Acid-Base and Solubility Equilibria in Our Bodies, Environment and Industry

Equilibrium is everywhere

Right now your blood is holding its pH between about 7.35 and 7.45 while your muscles make acid and your lungs breathe out CO₂. The ocean is absorbing about a quarter of the CO₂ that people release into the air each year. In Northern Ontario, many lakes near Sudbury that were damaged by acid rain have been recovering since smelter emissions were cut. All of these depend on the equilibria from this unit: Ka, buffers, Ksp and Le Châtelier's principle. In Equilibrium in Living Things, the Environment and Technology (unit Chemical Equilibrium) you explained several of these systems qualitatively with Le Châtelier's principle. This lesson adds the numbers from this unit (pH, buffer ratios and Ksp) and is about assessing their impact: using chemistry to explain real systems and to judge claims and decisions about them.

1. The blood buffer

The main buffer in blood plasma is the carbonic acid-hydrogen carbonate system:

CO₂(aq) + H₂O(l) ⇌ H₂CO₃(aq) ⇌ H⁺(aq) + HCO₃⁻(aq)

At body temperature, with dissolved CO₂ counted as the acid, the effective pKa is about 6.1. Normal blood has [HCO₃⁻] ≈ 24 mmol/L and dissolved CO₂ ≈ 1.2 mmol/L, a ratio of 20 : 1:

pH = 6.1 + log(24 ÷ 1.2) = 6.1 + log 20 = 7.4

That pH is more than one unit above pKa, so on paper this is a lopsided buffer. It works so well because the body controls both partners: the lungs remove CO₂ (within minutes) and the kidneys adjust HCO₃⁻ and excrete H⁺ (over hours to days). When exercise floods the blood with acid, the HCO₃⁻ absorbs it, and faster breathing removes the extra CO₂, shifting the equilibria to the left.

If breathing is too slow (for example during some lung diseases), CO₂ builds up, the equilibria shift right and pH falls (respiratory acidosis). Breathing too fast (hyperventilating) removes too much CO₂, and pH rises (respiratory alkalosis). Phosphate (H₂PO₄⁻/HPO₄²⁻) and proteins act as buffers too, especially inside cells.

2. Hemoglobin, teeth and kidney stones

Hemoglobin (recap). You studied Hb + O₂ ⇌ HbO₂, altitude and carbon monoxide in Equilibrium in Living Things, the Environment and Technology. The acid-base link is new here: muscles produce CO₂ and H⁺, and a lower pH makes hemoglobin release O₂ more readily (the Bohr effect), so oxygen goes where it is needed most. The blood buffer above keeps that pH change small everywhere else.

Tooth enamel. As you saw in Chemical Equilibrium, enamel is mostly hydroxyapatite:

Ca₅(PO₄)₃OH(s) ⇌ 5 Ca²⁺(aq) + 3 PO₄³⁻(aq) + OH⁻(aq)

When bacteria in plaque ferment sugars, they make acids. H₃O⁺ removes OH⁻ and PO₄³⁻ (both are bases), so the equilibrium shifts right and enamel dissolves (demineralization). Saliva, a buffer containing HCO₃⁻ and phosphate, raises the pH again so minerals can re-deposit. Fluoride replaces OH⁻ to form fluorapatite, Ca₅(PO₄)₃F, which is less soluble in acid; F⁻ is a much weaker base than OH⁻.

Kidney stones. The most common kind is calcium oxalate (Ksp ≈ 2.3 × 10⁻⁹). When urine becomes concentrated (for example with too little water intake), the ion product of Ca²⁺ and oxalate can exceed Ksp, and crystals can form and grow. Drinking more water lowers both concentrations, and therefore Q.

3. Lakes, acid rain and the Canadian Shield

Burning fossil fuels and smelting sulfide ores release SO₂ and NOₓ, which form sulfuric and nitric acids in the atmosphere. Normal rain has a pH of about 5.6 (from dissolved CO₂); acid rain is lower.

What happens next depends on buffering capacity. Lakes in areas with limestone (CaCO₃) contain plenty of HCO₃⁻ and CO₃²⁻, which neutralize added acid. Many lakes on the Canadian Shield, which covers much of Northern Ontario, sit on granite and other rocks that dissolve very little, so they have little HCO₃⁻ and are easily acidified. As the pH of a lake falls, fish and other organisms are harmed, and aluminium leached from soil (its compounds are more soluble in acid) adds to the damage.

Around Sudbury, decades of smelter emissions acidified many lakes and damaged the surrounding land. In 1960 the Sudbury smelters were the largest single source of sulfur emissions in the world, and an estimated 7000 lakes in the region were damaged by acid. Since then, SO₂ and metal emissions have been cut by more than 90 % (through smelter controls such as acid plants and process changes), and many lakes have recovered chemically, with fish such as lake trout returning. Recovery is not complete: calcium loss from soils, climate change and other stresses slow it in some lakes. The Sudbury regreening program spread crushed limestone on acidic, barren soils so that plants could grow again: an acid-base reaction (CaCO₃ + 2 H₃O⁺ → Ca²⁺ + CO₂ + 3 H₂O) used on a landscape scale.

4. Ocean acidification

The ocean absorbs CO₂, which forms carbonic acid. The average pH of surface seawater has dropped from about 8.2 before industrialization to about 8.05 today (8.04 in 2024), often rounded to "about 8.1". That sounds small, but it is a logarithmic scale: [H₃O⁺] has risen by about 30%. In 2025, scientists assessing Earth's "planetary boundaries" reported that ocean acidification has moved past its proposed safe limit. The extra H⁺ combines with carbonate ions, lowering [CO₃²⁻]. Corals, oysters, mussels, pteropods and many plankton build shells and skeletons of CaCO₃, which becomes harder to build and easier to dissolve when [CO₃²⁻] falls.

5. Cleaning up water with Ksp

Acid mine drainage forms when sulfide minerals exposed by mining react with air and water, making acidic water rich in dissolved metal ions. A common treatment is to add lime (Ca(OH)₂ or CaO). Raising the pH neutralizes the acid and pushes Q above Ksp for metal hydroxides such as Fe(OH)₃ and Cu(OH)₂, so the metals precipitate and can be settled out as sludge. Treatment plants choose the pH carefully, because each metal hydroxide precipitates over a different pH range (and some, such as aluminium hydroxide, redissolve if the pH goes too high).

Other examples from industry and medicine:

• Water softening: hard water contains Ca²⁺ and Mg²⁺. Adding lime and sodium carbonate precipitates CaCO₃ and Mg(OH)₂.
• Medical imaging: BaSO₄ is so slightly soluble that a patient can drink a suspension of it before an X-ray; very little toxic Ba²⁺ dissolves.
• Swimming pools: operators keep pH near 7.2 to 7.8, using buffers and acids or bases, so that chlorine disinfects well and CaCO₃ scale does not form.

Careers that use this unit: respiratory therapist, nurse, physician, dentist and dental hygienist, pharmacist, water and wastewater treatment operator, environmental technologist, chemical engineer, mining engineer, oceanographer and food scientist.