1. Ionic bonds and main group chemistry

2. Towards the noble gas configuration  Noble gases are unreactive – they have filled shells  Shells of reactive elements are unfilled  Achieve noble gas configuration by gaining or losing electrons Metals lose electrons – form positive ions Metals lose electrons – form positive ions Nonmetals gain electrons – form negative ions Nonmetals gain electrons – form negative ions

3. Lewis dot model  The nucleus and all of the core electrons are represented by the element symbol  The valence electrons are represented by dots – one for each  Number of dots in Lewis model is equal to group number (in 1 – 8 system)

4. The Octet Rule  All elements strive to become a noble gas, at least as far as the electrons are concerned.  Filling the outer shell – 8 electrons  Achieve this by adding electrons  Or taking them away

5. Predicting ion charges  s and p block elements are easy: charge = group number for cations charge = group number for cations charge = -(8 – group number) for anions charge = -(8 – group number) for anions

6. Less predictable for transition metals  Occurrence of variable ionic charge Cr 2+, Cr 3+, Cr 4+, Cr 6+ etc. Cr 2+, Cr 3+, Cr 4+, Cr 6+ etc.  4s electrons are lost first and then the 3d  Desirable configurations coincide with empty, half-filled or filled 3d orbitals Fe 2+ ([Ar]3d 6 ) is less stable than Fe 3+ ([Ar]3d 5 ) Fe 2+ ([Ar]3d 6 ) is less stable than Fe 3+ ([Ar]3d 5 )

7. Ionic size and charge  Loss of electrons increases the effective nuclear charge – ion shrinks  Gain of electrons decreases the effective nuclear charge – ion expands

8. Ionization energy  Energy required to remove an electron from a neutral gaseous atom  Always positive  Follows periodic trend Increases across period Increases across period Decreases down group Decreases down group  Removal of electrons from filled or half-filled shells is not as favourable [He]2s 2 [He]2s 2 2p 3 [He]2s 2 2p 4 [He]2s 2 2p 1

9. Higher ionization energies  Depend on group number  Much harder to remove electrons from a filled shell  Stepwise trend below illustrates this Partially filled – valence electrons Completely filled – core electrons

10. Electron affinity  Energy released on adding an electron to a neutral gaseous atom  Values are either negative – energy released, meaning negative ion formation is favourable negative – energy released, meaning negative ion formation is favourable Or zero – meaning can’t be measured and negative ions are not formed Or zero – meaning can’t be measured and negative ions are not formed  Addition of electrons to filled or half-filled shells is not favoured (e.g. He, N)  It is easier to add an electron to Na (3s 1 ) than to Mg (3s 2 )

11. Ionic bonding  Reaction between elements that form positive and negative ions Metals (positive ions) and nonmetals (negative ions) Metals (positive ions) and nonmetals (negative ions)  Neutral Na + Cl → ionic Na + Cl - [Ne]3s 1 + [Ne]3s 2 3p 5 = [Ne] + + [Ar] - [Ne]3s 1 + [Ne]3s 2 3p 5 = [Ne] + + [Ar] -

12. Stability of the ionic lattice  Simply forming ions does not give an energy payout: E i (Na) = 496 kJ/mol E i (Na) = 496 kJ/mol E a (Cl) = -349 kJ/mol E a (Cl) = -349 kJ/mol  Net energy investment required  Formation of a crystal lattice releases energy  The lattice energy is the energy released on bringing ions from the gas phase into the solid lattice  Depends on coulombic attraction between ions -U = κz 1 z 2 /d (κ = 8.99x10 9 JmC -2

13. Born-Haber cycle for calculating energy  The lattice energy can be obtained using other experimentally determined quantities and the energy cycle

14. Lattice energies follow simple trends  As ionic charge increases, U increases (U  z 1 z 2 )  As ion size decreases, U increases (U  1/d) U(LiF) > U(LiCl) > U(LiBr) U(LiF) > U(LiCl) > U(LiBr) U(NaI) < U(MgI 2 ) < U(AlI 3 ) U(NaI) < U(MgI 2 ) < U(AlI 3 )

15. The Octet Rule  Main-group elements undergo reactions which leave them with eight valence electrons Group 1 (ns 1 ) M + Group 1 (ns 1 ) M + Group 2 (ns 2 ) M 2+ Group 2 (ns 2 ) M 2+ Group 6 (ns 2 np 4 ) X 2- Group 6 (ns 2 np 4 ) X 2- Group 7 (ns 2 np 5 ) X - Group 7 (ns 2 np 5 ) X -  Works very well for second row (Li – F)  Many violations in heavier p-block elements (Pb 2+, Tl +, Sb 3+ )