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Dissolved CO₂ in seawater exists in equilibrium as carbonic acid, bicarbonate, and carbonate ions, forming a powerful buffering system that maintains ocean pH near 8.2. This carbonate equilibrium regulates the solubility of biogenic calcium carbonate shells and controls how much atmospheric CO₂ the ocean can absorb before experiencing acidification.
Your background in acid-base chemistry gives you the tools to understand this system: when an acid or base is added to a buffered solution, the buffer absorbs the perturbation and resists large pH changes. The ocean's carbonate system is the planet's largest natural buffer, and it operates through a series of linked equilibria that you can trace step by step.
When CO₂ from the atmosphere dissolves in seawater, it first forms dissolved CO₂ (sometimes written as CO₂(aq)). A small fraction of this reacts with water to form carbonic acid (H₂CO₃), which is a weak diprotic acid. Carbonic acid rapidly loses a proton to form bicarbonate (HCO₃⁻), and bicarbonate can lose another proton to form carbonate (CO₃²⁻). At the ocean's typical pH of about 8.1–8.2, the equilibrium overwhelmingly favors bicarbonate, which constitutes roughly 90% of the total dissolved inorganic carbon, with carbonate at about 9% and dissolved CO₂ at only about 1%. This distribution is a direct consequence of the pKa values of carbonic acid — at ocean pH, the first deprotonation is essentially complete while the second is only partial.
The buffering works because adding CO₂ to the system does not simply accumulate as dissolved gas — it is absorbed into the equilibrium. Additional CO₂ reacts with water and shifts the equilibrium toward more bicarbonate, consuming carbonate ions and releasing hydrogen ions in the process. The pH drops, but far less than it would in unbuffered water, because the enormous reservoir of bicarbonate and carbonate ions absorbs most of the perturbation. This is why the ocean has been able to absorb roughly 30% of anthropogenic CO₂ emissions without catastrophic pH collapse — the buffer is doing its job. However, each additional increment of CO₂ consumes carbonate ions, progressively weakening the buffer and making the ocean more sensitive to further additions. This declining buffer capacity is quantified by the Revelle factor, which measures how much the partial pressure of CO₂ changes relative to changes in total dissolved inorganic carbon.
The carbonate system has direct consequences for marine life. Many organisms — corals, foraminifera, coccolithophores, mollusks — build shells and skeletons from calcium carbonate (CaCO₃). The saturation state of seawater with respect to calcium carbonate depends on the concentration of carbonate ions: as CO₂ is added and carbonate ions are consumed, the water becomes less saturated and eventually undersaturated, meaning existing shells begin to dissolve. This connection between atmospheric CO₂, ocean chemistry, and biological calcification is the mechanistic basis of ocean acidification — not a shift to truly acidic conditions, but a measurable decline in pH and carbonate saturation that threatens calcifying organisms and the ecosystems that depend on them.