9.10: Carbon Dioxide and pH

Oxygen and carbon dioxide are involved in the same biological processes in the ocean, but in opposite ways; photosynthesis consumes CO₂ and produces O₂, while respiration and decomposition consume O₂ and produce CO₂. Therefore it should not be surprising that oceanic CO₂ profiles are essentially the opposite of dissolved oxygen profiles (Figure \(\PageIndex{1}\)). At the surface, photosynthesis consumes CO₂ so CO₂ levels remain relatively low. In addition, organisms that utilize carbonate in their shells are common near the surface, further reducing the amount of dissolved CO₂.

In deeper water, CO₂ concentration increases as respiration exceeds photosynthesis, and decomposition of organic matter adds additional CO₂ to the water. As with oxygen, there is often more CO₂ at depth because cold bottom water holds more dissolved gases, and high pressures increase solubility. Deep water in the Pacific contains more CO₂ than the Atlantic as the Pacific water is older and has accumulated more CO₂ from the respiration of benthic organisms.

Figure \(\PageIndex{1}\) Representative carbon dioxide profiles for the Pacific and Atlantic oceans (PW).

But the behavior of carbon dioxide in the ocean is more complex than the figure above would suggest. When CO₂ gas dissolves in the ocean, it interacts with the water to produce a number of different compounds according to the reaction below:

CO₂ + H₂O ↔ H₂CO₃ ↔ H⁺ + HCO₃^– ↔ 2H⁺ + CO₃^2-

CO₂ reacts with water to produce carbonic acid (H₂CO₃), which then dissociates into bicarbonate (HCO₃^–) and hydrogen ions (H⁺). The bicarbonate ions can further dissociate into carbonate (CO₃^2-) and additional hydrogen ions (Figure \(\PageIndex{2}\)).

Figure \(\PageIndex{2}\) The fate of dissolved carbon dioxide in the oceans. Most of the carbon ends up in the form of bicarbonate (PW).

Most of the CO₂ dissolving or produced in the ocean is quickly converted to bicarbonate. Bicarbonate accounts for about 92% of the CO₂ dissolved in the ocean, and carbonate represents around 7%, so only about 1% remains as CO₂, and little gets absorbed back into the air. The rapid conversion of CO₂ into other forms prevents it from reaching equilibrium with the atmosphere, and in this way, water can hold 50-60 times as much CO₂ and its derivatives as the air.

CO₂ and pH

The equation above also illustrates carbon dioxide’s role as a buffer, regulating the pH of the ocean. Recall that pH reflects the acidity or basicity of a solution. The pH scale runs from 0-14, with 0 indicating a very strong acid, and 14 representing highly basic conditions. A solution with a pH of 7 is considered neutral, as is the case for pure water. The pH value is calculated as the negative logarithm of the hydrogen ion concentration according to the equation:

pH = -log₁₀[H⁺]

Therefore, a high concentration of H⁺ ions leads to a low pH and acidic condition, while a low H⁺ concentration indicates a high pH and basic conditions. It should also be noted that pH is described on a logarithmic scale, so every one point change on the pH scale actually represents an order of magnitude (10 x) change in solution strength. So a pH of 6 is 10 times more acidic than a pH of 7, and a pH of 5 is 100 times (10 x 10) more acidic than a pH of 7.

Carbon dioxide and the other carbon compounds listed above play an important role in buffering the pH of the ocean. Currently, the average pH for the global ocean is about 8.1, meaning seawater is slightly basic. Because most of the inorganic carbon dissolved in the ocean exists in the form of bicarbonate, bicarbonate can respond to disturbances in pH by releasing or incorporating hydrogen ions into the various carbon compounds. If pH rises (low [H⁺]), bicarbonate may dissociate into carbonate, and release more H⁺ ions, thus lowering pH. Conversely, if pH gets too low (high [H⁺]), bicarbonate and carbonate may incorporate some of those H⁺ ions and produce bicarbonate, carbonic acid, or CO₂ to remove H⁺ ions and raise the pH. By shuttling H⁺ ions back and forth between the various compounds in this equation, the pH of the ocean is regulated and conditions remain favorable for life.