Oxygen and sulfur

Pour liquid oxygen between the poles of a strong magnet and the pale-blue liquid hangs there, bridging the gap; liquid nitrogen simply runs through. Oxygen molecules are attracted by the magnet — yet in the O=O Lewis structure every electron is paired, so there should be no magnetism at all.

The O=O bond is 121 pm long with a bond enthalpy of 498 kJ/mol. Oxygen is the most electronegative element after fluorine, and dioxygen is a strong oxidant, $E^\circ(\ce{O2/H2O}) = +1.229$ V; luckily most of its reactions are slow — otherwise the organic matter of our bodies would long since have burned.

In the stratosphere, ultraviolet light shorter than about 240 nm splits $\ce{O2}$ into atoms, which join $\ce{O2}$ to form ozone, $\ce{O3}$. Ozone is bent (116.8°, with two equal O–O bonds of 127.8 pm, between single and double), and it absorbs strongly from 200 to 310 nm (peak near 255 nm), shielding the ground from most UV-B and all UV-C. Ozone is itself a far stronger oxidant than oxygen, $E^\circ(\ce{O3/O2}) = +2.076$ V, and $\ce{2O3} \mathrel{\ce{->}} \ce{3O2}$ has $\Delta G^\circ = -326$ kJ: near the ground it is a harmful pollutant, and it is used to disinfect drinking water.

Across period 3 from sodium to chlorine the bonding in the highest oxides changes from ionic ($\ce{Na2O}$, MgO), through an ionic lattice with covalent character ($\ce{Al2O3}$) and a giant covalent network ($\ce{SiO2}$), to small molecules ($\ce{P4O10}$, $\ce{SO3}$, $\ce{Cl2O7}$) — and the acid–base character runs from basic through amphoteric to acidic:

• Basic oxides contain $\ce{O^2-}$, form hydroxides with water and react with acids: $\ce{Na2O} + \ce{H2O} \mathrel{\ce{->}} \ce{2NaOH}$. MgO does too, but $\ce{Mg(OH)2}$ is sparingly soluble, so the solution only reaches pH ≈ 10.4.
• Amphoteric $\ce{Al2O3}$ does not dissolve in water but dissolves in both acids and bases.
• Acidic oxides give oxoacids with water and salts with bases: $\ce{P4O10} + \ce{6H2O} \mathrel{\ce{->}} \ce{4H3PO4}$, $\ce{SO3} + \ce{H2O} \mathrel{\ce{->}} \ce{H2SO4}$, $\ce{Cl2O7} + \ce{H2O} \mathrel{\ce{->}} \ce{2HClO4}$. $\ce{SiO2}$ is insoluble in water but is attacked by hot concentrated alkali or molten sodium carbonate.

Behind the pattern is the polarity of the E–O bond. A metal has a low electronegativity, so the E–O bond breaks on the E side and releases $\ce{O^2-}$ (a base); a non-metal is electronegative, so in E–O–H it is the O–H bond that breaks, releasing $\ce{H+}$ (an acid).

In $\ce{H2O2}$ oxygen is at −1, between $\ce{O2}$ (0) and $\ce{H2O}$ (−2), so it can be an oxidant ($E^\circ(\ce{H2O2/H2O}) = +1.776$ V) and a reductant ($E^\circ(\ce{O2/H2O2}) = +0.695$ V). The first lies 1.08 V above the second, so hydrogen peroxide can oxidise itself — it disproportionates: $\ce{2H2O2} \mathrel{\ce{->}} \ce{2H2O} + \ce{O2}$, $\Delta G^\circ = -234$ kJ. Fortunately this is slow, so 3 % peroxide keeps in the medicine cabinet; add $\ce{MnO2}$, iodide or the enzyme catalase and oxygen gushes out.

Ordinary sulfur is built of crown-shaped $\ce{S8}$ rings: rhombic sulfur is stable below 95.3 °C and changes to monoclinic above it, which melts at about 119 °C. Heat the liquid to about 160 °C and the rings break open and link into long chains — the melt suddenly turns as thick as syrup; pour it into cold water and you get rubbery “plastic sulfur”.

Sulfur burns to $\ce{SO2}$ ($\Delta H^\circ = -297$ kJ/mol). $\ce{SO2}$ dissolves to give sulfurous acid ($\mathrm{p}K_{\mathrm{a}1} = 1.86$), one source of acid rain; it is also a reducing agent and a bleach. $\ce{SO3}$ is a trigonal-planar molecule that reacts violently with water to give sulfuric acid.

Sulfuric acid is one of the most-produced chemicals, and its key step is $\ce{2SO2} + \ce{O2} \mathrel{\ce{<=>}} \ce{2SO3}$: $\Delta H^\circ = -198$ kJ, $\Delta S^\circ = -188$ J/K. For an exothermic reaction that loses entropy, the lower the temperature the better the equilibrium — $K_p$ falls from about $10^5$ bar⁻¹ at 425 °C to about $10^2$ bar⁻¹ at 600 °C. But at low temperature the reaction is far too slow, so a $\ce{V2O5}$ catalyst is used, working between about 400 and 630 °C.

The difficulty is the heat released: in an adiabatic bed, every 1 % of $\ce{SO2}$ converted warms the gas by about 3 °C. The first bed takes gas in at about 420 °C, and the outlet soon climbs above 600 °C, reaching the equilibrium for that temperature at only about two-thirds conversion. So the converter is split into three or four beds, with the gas cooled back to 420–440 °C between them. After the third bed the $\ce{SO3}$ formed is absorbed (“double contact”); the gas left holds little $\ce{SO2}$, but its equilibrium is favourable again, and the fourth bed pushes the overall conversion above 99.7 %.

The $\ce{SO3}$ is not bubbled straight into water — that makes a fog of acid droplets that is hard to catch — but absorbed in 98 % sulfuric acid, which is then diluted. Concentrated sulfuric acid is a strong acid (second $\mathrm{p}K_\mathrm{a} = 1.99$), a powerful dehydrating agent (it chars sugar to black carbon), and, when hot, an oxidant that dissolves copper while being reduced to $\ce{SO2}$. It releases a great deal of heat on mixing with water, so always add acid to water — slowly, with stirring.