Boron, aluminium, carbon, silicon and the inert pair

Both are in group 14: in the laboratory $\ce{SnCl2}$ is used as a reducing agent because tin(II) readily becomes tin(IV), while the positive plate of a lead–acid battery is $\ce{PbO2}$ because lead(IV) is eager to become lead(II) again. Two neighbours in one group prefer opposite oxidation states. This lesson travels from electron-poor boron to oxygen-bound silicon, and ends by reading the “lower state down the group” pattern straight from electrode potentials.

Boron has only three valence electrons. In trigonal-planar $\ce{BF3}$ the boron has six electrons around it and an empty p orbital, which makes $\ce{BF3}$ a textbook Lewis acid: $\ce{BF3} + \ce{NH3} \mathrel{\ce{->}} \ce{F3B-NH3}$ and $\ce{BF3} + \ce{F-} \mathrel{\ce{->}} \ce{BF4-}$.

Boric acid $\ce{B(OH)3}$ is acidic not by giving up a proton but by accepting $\ce{OH-}$ from water: $\ce{B(OH)3} + \ce{H2O} \mathrel{\ce{<=>}} \ce{B(OH)4-} + \ce{H+}$, $\mathrm{p}K_\mathrm{a} = 9.24$ — so weak that it is used in eyewash.

The simplest borane, $\ce{B2H6}$, is even shorter of electrons: each B–H–B bridge holds three atoms together with just two electrons (a three-centre two-electron bond).

With $E^\circ(\ce{Al^3+/Al}) = -1.662$ V aluminium is a very reactive metal, yet aluminium pans and window frames last for decades: a fresh surface instantly grows a dense, tightly bound film of $\ce{Al2O3}$ only a few nanometres thick that seals the metal off from air and water (passivation). Anodising thickens and dyes this film.

Aluminium hydroxide dissolves in both acids and bases — it is amphoteric:

Aluminium’s affinity for oxygen is enormous: the thermite reaction $\ce{2Al} + \ce{Fe2O3} \mathrel{\ce{->}} \ce{Al2O3} + \ce{2Fe}$ releases 852 kJ, enough to melt the iron it makes. Because $\ce{Al2O3}$ is so stable, aluminium can only be won from its oxide by electrolysis.

• Diamond: every carbon is sp³ and bonded to four others (C–C 154 pm) in a three-dimensional network — the hardest natural substance, an electrical insulator, density 3.51 g/cm³.
• Graphite: flat layers of hexagonal rings of sp² carbon (C–C 142 pm within a layer) stacked 335 pm apart and held only by van der Waals forces. Delocalised π electrons let it conduct along the layers, and the layers slide, so it lubricates; density 2.26 g/cm³.
• Graphene: a single sheet of graphite, first isolated in 2004 (Nobel Prize in Physics, 2010).
• Fullerenes: cage molecules such as $\ce{C60}$, discovered in 1985 (Nobel Prize in Chemistry, 1996); rolled-up graphite sheets give carbon nanotubes.

At 25 °C and 1 bar graphite is the stable form: $\ce{C(graphite)} \mathrel{\ce{->}} \ce{C(diamond)}$ has $\Delta G^\circ = +2.9$ kJ/mol. Diamond is merely metastable — converting it means breaking very strong C–C bonds, so the rate is effectively zero.

CO and CO₂. Carbon monoxide is colourless, odourless and deadly: it binds haemoglobin more than two hundred times more strongly than oxygen does. Carbon dioxide is a greenhouse gas — see how it absorbs infrared in CO₂ and infrared light.

Why is CO₂ a gas while SiO₂ is quartz, melting at 1710 °C? Carbon is small, so the 2p–2p π overlap in C=O double bonds is good and one carbon is satisfied by two double bonds: a small molecule results. Silicon is large, its 3p orbitals overlap poorly with oxygen’s 2p, and Si=O double bonds do not pay; silicon prefers four single bonds to four oxygens, each of which bonds on to another silicon — an endless network of $\ce{SiO4}$ tetrahedra. All the silicates of the Earth’s crust are built from $\ce{SiO4}$ tetrahedra: isolated, in chains, in sheets (micas, clays) or in three-dimensional frameworks (quartz, feldspars, zeolites).

Elemental silicon also has the diamond structure, but its band gap is only 1.12 eV: it is a semiconductor. A trace of boron makes p-type silicon, a trace of phosphorus n-type, and joining the two gives diodes and transistors.

The highest oxidation state in groups 13, 14 and 15 is the group number minus ten (+3, +4, +5), but down each group the state two lower (+1, +2, +3) becomes steadily more stable: thallium’s common chloride is TlCl rather than $\ce{TlCl3}$; $\ce{PbCl4}$ is a yellow oil that decomposes into $\ce{PbCl2}$ and $\ce{Cl2}$ on gentle warming; bismuth(V), as sodium bismuthate, is a ferocious oxidant. It is as if the outer ns² pair turned “inert” and declined to bond.

The reason is an energy balance with two sides. First, 6s electrons are held unusually tightly — the 4f and 5d electrons filled in between shield poorly, and relativistic effects contract the 6s orbital (thallium’s second and third ionisation energies together are about 320 kJ/mol larger than indium’s). Second, bonds to bigger atoms are weaker, so forming two extra bonds no longer releases enough energy to pay for using the s electrons.