Lecture 2711/07/05

Ionic bond Ionic compounds Valence electrons are transferred from one atom to another Metal + non-metal NaCl Bonding

Na (s) + ½ Cl 2 (g)  NaCl (s) ΔH˚ = kJ/mol [Na + Cl - ] Ca (s) + ½ O 2 (g)  CaO (s) ΔH˚ = kJ/mol [Ca +2 O -2 ] Ionic compounds

Na(g)  Na + (g) + e - ΔE = I.E. = kJ/mol Cl (g) + e -  Cl - (g) ΔE = E.A. = -349 kJ/mol Na + (g) + Cl - (g)  [Na + Cl - ] (g) ΔE ion pair = -498 kJ/mol ΔE net = -351 kJ/mol

Energy of formation C = constant n + = number of positive charges of ion n - = number of negative charges of ion d = distance between the ion centers Energy of attraction between ions dependent on: Magnitude of the ion charges: Higher the ion charges  greater the attraction  ΔE larger negative value Distance between the ions: Larger the atomic size  larger the d  smaller the attractive force  ΔE is less negative

Lattice Energy ΔE lattice = energy of formation of 1 mole of a solid crystalline ionic compound when ions in the gas phase combine Cannot be measured directly Calculated value

ΔH˚ f = ΔH 1a + ΔH 1b + ΔH 2a + ΔH 2b + ΔH 3 Born-Haber Cycle: Used to calculate lattice energy ΔH 1a = kJ/mol ΔH 1b = kJ/mol ΔH 2a = kJ/mol ΔH 2b = +496 kJ/mol ΔH˚ f = kJ/mol ΔH 3 = ΔH lattice = -786 kJ/mol

Covalent Bonds Sharing of electrons Two non-metals Bond is a balance of attractive and repulsive forces Bond length is from lowest potential energy Too close – repulsion Too far – no effect

Valence electrons are distributed as: shared or bonding electrons or bond pairs unshared or non-bonding electrons or lone pairs HCl lone pair (LP) shared or bond pair Lewis Dot Structures

H 2 CO SO 3 C2F4C2F4C2F4C2F4 Double and triple bonds 1 line for each bond Commonly observed for C, N, P, O, and S

Octet Rule Normally the bonding electrons + the lone pair electrons = 8 There are a few exceptions H Predict # of bonds in molecule [(# electrons to fulfill octets of all atoms in molecule) – (# valence electrons in all atoms in molecule)] / 2

Count valence electrons H = 1 and N = 5 Total = (3 x 1) + 5 = 8 electrons  4 pairs Decide on the central atom Normally atom of lowest affinity for electrons Never H N is central Building a Lewis Dot Structure Ammonia, NH3 Form a single bond between the central atom and each surrounding atom H H H N H H H N Remaining electrons form lone pairs to complete octet Use lone pairs to create double or triple if needed to complete octet

Sulfite ion, SO 3 2-