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The ocean covers about 71% of the Earth’s surface and holds roughly fifty times as much carbon as the atmosphere. For millions of years its chemistry has been remarkably steady, with a pH a little above 8. That’s now changing faster than at any time in at least tens of thousands of years, and the cause is the same carbon dioxide that’s warming the planet.
This article explains the chemistry behind ocean pH, what’s happening to it, and why a change that sounds tiny matters so much.
Seawater is slightly alkaline
Surface seawater today has an average pH of about 8.05–8.1. That’s mildly alkaline. Before the Industrial Revolution, it was about 8.2.
Why isn’t seawater neutral? Because it contains dissolved hydrogencarbonate (HCO₃⁻) and carbonate (CO₃²⁻) ions, originally washed out of rocks by rivers over geological time. These ions are weak bases, and together with dissolved carbon dioxide they form a natural buffer system.
The ocean’s carbonate system
When carbon dioxide from the air dissolves in seawater, a chain of equilibria follows:
CO₂(g) ⇌ CO₂(aq) CO₂(aq) + H₂O ⇌ H₂CO₃ (carbonic acid) H₂CO₃ ⇌ H⁺ + HCO₃⁻ HCO₃⁻ ⇌ H⁺ + CO₃²⁻
At seawater’s pH of about 8.1, the dissolved inorganic carbon is distributed roughly as:
| Species | Share of dissolved inorganic carbon |
|---|---|
| Hydrogencarbonate, HCO₃⁻ | about 90% |
| Carbonate, CO₃²⁻ | about 9% |
| Dissolved CO₂ and H₂CO₃ | about 1% |
That’s consistent with the pH sitting between the two pKa values of carbonic acid in seawater (roughly 5.9 and 9.0 under seawater conditions). See polyprotic acids for how these stepwise equilibria work.
What adding CO₂ does
The ocean has absorbed roughly a quarter of all the carbon dioxide humans have released since the industrial era began. That’s slowed the rise of CO₂ in the atmosphere, but it’s changed the ocean’s chemistry.
When extra CO₂ dissolves:
- More carbonic acid forms.
- It releases H⁺ ions, so pH falls.
- Many of those H⁺ ions react with carbonate ions: H⁺ + CO₃²⁻ → HCO₃⁻. So carbonate concentration falls.
The overall result can be summarised as:
CO₂ + H₂O + CO₃²⁻ → 2HCO₃⁻
This is ocean acidification. The name doesn’t mean the ocean is becoming acidic (it’s still alkaline, above pH 7). It means it’s moving in the acidic direction.
Why a drop of 0.1 matters
A drop from 8.2 to about 8.1 sounds tiny. But pH is logarithmic. A drop of 0.1 pH units means:
[H⁺] increases by a factor of 10⁰·¹ = 1.26
That’s a 26% increase in hydrogen ion concentration (often rounded to “about 30%”) in roughly two centuries. It’s a larger and faster change than anything recorded in ice cores and marine sediments for at least hundreds of thousands of years. See the pH scale explained for why small pH changes mean large concentration changes.
Projections depend on future emissions. Under high-emission scenarios, surface ocean pH could fall to about 7.7–7.8 by 2100, which would mean roughly double to triple the pre-industrial H⁺ concentration. Under strong emission cuts, the decline would slow and stabilise.
How we know: long-term measurements
Direct, continuous measurements began in the late 1980s at ocean time-series stations such as Station ALOHA near Hawaii and the Bermuda Atlantic Time-series Study (BATS). They show seawater CO₂ rising in step with atmospheric CO₂, and pH falling by about 0.016–0.02 units per decade. Older values come from reconstructions using the chemistry of fossil shells and corals, which record the conditions they grew in.
Why falling carbonate is a problem
Many marine organisms build shells or skeletons of calcium carbonate: corals, oysters, mussels, clams, sea urchins, and tiny plankton such as coccolithophores and pteropods (“sea butterflies”).
They make it from calcium and carbonate ions:
Ca²⁺ + CO₃²⁻ → CaCO₃
When carbonate concentration falls, building shells takes more energy. If carbonate falls far enough, seawater becomes undersaturated in calcium carbonate, and existing shells and skeletons can begin to dissolve:
CaCO₃ + CO₂ + H₂O → Ca²⁺ + 2HCO₃⁻
Chemists describe this using the saturation state (Ω), the ratio of the actual ion product [Ca²⁺][CO₃²⁻] to its value at equilibrium with solid calcium carbonate. Above 1, shells tend to form; below 1, they tend to dissolve. Aragonite, the form of calcium carbonate used by corals and many molluscs, is more soluble than calcite and more vulnerable.
Effects on marine life
Research in laboratories and in naturally acidified waters (such as near underwater volcanic CO₂ seeps) shows:
- Corals grow more slowly and their reefs erode more readily, especially combined with warming-driven bleaching.
- Oyster hatcheries on the Pacific coast of North America experienced mass larval die-offs in the late 2000s linked to corrosive, CO₂-rich upwelled water. Some now monitor and buffer their seawater.
- Pteropods, a key food for fish such as salmon, show shell damage in more acidic waters.
- Some fish show changes in behaviour and sense of smell in high-CO₂ water in experiments, though results vary between studies.
- Some algae and seagrasses may benefit from extra CO₂ for photosynthesis.
The overall picture is that acidification adds stress to ecosystems already affected by warming, pollution and overfishing.
The ocean’s natural buffer is limited
Seawater’s carbonate buffer is powerful, which is why the pH has only dropped by about 0.1 so far despite enormous CO₂ uptake. But every CO₂ molecule absorbed uses up some carbonate. As carbonate declines, the buffer capacity falls, and the ocean becomes less able to absorb further CO₂. It’s the same principle as buffer capacity in a lab solution: add enough acid and any buffer weakens.
Over thousands of years, dissolving carbonate sediments on the sea floor and weathering of rocks on land will restore carbonate, but these processes are far too slow to keep up with today’s emissions.
Not the same everywhere
Ocean pH isn’t uniform. Cold water absorbs more CO₂ than warm water, so polar seas are acidifying faster and are already closer to the point where aragonite dissolves. Coastal waters swing more widely because of river inflow, upwelling of deep CO₂-rich water, and nutrient pollution that fuels algal blooms (when those algae die and decompose, they release more CO₂). Deep water is naturally more acidic than surface water because it has accumulated CO₂ from sinking, decomposing organic matter over centuries.
Can we reverse it?
The only lasting fix is reducing CO₂ emissions. Researchers are also exploring:
- Ocean alkalinity enhancement: adding crushed alkaline minerals (such as olivine or limestone) or alkaline solutions to seawater to raise pH and increase CO₂ uptake.
- Protecting seagrass and kelp, which take up CO₂ locally and may create small refuges of higher pH.
- Local buffering in shellfish hatcheries.
These are active research areas with open questions about scale, cost and side effects.
Key takeaways
- Seawater is slightly alkaline (pH about 8.1) because of its hydrogencarbonate–carbonate buffer.
- The ocean has absorbed about a quarter of human CO₂ emissions, lowering pH by about 0.1 since pre-industrial times: roughly a 26–30% increase in [H⁺].
- Dissolving CO₂ consumes carbonate ions, making it harder for corals and shellfish to build calcium carbonate.
- High-emission projections suggest surface pH could reach about 7.7–7.8 by 2100.
- For the underlying equilibria, see Ka and pKa.
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