Oxides of Sulphur & Allotropes of Sulphur

Why Sulphur Shows Allotropy

Allotropy is the existence of an element in two or more physical forms that differ in the arrangement of their atoms. Across Group 16 every member exhibits allotropy, but sulphur is the richest case: it forms numerous allotropes of which the yellow rhombic (α-sulphur) and monoclinic (β-sulphur) crystalline forms are the most important. The reason is sulphur's strong tendency to catenate — to form S–S bonds — which lets it close into puckered $\ce{S8}$ rings that can then pack in more than one way.

The two crystalline forms are not chemically different substances; both are built from identical $\ce{S8}$ molecules. They differ only in how those rings stack in the crystal lattice, which is exactly why they can interconvert simply by changing the temperature. Beyond the two crystalline solids, sulphur also gives non-crystalline (amorphous) and chain-like (plastic) forms, and in the vapour it fragments down to the small, oxygen-like $\ce{S2}$ molecule.

Rhombic vs Monoclinic Sulphur

The single most examined fact here is the transition temperature, 369 K (96.6 °C). Below it the rhombic form is the more stable allotrope; above it the monoclinic form is more stable. At exactly 369 K both forms coexist in equilibrium, which is the operational definition of a transition temperature.

$$\ce{Rhombic\ sulphur\ (\alpha)\ <=>[\ 369\ K\ ]\ Monoclinic\ sulphur\ (\beta)}$$

Rhombic sulphur is the form normally encountered at room temperature. Its crystals grow when a solution of roll sulphur in carbon disulphide ($\ce{CS2}$) is allowed to evaporate slowly. Monoclinic sulphur is made by melting rhombic sulphur in a dish, cooling until a crust forms, piercing two holes in the crust and pouring out the still-liquid interior — colourless, needle-shaped β-crystals are left behind. The comparison table below collects the data points NEET asks for.

The S8 Crown Ring

Both crystalline allotropes contain the eight-membered $\ce{S8}$ ring. The ring is not flat: it is puckered into a crown shape, in which alternate sulphur atoms lie in two parallel planes. The schematic below shows the staggered, zig-zag profile of the crown that distinguishes it from a planar octagon.

The crown geometry — with an S–S bond length of roughly 205 pm and an S–S–S angle near 107° — is energetically favoured because it lets all eight bond angles relax close to the tetrahedral value while keeping every sulphur two-coordinate. In addition to $\ce{S8}$, several ring modifications containing 6 to 20 sulphur atoms have been synthesised; in cyclo-$\ce{S6}$ the ring adopts a chair form rather than a crown.

Plastic Sulphur & Paramagnetic S2

Heating sulphur reveals the chemistry behind its allotropy. Rhombic and monoclinic sulphur both melt to a mobile yellow liquid. As the temperature climbs, the $\ce{S8}$ rings begin to open and link into long chains; the liquid darkens and becomes highly viscous, reaching maximum chain length and maximum viscosity around 470 K (≈200 °C). Above this the chains break down, $\ce{S8}$ rings re-form and the liquid thins again before boiling near 718 K.

If this near-boiling, chain-rich liquid is poured into cold water, the long sulphur chains are frozen in a random tangle, giving plastic sulphur — a rubbery, amorphous solid. It is metastable: on standing the chains snap and re-close into $\ce{S8}$ rings, so plastic sulphur slowly reverts to the stable rhombic form.

Sulphur Dioxide — Preparation

Burning elemental sulphur in air or oxygen gives sulphur dioxide together with a little (6–8%) sulphur trioxide. This is the most direct route and the entry point of the Contact Process for sulphuric acid.

$$\ce{S(s) + O2(g) -> SO2(g)}$$

In the laboratory $\ce{SO2}$ is conveniently generated by treating a sulphite with dilute sulphuric acid; alternatively, hot concentrated sulphuric acid is reduced by copper turnings. Industrially it arises as a by-product of roasting sulphide ores such as iron pyrites and zinc blende.

Sulphur dioxide is a colourless gas with a pungent, choking smell (the odour of burning sulphur). It is highly soluble in water, liquefies at room temperature under about 2 atm pressure, and boils at 263 K. After drying it is liquefied under pressure and stored in steel cylinders; liquid $\ce{SO2}$ is itself a useful non-aqueous solvent.

Structure of SO2

The $\ce{SO2}$ molecule is angular (bent). The central sulphur carries one lone pair, so the three electron domains (two bonds + one lone pair) adopt a trigonal-planar arrangement; lone-pair repulsion compresses the O–S–O angle to about 119°. Crucially, $\ce{SO2}$ is a resonance hybrid of two equivalent canonical forms, so the two S–O bonds are identical in length and order — each intermediate between a single and a double bond.

Properties of SO2

Sulphur dioxide is an acidic oxide and, in the presence of moisture, a reducing agent and a bleaching agent. As an acidic oxide it dissolves in water to give sulphurous acid and reacts with alkalis to give sulphites and hydrogen sulphites.

$$\ce{2NaOH + SO2 -> Na2SO3 + H2O}$$ $$\ce{Na2SO3 + H2O + SO2 -> 2NaHSO3}$$

Moist $\ce{SO2}$ is a fairly strong reducing agent because the $\ce{S(IV)}$ centre is readily oxidised to $\ce{S(VI)}$. It reduces iron(III) to iron(II) and, most importantly for the laboratory, decolourises acidified potassium permanganate — the standard test for the gas.

Its bleaching action is also reductive: in the presence of moisture the nascent hydrogen produced reduces and removes the colour of the substance. Because the bleaching works by reduction and not by oxidation, it is temporary — coloured material slowly regains its colour on standing in air as it is re-oxidised. This is why $\ce{SO2}$ is used only for delicate articles such as wool, silk and straw.

Uses of $\ce{SO2}$ follow directly from these properties: refining petroleum and sugar, bleaching wool and silk, acting as an anti-chlor, a disinfectant and a food preservative, and serving as the feedstock for sulphuric acid, sodium hydrogen sulphite and calcium hydrogen sulphite. In its reactions with water and alkalis, $\ce{SO2}$ behaves much like carbon dioxide.

Sulphurous Acid & SO3

Passing $\ce{SO2}$ through water gives an aqueous solution of sulphurous acid, $\ce{H2SO3}$ — a weak, dibasic and unstable acid that exists only in solution and cannot be isolated as the pure compound. The equilibrium lies well to the left.

$$\ce{SO2(g) + H2O(l) <=> H2SO3(aq)}$$

Catalytic oxidation of $\ce{SO2}$ by atmospheric oxygen over vanadium(V) oxide at about 720 K gives sulphur trioxide, $\ce{SO3}$ — the key step in the manufacture of sulphuric acid. The reaction is exothermic and reversible, so moderate temperature and pressure are used to balance yield against rate.

$$\ce{2SO2(g) + O2(g) ->[V2O5][720\ K] 2SO3(g)} \qquad \Delta_r H = -196.6\ \text{kJ mol}^{-1}$$

In gaseous $\ce{SO3}$ the sulphur is $sp^2$ hybridised and the molecule is trigonal planar, with three equivalent S–O bonds at 120° and the molecule a resonance hybrid. $\ce{SO3}$ is a strongly acidic oxide; it is absorbed into concentrated $\ce{H2SO4}$ to give oleum ($\ce{H2S2O7}$) rather than being added straight to water, because direct hydration produces a dense, corrosive acid mist.

$$\ce{SO3 + H2SO4 -> H2S2O7}\ (\text{oleum})$$

The downstream conversion of oleum and $\ce{SO3}$ into sulphuric acid is treated in the dedicated sibling note; the takeaway here is the trend across the two oxides — $\ce{SO2}$ (S in +4) is the reducing, acidic, bent oxide, while $\ce{SO3}$ (S in +6) is the strongly acidic, planar oxide that anchors industrial $\ce{H2SO4}$.