In any aqueous solution containing carbonate, the carbonate is distributed among three species: carbonic acid (H₂CO₃), hydrogen carbonate ions (HCO₃^-), and carbonate ions (CO₃^2-). This equilibrium is dependent on the pH of the solution; as the pH increases, the relative percentage of species will shift towards CO₃^2- and away from HCO₃^- and H₂CO₃. However, as the pH decreases, the equilibrium will shift towards H₂CO₃ with little to no carbon existing as CO₃^2-.

To show that increased ocean acidity reduces carbonate ion concentration, some fairly complex calculations are required. However, the evidence that increasing carbon dioxide concentration also lowers the percent of total species present as carbonate ion can be illustrated by applying Le Chatelier's principle to an additional equilibrium occurring in the ocean.

As carbon dioxide dissolves, the equilibrium shifts away from carbonate to hydrogen carbonate ions. This process causes the water to become less saturated with calcium carbonate. As the pH and concentration of dissolved carbon dioxide shift, so do the calcium carbonate saturation horizons.

Open up the pH vs. CO₂ learning tool and click "Show Graph". Use the CO₂ slider to explore how atmospheric CO₂ concentrations affect CO₃^2-, HCO ₃^-, and H ₂CO₃ concentrations.

Question: Use the CO₂ slider to change the atmospheric CO₂ to 720 ppm. At this pH, what percentage of the total carbon exists in the form of carbonate? What RCP projection is closest to this situation?