Questions: Ocean Acidification: Chemistry and Biological Impacts

Question 1

Ocean pH has dropped from 8.2 to 8.1 since preindustrial times. A skeptic argues this 0.1 unit change is negligible. Which response best captures why this change matters biologically?

• A. The change is negligible; biological systems can easily buffer a 0.1 pH unit shift
• B. A 0.1 unit drop represents a roughly 26% increase in hydrogen ion concentration and a significant reduction in carbonate ion concentration that impairs calcification
• C. The change matters because ocean pH has now crossed below 7.0, entering the chemically acidic range
• D. The change matters only because of its speed, not its absolute magnitude

Answer: (B)

Because pH is logarithmic, a 0.1 unit drop corresponds to a ~26% increase in [H+]. More importantly for biology, this pH drop goes hand-in-hand with a significant reduction in carbonate ion concentration [CO3²⁻], which directly lowers the saturation state Ω of calcium carbonate. For calcifying organisms, it is this drop in Ω — not pH per se — that makes shell-building harder. The ocean has not crossed pH 7; it remains alkaline, which is why 'ocean acidification' technically means 'becoming less alkaline,' though the chemical and biological consequences are real.

Question 2

Oyster larvae in a coastal upwelling zone are failing to form shells properly despite a pH of 7.9. What is the most likely biochemical explanation?

• A. A pH of 7.9 is so close to neutral that acid is directly dissolving the shells
• B. Upwelling brings deep, CO2-rich water to the surface, reducing carbonate ion concentration and dropping the saturation state (Ω) below the threshold for calcification
• C. Upwelling reduces water temperature, which inhibits the enzyme responsible for shell formation
• D. The larvae are absorbing too much bicarbonate, which blocks calcium carbonate precipitation

Answer: (B)

Upwelling zones bring cold, deep, CO2-rich water to the surface. This CO2 drives the carbonate equilibrium toward more bicarbonate and H+, consuming carbonate ions in the process. The result is a low carbonate saturation state Ω, which can drop below 1 — the point at which CaCO3 spontaneously dissolves rather than precipitates. Larval calcifiers are especially vulnerable because they must rapidly form their first shell with limited energy reserves. This mechanism explains the oyster larvae collapses observed in Pacific Northwest hatcheries in the 2000s.

Question 3

Ocean acidification refers to the ocean becoming chemically acidic, with pH dropping below 7.

• T. True
• F. False

Answer: False

The term 'ocean acidification' is technically misleading in common usage. Ocean pH remains above 7 (currently around 8.1), meaning seawater remains alkaline. 'Acidification' refers to the directional trend — pH is decreasing toward more acidic values — not to the absolute value crossing 7. The important consequences come not from the ocean becoming acid, but from the reduction in carbonate ion concentration that accompanies rising H+ levels, which reduces the saturation state for calcium carbonate and threatens calcifying organisms.

Question 4

When CO2 dissolves in seawater, it simultaneously lowers pH and reduces carbonate ion concentration — both effects together make calcification more difficult for shell-building organisms.

• T. True
• F. False

Answer: True

These are linked consequences of the same equilibrium reactions. CO2 + H2O → H2CO3 → HCO3⁻ + H+. The H+ produced then reacts with carbonate ions: H+ + CO3²⁻ → HCO3⁻. This consumes carbonate ions, reducing [CO3²⁻]. So adding CO2 simultaneously increases [H+] (lower pH) and decreases [CO3²⁻] (lower saturation state Ω). It is the reduction in [CO3²⁻] — reflected in Ω — that directly reduces the thermodynamic driving force for CaCO3 precipitation and makes shell-building energetically costlier.

Question 5

Why is carbonate saturation state (Ω) a more useful metric than pH alone for predicting whether calcifying organisms can survive in acidifying waters?

Model answer: Saturation state Ω = [Ca²+][CO3²⁻] / Ksp directly measures whether the water is thermodynamically favorable for CaCO3 precipitation. When Ω > 1, shell-building is favorable; when Ω < 1, existing shells dissolve. pH describes hydrogen ion concentration but does not directly quantify carbonate ion availability. Two water samples could have the same pH but different alkalinities, yielding different carbonate ion concentrations and different Ω values. For a calcifier, the relevant question is not 'how acidic is this water?' but 'can I precipitate calcium carbonate here?' — and Ω answers that question directly.

This distinction is practically important: pH alone can be misleading as an indicator of biological stress to calcifiers. Ω integrates the carbonate chemistry that actually governs calcification thermodynamics, making it the preferred metric in ocean acidification research and monitoring programs.