 s-electrons and 2-p electrons are all removed from the shell.

If the elements in group 14 form +2 ions, they will lose the p electrons leaving the s-electrons pair unused. For example, to form Pb2+ ions lead will lose the two 6p electrons but the 6s electrons will remain in its sub-energy level.

The inert pair effect shown in Pb2+ explains why the compounds of lead are predominantly ionic.

The compounds of group 14 with +4 oxidation state are more covalent than those with +2 oxidation state. The covalent character increases as a bigger anion is bonded to the group 14 element. For example, PbBr2 is more covalent than PbCl2. A big anion is more polarisable than a small anion.

Checking up 9.5

1. Explain the reason why carbon does not form CCl2 compound.

2. Explain the term”inert pair effect”. Give examples to illustrate.

3. Indicate which compound between PbCl2 and PbF2 is more covalent than an-other and justify.

4. Explain the reason why PbBr4 and PbI4 do not exist whereas PbCl4 exists

9.6 Uses of Group 14 elements

Activity 9.6

1. State 2 uses of carbon in daily use.

2. Diamond is considered to be very hard, predict its application.

3. Carbon forms hydrocarbons such as methane when combined with hydrogen. State where methane is obtained in Rwanda and indicate its uses.

4. Silicon is a semi-conductor. Predict its uses in electric gadgets.

5. Tin resists corrosion. Discuss where it can be used to prevent corrosion.

Group 14 elements are used in numerous applications. The wide applications are due to the fact that some of the elements in the group are non-metals, metals and metalloids.

Carbon uses:

•As a component of fuel for combustion as charcoal or coal.

•As the main component of crude oil and its derivatives used in our everyday life such: fuel, plastics, etc...

•As good chemical reducing agent used in extraction of metals (metallurgy).

.As a lubricant in moving parts of machines, to make electrodes, in lead pencils when mixed with clay. •Carbon isotope, C-14 isotope is used in archaeological dating.

•Diamond is used to make glass cutters, drilling devices and as abrasive for smoothing hard materials as precious gemstone in jewelry and ornamental objects; it is also a precious stone appreciated in jewelry.

Silicon uses:

•Silicon is used as a semi-conductor in transistors in electrical gadgets such as radios, computers, amplifiers etc.

.•Silicon in form of silicates is used in ceramics and in glass production.

•Silicon is also used in medicine to make silicone implants.

•Many rocks that we use for building our houses and other buildings are Silicates.

•Ferrosilicon alloy is used as a deoxidizer in steel manufacture.

•Silicon dioxide can be used to produce toothpastes and in semiconductors; silicon dioxide is the main component of sand, a raw material in the manufacture of glass.

Germanium uses:

•Germanium being a metalloid, is used in transistors in electrical gadgets such radios, computers, amplifiers etc..

Tin uses:

•Tin is used in plating steel sheets to resist corrosion; it is used for example to make canned tins to avoid the corrosion of the materials which are in contact with an acid medium.

Lead uses:

•Lead is used in making linings of vessels which are used in industrial production of sulphuric acid.

•Lead bricks alloyed with 4% antimony is used in radiation shielding

•Lead is used in accumulator plates of the batteries and as shielding materials against dangerous radiations such as X-rays, gamma rays etc.

Checking –up 9.6

1. Describe how diamond is used on a large scale.

2. Describe the materials that are manufactured using lead, tin and silicon as the main component.

3. Describe 3 compounds which contain:

i) Silicon

ii) Lead.

9.7 Define diagonal relationship

Activity 9.7

1. Discuss and compare the type of bonding in beryllium and aluminium elements.

2. Describe 2 chemical properties of oxides of beryllium and aluminium com-pounds.

3. Describe the solubility of lithium and magnesium in water.

4. Discuss the similarities in physical properties and chemical properties of BeCl2and AlCl3.

9.7.1.The diagonal relationship in groups1 & 2, 13 &14 elements

Diagonal relationships are similarities between pairs of elements in different groups which are adjacent to one another in the second and the third period of the periodic table.

These pairs are in Groups 1 and 2(Li/Mg), Groups 2 and 13(Be/Al) and Groups 13 and 14(B/Si). They exhibit similar properties; for example, boron and silicon are both semi-conductors, they form halides that are hydrolyzed in water and have acidic oxides.

NB: This relationship does not hold for allcertain groups; the only significant relationships are the ones mentioned above.

Such a relationship occurs because as we move across and descend in the periodic table, there are opposite effects.

On moving across a period of the periodic table from left to right, the size of the atoms decreases, and on moving downward a group, the size of the atoms increases.

If you look into those two movements in the first 3 positions of Period 2 of the periodic table, the result is: Li and Mg, Be and Al, B and Si, all couples in diagonal relationship, or diagonally adjacent resemble each other in their behavior. This resemblance, called diagonal relationship is attributed to the similarity of their size; for example the atomic radii of lithium and magnesium are similar (1.55×10-10m and 1.6×10-10m respectively).

Due to this relationship, those elements will show similar chemical properties and their compounds will have similar properties.

9.7.2. Diagonal Relationship of Li with Mg

At room temperature and pressure:

1. Li and Mg react with air to form only normal oxides, whereas Na forms a peroxide but metals below Na in the group, form superoxide.

2. The bicarbonates of Li and Mg do not exist in solid state, they exist in solution only.

3. Nitrates thermally decompose, to give oxides, nitrogen dioxide and oxygen:

4. Hydroxides of both Li and Mg decompose on heating to give oxides.

Hydroxides of both Li and Mg are weak alkalis.

NB: Hydroxides of other alkali metals are stable on heating while their nitrates

9.7.3. Diagonal Relationship of Be with Al

•Due to many similarities in their properties and reactions, aluminium and beryllium will be presented together even though they are in different groups of the periodic table.

Atomic radii: Be= 111 pm, Al= 143 pm

Ionic radii: Be2+= 31pm, Al3+ = 50 pm

Boiling point: Be = 2477oC, Al =2467oC

•Oxides: BeO and Al2O3 are both amphoteric, the rest of Group 2 oxides are only basic.

•Reaction with dilute acids: Be and Al are fairly resistant unless amalgamated or when they are very finely divided, other Group 2 metals readily react.

•Like aluminum, beryllium is not readily attacked by acids because of the presence of an oxide film on the surface of the metal.

• Reaction with conc. nitric acid: Be and Al are rendered passive, the rest of Group 2 metals readily form salts.

•Beryllium and aluminium have an appreciable covalent character of compounds (e.g. the chlorides are predominantly covalent).

9.7.4. Diagonal relationship between Boron and Silicon

Due to its small size and similar charge/mass ratio, boron differs from other group 13 members, but it closely resembles silicon, the second element of group 14 to exhibit diagonal relationship. Some important similarities between boron and silicon are given below:

•Both boron and silicon are typical non-metals that exist as non-metallic giant covalent 3D lattices, having high melting points and boiling points, they have almost the same densities (B=2.35gml–1 S=2.34 g//ml), low atomic volumes and bad conductor of electricity. However both elements are used as semiconductors.

•Both B and Si do not form cations but only form covalent compounds.

•Both B and Si form numerous volatile hydrides which spontaneously catch fire on exposure to air and are easily hydrolyzed.

•Both B and Si form weak acids like H3BO3 and H2SiO3when they react with water.

•B and Si are non-metals and their oxides are readily soluble in alkalis.

Checking-up 9.7

1. Write the chemical equation for the reaction between sodium hydroxide and:

a) Boron oxide

b) Silicon dioxide.

2. Be and Al belong to two different groups but their chlorides present the same properties; explain the origin of that resemblance.

3.Describe physical properties ofhydrides of boron and silicon and their behaviour in water.

4. Explain the acidity of solutions of beryllium chloride and aluminium chloride.

9.8. End unit assessment

I: Fill in the following statements with a missing word:

1. The arrangement of atoms in diamond structure is called....................

2.......................is the only element of group 14 whose chloride does not hydrolyse in water.

3.....................is a semi-metallic element of group 14 whose oxide reacts with HF acid only.

4.....................is the only element of group 14 that does not exist in various allotropic forms.

5......................is the only element of group 14 whose compounds in the oxidation state of +2 is more stable than that of +4.

II. Answer the following questions:

6. Write the equations for the reaction of decomposition of:a) Lead (II) hydroxideb) Tin tetrachloride

7. Explain the amphoteric character of tinby using appropriate equations of reaction.

8. Discuss the stability of +2 oxidation state as you move down in group 14 elements.

9. Explain the reason why the melting and boiling points of group 14 elements decrease down the group.

10. Explain by using appropriate equations, how you can confirm the presence of Sn2+ ions in a solution.

11. Discuss the origin of similarities of chemical properties of boron and silicon, use appropriate equations to emphasize the explanation.

12. Given the following three chloride, CCl4, SnCl2, SnCl4. Classify them in two groups of ionic and covalent compounds and justify your answer.

13. Complete the following chemical equation if possible and justify:

Key unit competency: Compare and contrast the properties of Group 15 elements and their compounds, in relation to their position in the Periodic Table.

Learning objectives

By the end of this unit, the learner will be able to:

•Describe the physical properties of Group 15 elements.

•Describe the variation in the metallic and non-metallic character of Group 15 elements.

•Explain recall the physical properties of the allotropes of phosphorus.

•Describe the chemical reactions of nitrogen compounds.

•Describe the impact of nitrogen oxides to the environment.

•Describe the industrial preparation of ammonia and nitric acid.

•Explain the reactions of nitric acid with metals and non-metals.

•Describe the chemical properties of phosphorus compounds.

•State the uses of the group 15 elements and its compounds

10.1. Physical properties of group 15 elements and the relative inertness of nitrogen

Activity 10.1

In pairs:

1. Assign the physical state for each of the elements in group 15

2. Explain what is meant by the term “metallic character”.

3. Classify each element in this group as metal, non-metal or metalloid

.4. Study the following figure carefully and answer the questions that follow

a. Identify the molecule represented in the figure.

b. What type of bond is there in the molecule?

c. Suggest if the bond is strong or weak.

10.1.1.Physical properties of group 15 elements

a. Physical stateThis group consists of nitrogen (7N), phosphorus (15P), arsenic (33As), antimony (51Sb), and bismuth (83Bi). This group consists of atoms having 5 electrons in their outer energy level. Two of the electrons are in the s sub-shell, with 3 unpaired electrons in the p sub-shell: ns2np3 (px1py1pz1). Group 15 Elements often form covalent compounds, usually with the oxidation numbers +3 or +5.

All the elements of this group are polyatomic. Dinitrogen is a diatomic gas while all others are solids.

•Nitrogen is a colorless gas and is the major component of air (78%).

•Phosphorus has allotropic forms and is avery reactive non-metallic element,

•Arsenic, antimony (their oxides are amphoteric) and bismuth (its oxide is basic) are solidin nature.

Table 10.1: Physical properties of group 15 elements

b. Metallic character

Down group 15 elements, the atomic radius increases which makes the outermost electron to be less attracted by the nucleus as you move down the group. Therefore, less energy is required to remove the outermost electron, which results in the increase in the metallic character down the group. This results also in decreasing of ionization energy down the group.

Nitrogen and phosphorous are non-metals, with the metallic properties first appearing in arsenic and increasing down the group. Arsenic and antimony are metalloids. Bismuth is a metal.

10.1.2. The relative inertness of nitrogen

The most striking aspect of nitrogen chemistry is the inertness of N2 itself. Nearly four-fifth of the atmosphere consists of nitrogen, and the other fifth is nearly all oxygen, a very strong oxidizing agent. Nevertheless, the searing temperature of a lightning bolt or very high temperature in a vehicle motor combustion is required for significant amounts of atmospheric nitrogen oxides to form. Thus, even though N2 is inert at moderate temperatures, it reacts at high temperatures with H2, Li, Group 2 members, B, Al, C, Si, Ge, O2, and many transition elements. In fact, nearly every element in the periodic table forms bonds to N, but at high temperature.

Nitrogen is found as a diatomic molecule (N2) in nature. There is a triple bond between the two nitrogen atoms. This explains the very high strength of the nitrogen-nitrogen triple bond (N≡N) that requires high energy, 942 kJmol-1, to be broken. For comparison, the single bond (N-N) and the double bond (O=O) require 247kJmol- and 498 kJmol- respectively.

Checking Up 10.1

1. Briefly describe how each of the following factors varies in group 15 elements:

a) Atomic radius

b) Electron affinity

c) Melting point

d) First ionization energy

2. Explain the following observations:

a) In group 15 of the periodic table, metallic character increases as you move down the group.

b) The atomic radii of two elements A and B from group 15 are 0.121nm and 0.141nm respectively. Identify the element with more metallic character. Justify your answer.

10.2. Reactions of group 15 elements

Activity 10.2

1. Explain what is meant by the term:

a. Hybridisation

b. Electronic configuration

c. Briefly explain what is meant by:

d. Acidic oxide

e. Basic oxide

f. Amphoteric oxide

2. Respond by True or False and justify

a. Nitrogen is the only member of the group that readily forms multiple covalent bonds.

b. Nitrogen is the most electronegative member of group 15 elements and the only one that can be involved in the formation of hydrogen bond

c. Nitrogen is the only group member that has an oxidation state -3.

d. Only nitrogen form oxides of the form: E2O3 and E2O5

All group 15 elements exhibit a common valency of three. They can complete their octet structure in chemical combination by gaining three electrons.

However, with the exception of nitrogen, group 15 elements have vacant d-orbitals which they use to expand their octet to form compounds with a valency of five. For instance phosphorous has a covalency of 5 due to availability of easily accessible empty d orbitals which can be used for sp3d hybridization that allows it to have 5 unpaired electrons. Consider phosphorous, atomic number 15.

Ground state electronic configuration of P: 1s2 2s2 2p6 3s2 3p3 3d0 (three unpaired electrons)

After sp3d hybridization, the 5 electrons previously in 3s23p3 are redistributed in the five hybrid orbitals sp3d:

Hence phosphorus can now accommodate 5 more electrons or create 5 covalent bonds.

The same applies for any element in the group, such as Bi, that participates in a five covalency bonding.

Reaction with oxygen

All group 15 elements form two types of oxides: E2O3 and E2O5. The oxide in the higher oxidation state of the element is more acidic than that of lower oxidation state.Their acidic character decreases down the group. Group 15 oxides of the form E2O3 of nitrogen and phosphorus are purely acidic, those of arsenic and antimony amphoteric and that of bismuth predominantly basic.

Nitrogen gas does not react with air under normal conditions. Nitrogen (II) oxide is produced when nitrogen combines with oxygen. This is further oxidized into nitrogen (IV) oxide:

Reaction with water

Elements in this group do not tend to react with water. Nitrogen gas does not react with water. Phosphorous does not react with cold water. However, phosphorous vapour reacts with steam at high temperatures.

Reaction with chlorine

White phosphorous readily ignites in chlorine to form the trichloride and in excess of chlorine, the pentachloride is produced. Red phosphorous behaves similarly when heated:

Reaction with hydrogen

All Group 15 elements form hydrides of the type EH3, where E = N, P, As, Sb or Bi. The hydrides show regular gradation in their properties. The stability of hydrides decreases from NH3 to BiH3 which can be observed from their bond dissociation energy. Consequently, the reducing character of the hydrides increases down the group. NH3 is only a mild reducing agent while BiH3 is the strongest reducing agent amongst all the hydrides of the Group 15. Basicity also decreases in the order:

Under normal conditions the nitrogen gas is very inert, but, Nitrogen reacts with Hydrogen at a high temperature (about 500oC) and extremely high pressure (200-1000atm) over a catalyst (consisting of finery divided iron (Fe) and aluminum oxides a promoter). These are conditions used in the manufacture of ammonia, NH3:

Reaction with metals

All group 15 elements react with metals to form their binary compounds exhibiting –3 oxidation state, such as, Ca3N2(calcium nitride), Ca3P2(calcium phosphide), Na3As2(sodium arsenide), Zn3Sb2(zinc antimonide) and Mg3Bi2(magnesium bismuthide).

Nitrogen combines with many metals to form metal nitride:

Checking Up 10.2

1. Group 15 elements form compounds by gaining three electrons, for example in the formation of ionic nitrides such as Li3N, Ca3N2, Mg3N2 and AlN, in which nitrogen forms N3-. Other large atoms do not form M3-. Explain why.

2. State how the stability of +5 oxidation state varies in group 15.

3. The atomic number of an element A is 33.

a. Write the electronic configuration of A using s,p,d, notation.

b. Write the molecular formula of all possible oxide of A.

c. i) State whether each oxide of A you have given in (b) is acidic, basic, neu-tral, or amphoteric and justify.

ii) Write the equation of reaction to illustrate your answer.

10.3. Ammonia and nitric acid

Activity 10.3 (a)

Experiment: Laboratory preparation of ammonia

Materials and chemicals

Round bottom flaskor hard glass test tube, U-tube, 3 corks, 10 grams of calcium hydroxide, gas jars, bent delivery tube and straight delivery tube, 5 grams of ammonium chloride on a watch glass and calcium oxide lumps.

Procedure

1. Set up the apparatus as shown in the diagram, with the chemicals indicated. Do not start heating yet.

2. When everything is in position, heat the hard flaskandcollect several gas jarsof ammonia. Cover each jar with a glass slip and keep the jars for other experiments.

Activity evaluation questions

1. Record your observations

2. Write a balanced equation of the reaction that take place.

10.3.1. Laboratory preparation of ammonia and nitric acid

a. Laboratory preparation of ammonia

Ammonia is a covalent compound, consisting of nitrogen bonded to three hydrogen atoms. It exists as a colourless gas at room temperature and it is naturally produced during the decaying of nitrogenous organic compounds such as proteins. Ammonia has a characteristic pungent odour.It is less dense than air and thus collected by upward delivery method. In the laboratory it is prepared by heating a mixture of any ammonium salt and an alkali.

Examples

10.3.2. Industrial production of ammonia and nitric acid

Production of ammonia (Haber process)

The Haber process combines nitrogen from the air with hydrogen derived mainly from natural gas (methane) into ammonia. The reaction is reversible and the production of ammonia is exothermic.

This reaction, which is reversible, requires special conditions for an optimum yield of ammonia

•Pressure : 200 –1000 atmospheres

•Temperature: 450- 500 ̊C

•Catalyst: finely divided iron (Fe) catalyst, aluminium oxide which is added to the catalyst to improve its performance (a promoter). It makes it more porous and this provides a high surface area to the reaction and potassium oxide

The ammonia formed can either be liquefied or dissolved in water to separate it from the mixture.

i. Uses of ammonia

Agricultural industries are the major users of ammonia. Ammonia and urea are used as fertilizer, as very valuable source of nitrogen that is essential for plant growth.

Ammonia and urea are used as a source of protein in livestock feeds for ruminating animals such as cattle, sheep and goats.

Ammonia can also be used as a pre-harvest cotton defoliant, an anti-fungal agent on certain fruits and as preservative for the storage of high-moisture corn.

The pulp and paper industry uses ammonia for pulping wood and as casein dispersant in the coating of paper.

The food and beverage industry uses ammonia as a source of nitrogen needed for yeast and microorganisms involved in the fermentation process.

ii. Environmental impact for industrial production of ammonia

Making ammonia using the Haber process requires a lot of energy, which usually involves burning fossil fuels. This releases carbon dioxide which causes global warming.

b. Production of Nitric acid (Ostwald’s process)

In the industrial manufacture of nitric acid a catalytic oxidation of ammonia to nitrogen (II) oxide,NO, is carried out then a further oxidation of nitrogen (II) oxide produces nitrogen (IV) oxide, NO2. Nitrogen dioxide is passed through water sprays in a steel absorption tower to produce nitric acid. The excess nitrogen monoxide is recycled back for more oxidation. Platinum is used as a catalyst. There are three steps:

Nitric acid is a strong acid, i.e. dissociates completely when dissolved in water.

i. Chemical properties of nitric acid

Activity 10.3 (b)

Investigating reactions of HNO3with metals

Caution

Work in a well ventilated roomAvoid big yield of the gas when HNO3 reacts with non-metals and metals. The gas is toxic.Concentrated acidis corrosive; much care must be taken in using

Materials and chemicals

Concentrated nitric acid, copper turnings, magnesium ribbon, aluminium powder, carbon (or charcoal), sulphur powder, red and blue litmus paper, spatula, source of heat.

Procedure

1. Collect a pyrex beaker or a clean test tube.

2. Pour 5 cm3 of concentrated HNO3 in the beaker or 2 cm3 of the same acid in the test tube.

3. Add a small piece, 2g to 4g of magnesium to the container.

4. Using wet blue and red litmus paper, check the nature of the gas produced.

5. Record your observations.

6. Repeat the experiment but at step (3) replace magnesium with copper turnings.

7. Record your observations.

This reaction explains why nitric acid is called an “oxidizing acid”; in opposition of acids that cannot dissolve copper called non-oxidizing acids, e.g. HCl(aq), H3PO4(aq).

This reaction can be used to distinguish between a gold sample and a copper sample; gold and copper have almost the same color. If a sample is gold, it does not dissolve in nitric acid; if it is copper, it dissolves in nitric acid to form a blue solution and a brown gas NO2.

Activity 10.3 (c):

Investigate reaction of concentrated HNO3 with non-metals. e.g carbon, and sulphur,

Materials and chemicals

Concentrated nitric acid, carbon (or charcoal), sulphur powder, red and blue litmus paper, spatula, source of heat.

Procedure

1. Have a pyrex beaker.

2. Put one spatula of carbon powder into a beaker; add about 5cm3 of concentrat-ed nitric acid

3. Record your observations.

4. If no reaction is taking place, gently heat the mixture.

5. Test the gas produced with a wet blue and red litmus paper.

6. Record your observations.

Procedure 2

1.Collect a pyrex beaker.

2. Measure one spatula of sulphur powder into the beaker. Add about 5cm3 of concentrated nitric acid.

3. Record your observations.

4. If no reaction is taking place, gently heat the mixture.

5. Test the gas given off with a wet blue and red litmus paper.

6. Record your observations.You can also replace the concentrated HNO3 with dilute nitric acid (1M) and record your observations.

Activity evaluation questions

1. Write balanced equations of the reactions taking place when concentrated nitric acid is used.

2. Write balanced equations of the reactions when dilute nitric acid (1M) is used.

3. Write the formula of the gas given off in all the experiments done.

4. Is the gas released acidic or basic? Explain your answer.

i. Uses of nitric acid

The main industrial use of nitric acid is the production of nitrates as fertilizers. Nitric acid is neutralized with ammonia to give ammonium nitrate.

Others uses are: manufacture of dyes, artificial fiber and drugs, making gun powder and explosives such as trinitrotoluene (TNT)

ii. Environmental impact of industrial production of nitric acid

The manufacture of nitric acid may constitute an environmental problem if appreciable quantities of nitrogen oxides, NOx, are rejected into atmosphere; because they form acid rains, and can contribute the greenhouse effect and global warming.

The main gaseous emissions from the Ostwald process include NO and NO2. Both gases contribute to photochemical smog, and therefore, careful attention must be paid to minimizingthe emission of those gases into the atmosphere. Note that Nitrogen (I) oxide is a significant greenhouse gas.

Checking Up 10.3

1. a. State whether NH3 is a Lewis acid or Lewis base

b. Explain your answer in (a) above

2. Explain briefly how ammonia is prepared in the laboratory.

3. How nitric acid is prepared in the laboratory? Use equation to illustrate your answer.

4. Ammonia is a gas used to manufacture various compounds.

a. State 3 physical properties of ammonia gas.

b. The manufacture of Ammonia in the Haber process is a key industry that is linked to other useful chemicals.

i) Write a balanced equation for the reaction that occurs in the Haber process.

ii) State the conditions of temperature, pressure and catalyst used in this process.

iii. With the help of balanced equations, outline the steps involved in the manufacture of nitric acid from ammonia

5. Write equations to show how ammonia reacts with the following:

a. Oxygen in the absence of a catalyst

b. Oxygen in the presence of a catalyst

c. Copper (II) oxide

d. Hydrochloric acid

6. Write equations to show how nitric acid reacts with the following substances:

a. Copper

b. Sulphur

c. Potassium hydroxide

10.4. Phosphorus and chemical properties of its compounds

Activity 10.4 (a)

Study the figure below and answer the questions asked

1. Explain what is meant by “allotropes” of an element

2. How red phosphorus and white phosphorus differ?

3. Write balanced equations to show how the following oxides are obtained from phosphorous:

a) P2O3

b) P4O104. How does phosphorous react with chlorine?

10.4.1. Allotropes of phosphorus

By definition, allotropy is a property exhibited by some elements to exist in multiple forms with different crystal structures. Allotropes are any two or more physical forms in which an element can exist. Phosphorus exists in two main allotropic forms:

•White phosphorus

•Red phosphorusand

a. White phosphorus (P4)

White phosphorous (P4) exists as molecules made up of four atoms in a tetrahedral structure. Thetetrahedral arrangement results in ring strain and instability. That is why it is very reactive.

When prepared, ordinary phosphorus is white, but it turns light yellow when exposed to sunlight. It is a crystalline, translucent, waxy solid, which glows faintly in moist air and is extremely poisonous.It ignites spontaneously in air at 34°C and must be stored under water. It is insoluble in water, slightly soluble in organic solvents, and very soluble in carbon disulfide. White phosphorus melts at 44.1°C, boils at 280°C. White phosphorus is prepared commercially by heating calcium phosphate with sand (silicon dioxide) and coke in an electric furnace. When heated between 230°C and 300°C in the absence of air, white phosphorus is converted into the red form.White phosphorus spontaneously takes fire in contact with air. White phosphorus is considered and has been used as a chemical weapon.

b. Red phosphorus

It is a micro-crystalline, nonpoisonous powder. It sublimates (passes from the solid state directly to the gaseous state) at 416°C. As shown in the figure below, its structure is made of a chain of P4 units, a kind of polymer: ...(P4)n...

10.4.2. Chemical properties of phosphorous compounds

10.4.3. Phosphoric acid

a. Laboratory preparation of phosphoric acid

In the laboratory phosphoric acid can be prepared by boiling a mixture of red phosphorus with 50% nitric acid in a flask fitted with a reflux condenser on a water bath till no more oxides of nitrogen are liberated. Iodine acts as a catalyst. Phosphoric acid is a weak acid.

b. Reaction of phosphoric acid (H3PO4)with metals and bases

Activity 10.4 (b)

Experiment: Investigating the presence of nitrate ion (NO3-)

Chemicals and apparatus needed are:

1. Freshly prepared iron (II) sulphate solution

2. Concentrated sulphuric acid

3. Test tubes

4. Droppers

Procedure:

1. To about 1cm3 of suspected solution of nitrate, add an equal volume of prepared iron (II) sulphate and tilt the tube.

2. Add cold concentrated H2SO4 slowly on the sides of the test tube.

Study questions

1. Record your observations

2. What is the name given to this test?

10.4.4. Identification of PO43- and NO3- ions

10.4 (b) Identification of NO3- ions

a. Brown ring test: the test carried out in Activity 10.4(b) is called “Brown ring test”. When the test is carried out as indicated, a brown ring is slowly formed between the two layers (acid and aqueous layers).

b. Addition of concentrated sulphuric acid: When a solid nitrate is mixed with about 1cm3 concentrated sulphuric acid and warmed with Cu metal, brown fumes of nitrogen dioxide are formed.

c. Boiling nitrates with powdered Devarda’s alloy (a mixture of copper, zinc and aluminium) with NaOH solution gives a colourless gas with a chocking smell. The gas is ammonia.

Identification of PO43- ions

Formation of immediate yellow precipitate of ammonium phosphomolybdate when a solution of phosphate (V), PO43-, is warmed with nitric acid and ammonium molybdate solution confirms a phosphate (V) ion in solution. Heptaoxodiphosphates (V), P2O74-, and trioxophosphates (III), PO3-3, also give yellow precipitates with the reagents but at a slow rate. The yellow precipitate is due to the hydrolysis of phosphates (V) on warming.

10.4.5. Uses of group 15 elements and compounds

a. Uses of Phosphorus

The main uses of Phosphorus are:

1. Red phosphorus is mainly used on making matches.

2. White phosphorus and zinc phosphate are mainly used as a poison for rats.

3. Phosphorus is used in making phosphor bronze, which is an alloy of copper and tin containing phosphorus.

4. Phosphorus is used for the preparation of phosphorous compounds like P2O5, PCl3,PCl5 and phosphoric acid. Phosphoric acid, in turn, is used in the manufacture of phosphates which are used as fertilizer: NPK (Nitrogen, Phosphate, and Potassi-um) fertilizer.

5. It is used in making incendiary (fire causing) bombs, tracer bullets and for produc-ing smoke screen.

b. Uses of Nitrogen and its compounds

1. Cryopreservation (conservation at low temperature) uses nitrogen to conserve egg, blood, sperm and other biological specimens at very low temperature.

2. Nitrogen gas is used in laboratory or hospital to create inert atmosphere

3. Nitrogen is a constituent of almost every major class of drugs, including antibiot-ics. In the form of nitrous oxide, nitrogen is used as a pharmaceutical anesthetic agent.

4. Ammonia and Nitrates are used as fertilisers to increase soil fertility.

5. Ammonia is used in the manufacture of nitric acid, a starting material for the pro-duction of chemical fertilizers such as ammonium sulphate, ammonium nitrate, and many other kinds of nitrate fertilizers.

6. Ammonia is used in the manufacture of aminobenzene dyes.

7. Nitric acid is used in making explosives such as trinitrotoluene (TNT), and in the production of nitrates.

8. The compound ammonium nitrate mixed with large quantity of salt in water can be used as a freezing mixture.

c. Uses of Arsenic

1. Arsenic is used in semi-conductor electronic devices

2. Arsenic is also used in various agricultural insecticides and poisons.

3. Metallic arsenic forms alloys with lead

Checking Up 10.4

1. What is another name given to yellow phosphorus?

2. Why is white phosphorus stored in water?

3. Write an equation to show how phosphoric acid reacts with calcium hydroxide

4. The chloride and oxide of phosphorus in the higher oxidation state are: PCl5 and P2O5.

a) Give the formulae of the chloride and oxide of phosphorus in a lower oxidation state.

b) Write a balanced equation for the reaction of PCl5 with water.

c) Write a balanced equation for the reaction of P2O5 with water.

d) 25cm3 of the resulting solution of the reaction between P2O5 and water reacted completely with 25cm3 of 0.6moldm-3 NaOH. Calculate the con-centration of the solution.

e) What is the name of the shape of PCl5? Give one of the bond angles in that shape.

10.1.1 Environmental problems of using chemical fertilizers of nitrates and phosphates

Nitric acid is mainly used in the manufacture of nitrates fertilizers. Excess use of nitrates as fertilizers is responsible of one type of pollutions of lakes and rivers called eutrophication.

Eutrophication results from the excessive richness of nutrients in a lake or a water body which causes a dense growth of plant life. When those water plants die and are decomposed, during the decomposition process that uses oxygen, they deplete the oxygen of the water body and render that water incapable of sustaining living aquatic organisms. In that case, the body of water is said to be dead (biologically).

The fraction of the nitrogen-based fertilizers which is not converted to be used by plants accumulates in the soil or lost as run-off. High application rates of nitrogen-containing fertilizers combined with the high water solubility of nitrate leads to increased runoff into surface water as well as leaching into groundwater, thereby causing groundwater pollution.

Ammonium nitrate fertilizers can cause soil acidification:

This may lead to decreases in nutrient availability which may be offset by liming. High levels of fertilizer may cause the breakdown of the symbiotic relationships between plant roots and mycorrhizal fungi.

Conclusion: Use of chemical fertilizers such as nitrates is good because it allows increasing of soil fertility. But we may be careful to avoid that excess of those fertilizers becomes a pollution problems for our soil, lakes and water bodies.

10.5 End unit assessment

Multiple choice questions

1.When heated NH3 is passed over CuO. The gas evolved may be one from the list below. Identify which one.a) N2    b) N2O     c) HNO3     d) NO2

2. When concentrated nitric acid is heated it decomposes to give one from the list below. Identify which one.   a) O2 and N2    b) NO     c) N2O5      d) NO2 and O2

3. Which of the following metal produces nitrous oxide with dilute HNO3

a) Fe     b) Zn     c) Cu      d) Ag

4. Which Nitrogen trihalides is least basic and why?

a) NF3      b) NCl3       c) NBr3       d) NI3

5. P4O6 reacts with water to give one of the following. Identify which one.

a) H3 PO3    b) H4 P2O7       c) HPO3     d) H3 PO4

6. Which one does not form complex as central atom in the list below; justify your answer:

a) N       b) P           c) As         d) Bi

7. Nitrogen dioxide is released by heating one of the following: identify which one.

a) Pb(NO3)2        b) KNO3      c) NaNO2       d) NaNO

Short and long answer Open questions

1. Though nitrogen exhibits +5 oxidation state, it does not form pentahalide. Give a reason.

2. PH3 has lower boiling point than NH3. Why?

3. a. Write the equation for the reaction to prepare ammonia gas from ammonium chloride and calcium hydroxide

b. i) Which other ammonia salt can be used in place of ammonium chloride?

ii) How can you show that ammonia is an alkaline gas?

iii) State three large scale uses of ammonia.

4. a. Describe how nitric acid reacts with copper (your answer should include equations for the reactions and clear observations).

b. Give two reactions that show how nitric acid can behave as an acid and as an oxidizing agent

c. State two uses of HNO3.

5. a. When phosphorous (P) burns in air, it reacts with oxygen to produce an oxide. Write a chemical equation for the reaction of phosphorous with oxygen.

b. Phosphorous pentachloride (PCl5) is one of the chlorides of phosphorous. Deduce the oxidation state of phosphorous atom in phosphorous pentachloride (PCl5) molecule.

c. Give two uses of phosphorous compounds.

Key unit competence:

The learner should be able to compare and contrast the chemical properties of the Group 16 elements and their compounds in relation to their position in the Periodic Table.

Learning objectives

By the end of this unit, the learner will be able to:

•Describe the physical properties of Group 16 elements.

•Describe the reactions between sulphur and oxygen.

•Describe the steps and conditions applied in the industrial preparations of sulphuric acid.

•Describe the chemical properties of sulphuric acid.

•Describe the properties of oxoanions.

•State uses of the Group 16 elements and compounds.

11.1. Physical properties of group 16 elements

Activity 11.1

Study carefully the following figure and answer the questions that follow

1. What do you observe on the above figure?

2. Using your own observation and knowledge, predict the physical state of each of the elements in the figure.

3. Which one of the following best defines the word “allotropes”?

a. Different structural forms of an element

b. A pair of substances that differ by H+c. Elements that possess properties intermediate between those of metals and nonmetals

d. Atoms of a given atomic number that have a specific number of neutrons

e. The different phases (solid, liquid or gas) of a substance

11. 1.1. Physical state and metallic character of group 16 elements

This group consists of oxygen (8O), sulphur (16S), selenium (34Se), tellurium (52Te), and polonium (84Po). It is also known as group of chalcogens. The name is derived from the Greek word for brass and points to the association of sulphur and its congeners with copper. Most copper minerals contain either oxygen or sulphur and frequently the other members of the group. Group 16 elements have the valence-shell electron configuration ns2np4.There is a steady change down the group from non-metallic to metallic properties.

Oxygen is a colourless gas, a non-metal. Sulphur is a white and waxy non-metal solid. Selenium has acidic oxides only. It is classified as a semi metal or metalloid. Tellurium has amphoteric oxides, so it is a semi-metal or metalloid element. Polonium is a silvery metallic solid.

The melting and boiling points increase with the increase in atomic number down the group. The large difference between the melting and boiling points of oxygen and sulphur may be explained using their atomicity. Oxygen exists as diatomic molecule (O2) whereas sulphur exists as polyatomic molecule (S8).

Table 11.1: Physical properties of group 16 elements

11.1.2. Allotropes of oxygen and sulphur

a) Oxygen

Oxygen is a gas at room temperature and is colorless, odorless, and tasteless. It is the most abundant element by mass in both the Earth’s crust and the human body. It is second to nitrogen as the most abundant element in the atmosphere (21%).

Oxygen has two allotropes, dioxygen (O2) and trioxygen (O3) (see in Activity 11.1). In general, O2 (or dioxygen) is the form referred to when talking about the elemental or molecular form because it is the most common form of the element. The O2 bond is very strong, and oxygen can also form strong bonds with other elements. However, compounds that contain oxygen are considered to be more thermodynamically stable than O2 itself.

Dioxygen is a colourless and odourless gas. It liquefies at 90 K (-183oC) and freezes at 55 K (-218oC). Oxygen atom has three stable isotopes: 16O, 17O and 18O. Dioxygen directly reacts with nearly all metals and non-metals except some metals (e.g., Au, Pt) and some noble gases. Its combination with other elements is often strongly exothermic which helps in sustaining the reaction. However, to initiate the reaction, some external heating is required as bond dissociation energy of oxygen-oxygen double bond is high (493.4 kJmol–1).

The other allotrope, ozone, O3, is a pale-blue poisonous gas with a strong odor. It is a very good oxidizing agent, stronger than dioxygen, and can be used as a substitute for chlorine in purifying drinking water without giving the water an odd taste. However, because of its unstable nature it disappears and leaves the water unprotected from bacteria. A layer of Ozone at very high altitude in the atmosphere is responsible for protecting the Earth’s surface from ultraviolet radiations; however, at lower altitudes it is poisoneous and becomes a major component of smog.

Destruction of the ozone layer

Human activities are rejecting in the atmosphere substances that have the capacity of destroying the ozone layer.The most destructrive substances are the chlorofluorocarbons (CFCs). That is why CFCs have been banned. We can contribute to the protection of the ozone layer by avoiding to use ozone destroying substances in our daily activities.

a. Sulphur

Sulphur occurs in native state, in volcanic lands and some sedimentary lands. Sulphur occurs as free element underground in some countries such as USA and Japan. Sulphur also occurs in many important metallic sulphides (0.05% of earth’s crust), such as lead sulfide, or galena, PbS; zinc blende, ZnS; copper pyrite, (CuFe)S2; cinnabar, HgS; stibnite, Sb2S3; and iron pyrite FeS2. It is also combined with other elements in the form of sulfates such as barite, BaSO4; celestite, SrSO4; and gypsum, CaSO4.2H2O, and it is present in the molecules of many organic substances such as

mustard, eggs, hair, proteins, oil of garlic, in petrol and coal.

Sulfur exists in a variety of forms called allotropes, which consist of several solid varieties, of which the most familiar are rhombic sulfur and monoclinic sulfur.

Rhombic sulphur (α-sulphur) (see figure in Activity 11.1)

Rhombic sulphur consists of yellow, translucent, octahedral crystals. It is stable below 96oC and it slowly changes to monoclinic at temperatures above 96oC. The temperature of 96oC is the transition temperature. It melts at 113oC. It has a density of 2.06 gcm-3

Monoclinic Sulphur (β-sulphur) (see figure in Activity 11.1)

Monoclinic Sulphur is stable above 96oC up to its melting point of 119oC. It consists of needle-shaped amber transparent crystals with density of 1.98gcm-3. Below 96oC, monoclinic sulphur changes to rhombic sulphur. This type of allotropy, in which one form changes to the other depending on the temperature, is called enantiotropy.

The most common naturally occurring allotrope, S8, cyclo-octasulphur, has a zigzag arrangement of the atoms around the ring, also called crown.

All forms of sulfur are insoluble in water, but the crystalline forms are soluble in carbon disulfide.

The most stable variety of the element is rhombic sulfur, a yellow, crystalline solid with a density of 2.06 g/cm3 at 20°C. Rhombic sulfur is slightly soluble in alcohol and ether, moderately soluble in oils and extremely soluble in carbon disulfide. When kept at temperatures above 94.5°C but below 120°C the rhombic form changes into monoclinic sulfur.

When ordinary sulfur melts, it forms a straw-colored liquid that turns darker with additional heating and then finally boils. When molten sulfur is slowly cooled, its physical properties change in accordance with the temperature, pressure, and method of crust formation.

Checking Up 11.1

1. True or false

a. Sulphur is a non-metal element

b. Rhombic sulphur is stable above 96oC

2. List the important sources of sulphur

3. Elements of Group 16 generally show lower value of first ionisation energy compared to the corresponding periods of group 15. Why?

11.2. Comparison of acidity and volatility of group 16 hydrides

Activity 11.2

1. In pairs, carry out research and write a note on the following terms:

a. Hydrides

b. The strength of an acid

d. A weak acid

e. A strong acid

2. With an example, explain what is meant by the term “hydrogen bond” and show how it is formed.

All group 16 elements form hydrides of the type H2E (E = O, S, Se, Te, Po). They are bent in shape. Group 16 hydrides, contray to group 1 & 2 hydrides, are covalent. Their acidic character increases from H2O to H2Te. The increase in acidic character down the group can be explained in term of increase of ionic character of the metallic hydrides, allowing easy dissociation in water to liberate H+ ion. The Acidic nature is in the order, H2O < H2S < H2Se < H2Te.

Owing to the decrease in energy of dissociation of H-E bond down the group, the thermal stability of hydrides also decreases from H2O to H2Po. Thermal stability is in the order: H2O > H2S >H2Se > H2Te > H2Po.

The main hydrides of group 16 are hydrogen oxide (H2O) known as water and hydrogen sulphide (H2S), the rotten egg smelling and poisonous gas.

Hydrogen oxide (H2O) is a liquid at room temperature because of strong intermolecular forces (hydrogen bond). Water is neutral (neither acid nor base).

Hydrogen sulphide(H2S) is a colorless gas with a repulsive smell similar to that of a rotten egg. It is also fairly soluble in water. It is one of the weakest acids known, this gas is extremely poisonous. H2O and H2S are both polar molecules therefore two molecules of H2O near each other, and two molecules of H2S near each other have dipole-dipole force between them. Yet the dipole-dipole force between two molecules of H2O is a special type of dipole-dipole force called Hydrogen bonding. And since Hydrogen bond is a stronger dipole-dipole force than that of a normal dipole-dipole bond the boiling point of water will be abnormally greater than that of hydrogen sulphide. That is why, water, with a molar mass of 18, is liquid at room temperature, whereas, hydrogen sulfide, with a higher molar mass of 34, is a gas at room temperature.

Checking Up 11.2

1. Give and explain the order of acidic nature of hydrides of group 16 elements.

2. H2S is less acidic than H2Te. Why?

3. Given two hydrides H2S and H2O. Which one is more thermally stable? Justify your answer.

4. H2O and H2S are hydrides of two elements belonging to the same group; yes H2O is liquid at room temperature, whereas H2S is gaseous. Explain why?

11.3 Preparation and properties of sulphuric acid

Activity 11.3

1. Do you know sulphuric acid?

2. If yes, what do you know about it?

3. How will you proceed if you are mixing sulphuricaci with water?

Experiment: Investigating the dehydrating power of Sulphuric acid on sugar

Materials and chemicals:

1. 70 grams granulated sugar, C12H22O11

2. 70 ml of concentrated (18M) sulfuric acid, H2SO4

3. 300 ml tall-form beaker4. 40cm stirring rod5. Paper towels

6. Disposable gloves or

7. 100 ml graduated cylinder

8. Tray9. 1 liter beaker

10. Sodium bicarbonate, NaHCO3

11. Spatula

Safety: Sulfuric acid is a very strong acid and is extremely corrosive to skin. Wear gloves and safety goggles. During the reaction, steam is generated. It is hot. It is recommended to work in a fume cupboard.

Procedure:

Spread some paper towels on the tray.1. Put sugar into 300 ml beaker. 2. Insert stirring rod into center of sugar.3. Put beaker on paper towels on the tray.4. Add 70 ml of sulfuric acid to the sugar and stir briefly.5. Stand about 1 - 2 meters away and wait for reaction to begin and observe what will happen.

Clean Up: You might want to incorporate part of the clean up procedure into the demonstration.

Remove black carbon column from the beaker and put it into a liter beaker with some sodium bicarbonate (hydrogen carbonate). With spatula, break the column of carbon into smaller pieces. Add a little water and set back on the tray. The foaming action is also exciting.

Neutralize any acid spills with sodium bicarbonate and wipe clean. Leave lecture hall clean for the next class.

Rinse all glassware and carbon chunks with lots of water. Carbon can be thrown away in trash.

Study questions

1. Record your observations

2. Write equation that take place in this experiment

11.3.1.Industrial preparation of sulphuric acid and its environmental impact

a. Preparation

Sulphuric acid is manufactured industrially using the contact process which is based on the catalytic oxidation of SO2 to SO3. In this process, there are five steps:

Step 1: Production of sulphur dioxide

Sulphur dioxide is obtained by either burning elementary sulphur or roasting metal sulphides in air in combustion chamber.

Step 2: Purification of gases

The SO2 produced in Step1 contains many impurities (e.g arsenic (III) oxide) and these impurities would poison the catalyst (platinum, Pt) by rendering it ineffective.The mixture is passed through a purifier chamber.

Step 3: Oxidation of SO2 to SO3

The purified gases (mixture of SO2 and O2) are dried and passed through converters containing a catalyst, vanadium (V) oxide V2O5 or fine divided platinum and at temperature about 

The reaction is exothermic and once the reaction has started no external heat is re-quired. The evolved heat maintains the temperature of the reaction

Step 4: Absorption of SO3

The SO3 is an anhydride of sulfuric acid, direct absorption in water gives a very vigorous reaction and forms droplets of sulphuric acid. To overcome this problem, SO3 is absorbed in the sulfuric acid in which it dissolves with ease in the absorption tower to form oleum or pyrosulphuric acid

Step 5: Dilution of oleum

The oleum formed is diluted with the correct quality of water to give concentrated sulphuric acid 98%.

b. Environmental impact of the industrial manufacture of sulphuric acid

Industrial emissions of sulfur oxides result into the increase of the concentration of sulfuric acid (H2SO4) and sulfurous acid (H2SO3) in the atmosphere. Those acids are present as particles or droplets which dissolve in clouds, fog, rain, dew, or snow, resultingin very dilute acid solutions. This constitutes a kind of pollution called acid rain. Acid rains cause acidification of soils and lakes, and destruction of some buildings, infrastructures and statues made in marble or other carbonaceous rocks.

11.3.2. Oxidising and dehydrating properties of sulphuric acid

a. Oxidizing properties

Hot concentrated Sulphuric acid is an oxidizing agent:

In other conditions, sulphuric acid is not considered as an oxidizing acid. When dilute, sulphuric reacts with other metals to form metal sulphates, and hydrogen gas, but it does’t attack copper.

b. Dehydrating properties

Concentrated sulphuric acid has a high affinity for water. It shows this property when it reacts with sugars, hydrated copper (II) sulphate, ethanol, oxalic acid and other organic matter. For example:

Because of this, concentrated sulphuric acid is used as a drying agent in laboratories.It is also this property that makes sulphuric acid destroy or burn any object containing cellulosic matter such as paper, clothes, skin, etc...). Due to this avidity for water, it is recommended, when mixing sulphuric acid and water, to pour slowly the acid in water, not the opposite.

Checking Up 11.3

1. a) Describe the Haber or Contact process for the manufacture of sulphuric acid.

b) Why is sulphur trioxide formed in this process not absorbed directly in water?

2. Concentrated sulphuric acid acts as a dehydrating agent. What does it mean?

3. Write equations to show how concentrated sulphuric acid reacts with:

a. Zinc

b. Magnesiumc. Carbon

11.4 Properties of oxoanions of sulphur

Activity 11.4 (a)

1. Use the library and/or internet to explain the following:

a. Oxidation

i) In terms of oxidation state

ii) In terms of electron transfer

b. Reduction

i) In terms of oxidation state

ii) In terms of electron transfer

c. Oxidizing agent

d. Reducing agent

Activity 11.4 (b)

2. An experiment for Heating hydrated copper(II) sulfate

Objectives:

Students remove the water of crystallisation from hydrated copper (II) sulfate by heating. Condensing in a test-tube collects the water. The white anhydrous copper (II) sulfate can then be rehydrated, the blue colour returns.

Apparatus and equipment (per group)

1. Two test-tubes

Activity questions

Record any observations made during the process and when the water was poured back onto the white copper (II) sulfate.

11.4.1. Reducing action of S2O32-

The thiosulfate anion reacts with iodine in aqueous solution, reducing it to iodide as it is itself oxidized to tetrathionate:

Thiosulfate ion reduces Iron (III) to iron (II) according to the following equation:

Thiosulfate ion is also used in analytical chemistry. It can, when heated with a sample containing aluminium cations, produce a mixture of whitish precipitate:

11.4.2. Reaction of SO32- with acids

The sulphites of Na+, K+, and NH4+ are soluble in water. Most of other sulphites are insoluble in water. However, due to the basic nature of SO32-, all sulphites dissolve in acidic solution.

In general, sulphite ion reacts with acids liberating sulphur dioxide according to the

ionic equation:

Examples:

11.4.3.Action of heat on sulphates (SO42-)The sulphates are normal salts of sulphuric acid and contain the sulphate ion, SO42-. Soluble sulphates are prepared by action of dilute acid on the metal, metal oxide, metal hydroxide,metal carbonate or hydrogencarbonate:

11.4.3.Action of heat on sulphates (SO42-)

The sulphates are normal salts of sulphuric acid and contain the sulphate ion, SO42-. Soluble sulphates are prepared by action of dilute acid on the metal, metal oxide, metal hydroxide,metal carbonate or hydrogencarbonate:

Insoluble sulphates (BaSO4 and PbSO4) are obtained by precipitation method:

Most sulphate salts do not decompose on heating. For instance, sodium sulphate, potassium sulphate, and calcium sulphate are not decomposedby heat.

Only certain sulphate salts are decomposed by heat when heated strongly. On heating, some sulphates decompose to give either sulphur trioxide or sulphur dioxide or both.

For instance, strong heating of green crystal iron (II) sulphate will release steam, sulphur dioxide, sulphur trioxide and leave behind a reddish solid iron (III) oxide residue. The steam released comes from the hydrated water of the crystallize salt:

Zinc sulphate, copper (II) sulphate, and iron (III) sulphate decompose when heated strongly to produce sulphur trioxide gas and a metal oxide.

Examples:

When ammonium sulphate is heated strongly, this white solid sublimate and is decomposed into ammonia gas and sulphuric acid vapour.

On heating, hydrated crystals of metal sulphates lose their water of crystallization and turn into a powder. They are then said to be anhydrous. Sometimes they may also lose their colour. Hydrated copper (II) sulphate is heated according to the following equation:

Checking Up 11.4

1. Write equations to show how thiosulfate ions reduce the following substances:

a. Iodine

b. Iron (III) ion

c. Aluminium ion

2. Write equations to show the action of heat on the following sulphates:

a. Zinc (II) sulphate

b. Iron (III) sulphate

c. Copper (II) sulphate

3. When hydrated copper II sulphate solid is heated in a boiling tube, a white solid Q and droplets of a colourless liquid P are observed.

a. Identify substances; liquid P and solid Q.

b. Explain the observation above.

Explain what would be observed if water is added to white solid Q.

11.5 Identification of sulphite and sulphate ions

Activity 11.5

Given a substance Y which contains one cation and one anion, identify the cation and the anion in Y. Carry out the following tests on Y and record your observations and deductions in the table below.

11.5.1. Identification of sulphate ion (SO42-)

a. Addition of barium nitrate solution

When an aqueous solution of Ba2+ ion is added to a solution containing SO42- ion, a white precipitate of barium sulphate insoluble in dilute nitric acid or hydrochloric acid is formed.

Use of barium chloride solution also gives a white precipitate insoluble in dilute hydrochloric acid. This reaction serves as the test for the sulphate ion in solution and also differentiates sulphate ions, SO42-, from sulphite ions, SO32-.

b. Addition of lead (II) nitrate solution

When an aqueous solution of Pb2+is added to a solution containing SO42-, a white precipitate of lead (II) sulphate is formed:

c. Addition of silver nitrate solution

When an aqueous solution of Ag+ is added to a solution containing SO42-, a white precipitate of silver sulphate:

11.5.2. Identification of sulphite ion (SO32-)

To an aqueous solution of any sulphite, dilute nitric acid is added followed by a few drops of barium nitrates or barium chloride solution. A white precipitate of barium sulphite forms which dissolves in the nitric acid or hydrochloric acid. The sulphite dissolves in acid with evolution of sulphur dioxide (seen as bubbles of a colourless gas).

Hence, sulphites give a white precipitate of barium sulphite soluble in acid while sulphates give a white precipitate insoluble in acid.

Checking Up 11.5

1. Describe a chemical test you would use to distinguish between the pairs of compounds below. In each case give the reagent, conditions if any and the ex-pected observation.

a. BaSO3 and BaSO4

b. BaSO4 and PbSO4

2. Carry out the following experiment to identify the cation and the anion in substance X provided by the teacher.

11.6 Uses of group 16 elements and compounds

Activity 11.6

1.What are the most known applications of Sulphur and Oxygen?

2.The following figures show how some elements play a big role in our daily lives. Observe carefully and answer the following questions:

a. What can you say about figure A, B, C and D?

b. Research in library or internet to find out the substance used in each of the figure.

c. Research in library or internet to find different uses of sulfuric acid.

11.6.1.Uses of oxygen

The first use of oxygen is in breathing and metabolism processes of all living organisms.

There are many other commercial uses for oxygen gas, which is typically obtained through fractional distillation of air. It is used in all operations involving combustion as the active component of air.

It is used in the manufacture of iron, steel, and other chemicals. Oxygen is also used as an oxidizer in rocket fuel, and for medicinal purposes. Mixture of oxygen and ethyne (oxyacetylene) is used for welding and metal cutting.

11.6.2. Uses of sulphur

The main use of Sulphur is the manufacture of sulphuric acid.Sulphur is also used in vulcanization of rubber, a chemical process for converting natural rubber or related polymers into more durable and pressure resisting materials by heating them with sulfur or other equivalent curatives or accelerators. These additives modify the polymer by forming cross-links (bridges) between individual polymer chains, making the final product very hard and resistant to pressure and other conditions.Sulphur is an ingredient in the manufacture of dyes, fireworks and other sulphur compounds.

11.6.3. Uses of sulphuric acid

Sulphuric acid is a very important industrial chemical. It used to be called the giant of chemical industry. It is used in the manufacture of hundreds of other compounds in many industrial processes.

•The bulk of sulphuric acid produced is used in the manufacture of fertilisers (e.g., ammonium sulphate, superphosphate).

•Sulfuric acid is also used in many other applications such as in: metallurgical industry, storage batteries, chemistry laboratories, etc....

11.6.4. Uses of hydrogen sulphide (H2S)

Some nuclear power plants use hydrogen sulphide to produce heavy water. Farmers use H2S as an agricultural disinfectant. In Analytical Chemistry, H2S is used in selective precipitation of metal sulfides

Checking Up 11.6

1. Give at least 3 uses of oxygen

2. Mention three areas in which H2SO4 plays an important role

3. Briefly explain what is meant by “vulcanization”

4. What property makes sulphuric acid a dangerous corrosive product?

5. How do you proceed when you mix sulphuric acid with water

11.7. End unit assessment

Multiple choice questions

1. Which one(s) of the following substancs is not easily soluble in water and why?

(A) H2        (B)O2       (C) SO2         (D) CO2

2. Select the molecular formula of sulphur in the following list at room temperature:

(A) S      (B) S2            (C)S4        (D) S8

3. Choose the correct answer. All the elements of Group 16 are:

(A) Non - metal         (B) Matalloids       (C) Radioactive        (D) All have allotropic forms

4. The triatomic species of element oxygen is known as:

(A) Azone      (B) Polyzone       (C) Trizone      (D) Ozone

5. Which of the following acts as pickling agent:

(A) HNO3     (B) H2SO4     (C) HNO2

6. Which of the following is not suitable for use in desiccators to dry substance and why?

(A) ConcH2SO4      (B) Na2SO4      (C) CaCl2     (D) P4O10

7. Which of the following methods are you going to use to test presence of water in a liquid:

(A)Taste      (B) Smell      (C) Use of litmus paper       (D) Use of anhydrous CuSO4

8. When sulphur is boiled with Na2SO3 solution, the compound formed is:

(A) Sodium Sulphide    (B) Sodium Sulphate    (C) Sodium persulphate     (D) Sodium thiosulphate

Short and long answer open questions

9. State two ways in which oxygen differs from other group 16 elements.

10. State two sources and two uses of oxygen.

a. Name two crystalline allotropes of sulphur.

b. Write equation to show how sulphur reacts with oxygen, iron and sulphu-ric acid.

c. Write two equations to illustrate the reducing property of hydrogen sul-phide

12. Sulphur is used to make sulphuric acid. The reactions in the manufacture of sulphuric acid by the Contact Process are shown below.

a. Give a large scale source of the element sulphur

b. State another use of sulphur dioxide

c. How is sulphur changed into sulphur dioxide?

d. Write a complete chemical equation for reaction 3.

e. Write a chemical equation for reaction 4.

Key unit Competence:

To be able to: Compare and contrast the chemical properties of the Group 17 elements and their compounds in relation to their position in the Periodic Table.

Learning objectives

By the end of this unit, the learner will be able to:

•Prepare and test halogens.

•Perform experiments to prepare and test chlorine, bromine and iodine.

•Relate the oxidizing power of Group17 elements to their reactivity.

•Relate the acidity strength of oxoacids to the number of oxygen atoms combined with the halogen.

•Compare the reactions of the halogens with cold dilute sodium hydroxide and hot concentrated sodium hydroxide solutions.

•State the uses and hazards of halogens and their compounds.

•Test for the presence of halides ions in aqueous solutions

•State the natural occurrence of halogens•Describe the extraction methods of halogens

•Explain the trends of physical and chemical properties of Group 17 elements down the group

•Describe the trends in strength acidity, volatility and reducing power of halogens hydrides

•Describe the chemical properties of chlorates, iodates, perchlorates and periodates

12.1 Natural occurrence of halogens

Group 17 elements are called halogens. They include: Fluorine (9F), chlorine (17Cl), bromine (35Br), iodine (53I) and astatine (85At). Although astatine is a member of group 17 elements, its chemistry is of little importance in this unit because all of its known isotopes are radioactive. The longest-lived isotope of Astatine (At-210) has a half-life of only 8.3 hours. All these elements have 7 electrons in their outermost shell (ns2np5) which is one electron short of the next noble gas.

Activity 12.1

1. (i).Can halogens occur freely in nature?

(ii).Give explaination to your answer in (i)

2. Which of the halogen compounds is the most abundant in nature?

All halogens occur in sea water as halide ions. Halogens also occur in minerals such as, NaCl (rock salt), CaF2, Ca(PO4)3F and NaAlF6 . Chlorides also occur in the Great Salt Lake and the Dead Sea, and in extensive salt beds that contain NaCl, KCl or MgCl2. Hydrochloric acid is present in the human stomach to help in the digestion process of food.

Bromine compounds occur in the Dead Sea and underground brines. Iodine compounds are found in small quantities in chile saltpeter, underground brines.

The best sources of halogens (with the exception of iodine) are halide salts. It is possible to oxidize the halide ions to produce the diatomic molecules by various methods, depending on the ease of oxidation of the halide ion. Fluoride is the most difficult to oxidize, whereas iodide is the easiest.

Checking up 12.1

1. State 2 locations where chlorine can be found in nature.

2. Write the chemical formulae of the compounds of halogens in nature.

3. Give the name of one lake in Rwanda where salt is abundant in water.

4. Explain the separation method you can use to get the salt crystals from the water.

5. Fluoride ion is the most difficult to oxidise into fluorine, whereas iodide ion is the easiest. Explain why.

12.2. Preparation methods of halogens

Activity 12.212.2 (a). Preparation of Cl2 and I2

1. Dissolve a small quantity of potassium iodide (2g) in a beaker containing 120 ml of water. Keep this solution.

2. Pour 20 ml of a 0.02 mole/litre potassium permanganate solution into a flat-bottomed flask.

3. Pour concentrated hydrochloric acid into potassium permanganate in the flat-bottomed flask (20 ml of a 2 mol/litre HCl).

4. Pour potassium iodide (10 ml of the potassium iodide solution) in the boiling tube.

5. Direct the delivery tube into a boiling tube containing potassium iodide solution as shown in the set-up diagram.

6. Heat the flask until there is a colour change in the boiling tube. Record the observations.

7. Add few drops of starch indicator to the resultant solution in the boiling tube

8. Record the observations.

Repeat procedures steps 1 to 7 but this time use solutions of KBr or NaBr instead of KI in the boiling tube to prepare bromine and chlorine.

Activity 12.2 (b) Preparation of bromine and iodine

1. Put 0.5 gram of MnO2 in a round bottomed flask.

2. Pour concentrated NaBr solution (5 ml of a 0.1 mol/litre) in the round bottomed flask.

3. Pour 5 ml of 1 mol/litre HCl solution in the round bottomed flask mixture.

4. Connect the apparatus to a delivery tube using a rubber stopper.

5. Heat the round bottomed flask mixture.

6. Direct the delivery tube in a solution of AgNO3 in a test tube.

7. Note the observable changes.

Activity 12.2. (c): Electrolysis of concentrated NaCl solution

a) Put 1 g of NaCl in a beaker.

b) Add water and stir using a glass rod until all the salt dissolves.

c) Pour the solution in an electrolyser.

d) Connect the electrolyser to the source of direct current and switch on.

e) Dip a test tube full of water in the NaCl solution in inverted position from above each electrode.

f ) Put 2 a little phenolphthalein indicator in the solution under each test tube.

g) Record the observations that take place for 5 minutes.

Apparatus set-up: Electrolysis of concentrated NaCl solution

12.2.1. Chlorine

Most commercial chlorine is obtained by electrolysis of chloride ions in aqueous solutions of sodium chloride or molten NaCl.

The reactions that take place are shown by the following chemical equations:

At the cathode, reduction reaction that takes place depends on the state of NaCl:If molten NaCl is used,

Sodium can be collected in metallic form or dissolved in water to produce NaOH solution:

If salt water or brine is used: hydrogen will form at the cathode, not sodium:

It is feasible to prepare chlorine, bromine and iodine in the laboratory by the chemical oxidation of the halide ion in acid solution with strong oxidizing agents such as manganese dioxide (MnO2) or sodium dichromate (Na2Cr2O7).

The reaction with manganese dioxide is represented by the equation given below:

Where X= chlorine, bromine or iodine. The halide used here may be NaX or KX

12.2.2. Fluorine

The common method for preparing fluorine is electrolysis. The most common electrolysis procedure is to use a molten mixture of potassium hydrogen fluoride, KHF2, and anhydrous hydrogen fluoride. During the process, fluoride ion, F- is oxidized to F2, whereas hydrogen ion, H+, is reduced to H2 .

Electrolysis causes HF to decompose, forming fluorine gas at the anode and hydrogen at the cathode. The two gases are separated to prevent their explosive recombination to reform hydrogen fluoride, HF.

12.2.3. Bromine

The industrial production of bromine involves the oxidation of bromide ion by chlorine:

Chlorine is a stronger oxidizing agent than bromine.

12.2.4. Iodine

Iodine can be obtained by oxidation of iodide ions by chlorine gas or another halogen that is higher than iodine in the group.

The industrial production of iodine can be done using the reduction of sodium iodate, NaIO3. The reaction is carried out using sodium hydrogen sulfite, NaHSO3, as reducing agent:

Checking up 12.2

1. Describe the preparation of chlorine in the laboratory.

2. How is electrolysis used to produce chlorine industrially?

3. Explain the method used to produce bromine.

4. Explain the hazards that may be encountered during the production of fluorine from its ores.

5. Say if the following statement is correct or wrong and justify your answer: “Bromine can be produced by oxidation of bromide by chlorine and vice-versa”.

12.3. Trends of physical and chemical properties

Activity 12.3: Testing a few properties of halogens

1. Dissolve about 0.5g of sodium chloride crystals in a test tube containing 20cm3of water. Keep this solution.

2. Pour 4 ml of the NaCl solution in 4 ml of a solution of conc H2SO4 in a test tube.3. Heat the above mixture in (NaCl and H2SO4 ) on a Bunsen burner flame.

4. Test the gas evolved with blue and red litmus papers.

5. Repeat the above procedures but using NaBr instead of NaCl.

6. Repeat the above procedures but these times use KI instead of NaCl.

7. Test the presence of any evolved gas with litmus papers.

8. Test the resultant solution of the mixture of KI and concentrated H2SO4 with a solution of starch indicator.

7. Record the observations

8. Interpret the results of the reactions in the experiment and draw a conclusion of the reactions.

Observations: The colourless solution of KI(aq) gives red or reddish mixture when chlorine gas is bubbled in it.

The red mixture turns the colourless starch indicator solution to dark–blue. On standing for a while, a black or deep violet residue settles.

Interpretations

Chlorine is liberated by the reaction between 2M hydrochloric acid and potassium permanganate solution. Chlorine displaces iodine from potassium iodide solution, which dissolves in water to give a dark-red solution, and turns starch indicator dark-blue. The greyish-black residue is due to the formation of Iodine solid.

12.3.1. Physical properties of group 17 elements down the group

a. Appearance

At room temperature and pressure, fluorine and chlorine are gases, bromine is a liquid and iodine is a solid. The table below shows the physical properties of halogens.

Table12.1: Physical properties of halogens

As shown in the table above, the boiling and melting points of halogens increase as you move down the group due to increase of the atomic mass which increases the strength of van der Waals forces.

a. Solubility in water

•Fluorine and chlorine dissolve in cold water.

•Bromine and iodine are sparingly soluble in polar solvents like water and they form coloured solutions. However they are soluble in organic solvents

like carbon tetrachloride, carbon disulphide and hydrocarbons.

•Halogen-halogen bond dissociation energy decreases from chlorine downwards as shown in the following trend:

The reason why the trend is in this order is that in a smaller molecule such

as Cl-Cl, the bond length is shorter so the internuclear attraction becomes

bigger and therefore more energy is required to break the bond Cl-Cl.

•Fluorine, F2 presents an exception because its energy of dissociation F-F is far smaller than that of Cl-Cl (opposite to the general trend from Cl2 to I2) because there is a repulsion between the lone pairs of electrons of fluorine atoms in the molecule due to the small size of fluorine.

12.3.2. Trend in chemical properties of group 17 elements down the group

Due to their valence electronic structure of ns2 np5 halogens gain easily one electron to complete the octet structure, that is why they exhibit an oxidation state (-1) in most of their compounds. They are generally considered as good oxidizing agents.

All the halogens are highly reactive; they are the most reactive group of non-metals in the periodic table. They react with metals and non-metals to form halides. But in Group 17 elements (halogens), the reactivity of halogens decreases down the group in the order: F2> Cl2>Br2 >I2. Hence each halogen displaces those below it from their salts.

The ease of capturing an electron indicates the strength of oxidising nature of halogens. F2 is the strongest oxidising halogen and it oxidises other halide ions in solution and even in the solid phase. In general, a halogen oxidises halide ions of higher atomic number. The halogens decrease in strength of oxidising power as we descend the group.

Examples:

In addition to the normal oxidation state -1, chlorine, bromine and iodine exhibit +1, +3, +5 and +7 oxidation states. The higher oxidation states is achieved when the halogen is in combination with a small and highly electronegative atom such as fluorine and oxygen. Example: Interhalogens, oxides and oxoacids. The oxidation states of+1, +5, +7 occur in the oxides and oxoacids of chlorine, bromine and iodine

Compounds where halogens have positive oxidation states are generally very reactive, because they tend to be reduced to the oxidation state -1, the most stable oxidation state of halogens.

For example, perchlorate, ClO4-, is a powerfull oxidizing agent due the the presence of Cl(VII) species that is eager to capture electrons and get the oxidation state -1.

Being the most electronegative element, fluorine exhibits only –1 oxidation state.

a)Reactions of halogens with water

Fluorine exhibits anomalous behaviour in many properties. The anomalous behaviour of fluorine is caused by its small size and very high electronegativity.

Fluorine reacts vigorously with water to produce oxygen O2 and hydrofluoric acid HF.

Chlorine dissolves slightly in water at room temperature (25oC). At 800C and above, Cl2 gas is insoluble in water. Chlorine reacts with water to produce hydrochloric acid, HCl and hypochlorous acid, HClO:

Hypochlorous acid is a bleaching agent.Bromine dissolves in water; it reacts with water slowly to form hydrogen bromide HBr and hypobromous acid, HBrO:

Solubility of iodine in water is very low. Iodine does not react with water. Table 12.2 compares the oxidizing power of halogens when a given halogen is mixed

with an aqueous solution of another halide.

Checking up 12.3

1 State different hydrogen halides that can be formed.

2. What is the difference between hydrogen halides and hydrohalic acids?

3. Describe the anomalous behaviour of fluorine in the group (explain at least 3 differences from the rest of the halogens).

4. Write the reaction equation that takes place between:

a) Solid chloride with concentrated sulphuric acid. b)Bromine in hot concentrated sodium hydroxide solution.

5. Briefly explain the trend in volatility of hydrogen halides as you move down the group.

6. Explain why the hydrogen halides acidity increases in the order: HF < HCl < HBr <HI

7. Explain the trend in solubility of halogens in water as you move down the group.

8. You have two test tubes; one contains water, the other contains a solution of chlorine in water, and you are asked to identify them. What test are you going to do in order to identify the two test tubes?

Table 12.2: Comparison of oxidising power of halogens

b) Reaction of halogens with hydroxide:

• Cl2 reacts with excess cold dilute OH- and the reaction equation is:

With hot concentrated OH-, chlorate ClO3- , is formed:

The same reactions occur for Br2 and I2.

• F2, unlike other halogens, does not form oxosalts with alkalis and it does not form oxoacids as shown in the following reactions:

c) Reaction with metals:

Halogens react with metals to form salts, metal halides. Example: bromine reacts with magnesium to form magnesium bromide

The preparation of certain anhydrous metal halides is carried out using this method of reacting a halogen and a metal. Examples: MgCl2 (white solid), AlCl3 (white solid), FeCl3 (brown solid), FeBr3 (dark red solid) etc... are halides that can be prepared in this way.

Generally metal halides are ionic, the most ionic being the ones where the metal has low oxidation states (+1, +2).

For the same metal in the same oxidation state, the ionic character of the metal halides decreases down the group:

MF >MCl > MBr > MI; whereby M is a monovalent metal.

If a metal exhibits more than one oxidation state, the halide in a higher oxidation state will be more covalent than the one with a lower oxidation state.

Example: SnCl4 and PbCl4 are more covalent than SnCl2 and PbCl2 respectively.

d) Reaction with non-metals

Halogens react with hydrogen gas to form hydrogen halides but the affinity for hydrogen decreases from fluorine to iodine.Hydrogen halides dissolve in water to form hydrohalic acids. The acidic strength of these acids varies in the order:

HF < HCl < HBr < HI

The stability of these halides decreases down the group due to the decrease in (H–X) bond strength order:

Halogens react with oxygen to form many oxides such as, Cl2O, Cl2O6, etc.. In those oxides, except for fluorine, halogens have positive oxidation states due to the high electronegativity of oxygen; this makes them unstable and very reactive. Fluorine forms compounds with oxygen where it has a negative charge; they are rather oxygen fluoride; e.g. OF2With phosphorus, all halogens react with phosphorus to form phosphorus (III) halides of the form PX3. The reaction equation of phosphorus with bromine is represented in a general equation as:

In excess chlorine or bromine, phosphorus reacts to form phosphorus (V) chloride or bromide:

The reaction between phosphorus (III) chloride and chlorine to form phosphorus (V) chloride is reversible:

Halogens combine with other halogens to form various compounds known asinterhalogensof the types XY, XY3, XY5 and XY7, whereby X is a larger size halogen and Y is smaller size halogen.These compounds are easiest to form when Y is fluorine.

d) Bleaching property

Bleaching is the process of removing stains or colours in fabrics, especially by the use of chemical agents such as halogens, chlorine and bromine. The bleaching action of chlorine is an oxidizing action of hypochlorous acid, HOCl:

The hypochlorous acid oxidizes the dye material to form a colourless substance:

Chlorine bleaches ordinary inks which are solutions of dyes.

Bromine has a similar action but with a much less vigorous action.

Iodine has no reaction with water and therefore it is not a bleaching agent.

12.4. Preparation of Hydrogen halides

Activity 12.4.(a)

Laboratory preparation of Chlorine

Reactants:Sodium Chloride and sulfuric acid

procedure:

1.Put 50g of NaCl in round bottomed flask

2.Pour conc sulfuric acid throuth the filter fannel and heat

3.The liberated gas is pased throuth the concentrated sulfuric acid

4.Collect the gas by the Downward delivery

Note:The lower end of the thistle funnel must be dippen in acid,or you can use the funnel with Syphon.

Halogens form hydrogen halides (HF(g), HCl(g), HBr(g) and HI(g), these hydrogen halides form a series of acids when they are dissolved in water. When in aqueous solution, the hydrogen halides are known as hydrohalic acids. Hydrohalic acids are strong acids, except HF(aq) that is a weak acid.

Hydrogen halides are formed by reacting a solid halide (such as NaBr) with concentrated phosphoric (V) acid (H3PO4):

The hydrogen halides fume on contact with air and with moisture, they react to form acids:

HF and HCl can be prepared by using concentrated sulphuric acid, H2SO4, a less expensive method:

Similar equation for NaF.

Concentrated sulphuric acid is an oxidising agent and can’t be used to prepare HBr or HI. It is observed that as HBr or HI are formed, they are oxidised by sulphuric acid:

HBr and HI are prepared by hydrolysis of phosphorus trihalides.

Checking up 12.4

1. Explain why fluorine is the most oxidising element in group 17 and of all chem-ical elements?

2. Explain why the acidity of group 17 hydrohalides increases down the group.

3. Explain why HF does not follow the general trend of volatility but instead it has a higher boiling point than HCl.

4. Describe the natural state and appearance of halogen elements at room tem-perature and pressure.

5.You have two unknown sample solutions A and B. You are told that one is a solution of NaCl(aq), the other is a solution of NaI(aq). You are asked to identify them. When you add a solution of Fe3+(aq) which has a yellow colour, to both sample solutions, you get the following results: (a) Adding Fe3+(aq) to A gives a yellow solution; (b) Adding Fe3+(aq) to B gives a solution with a complicated mixture of colours between green and violet. Question: Which solution is in A, which solution is in B? Justify.

12.5. Trends in strength of acidity, volatility and reducing power of hydrogen halides

12.5.1. Acid strength

The acid strength is a measure of how an acid dissociates in water into its ions. Strong acids dissociate completely into their ions, whereas a weak acid dissociates partially into its ions.

Examples:

•Hydrochloric acid is a strong acid:

•Hydrofluoric acid is a weak acid:

The acid strength of hydrogen halides (HX) increases down the group from HF to HI:

Because of the very electronegativity and small size of F, it forms very strong H-F bond. In water, it is slightly dissociated to give few H+ ions in solution; therefore it is a weak acid.

This is due to the strong bond H-F; in other words, the bond strength HX decreases down the group.

12.5.2. Volatility

Volatility is the state of having a low boiling point to evaporate easily. The boiling points of hydrogen halides generally increase down the group. This is because the van der Waals’ forces increases with size (the forces increase with increase of the surface of contact between molecules).

The high electronegativity of fluorine atom is the root cause of the strong hydrogen bonds. It is liquid at room temperature while other hydrogen halides are gases.

The trend in volatility:

HF is an exception to the general trend (it has a high boiling point of 19.9 0C). It is almost liquid at room temperature while the rest are gases because there are strong hydrogen bonds between HF molecules. For the other Hydrogen halides, the boiling point increases with increasing molecular mass.

12.5.3. Reducing power

The reducing power of hydrogen halides is dependent on the ability of their anions (halide ions) to liberate electrons.

It becomes easier to oxidise the hydrogen halides as we go down the group because the electron on the halide ion becomes less attracted to the nucleus as we move down the group.

Hence the order of reducing power is: I-> Br- > Cl-> F-

Given that halogens are the best oxidizing agents since they tend to capture electrons from other elements to achieve the octet electronic configuration, the halide ions are not expected to be good reducing agents.

However, an iodide ion, due to its big size, can easily liberate an electron and therefore act as a mild reducing agent.

Hence HI is the strongest reducing agent of all hydrogen halides.

HI can reduce Cl2, Br2, O2, Fe3+ salts, etc... to produce molecular iodine, I2:

That is why when a colorless aqueous solution of HI is exposed to air, it turns brown because of the presence of I2.

Concentrated H2SO4 oxidises Br- and I- to form Br2 and I2 elements respectively.

Concentrated H2SO4 cannot oxidise F- and Cl- ions.

12.5.4. Tests for halide ions in aqueous solution

Test of substance X with an unknown anion

Activity 12.5

Identification of ions:

i) You are provided with a solution of X substance.

ii) Put 1 ml of X solution in each of the 4 test tubes.

iii) Add in each test tube the reagent solutions as indicated in the table be-low.

iv) Note down the observations for interpretation later in each test.

Halide ions are detected by precipitation reactions using silver nitrate solution AgNO3 or lead (II) nitrate Pb(NO3)2.The ionic equations below illustrate the general equation reactions:

i) Testing using AgNO3 and ammonia solution, NH3

The first step is to dissolve 1 ml sample solution to be tested in about 2 ml of dilute nitric acid.The table given below shows the observable changes when we add the silver nitrate solution, AgNO3(aq), followed by ammonia solution, NH3(aq)

The addition of NH3 ammonia solution is called ‘confirmatory test”, i.e. to confirm the first one.

Table 12.5: Tests of halides

If the test uses Pb2+(aq) and Cl- ions, heating the precipitate, it will dissolve; PbCl2(s) is insoluble in cold water, but soluble in hot water.

g) Lead nitrate Pb(NO3)2 solution

Chloride, Cl- ion: a white precipitate solution is formed in cold water, the precipitate dissolves when it is hot water

Checking up 12.5

1. You are given 1ml of solution Y in which you add 1 ml of AgNO3 solution and there is formation of a white precipitate that dissolves on adding 1 ml of ammonia solution. Deduce the anion present in solution Y.

2. Write a balanced equation for the reaction of: a) Decomposition of KClO4b) Oxidation of Cl- by ClO3-in acid medium.

3. Explain how ClO3- is prepared in the laboratory from chlorine.

4. Explain why fluorine forms only 1 oxoacid (HFO).

12.6. Chemical properties of chlorates, iodates, perchlorates and periodates

Activity 12.6 (a)

1. Put 0.5g of solid KClO3 in a pyrex test tube.

2. Using a test tube holder, heat the KClO3 strongly while placing a glowing splint on the test tube outlet.

3. Note the observation as heating takes place.

4. Dissolve a small portion of the solid residue in water.

5. Put a little solid residue in 1 ml of aqueous solution of AgNO3.

6. Record the observed changes.

7. Deduce the chemical changes that took place. a)While heating b) On adding the residue to AgNO3 solution.8. Write the equations for the reaction that took place in 7 (a) and 7(b).

Activity 12.6 (b)

1. Dissolve 0.5g of KIO3 solid in water and make 20 ml of solution.

2. Add 20 ml of a 1mol/litre solution of HCl acid to the solution of KIO3.

3. Pour 2 ml of the above mixture in a test tube.

4. Add 1 ml of a solution of KI to the solution containing KIO3 and HCl acid.

5. Note down the changes observed.

6. Identify the reagent compound that is oxidizing and the one that is reducing in the above reactions.

Due to the high electronegativity and small size, fluorine forms only one oxoacid, HOF known as fluoric (I) acid or hypofluorous acid. Other halogens form many oxoacids. Most of them cannot be isolated in pure state. They are stable only in aqueous solutions but they form stable salts.

Halogens normally form 4 series of oxoacids, namely: Hypohalous acids (HXO), halous acids (HXO2 ), halic acids (HXO3) and perhalic acids (HXO4).

The difference betweem oxoacids is the oxidation state of the halogen in each oxoacids or halates.

Table 12.3: Oxohalic acids

a) Acid character

Table 12.3 shows that the acid strength increases with increasing oxidation state of the halogen.

Example: Among of the four oxoacids of chlorine, the acidity increases as follows:

HClO < HClO2 < HClO3 < HClO4

In other words the acidity increases with the number of oxygen atoms forming bonds with the halogen atom.

Example: Potassium chlorate decomposes at 400oC to give potassium perchlorate and potassium chloride.

At higher temperatures, the perchlorate decomposes to give the chloride and oxygen.

Potassium bromate and iodate give similar products on strong heating. No perbromate and periodate is formed during the heating. The halates are also strong oxidizing agents as their corresponding acids behave. Potassium chlorate (V) oxidizes hot concentrated hydrochloric acid to an explosive bright yellow gaseous mixture

Checking up 12.6

1. Write a balanced equation for the reaction of:

a) Decomposition of KClO4

b) Oxidation of Cl- by ClO3-in acid medium.

2. Explain why fluorine forms only 1 oxoacid (HFO).

12.7. Uses of halogens and their compounds

Activity 12.7

1. When you want to eat food, salt is dissolved in it. Indicate the chemical composition of table salt and its natural occurrence.

2. Chlorine compounds are used in the treatment of water. Explain how chlorine re-acts to be a good disinfectant in water treatment.

3. a)Write the observations of the phenomenon that takes place when electrolysis of a concentrated solution of chlorine is carried out in the laboratory.

b) Deduce the product of reaction that is formed at the anode.

12.7.1. Uses of Halogens

Halogens and their compounds have many applications and uses in different domains:

•Fluorine is used in production of synthetic fibres. It is also used as an ingredient in toothpastes as well as in the manufacture of HF.

•NaCl is the main food seasoning table and kitchen salt; it is also used in many industrial processes such as: NaOH production, soap manufacturing, etc...

•Chlorine is used to synthesise products for bleaching clothes, papers, etc. it is used as an antiseptic and as fungicide.

•Chlorine is used in water treatment as well as in the manufacture of plastics such as Polychloroethene (polyvinyl chloride: PVC).

•Chlorine is used to make DDT (Dichlorodiphenyltrichloroethane: a banned chemical), and other chlorinated aromatic compounds used as pesticides.

•A Chlorine compound, HCl is important in in the human stomach for digestion.

•Halogen elements are used to manufacture polymers.

•Bromine is used in photographic industry (film manufacture).

•Iodine is used in food in form of iodised salt and in drugs to fight against goitre and to kill bacteria in wounds, etc.

12.7.2. Hazards caused by group 17 elements

Bromine effects

•On heating, toxic fumes are formed.

•Reacts violently (explosively) with many compounds.

•Attacks plastics, rubber and coatings.

Chlorine effects

•It reacts violently with many compounds like ammonia and may cause fire and explosion.

•It attacks many metals in the presence of water.

•It attacks plastics, rubber and coatings.

Chlorine oxide effects

•It may explosively decompose when it encounters shock and friction then it may explode on heating.

•It reacts violently with mercury, phosphorus, sulphur, etc causing fire and explosion hazard.

Fluorine effects

•It reacts violently with water to produce toxic and corrosive vapours: ozone and hydrogen fluoride.

•It reacts violently with ammonia, metals, oxidants, etc, to cause fire and explosion.

Hydrogen bromide effects

•It reacts violently with strong oxidants and many organic compounds to cause fire and explosion.

•It attacks many metals forming flammable hydrogen gas.

Hydrogen fluoride effects

•It reacts violently with many compounds causing fire and explosion.

•On contact with air, it emits corrosive fumes which are heavier than air.

•It attacks glass and other silicon-containing compounds.

SF6 effects

•The substance decomposes in a fire to produce toxic fumes of sulphur oxides and hydrogen fluoride •When it is heated, there is formation of toxic fumes.

CFCs: Chlorofluorocarbons

Chlorofluorocarbons are hydrocarbons (alkanes) where some or all of the hydrogen atoms have been replaced by chlorine and fluorine atoms. Most of those compounds are stable and unreactive at high temperature. That is why they were used as aerosol propellants, refrigerants, and solvents.

One example of CFC is Freon 12 (CCl2F2) that was used in refrigerators.Because of their chemical inertness, CFCs can diffuse unchanged into the upper atmosphere up to the ozone layer (10-15 km). There, photochemical reaction cause them to break down into radicals, .Cl. Radicals, being very active chemical species, react with ozone to form ordinary oxygen, hence destruction of ozone layer. For this reason, their use has been discourage or banned all over the World.

Checking up 12.7

1. Indicate at least 3 uses of chlorine compounds in daily life.

2. State at least 2 uses of bromine and its compounds in daily life.

3. Why do we use iodised salt as table and cooking salt?

4. Discuss the importance of halogen compounds in biological area.

5. Give at least 2 negative effects of halogens compounds

12.8. End of unit assessment

1. Chlorine is used in the preparation of bleaching agents and iodine.

(a) Write the reaction of Chlorine with water and mention the bleaching group

(b) Bromine is extracted from sea water by oxidising bromideions with chlo-rine gas.

(i) Write the ionic equation for this reaction.

(ii) Explain why chlorine is strong oxidising agent compared to bromine

2. Aqueous silver nitrate can be used as a test for halide ions. A student decided to carry out this test on a solution of magnesium chloride. The bottle of mag-nesium chloride that the students used showed the formula MgCl2.6H2O. The student dissolved a small amount of MgCl2.6H2O in water and added aqueous silver nitrate solution.

a)What would the student expect to see after adding aqueous silver nitrate?

b) Write an ionic equation for this reaction. Include state symbols.

c) Using aqueous silver nitrate, it is sometimes difficult to distinguish between chloride, bromide and iodide ions.What test are you going to use to distinguish between them? Explain

d) When carrying out halide test with aqueous silver nitrate, it is important that distilled water is used for all solutions, rather than tap water.

Suggest why it is done this way.

1. Chlorine gas is bubbled through an aqueous solution of bromide ions and alsothrough an aqueous solution of iodide ions. What will be observed? Explain the observation by chemical equations

2. A student carried out experiments using chlorine gas (Cl2) with sodium hydrox-ide solution and the equation for this reaction is shown below:

The student repeated the procedure in but with hot concentrated sodium hydroxide. A different reaction took place in which sodium chlorate (V) was formed instead of NaClO.

Write the equation for the reaction to produce sodium chlorate (V).

Key unit competence: Compare and contrast the properties of the group 18 elements in relation to their position in the periodic table.

Learning objectives

By the end of this unit, I will be able to:

•State the physical properties of the Group 18 elements.

•Explain the lack of reactivity of the group 18 elements.

•Associate chemical inertia of the group 18 elements to their full valence shell.

•Recognize the importance of noble gases or group 18 elements in the daily life.

Introductory activity

Make a research to find out the type gas :Inside the Bulb, in balloon, responsible for different colors dispayed by this house (or in advertising sings)

Group 18 consists of six elements: helium(2He), neon(10Ne), argon(18Ar), krypton(36Kr), xenon(54Xe) and radon( 86Rn).They are all gases and are characterized by filled s and p valence orbitals (valence electron configuration of 2s2 for helium and ns2np6 for the others).This structure has been qualified as ”the octet electronic structure”; octet means eight. This electronic structure, also called octet rule gives stability to the elements of Group 18.

Because of this, these elements are very unreactive. They form very few compounds. Because of this they are termed noble gases.

Other chemical elements tend to react, to combine with other elements or with themselves in order to fulfill the octet rule, i.e. to acquire the same electronic structure as the nearest noble gas.

13.1. Occurrence and physical properties of noble gases

Activity 13.1

The air is composed of a mixture of gases including water vapour.

i)Make a research (with any documentation) to identify its components and ar-range them according to their abundances (Component1> Component 2, etc...)

ii)Show how these components can react each other if possible •If not possible, justify your answer.

iii)Explain how neon lamp works.

13.1.1. Occurrence

•All the noble gases except radon occur in the atmosphere. Their total atmospheric abundance in air is 0.03%; argon is the major component.

• Helium and sometimes neon are found in minerals of radioactive origin e.g., pitchblende, monazite, cleveite.

.Xenon and radon are the rarest elements of the group. Radon is obtained as a decay product of 226Ra.

13.1.2. Physical properties

•All the noble gases are monoatomic.

•They are colourless, odourless and tasteless.

•They are sparingly soluble in water.

•They have very low melting and boiling points.

•Due to stable electronic configuration these gases exhibit very high ionisation energy. The ionization energy decreases down the group with increase in atomic size.

•Since they have stable electronic configurations, they have low tendency to accept or give up electron.

Helium has the lowest boiling point of any known substance: ~ 267 – 2690C

Checking up 13.1

Question:

(i) Using dot and cross diagram, show how neon(z =10) is considered as a noble gas.

(ii) Why do we classify helium in Group 18?

13.2. Reactivity and inertness of noble gases

Activity 13.2

1. Through a research, can you find any compound made of one of the rare gas?

2. Do you think that noble gases can react with one of these reagents in normal conditions?

Although group 18 elements used to be called inert gases, it has been discovered that some of them, in appropriate conditions can react and form compounds.

But up to now, of the group 18 elements, only krypton and xenon have been shown to form chemical compounds. The first of these was prepared in 1962 by Neil Bartlett,

an English chemist who made the ionic compound XePtF6, which contains the ions Xe+ and PtF6-:

It has also been discovered that xenon tetrafluoride can be prepared by reacting xenon and fluorine gas in a nickel reaction vessel at 400oC and 6 atmospheres:

Xenon tetrafluoride forms stable colorless crystals. Two other xenon fluorides, XeF2and XeF6, were synthesized by the group of Argonne, and a highly explosive xenon oxide (XeO3) was also found. The xenon fluorides (XeF2, XeF4 and XeF6 ) are colourless crystalline solids and sublime readily at about 25oC. They are powerful fluorinating agents.

They are readily hydrolysed even by traces of water, for example:

In the past years other xenon compounds have been prepared, for example, XeO4(explosive), XeOF4, XeOF2, and XeO3F2. These compounds contain discrete molecules with covalent bonds between the xenon and the other atoms. A few compounds of krypton, such as KrF2 and KrF4, have also been observed. There is evidence that radon also reacts with fluorine, but the radioactivity of radon makes its chemistry very difficult to study.

Checking up 13.2

Question: Explain why in some applications such as air balloons, helium preferred to hydrogen?

13.3. Uses of noble gases

Activity: 13.3

Do a research (with any documentation) to find how each noble gas has been discovered and its uses?

Helium

•Helium is a non-inflammable and light gas. Hence, it is used in filling balloons for meteorological observations, replacing the flammable hydrogen gas.

•It is also used in gas-cooled nuclear reactors.

•Liquid helium (B.P:-267.8oC) finds use as cryogenic agent for carrying out various experiments and conservation at very low temperatures.

Neon

•Neon is used in advertising signs, it glows when electricity is passed through it. Different coloured neon lights can be made by coating the inside of the glass tubes with colored chemicals.

•Neon bulbs are more used in our daily life.

Argon

•It is used in light bulbs. The very thin metal filament inside the bulb would react with oxygen and burn away if the bulb were filled with air instead of argon.

•Argon is used mainly to provide an inert atmosphere in high temperature metallurgical processes (arc welding of metals or alloys).

•It is also used in the laboratory for handling substances that are air-sensitive.

Krypton

Krypton is used in lasers. Krypton lasers are used by surgeons to treat certain eye problems and others. It is used in light bulbs designed for special purposes.

Xenon

Xenon is used in fluorescent bulbs, flash bulbs and lasers. Xenon emits an instant, intense light when present in discharge tubes. This property of xenon is utilized in high-speed electronic flash bulbs used by photographers.

Radon

Radon is radioactive and is used in medicine as a source of gamma rays. The gas is sealed in small capsules, which are implanted in the body to destroy malignant (e.g., cancerous) growths.

13.4. End unit assessment

1. a) Give a reason why the first ionization energies of noble gases are very high.

b) State one use of neon and give a reason to support your answer.

c) State and explain the trend in atomic radius among noble gases.

d) Why are noble gases unreactive?

e) Explain why the value of the first ionisation energy of neon is higher than that of sodium.

2. Explain why Group 18 elements are rare on Earth?

3. The discovery of compounds of noble gases has been done, up to date, with Xe and Kr, not with He or Ne. Can you suggest a probable reason?

Key unit competency: Compare and contrast the properties of the Period 3 elements and their compounds in relation to their positions in the Periodic Table.

Learning objectives:

By the end of this unit I will be able to:

•Compare the physical properties of the Period 3 elements.

•Describe the nature of the oxides of the Period 3 elements and the type of bonding in their chlorides, oxides and hydrides.

•Relate the physical properties of the Period 3 elements to their position in Periodic Table.

•Relate the physical properties of compounds of the Period 3 elements to their nature of bonds across the period.

14.1. Physical Properties of the Period 3 elements

Activity 14.1

1. Write the electronic configuration of the following elements in terms of s, p, d and f...(i) Sodium (ii) Magnesium (iii) Aluminium (iv) phosphorous (v) sulphur

2.Considering the electronic configuration of magnesium and Aluminium, phos-phorus and sulphur. How do you expect their ionization energies to vary?

3. How do you expect the general trend in ionization energy, electron affinity, melting and boiling point, electronegativity to vary for the elements in the pe-riod 3?

4. Considering the electronic configuration of magnesium and Aluminium, phos-phorus and sulphur. What can you say about them, how do you expect their ionization energies to vary?

(a) Variation of First ionization energies (IE) of Period 3 elements

First ionization energy generally increases across Period 3 from left to right. However, it drops at aluminium and Sulphur (table 14.1 and Fig.14.1). This can be exaplained in term of more stable electronic structures of the two elements after loosing 1 electron:

Table 14.1: Variation of first ionization energies of period 3

Going across Period 3, there are more protons in each nucleus so the nuclear charge in each element increases. Therefore the force of attraction between the nucleus and outer electron is increased, and there is a negligible increase in shielding because each successive electron enters the same energy level. So apart from the two exceptions mentioned above, the first ionization energy increases from left to right in the period (figure 14.1).

b. Variation of atomic radius of Period 3 elements

Table 14.2: Variation of atomic radius of period 3 elements

Going across Period 3 from left to right, the number of protons in the nucleus increases so, the nuclear charge increases. There are more electrons, but the increase in shielding is negligible because each extra electron enters the same principal energy level. Therefore, the force of attraction between the nucleus and the electrons increases. So the atomic radius decreases as indicated in the Figure 14.2 and table 14.2.

(c)Variation of electronegativity of Period 3 elements

Table 14.3: Variation of electronegativity of period 3 elements

Going across Period 3 from left to right, electronegativity increases almost linearly due to the nuclear charge increase as atomic radius decreases. There are more electrons, but the increase in shielding is negligible because each extra electron enters the same principal energy level so electrons will be more strongly attracted to the nucleus.

You might expect argon (with 18 electrons) to be the most electronegative element in Period 3, but its outer energy levels are full. Therefore, it does not form covalent bonds with other atoms, so it is given an electronegativity value of zero.

d. Variation of melting and boiling points in Period 3

Melting and boiling points generally increase going from sodium to silicon, then decrease going to argon with a “jump” at Sulphur (Fig 14.4 and Table 14.4).

Table 14.4: Variation of melting and boiling points of period 3 elements

Sodium, magnesium and aluminium are all metals. They have metallic bonding, in which positive metal ions are attracted to delocalized electrons. Going from sodium to aluminium, the charge on the metal ions increases from +1 to +3 through magnesium at +2, the number of delocalized electrons increases, so the strength of the metallic bonding increases and the melting points and boiling points increase.

Silicon has giant covalent bonding structure. It has a giant lattice structure similar to that of diamond, in which each silicon atom is covalently-bonded to four other silicon atoms in a tetrahedral arrangement. This extends in three dimensions to form a giant macromolecule and this explains the very high melting point and boiling point.

Phosphorus, sulphur, chlorine and argon, are all non-metals, and they exist as small, separate molecules. Phosphorus, sulphur and chlorine exist as simple molecules. Argon exists as separate atoms (monatomic). When these these elements melt or boil, it is the van der Waals’ forces between the molecules which are broken; being weak they need little energy to overcome them.This explains why their melting and boiling points are low. However, Sulphur has a higher melting point and boiling point than the other three because

Phosphorus exists as P4 molecules, Sulphur exists as S8 molecules, chlorine exists as Cl2 molecules, and argon exists as individual Ar atoms.

The strength of the van der Waals’ forces decreases as the size of the molecule decreases, so the melting points and boiling points decrease in the order S8> P4> Cl2>Ar

e. Variation of electrical conductivity of Period 3 elements

Electrical conductivity increases going across Period 3, left to right, from sodium to aluminium, then decreases sharply to silicon as indicated by the graph 14.5 below.

Table 14.5: Variation of relative electrical conductivity of period 3 elements

The delocalized electrons are free to move and carry charge. Going from sodium to aluminium, the number of delocalized electrons increases, there are more electrons which can move and carry charge so the electrical conductivity increases.

Silicon is called a semi-conductor because at higher temperatures more electrons are promoted to the higher energy levels so there are more delocalized electrons to move and carry charge.

Phosphorus, sulphur and chlorine, the outer electrons are not free to move and carry charge because they are held strongly in covalent bonds. In argon (mono atomic) the outer electrons are not free to move and carry charge because they are held strongly in a stable third energy level and this explains their zero electrical conductivity.

f. Variation of metallic character of period 3 elements

Metallic character decreases as you move across a period 3 in the periodic table from left to right. This occurs as atoms more readily accept electrons to fill the valence shell than lose them. Note that as the metallic character decreases across the period, the reducing power decreases whereas oxidizing power increases.

g. Variation of electron affinity across period 3 elements

The electron affinity [EA] is the energy change for the process of adding an electron to a gaseous atom to form an anion (negative ion).

As we might predict, it becomes easier to add an electron to an atom across the period from left to right as the effective nuclear charge of the atoms increases. As we go from left to right across a period 3, EAs tend to become more negative, i.e. ability to acquire electrons increases.

The exceptions found among the elements of group 2 (2A), group 15 (5A), and group 18 (8A) can be understood based on the electronic structure of these groups.

Magnesium and phosphorous have anomalous electron affinity, Magnesium has a positive EA while phosphorous is less negative. This is due their electron arrangement, where in magnesium the electron should be added to 2p orbital which is a less stable structure compared to the electronic structure of the atom [Ne]3s2. Similarly in phosphorous the electron should be added to 3p orbital which is half filled and thermodynamically stable.

Checking up 14.1

1.Explain the variation of the following terms as applied in period 3 of the periodic table:

(i) Ionization energy, Electronegativity, )

(ii)Explain the anomalous behavior indicated by magnesium and phospho-rous in graph 14.1 above

2.The table below shows the melting points of the period 3 elements except for silicon:

(a)Explain in terms of bonding why the melting point of magnesium is higher than that of sodium.

(b) Predict the approximate melting point of silicon.

(c) Explain why chlorine has a lower melting point than sulphur.

(d) Explain the variation of metallic character, electronegativity, atomic radii ,first ionization energy, melting and boiling points, electron affinity and electrical conductivity across the period

14.2. Chemical properties of period 3 elements

Activity 14.2

a.Experiment to investigate the action of water on period 3 elements

Materials /apparatus

Water , test tubes, a piece of sodium metal, aluminium powder, magnesium ribbon/powder, red litmus paper/ universal indicator

Procedure:

1. Put about 5cm3 of water in each of the three test tubes arranged in a test tube rack

2. Add a small piece of Mg, Na and Al in each of the three test tubes in 1.

3. Record your observations in a suitable table.

Study questions:

1. What do you say about your observations made in experiment above.

2. Write equation for the reaction that occurs in each test tube in procedure 2.

(b) Experiment to investigate the action of heat on period 3 elements

Materials /apparatus:

Water , test tubes, a piece of sodium metal, aluminium power/sheet, magnesium ribbon/powder, phosphorous and sulphur powder, universal indicator , pair of tongs, source of heat

Procedure:

1. Hold a piece of magnesium ribbon on a Bunsen flame and record you observa-tion.

2. Repeat experiment 1 for sodium, aluminium, phosphorous and sulphur and record your observation in each case.

3. For each of the products formed i.e. for metal oxides formed, add water and dip a litmus paper to test their nature.

Note: if the oxide is gaseous hold a piece of litmus paper on the mouth of the test tube.

Study questions:

1.Write equations to show how the metals react with oxygen.

2. What would you expect to observe when the metal is burned in oxygen.

a) Reaction with water

Reactivity with water generally decreases across the period from left to right because there is a decrease in metallic properties.

i) Sodium reacts vigorously with cold water to form sodium hydroxide and hydrogen gas.

ii) Magnesium does react slowly with cold water to form magnesium hydroxide solu-tion, but with steam the reaction is faster to form magnesium oxide and hydrogen gas.

iii) Aluminium reacts with steam very slowly to form aluminium hydroxide and hy-drogen gas.

Aluminum powder heated in water vapor produces hydrogen and aluminum oxide. The reaction is relatively slow because of the existing strong aluminum oxide layer on the metal, and the build-up of more oxide during the reaction.

iv)Silicon reacts a little bit faster with steam at red hot to produce silicon dioxide and hydrogen.

v. Phosphorus and sulfurhave no reaction with water

.vi)Chlorine dissolves in water to give a green solution. A reversible reaction produces a mixture of hydrochloric acid and chloric(I) acid (hypochlorous acid).

In the presence of sunlight, the hypochlorous acid slowly decomposes to produce more hydrochloric acid, releasing oxygen gas:

There is no reaction between of argon and water that is known.

b. Reaction with oxygen

Elements must be heated to react with oxygen; however, dry white phosphorus can ignite spontaneously and that is why it is stored under water. The reactivity depends much on the state of subdivision

Sodium and magnesium: Sodium and magnesium burnvigorously in oxygen to form ionic white oxides

Aluminium: sheets of aluminium get slowly coated with thin oxide layer of aluminium oxide.

c. Reaction with NaOH

d. Reaction with HCl.

e. Reaction with Hydrogen

f. Reducing and oxidizing power

Elements on the left of the period three are metal.They react by losing their valance electrons; hence they are good reducing agents.Their reducing power decreases from Na to Al.

The elements on the right of the period three are nometals. They react by gaining or sharing their valance electrons; they are good oxidising agents.Their oxidising power increases from Si to Cl.

In general the oxidizing power increases from left to right and the reducing power decreases from left to right across period 3 due to decrease of atomic size that affects the ionization energy and electronegativity.

Checking up 14.2

1. Describe the nature of hydrogen compounds of period 3.

2. Explain why the reducing power of period 3 elements decreases across the period.

14.3. Compound of period 3 elements

The oxides of period 3 elements:

Activity 14.3(a)

1. Write the formulae of the oxides of period 3 elements

2. What did you consider when writing the formulae of the oxide in 1 above?

3. How do you expect the oxides to behave in water? Explain your answer.4. Suggest the trend of acid- base character of the oxides of period 3

The metallic oxides on the left of the period adopt giant structures of ions. In the middle, silicon forms a giant covalent oxide (silicon dioxide); the elements on the right form simple molecular oxides with simple structures. The intermolecular forces binding one molecule to its neighbors are van der Waals dispersion forces or dipole-dipole interactions.

Physical properties of the oxides of period 3 elements

Melting and boiling points: the metal oxides and silicon dioxide have high melting and boiling points because a large amount of energy is needed to break the strong bonds (ionic or covalent) operating in three dimensions. The oxides of phosphorus, sulfur and chlorine consist of individual molecules.

Electrical conductivity: None of the oxides above have any free or mobile electrons, indicating that none of them will conduct electricity when solid.

Phosphorus has two common oxides, phosphorus (III) oxide, P4O6, and phosphorus(V) oxide, P4O10.

Phosphorus (III) oxide is a white solid, melting at 24°C and boiling at 173°C.

Phosphorus(V) oxide is also a white solid, which sublimes at 300°C. In this case, the phosphorus uses all five of its outer electrons in the bonding.

Sulphur has two common oxides, sulphur dioxide (sulphur(IV) oxide), SO2, and

sulphur trioxide (sulphur(VI) oxide), SO3.

Sulphur dioxide is a colorless gas at room temperature with an easily recognized pungent smell. It consists of simple SO2 molecules.

Pure sulphur trioxide is a white solid with a low melting and boiling point. It reacts very rapidly with water vapor in the air to form sulfuric acid. Under laboratory

conditions, it forms a white sludge which fumes dramatically in moist air, forming a fog of sulfuric acid droplets.

Chlorine forms several oxides. Two are considered here: chlorine(I) oxide, Cl2O, and

chlorine(VII) oxide, Cl2O7.

Chlorine (I) oxide, Cl2O, is a yellowish-red gas at room temperature. It consists of simple, small molecules.In chlorine (VII) oxide, Cl2O7, the chlorine losesall of its seven valence electrons in bonds with oxygen. This produces a molecule much larger than chlorine (I) oxide, suggesting higher melting and boiling points. Chlorine (VII) oxide is a colorless oily liquid at room temperature.

Checking up 14.3(a)

1. The table below shows oxides of period 3 in the periodic table

(a)Copy and complete the table using the following guidelines

(i)Complete the ‘bonding’ row using only the words ionic or covalent

(ii) Complete the ‘structure’ row using simple molecular,giant ot lattice

(iii) Explain in terms of forces, the difference between the melting points of MgO and SO3.

b. The oxides Na2O, Al2O3, SO3, were each separately added to water. For each oxide, construct a balanced equation for its reaction with water

Acid-base Behavior of the Oxides

Activity 14.3 (b)

1. Classify the oxides in terms physical states of the oxides of period 3.

2. How do you expect the oxides react with water, acids, and sodium hydroxide.(use equations to justify your answer)

3.(a) Predict the nature of oxides of period 3 elements when dissolved in water.

(b)What would you expect to observe when both blue and red litmus papers are dropped into each of the solutions formed in question (2) above in water.

Acidity increases from left to right, ranging from strongly basic oxides on the left to acidic ones on the right, with an amphoteric oxide (aluminum oxide) in the middle.

Checking up 14.3(b)

1.Consider the following oxides: CaO, Al2O3 , Na2O , MgO , P2O5 , SO2 SiO2.

(a)State which of the oxides arei)Basic (ii) acidic (iii) neutral (iv)Amphoteric

(b)Describe briefly how you could test to see whether a solid oxide is basic , acidic , neutral or Amphoteric 2. (a)In each of the following species state whether the metal to non metal bond is ionic or covalent and explain why : Na2O , Al2O3

(b) An uknown inorganic oxide is a white solid of melting temperature 17100C and is uncreative towards water.(i)State and explain whether or not this is sufficient information to deduce whether the bonding is ionic or covalent.(ii) Outline one additional test which could be applied to support your conclusion.

14.3.2. Chlorides of Period 3 Elements

Activity 14.3(c)

The table below gives some data about the chlorides of elements of period 3

(a) Explain why the boiling point of NaCl is higher than that of MgCl2.

(b) Make a research and explain why the pH of NaCl is 7 and that of AlCl3 is 2.

The chlorides of interest are given in the table below:

Sulphur forms three chlorides, but S2Cl2 is the most common. Aluminum chloride may be found, in certain conditions, as a dimmer, Al2Cl6. Phosphorous forms two chlorides, PCl3 and PCl5.

The nature and structure: Sodium chloride and magnesium chloride are ionic and consist of giant ionic lattices at room temperature. Aluminum chloride exhibits covalency characters. Aluminium ion has high charge density; due to this, the electron cloud of the chloride is distorted toward the aliminium ion, Al3+, impacting an appreciable covalent property to the bond.

The other chlorides are simple covalent molecules.

Melting and boiling points of Period 3 chloride

Sodium and magnesium chlorides are solids with high melting and boiling points because of the large amount of heat which is needed to break the strong ionic attractions.

Aluminum chloride and phosphorus(V) chloride are solids with relatively low melting and boiling points. The remaining chlorides are volatile liquids due to the weak van der Waals that hold their molecules together.

Electrical conductivity: solid chlorides do not conduct electricity because the ions are not free to move

Sodium, magnesium and aluminium chlorides are ionic and so will conduct electricity when they are molten or in aqueous solution. The rest of the chlorides do not conduct either in solution or molten state due to absence of ions.

Reactions with water

The chlorides of more electropositive metals (NaCl and MgCl2) dissolve in water to form neutral solutions. The other chlorides (covalent) hydrolyze or react with water to form acidic solutions. As shown below:

The greater positive charge on aluminium attracts electrons in the water molecules quite strongly, making O-H bond of water molecule in the complex ion so weak to dissociate and liberate H+ ion.

Checking up 14.3(c)

1. a. Distinguish between dissolving and hydrolysis.

b. Name one chloride that dissolves in water, and one chloride that undergo hydrolysis.

c. State how the bonding in the chlorides changes on crossing the second and third periods from left to right

d. Suggest two ways how you would know that a reaction has taken place when a few drops of water is added to silicon (iv) chloride.

2. a. Write an equation that shows thata solution of beryllium chloride in water is acidic.

b. Predict the shape of PCl3 molecule and suggest a likely bond angle.

c. Iron (iii) chloride, FeCl3, forms dimmers in the gas phase similar to those of aluminium chloride. Draw the likely structure of these two dimers.

4. Carbon and silicon are in the same group of the periodic table and forms chloride of CCl4 and SiCl4 respectively. However, CCl4 does not undergo hydrolysis whereas SiCl4 hydrolyses in water to form white fumes of the gas.

14.3.3. The hydrides of period 3 elements

Activity 14.4(d)

1. Period 3 elements from sodium to chlorine form different hydrides of different bond nature, physical properties and structure.

(a) Write the formula of the hydrides formed by period 3 elements.

(b) Predict the nature of bonding based on your knowledge of periodicity of ele-ments in the periodic table.

(c ) Basing on the nature of bonding predicted in (b) above. How would you ex-pect their boiling and melting point vary across the period?

(d) Predict the nature of solutions formed by hydrides when dissolved in water. What would you expect to observe if red and blue litmus papers were separate-ly dropped into each solution?

Hydrides are commonly named after binary compounds that hydrogen forms with other elements of the periodic table. Hydride compounds in general form with almost any element, except a few noble gases. The common hydrides of period 3 elements are as shown in the table 14.5 below

The hydrides above are examples of period 3 elements with some of their properties summarized in the tabl. As we can see the hydrides of period 3 vary from ionic hydride such as NaH at the left side to polar covalent hydride such as HCl at the right side of the period.

Checking up 14.3(d)

1. The table below gives some properties of oxides of period 3 in the periodic table.

a. Explain the trend in the melting points of the above oxides and their origin.

b. Write the equations to show the reaction of

i. SO3 with water

ii. Na2O

c. One of the above oxides has acidic and basic character.i. Identify that oxide from the above table ii. Using equations explain the acidic and basic character of the oxide identified in (c) (i).

2. This question concerns the following oxides: Na2O, MgO, SiO2, and SO3.From the list above identify the oxide that best fits the description given:

i. An oxide that is insoluble in water.

ii. An oxide that has simple molecular structure at room temperature and pressure.

iii. An oxide that reacts with water forming a strongly alkaline solution.

iv. An oxide that is slightly soluble in water forming weak alkaline solution.

3. There is a link between the properties of the oxides of the Period 3 elements and their structure and bonding. The table below shows the melting points of the oxides of some Period 3 elements.

a. In terms of crystal structure and bonding, explain in each case why the melting points of sodium oxide and silicon dioxide are high.

b. Predict whether the melting point of lithium oxide is higher than, the same as, or lower than the melting point of sodium oxide and explain your prediction.

c. Phosphorus (V) oxide has a lower melting point than sodium oxide.

i. State the structure of and bonding in phosphorus (V) oxide.

ii. Explain why the melting point of phosphorus(V) oxide is low.

d. Samples of phosphorus(V) oxide and sodium oxide were reacted with water. In each case, predict the pH of the solution formed and write an equation for the reaction.

4. Sodium chloride is a high melting point solid which dissolves in water to make a colorless solution. Silicon (IV) chloride is a liquid at room temperature which fumes in moist air, and reacts violently with water.

a. Draw a diagram to show the arrangement of the particles in solid sodium chloride, making clear exactly what particles you are talking about.

b. Explain why this arrangement leads to a high melting point.

c. Draw a simple diagram to show the structure of silicon (IV) chloride, and explain why silicon (IV) chloride is a liquid at room temperature.

d. Why is there such a big difference between the chlorides of sodium and silicon?

e. Briefly describe and explain the difference in electrical conductivity between sodium chloride and silicon (IV) chloride in both solid and aqueous molten state.

f. Write an equation to show what happens when silicon (IV) chloride reacts with water.g. Name another Period 3 chloride which behaves similarly to sodium chloride, and one which behaves similarly to silicon (IV) chloride.

5. With the help of equation describe how the hydrides of period 3 react with water.

14.5. End unit assessment

1. Use the information in the following table to explain the statements below

(a)What is the trend in atomic radius? Explain the origin of that trend.

(b) The ionic radii of Na+, Mg2+ and Al3+ are less than their respective atomic radii , whereas the ionic radii of Cl- and S2- are greater than their respective atomic radii.Compare the atomic radii with the ionic radii and explain what your observation

(c) What trend do you observe in the 1st ionization energy? Explain the origin of that trend?

d. The first ionization energy of Al is less than that for Mg; why?

2. The table below shows the melting points of the period 3 elements except for silicon.

a. Explain in terms of bonding why the melting point of magnesium is higher than that of sodium

b. State the type of bonding between atoms in the element silicon and name the type of structure which silicon forms.

c. Predict the approximate melting point of silicon.

d. Explain why chlorine has a lower melting point than sulphur.

e. Predict the approximate melting point of potassium and give one reason why it is different from that of sodium.

3. The elements Sodium, Magnesium, silicon, phosphorous and chlorine are members of the third period of the periodic tablea.

a.i. Write down the formula of the principal oxides and chlorides of the elements listed above and in each case indicate the type of bonding.

ii. Explain what happens when each of the above oxides and chloride is added to water and indicate whether the resultant solution will be acidic, basic or neutral.

c. The melting points of Mg , Si and S are 6500C, 14230C respectively. Explain the differences in the melting points of the elements.

d. Name the type bonding that exists in the hydrides of the elements Sodium, Phosphorous and sulphur and write the equations to show the reactions if any of the hydrides with water.

4. Choose from the elements: Sodium, magnesium aluminium, silicon, phospho-rous, chorine and argon a. List the elements that react readily with cold water to form alkaline solutions. And write the equations for the reactions.

b. List the hydrides that have hydrides with low boiling points/temperatures and explain why.

c. List the elements that form nitrates and write the formulae of nitrates.

d. What is the most ionic compound that can be formed by the combination of two of these elements.

e. Which element has both metallic and non metallic properties?

f. Name the elements that normally exist as molecules.

5.a. Describe the formation of Al2Cl6 dimers from AlCl3 monomers in terms of or-bital hybridization

b. i. Which of the oxides listed above form oxide of the type X2O3

ii.Describe briefly how you would prepare each of the oxide and give the equations for the reactions involved.iii. What would be observed if each of the above oxide in (b)

(ii) was reacted with sodium hydroxide? Write equation for the reactions involved in each case

 Key unit competency: Deduce how concentration, pressure, catalyst and temperature affect the chemical processes in industry.

Learning Objectives

By the end of this unit, students will be able to:

•Distinguish between complete and reversible reactions.

•Explain dynamic equilibrium.

•State the characteristics of dynamic equilibrium.

•Explain the factors that affect the position of the equilibrium in a reversible reaction.

•Apply Le Châtelier’s principle to explain the effects of changes in the temperature, concentration and pressure on a system in equilibrium.

•Compare and contrast theoretical and actual optimal conditions in the industrial processes.

•Relate the effect of concentration, temperature, pressure and catalyst to the amount of products in the manufacturing industries.

•Recognize the importance of Le Châtelier’s principle in Haber and Contact processes.

Introductory activity

Observe the figures bellow.

1. What is the difference between the figure a and b?

2. What conditions for the figure b to be like figure a?

3. What will happen if one person on the right side leaves his/her group in figure C?

The above figure shows that when two teams pull on a rope with equal force. The resulting force is equal in magnitude and equal to zero and the rope does not move, the system is said to be in equilibrium. Students in figure (a) represent a system in equilibrium. The equal and opposite forces on both ends of the seesaw are balancing. If, instead one force is greater in magnitude than the other, the system is not in equilibrium [ figure (b)]

In chemistry, a chemical reaction is a process where old bonds are broken and new bonds are formed. For a chemical reaction to take place, two or more substances called reactants are interacted. In general, when reactants collide with sufficient energy and in a proper orientation, the products are formed. Many chemical reactions proceed to a certain extent and stop. In some cases, reactants combine to form products and the products also start combining to give back the reactants. When such opposing processes take place at equal rates, no reaction appears and it is said that a state of equilibrium has reached.

15.1. Difference between complete and incomplete reactions (irreversible versus irreversible reactions)

Activity 15.1

1. Write any two equations of your choice to show a reaction that undergo com-pletion.

2. Write any two equations of your choice to show a reaction that does not go completion

A chemical reaction can proceed in either non-reversible (irreversible or complete) or reversible reaction.

During chemical processes, many chemical reactions do not undergo completion but instead they attain a state of chemical equilibrium. Chemical reaction can proceed in either non-reversible (irreversible or complete) or reversible reaction.

A non-reversible reaction is a reaction which proceeds in only one direction, in other words,the reactants are completely transformed into products.

In reversible reactions, both the forward and reverse reactions occur at different rates at the beginning but approach the same rate at equilibrium state.

Reversible reactions are indicated by placing two half arrows pointing in opposite directions between the reactants and products. Note that all reactions at equilibrium can be forced to go in either direction depending on change of external factors.

Reversible reactions are indicated by placing two half arrows pointing in opposite directions between the reactants and products.The forward reaction is indicated by an arrow oriented from left to right and the reverse reaction from right to left.

Checking up 15.1

1. Identify the reactions which are non-reversible and reversible by writing the words irreversible and reversible.

15.2. Concept of equilibrium (dynamic equilibrium) and its characteristics

Activity 15.2

1. Explain the terms used in equilibrium reactions.

(a)Reversible reaction

(b) equilibrium state

(c) dynamic equilibrium

(d) position of equilibrium.

2. Suggest and explain the characteristics of dynamic equilibrium and how it can be attained.

3. Learners should do a tug-of-war game outsidetheclassroom and comment on the game.

4. In a given Hotel, clients enter others leave. At a certain moment if the number of leavers and arrivals is equal, the number of the clients in the Hotel doesn’t change.

i. Has the movements of clients coming in and out stopped?

ii. How can you qualify that status?

iii. How can you compare this with chemical equilibrium?

15.2.1. Concept of equilibrium reactions

When a chemical reaction takes place in a container which prevents the entry or escape of any of the substances involved in the reaction, the quantities of these components change as some are being consumed and others are being formed at the same time.

Chemical equilibrium is the state at which the rate of forward reaction becomes equal to the rate of backward reaction.

At the initial state, the rate of forward reaction is greater than the rate of backward reaction. However as the products are formed, the concentration of reactants decreases and the concentration of products increases.

The state of chemical equilibrium can be shown graphically as follows in Figures 15.4(a) and 15.(b)

At the equilibrium, the rate of formation of products is equal to that offormation of reactants.

Consider the reaction AB; the figure15.4 (b) indicates how the concentration of A decreases while that of B increases for the reaction. The dotted vertical line indicates the time when the concentrations of A and B are no longer changing.

If the reversible reaction is carried out in a closed system, the reaction is said to be in the equilibrium state when the forward and backward reaction occur simultaneously at the same rate and the concentrations of reactants and products do not change with time (Figure 15.4 b).

At this point, the rates of forward and reverse reactions are the same and the system is said to have reached a state of dynamic equilibrium.

A dynamic equilibrium is a process where the forward and reverse reactions proceed at the same rate;at that moment the concentrations of reactants and products remain constant (do not change).

However, in dynamic equilibrium, even if the concentrations of reactants and products do not change, it does not mean that the reaction has stopped. Rather, the reaction is proceeding in a way that it keeps the concentrations unchanged (the net change is zero).

There are two types of chemical equilibrium: homogeneous and heterogeneous equilibria.

In a homogeneous equilibrium, all the reactants and the products are in the same phase.

In heterogeneous equilibrium, the reactants and the products are present in different phases. This is the case of an aqueous solution in which the ions combine to produce a slightly soluble solid that forms a precipitate or an equilibrium reaction where solid and gaseous phases are present.

15.1.2. Characteristics of a system in a dynamic equilibrium

A chemical equilibrium is a process where reactants are converted into products and products can react to give back the reactants at equal rate. The characteristics of a system in dynamic equilibrium are the following;

1. The rate of the forward reaction is equal to the rate of the reverse reaction,

2. Microscopic processes (the forward and reverse reactions) continue in a balance which yields no macroscopic changes ( nothing appears to be happening),

3. The system is closed and the temperature is constant and uniform throughout the process,

4. The equilibrium can be approached from the left (starting with reactants) or from the right (starting with products).

Checking up 15.2

1. Briefly explain the characteristics of reactions at equilibrium

2. Compare the homogeneous and heterogeneous reactions using specific examples.

3. By giving an example, describe the term dynamic equilibrium.

4. When does a reaction attain equilibrium state?

5. Using a graph and specific examples, explain what happens during a reaction before, at and after the equilibrium has been attained.

15.3. Factors that affect the reactions in equilibrium and Le Châtelier’s principle

Activity 15.3(a)

1. Around 1908-1909 a young German research chemist, Fritz Haber, had discov-ered that nitrogen and hydrogen would form an equilibrium mixture contain-ing ammonia.

(a) Write a balanced equation for the formation of ammonia.

(b)Haber’s experiment yielded an equilibrium mixture containing only 8% by vol-ume of ammonia. What conditions of temperature and pressure does Le Châtel-ier’s principle predict for maximum yield of ammonia at equilibrium?

(c) Why do you think Haber employed the catalyst accompanied with promoters and heat exchanger in his equipment?

2. How is Le Châtelier’s principle used to explain the conditions that affect the equilibrium reactions?

Many industrial processes involve reversible reactions. It is important to understand how the variation of conditions can affect the composition of a chemical equilibrium. Some reactions to take place involve some conditions. For example, the rate of a chemical reaction depends on factors that affect the reaction.

Different factors which can affect the chemical equilibrium include:

1. Temperature

2. Pressure

3. Concentration of reactants and products

The effect of the above-mentioned factors on chemical equilibrium can be explained by the Le Châtelier’s Principle.