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9.10: Carbon Dioxide and pH

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    13012
  • 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.

    figure5.5.1b.png
    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+ + CO32-

    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 (CO32-) and additional hydrogen ions (Figure \(\PageIndex{2}\)).

    figure5.5.2.png
    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.

     

     

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