The Carbonate System in Aquarium Water: CO2, Bicarbonate and Buffering

How dissolved CO2, carbonic acid, bicarbonate and carbonate interconvert, what the pKa values 6.35 and 10.33 mean, and why alkalinity (KH) keeps aquarium pH from swinging.

The chain of equilibria

Carbon dioxide that dissolves in water, whether from the air, from fish and bacterial respiration or from injected gas, enters a sequence of linked reversible reactions. Dissolved CO2 combines with water to form carbonic acid (H2CO3); carbonic acid releases a hydrogen ion to become bicarbonate (HCO3-); bicarbonate can release a second hydrogen ion to become carbonate (CO3 2-). Written as one chain: CO2 + H2O ⇌ H2CO3 ⇌ H+ + HCO3- ⇌ 2H+ + CO3 2-. Because every step can run in either direction, adding or removing any one member shifts all the others and changes the hydrogen ion concentration, which is what pH measures (pH is the negative logarithm of the molar hydrogen ion concentration).

The first step lies far to the left. At 25 °C the ratio of true carbonic acid to dissolved CO2 is only about 1.7 × 10-3 in pure water and about 1.2 × 10-3 in seawater, so almost all of the 'carbonic acid' in a tank is in fact dissolved CO2 gas. For practical purposes the system is described by dissolved CO2, bicarbonate and carbonate, whose shares always add up to 100 % of the dissolved inorganic carbon.

pKa values and which species dominates

Each dissociation step has an acid constant, usually given as pKa. For the hydrated (CO2 plus H2CO3) form, the commonly quoted apparent values at 25 °C are pKa1 = 6.35 and pKa2 = 10.33; a Bjerrum plot, which graphs the share of each species against pH, places the crossover points at roughly 6.4 and 10.3. At a pH equal to pKa1 dissolved CO2 and bicarbonate are present in equal amounts; at a pH equal to pKa2 bicarbonate and carbonate are equal. The constants vary substantially with temperature and salinity, so the curves for seawater sit at different positions from those for fresh water.

• Below about pH 6.4: dissolved CO2 is the largest fraction.
• Between about pH 6.4 and 10.3, which includes the pH 6.5–9.0 range recommended for aquaculture: bicarbonate dominates.
• Above about pH 10.3: carbonate dominates.
• Below pH 8.5, carbonate is less than 1 % of bicarbonate in most natural waters.

Buffering capacity

A buffer is a weak acid together with its conjugate base. When acid is added, the base takes up the hydrogen ions; when base is added, the weak acid releases hydrogen ions to neutralise it, so pH changes only a little until one partner is used up. Buffer capacity peaks where pH equals pKa and falls to about one third of that maximum at pKa ± 1, which is why the useful range of a buffer pair is usually taken as pKa ± 1. The CO2/bicarbonate pair is the same buffer that holds human blood between pH 7.35 and 7.45.

Alkalinity is the measure of this acid-neutralising capacity: the amount of acid water can absorb before reaching a designated pH. It is determined by titration, in which bicarbonates neutralise the added acid first and the pH then falls steeply once they are exhausted. In fresh water carbonates and bicarbonates are the most common and important bases, with hydroxides, phosphates and borates contributing smaller amounts; in seawater borate adds a small share. Carbonate counts twice, because each carbonate ion can neutralise two hydrogen ions. Alkalinity is reported in mg/L as CaCO3 or in meq/L.

Why KH stabilises pH

Carbonate hardness (KH, from the German Karbonathärte) measures the carbonate and bicarbonate content of water; 1 °dKH equals 17.848 mg/L as CaCO3. Strictly, carbonate alkalinity counts carbonate twice while carbonate hardness counts it once, but below pH 8.5 the difference is under 1 %, so in aquarium practice a KH reading is effectively a carbonate alkalinity reading.

In water with moderate to high alkalinity, pH sits near neutral to slightly basic (7.0 to 8.3) and does not fluctuate widely: more CO2 or other acid is needed to move pH because more base is available to absorb it. In poorly buffered water, respiration at night can push pH dangerously low, and intense photosynthesis by day can push it dangerously high. Plants and algae remove CO2 and can split bicarbonate to obtain CO2, releasing carbonate, which reacts with water to form hydroxide; in dense blooms this can drive pH above 9. Adding CO2 lowers pH without changing alkalinity, which is why a tank's KH stays steady while its pH cycles daily with light and dark.

Where alkalinity comes from and how it is lost

In nature, CO2-bearing rain and groundwater dissolve calcitic limestone (CaCO3) or dolomite (CaMg(CO3)2), producing calcium and magnesium bicarbonate; this raises alkalinity, pH and hardness together. Water draining granite catchments carries far less alkalinity. Hardness (calcium and magnesium) and alkalinity are separate measurements that are often confused because both are reported as mg/L CaCO3: where sodium bicarbonate supplies the alkalinity, water can have low hardness and high alkalinity. Above pH 8.3, calcium and photosynthetically produced carbonate can precipitate as limestone, removing alkalinity from the water.