Buffer uses and ocean acidification

Many chemical and biological processes only work within a narrow range of pH, and buffers keep them there. The oceans are buffered too, by carbonate and hydrogencarbonate ions, but uptake of additional atmospheric carbon dioxide changes the carbonate balance and lowers pH. This lesson describes some uses of buffers and explains ocean acidification.

Uses of buffers

• Blood normally stays near pH 7.4. The carbonic acid/hydrogencarbonate system is one important buffer, H₂CO₃/HCO₃⁻. Protein buffers also contribute, and lungs and kidneys regulate the system. Enzyme activity and haemoglobin’s oxygen binding depend on pH.
• pH meters are calibrated with buffer solutions of known pH, because they resist small acid or base additions. Their reference pH is specified at a given temperature; carbon dioxide uptake and contamination can still change it.
• Biochemical experiments, food, cosmetics and medicines use buffers wherever a reaction or product must stay at a set pH.

In each case the buffer works only while both members of its conjugate pair remain in large amounts. It reduces a change in pH; it cannot remove an unlimited amount of acid or base.

The carbonate buffer in the oceans

Seawater contains hydrogencarbonate ions and carbonate ions, a conjugate acid–base pair:

• Added acid is removed by carbonate ions: CO₃²⁻(aq) + H⁺(aq) → HCO₃⁻(aq)
• Added base is removed by hydrogencarbonate ions: HCO₃⁻(aq) + OH⁻(aq) → CO₃²⁻(aq) + H₂O(l)

This limits pH changes rather than fixing every seawater sample at one value.

Carbon dioxide and ocean acidification

Some of the additional atmospheric carbon dioxide dissolves in the oceans. In water it forms carbonic acid, a weak acid:

As the concentration of carbon dioxide in the air rises, more dissolves, these equilibria shift to the right, and [H⁺] rises. Carbonate ions remove some of the extra hydrogen ions, so overall

The buffer absorbs most of the extra acid, but every carbonate ion that reacts is used up. The ratio of carbonate to hydrogencarbonate falls, so the pH still falls, only by less than it would without the buffer. Since the industrial revolution, global mean surface-ocean pH has fallen from about 8.2 to 8.1, which means about 26% more hydrogen ions. NOAA describes these global averages and their carbonate consequences. Local samples vary.

This fall is called ocean acidification. Seawater is still alkaline: acidification describes the direction of change, not a pH below 7.

Lower carbonate concentration matters to living things. Corals, shellfish and plankton build shells and skeletons of calcium carbonate:

With fewer carbonate ions, this equilibrium shifts to the left, so building shells becomes harder and thin shells may begin to dissolve.

Reading a pH difference

A pH fall of d gives a final-to-initial hydrogen-ion ratio of 10^d. For example, a 0.10 fall gives 10^(0.10) = 1.26, a 26% increase. pH alone measures the free hydrogen-ion concentration, rather than all acid absorbed by the buffer.

Common mistakes

• Saying ocean acidification means seawater has become acidic. It is still alkaline, but less so.
• Saying a buffer keeps the pH constant. It only reduces the change, and only while both members of the pair remain.
• Mixing up the members of the carbonate pair: carbonate removes added acid; hydrogencarbonate removes added base.
• Linking falling carbonate to “acid dissolving shells” without the equilibrium Ca²⁺ + CO₃²⁻ ⇌ CaCO₃.
• Writing carbon dioxide as fully converted to carbonic acid. Hydration and acid dissociation are equilibria, rather than complete conversion of all carbon dioxide to carbonic acid.