Blood, rain and the ocean

Blood is normally held within a narrow range, typically quoted as pH 7.35 to 7.45 for arterial blood, mainly by the bicarbonate buffer, carbonic acid (dissolved CO₂) and bicarbonate ion^[1]. Its pK of 6.1 is more than a unit away from 7.4, which would make it a poor buffer in a closed bottle; it works because the body controls both partners, the lungs by breathing out CO₂ and the kidneys by keeping or excreting bicarbonate. The same chemistry explains why unpolluted rain has a pH of about 5.6, and why the surface ocean, now near pH 8.1, has become measurably more acidic as it absorbs carbon dioxide from the air.

A buffer that breathes§

Carbon dioxide dissolves in blood and reacts with water to form carbonic acid, which gives up a proton to make bicarbonate: CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻. Physiologists write the Henderson–Hasselbalch equation for it with two practical substitutions. The acid is measured as the partial pressure of CO₂, P_CO₂, in mmHg, multiplied by its solubility, 0.03 mmol/L per mmHg; and the constant is an apparent pK′ of 6.1 for plasma at body temperature^[1][2]:

With the typical textbook values, 24 mmol/L of bicarbonate and 40 mmHg of CO₂ (1.2 mmol/L dissolved), the ratio is 20 to 1 and the pH is 6.1 + log 20 = 7.40^[1]. established That is the point drawn in the diagram at the top of this page.

Twenty to one is a lopsided buffer. As a closed solution, 25 mmol/L of this pair would have a buffer capacity at pH 7.4 of only about 0.003 mol/L per pH unit, a twentieth of the acetate buffer in Chapter 4. The reason it works is that the acid partner is a gas. Every hydrogen ion that bicarbonate absorbs becomes carbonic acid, then CO₂, which the lungs breathe out. Instead of accumulating, the acid leaves. The body sets the denominator of equation (5.1) with every breath, and the numerator over hours to days through the kidneys, which excrete non-volatile acids and reclaim bicarbonate^[1].

You can put a number on how much the open system gains. If the lungs hold the CO₂ fixed, then [HCO₃⁻] = K′·0.03P_CO₂/[H⁺], and every tenfold fall in [H⁺] needs ten times as much bicarbonate. Differentiating gives a buffer capacity of ln 10 × [HCO₃⁻], about 2.3 × 0.024 = 0.055 mol/L per pH unit: some twenty times the closed-bottle figure, and about as much as the 0.1 M acetate buffer of Chapter 4 at its best. illustrative The weak spot is the other side of the bargain. A buffer whose acid can leave is only as good as the breathing that removes it.

The diagram shows the two ways the pH can be pushed away from 7.4, and two ways back. If CO₂ rises to 60 mmHg because breathing is inadequate, the point slides along the bicarbonate line to pH 7.22. If an acid load from metabolism uses up bicarbonate, the point falls straight down: at 14 mmol/L of bicarbonate the pH is 7.17. Then press the last button. Breathing harder lowers the CO₂ to 28 mmHg and brings the ratio, and the pH, back to 7.32, closer to 7.4 although the bicarbonate is still low. Clinicians call this respiratory compensation of a metabolic acidosis, and the kidneys provide the slower compensation in the other direction^[1].

The two controls work on very different clocks. Breathing changes the CO₂ within minutes, so respiratory compensation is fast but partial. The kidneys change the bicarbonate over hours to days, so a person with long-standing lung disease and a high CO₂ often has a pH close to the reference range, because their bicarbonate has climbed to match^[1]. On the diagram that is a point well above the textbook one, sitting on a high-CO₂ isobar but still near the shaded band.

Why clean rain is acidic§

Rain falling through clean air is not pH 7. It has a pH of about 5.6, because carbon dioxide dissolves in it and forms carbonic acid^[3]. This is a calculation the site’s equilibrium engine can do from first principles. Henry’s law gives the dissolved CO₂: about 0.034 mol/L for every atmosphere of CO₂ partial pressure at 25 °C^[4]. In July 2026 the air held 426.1 parts per million of CO₂, as a global monthly mean^[5], so the rain contains 1.5 × 10⁻⁵ mol/L of dissolved CO₂. The carbonic acid constants^[6] and the charge balance [H⁺] = [HCO₃⁻] + 2[CO₃²⁻] + [OH⁻] then fix the pH.

At today’s CO₂ the model gives pH 5.59, matching the EPA’s “about 5.6”. established Acid rain is something else: rain with sulfuric and nitric acids formed from sulfur dioxide and nitrogen oxides, mostly from burning fossil fuels, and typically between pH 4.2 and 4.4^[3]. As Chapter 1 showed, that is about twenty times the hydrogen-ion concentration of clean rain.

A tenth of a pH unit in the ocean§

The ocean absorbs about 30 % of the carbon dioxide released into the atmosphere, and the pH of surface ocean water has fallen by about 0.1 since the industrial revolution began, to an average of about 8.1 today^[7]. established NOAA describes that as an increase in acidity of approximately 30 %. The arithmetic is worth doing, because it is where a logarithmic scale fools intuition: a fall of exactly 0.10 multiplies [H⁺] by 10^0.1 = 1.26, an increase of 26 %; a fall of 0.11 gives 29 %, and 0.12 gives 32 %. A change that looks like a rounding error on the scale is a quarter to a third more hydrogen ions.

Why the sea sits at 8.1, not 5.6§

Rain and seawater dissolve CO₂ from the same air, yet one ends up two and a half pH units more basic than the other. The difference is what else is in the water. Rock weathering has, over geological time, delivered more strong-base cations (sodium, magnesium, calcium, potassium) to the sea than strong-acid anions (chloride, sulfate) to balance them. The charge balance of Chapter 2 has to close, so the shortfall of negative charge is made up by bicarbonate, carbonate and borate ions. That excess is called the total alkalinity, about 2,300 micromoles per kilogram of surface seawater^[8][9]:

Alkalinity is literally a titration result. Oceanographers measure it by titrating a weighed sample of seawater with hydrochloric acid past the carbonic-acid end point near pH 4.5, the second equivalence point of the carbonate titration in Chapter 3^[9]. established Rain has no alkalinity at all, so its pH is set by the CO₂ alone. Seawater has a large store of base, and the CO₂ only nudges the pH that the base has already set.

Does that store of base make the ocean slower to acidify than rain? The figure answers it with the same chemistry the rest of this site uses, now with constants measured in seawater: the solubility of CO₂ (Weiss 1974), the two carbonic-acid constants (Lueker, Dickson and Keeling 2000), boric acid (Dickson 1990) and water (Millero 1995), as compiled in the standard oceanographic handbook^{[10][11][12][13][9]}. Before you look, guess.

Between 280 ppm, the air before industrialization, and today’s level, the model ocean falls from pH 8.18 to 8.03 while clean rain falls from 5.68 to 5.59. Per doubling of CO₂ the ocean drops about 0.25 and the rain 0.15. illustrative The reason is in the readout under the figure. Most of the extra CO₂ that enters seawater does not stay as carbonic acid; it reacts with carbonate ion,

and the carbonate ion concentration falls by about a quarter between 280 ppm and today in the model. illustrative Because the alkalinity of equation (5.2) is fixed and mostly made of bicarbonate, [H⁺] ≈ K₁[CO₂]/[HCO₃⁻] grows almost in proportion to the CO₂, with only borate and the remaining carbonate softening it. In rain, by contrast, [H⁺] follows the square root. In the language of Chapter 4, the ocean is a buffer whose base partner is being spent, and it is the partner that is running out, not the pH, that worries marine biologists: shell-building organisms such as corals and some plankton need carbonate ions to make calcium carbonate^[7].

The model is an equilibrium calculation, and the real ocean is not quite at equilibrium. Its fall of 0.15 since 1800 is a little larger than NOAA’s rounded “about 0.1”: surface water lags the air, mixes with older water from below, and differs from region to region and season to season by more than the trend itself. That is why the measured decline comes from long time series rather than from a formula. At Station ALOHA in the central North Pacific near Hawaii, nearly twenty years of measurements showed surface pH falling by 0.0019 ± 0.0002 per year, a rate indistinguishable from what equilibration with the rising CO₂ in the air predicts^[14]. established The figure agrees: Mauna Loa’s annual mean CO₂ rose from about 352 ppm in 1988 to 384 ppm in 2007^[15], and over that rise the model ocean falls by 0.032, or 0.0017 per year. The same study found a strong seasonal cycle on top of the trend, driven by temperature, mixing and photosynthesis, and larger changes at some depths below the surface mixed layer.

The thread through all three examples is the one that runs through the whole site. pH measures the concentration of hydrogen ions on a logarithmic ruler; that concentration is set by equilibrium constants and the balance of charges; and a buffer, whether a beaker of acetate, a bloodstream or an ocean, resists change only as long as one of its partners has something left to give.