Silicon chemistry and silicones

Silicon is in Group 4, the same group as carbon, so its atoms also form four covalent bonds. Most of the silicon on Earth is locked up as silicon dioxide (sand and quartz), where every silicon atom is bonded to four oxygen atoms in a giant covalent structure. That is why sand has such a high melting point.

Carbon can form long chains with itself, which is the basis of all the alkanes and polymers you study. Silicon cannot do this well: Si–Si bonds are weak and break easily. Instead, chemists make chains where silicon and oxygen atoms alternate, with small carbon groups attached to the silicon. These polymers are called silicones.

Silicones have a useful mix of properties:

• they stay soft and flexible over a wide range of temperatures;
• they repel water;
• they are unreactive and non-toxic, so they are safe in kitchens and in medicine.

So silicone is used for oven gloves, cake moulds, waterproof sealant around baths, and medical tubing. Do not confuse silicon (the element, used in computer chips) with silicone (the polymer).

Silicon and carbon are both in Group 14, but three differences change their chemistry.

Bond strengths. The Si–Si bond (about 226 kJ mol^{-1}) is much weaker than C–C (347 kJ mol^{-1}), while Si–O is stronger than C–O. Silicon chemistry therefore drifts towards Si–O bonds whenever oxygen or water is available. Silane, SiH_4, catches fire on contact with air; methane needs a spark.

Size and electronegativity. Silicon is larger and less electronegative than carbon (1.9 against 2.6), so in an Si–Cl or Si–C bond the silicon carries a sizeable \delta+ and there is room around it for a nucleophile to approach. Tetrachloromethane sits unchanged in water, but tetrachlorosilane fumes in moist air as it hydrolyses to silica and HCl.

Making silicones. Industrial silicone chemistry starts from the direct process, where chloromethane is passed over hot silicon with a copper catalyst to give mainly dichlorodimethylsilane, Me_2SiCl_2. Water then replaces both chlorines:

The silanediol molecules condense with one another, losing water each time an Si–O–Si link forms. This is condensation polymerisation:

Chemists control the product by mixing in other chlorosilanes. Me_3SiCl has only one reactive Si–Cl, so it caps the end of a chain and keeps it short (silicone oils). MeSiCl_3 has three, so it creates branch points and cross-links, giving rubbers and resins.

Why silicon behaves differently

Si–C bonds (about 1.87 Å) are longer than C–C bonds, and silicon is electropositive relative to carbon, so the C–Si bond is polarised towards carbon. Two consequences dominate synthetic use. First, silicon is attacked by hard nucleophiles, especially O and F, through a five-coordinate silicate intermediate or transition state. Second, the Si–F bond (roughly 540 kJ mol^{-1} for bond dissociation) is even stronger than Si–O, so fluoride is the reagent of choice for pulling silicon off oxygen.

Silyl ethers as protecting groups

An alcohol reacts with a chlorosilane and a mild base (imidazole in DMF is the classic recipe) to give a silyl ether. Bulk around silicon sets the stability: TMS ethers fall off during aqueous work-up, TBS (tert-butyldimethylsilyl) ethers survive chromatography and most bases, and TIPS and TBDPS are tougher still towards acid. This graded stability lets a chemist protect a primary alcohol with TBS, a secondary one with TIPS, and remove them one at a time. Tetrabutylammonium fluoride (TBAF) in THF cleaves almost all of them, driven by Si–F formation.

Peterson olefination

An α-silyl carbanion such as Me_3SiCH_2Li adds to an aldehyde or ketone to give a β-hydroxysilane. Treat that with base (KH) and the alkoxide attacks silicon intramolecularly, so elimination is syn. Treat it with acid and the protonated OH leaves while the C–Si bond lines up anti, so elimination is anti. One diastereomer of the β-hydroxysilane can therefore be steered to either the E or the Z alkene, something the Wittig reaction cannot offer, and the by-product is a volatile siloxane.

Brook rearrangement

When an alkoxide sits next to a carbon carrying a silyl group, the oxygen attacks silicon and the silyl group migrates from carbon to oxygen, leaving a carbanion. The new Si–O bond pays for the loss of the C–Si bond and the formation of a less stable anion. The reverse (retro-Brook) wins when the carbanion would be very unstable, for example an unstabilised aryl anion.

Sol-gel silica and hydrosilylation

Tetraethyl orthosilicate, Si(OEt)_4, hydrolyses and condenses in ethanol/water with acid or base catalysis to build an Si–O–Si network at room temperature. Gentle drying gives coatings, monoliths or (via the Stöber method) monodisperse silica spheres. Addition-cure silicone rubbers rely on hydrosilylation: an Si–H group adds across a vinyl group on another chain, catalysed by platinum complexes such as Karstedt's catalyst.

Silicone manufacture runs at the scale of millions of tonnes a year, and the platinum used to cure silicone rubbers by hydrosilylation is a significant cost because it stays in the product. Much current work therefore targets base-metal catalysts (iron, cobalt and nickel complexes with redox-active or pincer ligands) that match platinum's anti-Markovnikov selectivity without isomerising the alkene. Mechanistic studies separate the Chalk–Harrod and modified Chalk–Harrod pathways for platinum from σ-bond metathesis and radical routes for earlier metals, and copper hydrides give a non-redox route for hydrosilylating C=O and C=N bonds enantioselectively.

In medicinal chemistry, the carbon–silicon switch replaces a quaternary carbon with silicon to tune lipophilicity, metabolism and conformation; sila-analogues of haloperidol, for example, avoid a toxic pyridinium metabolite. Silanols are also explored as hydrogen-bond donors in place of alcohols.

Fundamental chemistry keeps finding new low-coordinate silicon. Stable disilenes (Si=Si) were first isolated in 1981, and isolable silylenes, silylium ions and silicon-centred radicals now serve as ligands, Lewis acids and even catalysts in their own right. Fluoride-free Hiyama–Denmark couplings use silanolates so that silyl protecting groups elsewhere survive.

Two environmental questions are open. Cyclic siloxanes such as D4, D5 and D6 are persistent and have been restricted in EU consumer products, which pushes formulators to alternatives. And because silicones do not biodegrade, chemical recycling (catalytic depolymerisation back to cyclic monomers or to alkoxysilanes) is an active area aiming to close the loop on silicone waste.