Group 13, 14, 15

[Lewis acids] -empty p orbital -electron poor -6 valence e- -> want to accept e- from Lewis base to make stable octet

monomeric trigonal planar more volatile than Al analogue B-X bond length increases F -> I (anion gets weaker)

B-X have 𝛑 character sp2 hybridised when “L” coordinates, the change in geometry (sp2 -> sp3) means 𝛑 interactions are now lost - endothermic reorganisation L->B bond formation = exothermic F = strongest stabiliser (smaller ∴ better overlap) ∴ bigger energy deficit on loss of B-F 𝛑 bond character

Al2M6 (dimer) volatile, reactive, liquid, flammable in air similar to Al2Cl6 - no l.p/filled orbitals to donate e- density into Al empty p orbital (lower stability)

CF4, CCl4 = inert SiCl4 = readily hydrolysed by water SiCl4 + 2H2O -> SiO2 + HCl Si = bigger than C - more prone to nucleophilic attack Sn, Pb - EX2 more stable than EX4 especially Pb as it’s at bottom of group - inert pair effect SnCl2 - reducing agent SnCl4 - liquid (freezes at -33°C - molecular rather than ionic) + tetrahedral less reactive than group I metals

[CO] -forms in limited O2 supply -binds to Fe in haemoglobin - prevents oxygen -transport from lungs to cells [CO2] -water soluble

PbO2 strong oxidising agent Pb(IV) = easily reduced to Pb(II) **Tin behaves similarly

lone pair lewis base e- rich

N2O, NO, N2O3, NO2, N2O4, N2O5

[oxidation states = +3, +5] P4O6 - P(III) -acidic -P4O6 + H2O -> 4H3PO3 P4O10 -acidic -P4O10 + 6H2O -> 4H3PO4

[protons from OH = source of acidity - low pKa; easily released from oxygen] H3PO3 = phosphonic acid -diprotic -tetrahedral H3PO4 = phosphoric acid -triprotic -tetrahedral

[NF3] -thermodynamically stable gas -no reaction with water, acid or bases -less basic than NH3 [NCl3] -less polarisation - more reactive -NCl3 + 3H2O -> NH3 + 3HOCl [NBr3 + NI3] -weaker bonds -more reactive -diffuse halide, poor orbital overlap

[PF3] -gas -hydrolysed in water -reacts faster in alkali -> PF3 + 3OH- -> H3PO3 + 3F- -binds to haemoglobin - toxic [PCl3] -PCl3 + H2O -> H3PO3 + 3HCl

[NF5] -not known -N can’t achieve +5 oxidation state with halides [PF5 + AsF5] -gas [SbF5] -liquid -F bridges even in gas phase (Sb > As) increased coordination number

NH3 - hydrogen bonding raises b.p. [PH3 -> BiH3] -Van der Waals forces only -weakening E-H bond strength

incompatible overlap of E and H1s orbitals decreases bond strength leads to increased acidity poorer overlap of E-H orbitals decreases down group results in long bonds less e- repulsion compared to N-H

[decreases] -less s-p mixing -lone pairs held in orbitals close to nucleus -s-type = not available for bonding -heavier EH3 are less basic (unable to abstract H+) -PH3 = more likely to release H+ -NH3 = less likely to release H+