Ionic bonds and main group chemistry

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 • 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 = -(8 – group number) for anions

6. Less predictable for transition metals • Occurrence of variable ionic charge • Cr2+, Cr3+, Cr4+, Cr6+ etc. • 4s electrons are lost first and then the 3d • Desirable configurations coincide with empty, half-filled or filled 3d orbitals • Fe2+ ([Ar]3d6) is less stable than Fe3+ ([Ar]3d5)

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. [He]2s22p3 [He]2s22p4 [He]2s2 [He]2s22p1 Ionization energy • Energy required to remove an electron from a neutral gaseous atom • Always positive • Follows periodic trend • Increases across period • Decreases down group • Removal of electrons from filled or half-filled shells is not as favourable

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 • 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 (3s1) than to Mg (3s2)

11. Ionic bonding • Reaction between elements that form positive and negative ions • Metals (positive ions) and nonmetals (negative ions) • Neutral Na + Cl → ionic Na+Cl- • [Ne]3s1 + [Ne]3s23p5 = [Ne]+ + [Ar]-

12. Stability of the ionic lattice • Simply forming ions does not give an energy payout: • Ei(Na) = 496 kJ/mol • Ea(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 =κz1z2/d (κ = 8.99x109 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  z1z2) • As ion size decreases, U increases (U  1/d) • U(LiF) > U(LiCl) > U(LiBr) • U(NaI) < U(MgI2) < U(AlI3)

15. The Octet Rule • Main-group elements undergo reactions which leave them with eight valence electrons • Group 1 (ns1) M+ • Group 2 (ns2) M2+ • Group 6 (ns2np4) X2- • Group 7 (ns2np5) X- • Works very well for second row (Li – F) • Many violations in heavier p-block elements (Pb2+, Tl+, Sb3+)