Silicon Compounds — Silicates, Silicones, Zeolites

Silicon Dioxide (Silica)

Silicon is the second most abundant element in the earth's crust (about 27.7% by mass), occurring almost entirely as silica and silicates. Silicon dioxide, $\ce{SiO2}$, commonly called silica, is found in several crystallographic forms — quartz, cristobalite and tridymite — which are interconvertible at suitable temperatures.

Structurally, $\ce{SiO2}$ is a covalent, three-dimensional network solid. Each silicon atom is bonded tetrahedrally to four oxygen atoms, and each oxygen atom is in turn bonded to two silicon atoms, so every corner of a tetrahedron is shared with a neighbouring tetrahedron. The result is a giant molecule: the entire crystal can be viewed as one continuous covalent lattice in which eight-membered rings of alternating silicon and oxygen atoms repeat through space.

Because of the very high $\ce{Si-O}$ bond enthalpy, silica in its ordinary form is almost non-reactive. It resists attack by halogens, dihydrogen, and most acids and metals even at elevated temperatures. The two reagents that do attack it are hydrofluoric acid and molten alkali:

Quartz, an ordered crystalline form of silica, is extensively used as a piezoelectric material: it has made possible extremely accurate clocks, modern radio and television broadcasting, and mobile radio communication. Silica gel serves as a drying agent and as a support for chromatographic materials and catalysts, while kieselguhr (an amorphous form of silica) is used in filtration plants.

Why SiO₂ Is a Solid but CO₂ a Gas

Carbon and silicon belong to the same group and both form a dioxide of formula $\ce{MO2}$, yet their physical states could not be more different. $\ce{CO2}$ is a gas that solidifies (as dry ice) only at low temperature, whereas $\ce{SiO2}$ is a hard, high-melting solid. The explanation lies entirely in bonding, not in formula.

Carbon is small and has its valence orbitals close to the nucleus, so it readily forms strong $p\pi$–$p\pi$ multiple bonds with oxygen. Carbon dioxide therefore exists as discrete, linear $\ce{O=C=O}$ molecules in which carbon is $sp$ hybridised; both $\ce{C=O}$ bonds are of equal length (115 pm) and the molecule has no dipole moment. These small molecules are attracted to one another only by weak van der Waals forces, which are easily overcome — hence $\ce{CO2}$ is a gas.

Silicon, in contrast, is larger; its valence orbitals are more diffuse and do not overlap effectively with the small, compact 2p orbitals of oxygen. It therefore cannot form $p\pi$–$p\pi$ double bonds with oxygen. To satisfy its valence, each silicon instead forms four single $\ce{Si-O}$ bonds, and every oxygen bridges two silicons, propagating the network endlessly in three dimensions. Melting $\ce{SiO2}$ requires rupturing strong covalent bonds throughout the lattice, so it is a hard solid with a very high melting point.

The SiO₄ Tetrahedron Building Block

Every silicate structure — and silica itself — is built from one repeating unit: the $\ce{SiO4^4-}$ tetrahedron. A silicon atom sits at the centre, bonded to four oxygen atoms placed at the corners of a tetrahedron. The silicon uses $sp^3$ hybrid orbitals that overlap with 2p orbitals of the oxygens. In the discrete orthosilicate anion all four oxygen atoms carry negative charge, giving the formula $\ce{SiO4^4-}$.

The richness of silicate chemistry comes from how these tetrahedra connect. They may stay as discrete units, or join through corners by sharing one, two, three or all four oxygen atoms. When all four corners of every tetrahedron are shared, the structure becomes a three-dimensional network — exactly the situation in pure silica, where the framework is electrically neutral because no terminal oxygens remain to carry charge.

Silicates — Linking the Tetrahedra

A large number of silicate minerals exist in nature; familiar examples are feldspar, zeolites, mica and asbestos. In every one, the $\ce{SiO4^4-}$ tetrahedron is the basic structural unit. Either a discrete unit is present, or several units are joined via corners by sharing 1, 2, 3 or 4 oxygen atoms per tetrahedron. The way they link determines the macroscopic structure and properties.

Whenever a tetrahedron does not share all four oxygens, the unshared oxygen atoms retain a negative charge. The overall negative charge on the silicate framework is then neutralised by positively charged metal cations sitting between the tetrahedra. Two important man-made silicates worth remembering for direct recall are glass and cement.

Silicones — Organosilicon Polymers

Silicones are a group of organosilicon polymers that have $\ce{(R2SiO)}$ as the repeating unit, where R is an alkyl or aryl group; the polymer is written $\ce{(R2SiO)_n}$. The backbone is an alternating chain of silicon and oxygen atoms, with the alkyl or aryl groups attached to each silicon by covalent bonds.

The starting materials are alkyl- or aryl-substituted silicon chlorides, $\ce{R_{n}SiCl_{(4-n)}}$. When methyl chloride reacts with silicon in the presence of a copper catalyst at about 573 K, a mixture of methyl-substituted chlorosilanes — $\ce{MeSiCl3}$, $\ce{Me2SiCl2}$ and $\ce{Me3SiCl}$, with a little $\ce{Me4Si}$ — is formed. Hydrolysis of dimethyldichlorosilane, $\ce{(CH3)2SiCl2}$, followed by condensation polymerisation, yields straight-chain polymers:

The chain length is controlled by adding $\ce{(CH3)3SiCl}$, which caps the ends and stops further growth; the trifunctional $\ce{RSiCl3}$, by contrast, introduces cross-linking and yields two-dimensional or network structures. Because the silicon-oxygen chain is surrounded by non-polar alkyl groups, silicones present a hydrophobic surface and are water-repellent. They have high thermal stability, high dielectric strength, and resistance to oxidation and chemicals.

Zeolites — Aluminosilicate Frameworks

A zeolite is created from the three-dimensional network of silicon dioxide by a simple substitution: if a few silicon atoms are replaced by aluminium atoms, the overall structure — now called an aluminosilicate — acquires a negative charge. Cations such as $\ce{Na+}$, $\ce{K+}$ or $\ce{Ca^2+}$ balance that charge, sitting loosely in the cavities of the framework. Feldspar and zeolites are the standard examples.

The substitution matters because $\ce{Al}$ has one fewer valence electron than $\ce{Si}$. Replacing a neutral $\ce{SiO2}$ framework site with aluminium leaves a deficiency that must be neutralised by an external cation. Those external cations are only weakly held, which is the structural origin of both major applications of zeolites.

Hydrated zeolites are used as ion-exchangers in the softening of hard water, and they serve as catalysts in petrochemical industries for the cracking of hydrocarbons and for isomerisation. The most quoted single example is ZSM-5, a type of zeolite used to convert alcohols directly into gasoline — a fact that has appeared verbatim in NEET.

Silica · Silicate · Silicone · Zeolite

Because these four classes share the silicon-oxygen motif yet differ sharply in structure and use, NEET frequently tests them as a contrast set. The following table consolidates the bonding, structure and applications of each.