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)