Equilibrium at Work: Industry, Body and Environment

Sulfuric acid is made in larger amounts than any other industrial chemical, mainly for phosphate fertilizers. Sulfur is burned to ${\text{SO}}_2$ (or ${\text{SO}}_2$ is captured from smelter gases), which is oxidized to ${\text{SO}}_3$ in the step above. Le Chatelier's principle says a low temperature would give the most ${\text{SO}}_3$, since the reaction is exothermic, but the rate would be far too slow. As in the Haber process, plants compromise at about 400 to 450 °C and add a catalyst, vanadium(V) oxide, ${\text{V}}_2{\text{O}}_5$, which speeds the approach to equilibrium without changing K. High pressure would also favour the side with fewer moles of gas, but at this temperature conversion already exceeds 99% at only 1 to 2 atm when the gas passes through several catalyst beds with cooling in between, so expensive high-pressure equipment is not worth it. Excess air supplies extra ${\text{O}}_2$, shifting the equilibrium right. ${\text{SO}}_3$ is absorbed into concentrated sulfuric acid rather than water, because with water it forms a corrosive mist; the product is diluted afterward. In Sudbury, nickel smelters now capture the ${\text{SO}}_2$ that once caused acid rain and turn it into sulfuric acid.

About four out of five kidney stones contain calcium, most often calcium oxalate, ${\text{CaC}}_2{\text{O}}_4$ ($K_{sp}\approx 2.3\times 10^{-9}$). Urine is often supersaturated with it ($Q>K_{sp}$); natural inhibitors such as citrate slow crystal growth, but if Q stays high, crystals grow into stones. Every prevention strategy lowers Q. Drinking more water dilutes both ions, cutting oxalate-rich foods lowers $[{\text{C}}_2{\text{O}}_4^{2-}]$, and potassium citrate binds ${\text{Ca}}^{2+}$ in a soluble complex, lowering the free calcium ion concentration. Uric acid stones are a pH problem: uric acid is a weak acid, and in acidic urine the equilibrium favours the poorly soluble neutral form, so making urine less acidic dissolves them. Gallstones follow the same logic: most form when bile holds more cholesterol than its bile salts can keep dissolved.

Each hemoglobin molecule binds up to four ${\text{O}}_2$ molecules reversibly. In the lungs, where oxygen is plentiful, the equilibrium shifts toward oxyhemoglobin; in active tissues, where oxygen is scarce, it shifts back and releases ${\text{O}}_2$. Carbon monoxide binds the same sites about 200 times more strongly than oxygen, so even a small concentration of CO ties up a large share of hemoglobin as carboxyhemoglobin, starving tissues of oxygen. Treatment is Le Chatelier's principle: breathing pure or pressurized oxygen raises $[{\text{O}}_2]$ and shifts the equilibrium below to the left, freeing hemoglobin. Ontario has required CO alarms in homes with fuel-burning appliances since 2014.

Medicines use acid-base equilibria too. Aspirin is a weak acid ($\text{p}{K}_a\approx 3.5$). Buffered aspirin includes bases such as calcium carbonate or magnesium oxide, which neutralize some stomach acid around the tablet, so it dissolves faster and may irritate the stomach lining less. Antacids raise stomach pH the same way, and many liquid medicines and eye drops contain buffers so their pH stays stable on the shelf and comfortable in the body.

Groundwater flowing through limestone dissolves it: ${\text{CaCO}}_3(s)+{\text{CO}}_2(aq)+{\text{H}}_2\text{O}\rightleftharpoons {\text{Ca}}^{2+}+2{\text{HCO}}_3^{-}\text{.}$ Water rich in ${\text{Ca}}^{2+}$ and ${\text{Mg}}^{2+}$ is hard. Heating drives ${\text{CO}}_2$ out of solution, shifting the equilibrium left, so ${\text{CaCO}}_3$ precipitates as scale in kettles and water heaters. Water softeners swap ${\text{Ca}}^{2+}$ for ${\text{Na}}^{+}$ by ion exchange, or add carbonate so the common ion precipitates ${\text{CaCO}}_3$. Industrial wastewater is cleaned of heavy metals the same way: adding lime raises $[{\text{OH}}^{-}]$ until metal hydroxides such as ${\text{Cu(OH)}}_2$ ($K_{sp}\approx 2.2\times 10^{-20}$) precipitate, leaving tiny dissolved concentrations. Metal sulfides have even smaller Ksp values, so sulfide precipitation removes metals more completely. The pH must be controlled, because some hydroxides redissolve as complex ions in strongly basic solution.

The oceans absorb about a quarter of the ${\text{CO}}_2$ humans emit. Dissolved ${\text{CO}}_2$ forms carbonic acid, which releases ${\text{H}}^{+}$: ${\text{CO}}_2+{\text{H}}_2\text{O}\rightleftharpoons {\text{H}}^{+}+{\text{HCO}}_3^{-}$. Much of that ${\text{H}}^{+}$ then combines with carbonate ions, ${\text{H}}^{+}+{\text{CO}}_3^{2-}\to {\text{HCO}}_3^{-}$, lowering $[{\text{CO}}_3^{2-}]$. Surface ocean pH has fallen from about 8.2 before industrialization to about 8.1. For corals, oysters and plankton that build shells from calcium carbonate, lower $[{\text{CO}}_3^{2-}]$ means a smaller Q for ${\text{CaCO}}_3\rightleftharpoons {\text{Ca}}^{2+}+{\text{CO}}_3^{2-}$, so building shells costs more energy, and where Q falls below Ksp, shells begin to dissolve.