Capstone: the chemistry of ocean acidification

This final lesson applies the whole course to one real system: the carbon dioxide–carbonate equilibria that link the atmosphere, the oceans and the shells of marine life. As people burn fossil fuels, the concentration of carbon dioxide in the air has risen from about 280 parts per million before industrialization to roughly 425 to 430 ppm in 2025, according to measurements at Mauna Loa by NOAA and the Scripps Institution of Oceanography. The oceans have absorbed roughly a quarter to a third of the carbon dioxide released by human activity, slowing climate change but changing seawater chemistry. Every step of that story is chemistry from this course: Henry's law (lesson 4), equilibrium and Le Châtelier's principle (lesson 7), solubility (lesson 8), weak acids (lesson 9), buffers (lesson 10), thermodynamics (lessons 1 to 3) and kinetics (lessons 5 and 6). The lesson ends with an integrated problem set.

From air to water: Henry's law

Carbon dioxide dissolves in water in proportion to its partial pressure (Henry's law, lesson 4). At 25 °C the Henry's law constant for CO₂ in fresh water is about 0.034 mol/(L·atm); it is somewhat lower in seawater and higher in cold water, which is one reason cold polar seas absorb carbon dioxide readily.

Dissolved CO₂ reacts with water to form carbonic acid, a weak diprotic acid (lesson 9). Only a small fraction exists as true H₂CO₃; by convention the dissolved CO₂ and H₂CO₃ are counted together and written H₂CO₃* or CO₂(aq). The two ionization steps produce bicarbonate and carbonate.

These are the fresh-water values at 25 °C. In seawater, the high concentration of other ions changes the effective constants, which oceanographers measure directly; the apparent values at 25 °C and typical salinity are about pK₁* = 5.86 and pK₂* = 8.92.

Why rain is acidic

The first ionization alone sets the pH of pure water in contact with air. This is a weak-acid calculation with one twist: the concentration of dissolved CO₂ is fixed by the atmosphere, not by how much acid was added.

Speciation: which form dominates?

In a solution containing a fixed total amount of dissolved inorganic carbon (DIC = [CO₂(aq)] + [HCO₃⁻] + [CO₃²⁻]), the fraction present in each form depends only on pH. Below pK₁ dissolved CO₂ dominates; between pK₁ and pK₂ bicarbonate dominates; above pK₂ carbonate dominates. The fractions follow from the two equilibrium expressions.

Surface seawater sits near pH 8.1, where about 85 to 90 % of the dissolved inorganic carbon is bicarbonate, most of the rest is carbonate, and less than 1 % is dissolved CO₂. This is why the ocean can hold about fifty times as much carbon as the atmosphere.

The ocean as a buffer

Seawater is buffered (lesson 10), mainly by the bicarbonate–carbonate system along with borate. When carbon dioxide dissolves, the H⁺ it produces is largely consumed by carbonate ions, forming more bicarbonate. Adding the reactions shows the net effect.

The buffer keeps the pH from falling as fast as it would in unbuffered water, but it does so by using up carbonate. Since the industrial revolution, average surface-ocean pH has fallen by about 0.1 unit, from roughly 8.2 to 8.1. Because pH is logarithmic, that is about a 26 % increase in hydrogen ion concentration (10^0.1 = 1.26). Projections under high-emission scenarios suggest a further fall of 0.3 or more by 2100.

# Simplified model: carbonate speciation in seawater at constant total dissolved inorganic carbon
# (apparent constants at 25 C, salinity 35). Real ocean calculations use measured alkalinity too;
# the numbers here illustrate the trend and are not predictions.
K1, K2 = 10 ** -5.86, 10 ** -8.92
DIC = 2.0e-3            # mol/kg, typical surface value (illustrative)
Ca = 0.0103             # mol/kg calcium in seawater
Ksp_arag = 6.65e-7      # apparent Ksp of aragonite at 25 C, (mol/kg)^2

print(f"{'pH':>4s} {'CO2(aq) %':>10s} {'HCO3- %':>8s} {'CO3 2- %':>9s} {'CO3 2- umol/kg':>15s} {'Omega':>6s}")
for pH in (8.2, 8.1, 8.0, 7.8):
    H = 10 ** -pH
    d = 1 + H / K2 + H * H / (K1 * K2)
    a2 = 1 / d
    a1 = (H / K2) / d
    a0 = (H * H / (K1 * K2)) / d
    co3 = a2 * DIC
    omega = Ca * co3 / Ksp_arag
    print(f"{pH:4.1f} {100 * a0:10.2f} {100 * a1:8.1f} {100 * a2:9.1f} {co3 * 1e6:15.0f} {omega:6.2f}")

  pH  CO2(aq) %  HCO3- %  CO3 2- %  CO3 2- umol/kg  Omega
 8.2       0.38     83.7      15.9             319   4.94
 8.1       0.50     86.4      13.1             262   4.05
 8.0       0.64     88.7      10.7             213   3.30
 7.8       1.06     92.0       7.0             140   2.16

Shells, saturation and solubility

Corals, oysters, mussels, sea urchins and many plankton build shells and skeletons of calcium carbonate, as calcite or aragonite (a slightly more soluble crystal form). Whether seawater favors building or dissolving these minerals depends on the ion product compared with K_sp (lesson 8), expressed as the saturation state.

When Ω > 1, seawater is supersaturated and shells form easily; when Ω < 1, it is undersaturated and unprotected carbonate tends to dissolve. Calcium is plentiful and nearly constant in seawater, so Ω tracks the carbonate concentration. As the code shows, a fall of a few tenths of a pH unit removes a large share of the carbonate and lowers Ω markedly, even in this simplified model. Organisms can build shells below the thermodynamic threshold by controlling the chemistry at their growth sites, but doing so costs energy, and larval shellfish are especially vulnerable.

The effects are already measurable in the northeast Pacific, where upwelling brings naturally CO₂-rich deep water to the surface and human emissions add to it. Oyster hatcheries in Oregon and Washington suffered heavy larval losses during corrosive-water episodes in the late 2000s; many now monitor seawater chemistry continuously and buffer their intake water, with support from NOAA's Ocean Acidification Program. Shellfish growers in British Columbia faced similar problems.

Temperature, kinetics and the long view

Thermodynamics and kinetics shape the story over different time scales. Warming lowers the solubility of CO₂ (lesson 4), so a warmer ocean absorbs less, a feedback that leaves more carbon dioxide in the air. Calcium carbonate is unusual in being less soluble in warm water, which is why carbonate sediments accumulate in warm, shallow tropical seas and dissolve in the cold, high-pressure deep ocean, below a depth called the carbonate compensation depth. Dissolving CO₂ into surface water is fast, taking months; mixing it into the deep ocean takes centuries; and neutralizing it by dissolving carbonate sediments and weathering rocks takes thousands to hundreds of thousands of years. That is why ocean acidification, once it happens, is effectively irreversible on human time scales, even if emissions stop.

Chemistry also points to responses. Cutting emissions addresses the cause. Researchers are testing ocean alkalinity enhancement, adding finely ground minerals or alkaline solutions to seawater to raise its pH and absorb more CO₂, with field trials off Nova Scotia and elsewhere; the approach raises questions about side effects and verification that are still being studied. Seaweed and seagrass beds raise the pH of nearby water as they photosynthesize, offering local refuges for shellfish.

Integrated problem set

The problems below draw on every lesson. Work them before reading the solutions; they are typical of the questions a final examination in general chemistry asks.

The code below runs a final integrated calculation: the pH of clean rain at different atmospheric CO₂ levels, from preindustrial times to a high-emission future, combining Henry's law with the weak-acid equilibrium.

from math import log10, sqrt

kH, Ka1 = 0.034, 4.3e-7       # mol/(L atm) and first ionization constant at 25 C
for label, ppm in (("preindustrial", 280), ("1960", 317), ("2025", 425), ("doubling", 560), ("high-emission 2100", 900)):
    co2 = kH * ppm * 1e-6
    H = sqrt(Ka1 * co2)
    print(f"{label:19s} {ppm:4d} ppm: [CO2(aq)] = {co2:.2e} M, rain pH = {-log10(H):.2f}")

preindustrial        280 ppm: [CO2(aq)] = 9.52e-06 M, rain pH = 5.69
1960                 317 ppm: [CO2(aq)] = 1.08e-05 M, rain pH = 5.67
2025                 425 ppm: [CO2(aq)] = 1.45e-05 M, rain pH = 5.60
doubling             560 ppm: [CO2(aq)] = 1.90e-05 M, rain pH = 5.54
high-emission 2100   900 ppm: [CO2(aq)] = 3.06e-05 M, rain pH = 5.44

Clean rain becomes only slightly more acidic as CO₂ rises, because the ocean and land buffers do not apply to a raindrop and [H⁺] varies only as the square root of P(CO₂). The ocean's story is different and more serious, because it is the carbonate concentration, not pH alone, that matters to marine life. Working through a real system like this, connecting enthalpy, entropy, rates, equilibria, solubility, acid–base chemistry and electrochemistry, is exactly what the physical, analytical and environmental chemistry courses that follow (CHEM-211, CHEM-271) build on.