1. Chapter 7
Ionic Bonding
7.1 Ionic Bonds: Donating and Accepting Electrons
7.2 Energetics of Formation of Ionic Compounds
7.3 Stoichiometry of Ionic Compounds
7.4 Ionic Crystals
7.5 Ionic Radii

2. Chapter 7 Ionic Bonding (SB p.180)
When a piece of sodium metal is allowed to react with a jar of chlorine gas …... +
1:1 ratio of Na+ and Cl- e- e- e- e-

3. Formation of ionic bond between sodium atom and chlorine atom
Chapter 7 Ionic Bonding (SB p.180)
Formation of ionic bond between sodium atom and chlorine atom Cl Na
Sodium atom Na
1s22s22p6
Chlorine atom Cl
1s22s22p63s23p5

4. linked up together by ionic bond
Chapter 7 Ionic Bonding (SB p.180)
Formation of ionic bond between sodium atom and chlorine atom + -
Sodium ion Na+
1s22s22p6
Chloride ion Cl-
1s22s22p63s23p6 Cl Na
linked up together by ionic bond

5. Ionic Bonds: Donating and Accepting Electrons
7.1 Ionic Bonds: Donating and Accepting Electrons (SB p.181)
Ionic Bonds: Donating and Accepting Electrons

6. Donating and Accepting Electrons
7.1 Ionic Bonds: Donating and Accepting Electrons (SB p.181)
Donating and Accepting Electrons
Ionic bonds are the strong non-directional electrostatic forces of attraction between oppositely charged ions.

7. Donating and Accepting Electrons
7.1 Ionic Bonds: Donating and Accepting Electrons (SB p.181)
Donating and Accepting Electrons + –
Internuclear distance

8. Donating and Accepting Electrons
7.1 Ionic Bonds: Donating and Accepting Electrons (SB p.181)
Donating and Accepting Electrons + – – +
Cationic radius (r+)
Anionic radius (r-)
Internuclear distance
Internuclear distance = r+ + r-

9. Electron transfer from a magnesium atom to two chlorine atoms
7.1 Ionic Bonds: Donating and Accepting Electrons (SB p.182)
Electron transfer from a magnesium atom to two chlorine atoms
Electron transfer from two lithium atoms to an oxygen atom.

10. Energetics of Formation of Ionic Compound
7.2 Energetics of Formation of Ionic Compounds (SB p. 183)
Energetics of Formation of Ionic Compound
macroscopic level
 Hf ø
Na(s) ½Cl2(g)  NaCl(s)
Actually passing through many steps at the molecular level
microscopic level

11. Standard Enthalpy Change of Atomization (H atom) ø
7.2 Energetics of Formation of Ionic Compounds (SB p. 184)
Standard Enthalpy Change of Atomization (H atom) ø
The enthalpy change when one mole of gaseous atoms is formed from its elements in the defined physical state under standard conditions.
Na(s) Na(g) H atom [Na(s)] = kJ mol-1 ø
1/2 Cl2(g) Cl(g) H atom [1/2Cl2(g)] = +121 kJ mol-1 ø
Questions:
Why are the changes endothermic?
What type of bond is broken in each case?

12. Ionization Enthalpy (H I.E.) ø
7.2 Energetics of Formation of Ionic Compounds (SB p. 184)
Ionization Enthalpy (H I.E.) ø
The amount of energy required to remove one mole of valence electrons from one mole of atoms or ions in the gaseous state.
Na(g) Na+(g) + e H I.E [Na(g)] = +494 kJ mol-1 ø
Mg(g) Mg+(g) + e H I.E [Mg(g)] = +736 kJ mol-1 ø
Mg+(g) Mg2+(g) + e- H I.E [Na(g)] = kJ mol-1 ø
Questions:
Why are the changes endothermic?

13. Electron Affinity (H E.A.) ø
7.2 Energetics of Formation of Ionic Compounds (SB p. 185)
Electron Affinity (H E.A.) ø
The energy change when one mole of electrons is added to one mole of atoms or ions in the gaseous state.
First electron affinity
O(g) + e O-(g) H E.A [O(g)] = kJ mol-1
Second electron affinity
O-(g) + e O2-(g) H E.A [O(g)] = kJ mol-1 ø
Questions:
Why may E.A. have -ve or +ve values?

14. Lattice Enthalpy (H L.E.) ø
7.2 Energetics of Formation of Ionic Compounds (SB p. 185)
Lattice Enthalpy (H L.E.) ø
The energy change when one mole of an ionic crystal is formed from its constituent ions in the gaseous state under standard conditions
Na+ (g) + Cl-(g) NaCl(s) H lattice [Na+Cl-(s)] ø + –

15. ø Why can’t L.E. be determined directly from experiments? Questions:
7.2 Energetics of Formation of Ionic Compounds (SB p. 185)
Na+ (g) + Cl-(g) NaCl(s) H lattice [Na+Cl-(s)] ø
+ve or -ve?
L.E. can be determined indirectly by either: (1) calculations basing on the knowledge of electrostatics in Physics (assuming ions are point charges); or (2) calculations basing on Hess’s Law. + – + – + –
Questions:
Why can’t L.E. be determined directly from experiments?

16. Born-Haber Cycle for the formation of sodium chloride
7.2 Energetics of Formation of Ionic Compounds (SB p. 186)
Born-Haber Cycle for the formation of sodium chloride
Hatom[Na(s)]
HI.E.

17. 7.2 Energetics of Formation of Ionic Compounds (SB p. 187)

18. ø ø ø ø ø ø ø ø By Hess’s law, ΔHf [NaCl(s)]
7.2 Energetics of Formation of Ionic Compounds (SB p. 187)
By Hess’s law,
ΔHf [NaCl(s)]
= ΔHatom[Na(s)] + ΔHI.E.[Na(g)] + ΔHatom[Cl2(g)]
+ ΔHE.A.[Cl(g)] + ΔHlattice [NaCl(s)]
i.e. ΔHf [NaCl(s)]
= (-364) +ΔHlattice [NaCl(s)]
ΔHlattice [NaCl(s)]
= ΔHf [NaCl(s)] +[ (-364)]
= [ (-364)] = -711 kJ mol-1 ø ø ø ø ø ø ø ø

19. Stoichiometry of Ionic Compounds
7.3 Stoichiometry of Ionic Compounds (SB p. 189)
Stoichiometry of Ionic Compounds
Stoichiometry is the simplest ratio of the atoms bonded together in a compound.
How can the stoichiometry of an ionic compound be determined?

20. In Terms of Electronic Configuration
7.3 Stoichiometry of Ionic Compounds (SB p. 189)
In Terms of Electronic Configuration
Example
magnesium chloride
Elements involved
Mg (Group II) Cl (Group VII)
Ions formed
Mg Cl- 2 1
Ratio of ions
Chemical formula
Mg2+(Cl-) or MgCl2

21. Ionic Crystals Structure of Sodium Chloride Unit cell of NaCl
7.4 Ionic Crystals (SB p. 193)
Ionic Crystals
Unit cell of NaCl
Structure of Sodium Chloride
Co-ordination number of Na+ = 6
6:6 co-ordination
Co-ordination number of Cl- = 6

22. 7.4 Ionic Crystals (SB p. 193)
A unit cell is the smallest basic portion of the crystal lattice that, when repeatedly stacked together at various directions, can reproduce the entire crystal structure.

23. Ionic Crystals Question
7.4 Ionic Crystals (SB p. 193)
Ionic Crystals
corner(Cl-)
face(Cl-)
edge(Na+)
Question
Determine the number of Na+ and Cl- in a unit cell of sodium chloride respectively.

24. 7.4 Ionic Crystals (SB p. 193)
Diagram showing the two inter-penetrating face-centred cubic structure of Na+ and Cl- ions

25. Ionic Crystals Structure of Caesium Chloride (CsCl)
7.4 Ionic Crystals (SB p. 194)
Ionic Crystals
How to describe the structure?
Structure of Caesium Chloride (CsCl)
Co-ordination number of Cs+ = 8
8:8 co-ordination
Co-ordination number of Cl- = 8

26. 7.4 Ionic Crystals (SB p. 194)
The structure is actually two inter-penetrating simple cubic structure of Cs+ and Cl- ions

27. Some simple ionic structures
7.4 Ionic Crystals (SB p. 195)
Some simple ionic structures
Type of structure
Examples
Radius Ratio
(r+ : r-)*
Coordination
Sodium
chloride
Na+Cl-, Na+Br-,
K+Cl-, K+Br-
< 0.732
> 0.414
6:6
Caesium
Cs+Cl-,
Cs+Br-,
Cs+I-
> 0.732
8:8

28. Ionic Radii The technique of X-ray diffraction Photographic plate
7.5 Ionic Radii (SB p. 196)
Ionic Radii
X-ray
Photographic plate
The technique of X-ray diffraction

29. Electron density map Electron density map found by X-ray diffraction
7.5 Ionic Radii (SB p. 196)
Electron density map
Electron density map found by X-ray diffraction

30. Size of ion vs size of atom
7.5 Ionic Radii (SB p. 197)
Comparing relative atomic radii of some elements with the ionic radii of the corresponding ions.
Size of ion vs size of atom

31. Size of cation < size of atom
7.5 Ionic Radii (SB p. 197)
Size of cation < size of atom
Reasons: (1) The number of electron shell decreases (2) No. of protons > no. of electrons (p/e ratio increases) The nuclear attraction is more effective to cause a contraction in the electron cloud.
Size of anion > size of atom
Reasons: (1) Repulsion between newly added electron(s) with other electrons (2) No. of protons < no. of electrons (p/e ratio decreases) The nuclear attraction is less effective and there is an expansion of the electron cloud.

32. 7.5 Ionic Radii (SB p. 198)
Variation of ionic radii of the first 20 elements in the Periodic Table
isoelectronic ions
Why ionic radius decreases along the isoelectronic series?

33. Isoelectronic ions are ions with the same number of electrons.
7.5 Ionic Radii (SB p. 198)
Isoelectronic ions are ions with the same number of electrons.
The following are examples of isoelectronic series:
1. H-, Li+, Be2+, B N3-, O2-, F-, Na+, Mg2+, Al P3-, S2-, Cl-, K+, Ca2+

34. Reason
Isoelectronic ions have the same number of electrons. An increase in the number of protons implies an increase in the p/e ratio which leads to a contraction of the electron cloud.

35. The END