Storing Carbon Dioxide

• The solubility of gases decreases as temperature increases; this is illustrated in the following graph which shows the solubility of carbon dioxide in water:
• There is also a relationship between pressure and the solubility of gas; the higher the pressure, the greater the volume of gas dissolves in the solution.
• Carbon dioxide is put into drinks under considerable pressure in order to make them fizzy.
• Therefore a high pressure and low temperature will help keep your can of coke fizzy.
• These conditions also favour carbon dioxide dissolving in the oceans, thus removing it from the atmosphere, helping to maintain a constant environment on Earth.
• The exchange of CO₂ between the atmosphere and the ocean takes place rapidly; much faster than if water was left to rest in a glass.
• The uptake of CO₂ by the oceans is made much more efficient by the effects of marine life; for example, CO₂ is taken up by phytoplankton (for photosynthesis), the phytoplankton is then eaten and metabolised by animals which release CO₂ into the water.
  
• Carbon dioxide is much more polar than O₂ and N₂; this is related to the fact that it has polar bonds which can interact with the dipolar water molecules.
• At the surface of the oceans, the following equilibrium is taking place:
  CO_{2 (g)} CO_{2 (aq)}
  
• Some of the carbon dioxide then reacts with water:
  CO_2(aq) + H₂O_(l) H⁺_(aq) + HCO₃^-_(aq)
  HCO₃^-_(aq) H⁺_(aq) + CO₃^2-_(aq)
  
• This removes some of the aqueous carbon dioxide from the water, and thus shifts the position of equilibrium (in the reaction of CO₂ dissolving in water) to the right.
• Much of the carbon dioxide we pump into the atmosphere is absorbed by the oceans in this way (probably around 35-50% is removed in this way).
• The aborption of CO₂ into the oceans is a continual process; the surface water that is rich in CO₂ sinks to the deeps of the oceans, where it is stored for hundreds of years.

Dissolving calcium carbonate

The overall equations listed previously can be combined into one equation:
CO_2(aq) + H₂O_(l) 2H⁺_(aq) + CO₃^2- _(aq)

• According to Le Chatelier’s principle, removing the H⁺ ions will result in more CO₂ dissolving into the ocean.
• One way to remove H⁺ ions would be to add a base; however this is not a very practical proposition where the oceans are concerned.
• There are, however, many marine organisms that are able to remove CO₃^2- ions from sea water; they use the ions to form protective shells of calcium carbonate; CaCO₃.
• Therefore by forming calcium carbonate shells, the small marine organisms are doing their part to combat the increasing levels of carbon dioxide in the atmosphere.
• Billions of years ago, the atmosphere contained much more carbon dioxide; as a result of this, shell production flourished.
• Limestone cliffs are the remains of the calcium carbonate shells formed by the marine organisms billions of years ago.
• Calcium carbonate is insoluble in salt water at the surface of the oceans, and is therefore a useful product for making protective shells.
• However, calcium carbonate does slightly dissolve in pure water:
  CaCO_3(s) + aq Ca²⁺_(aq) + CO₃^2- _(aq)
  
• The solubility product for calcium carbonate is 5 x 10^-9 mol² dm^-6.
• When the Ca²⁺ ions and CO₃^2- ions mix in solution, one of three things can occur:
  1. Calcium carbonate precipitates out of the solution; this occurs when the value of [Ca²⁺] x [CO₃^2-] is greater than the solubility product.
  2. The dissolved ions are in equilibrium with the calcium carbonate; this occurs when [Ca²⁺] x [CO₃^2-] is equal to the solubility product.
  3. The ions remain in solution; this occurs when [Ca²⁺] x [CO₃^2-] is less than the solubility product.
• At the surface of the oceans, the Ca²⁺ and CO₃^2- ions are at high concentrations, and so the calcium carbonate is insoluble.
• However, deeper down in the ocean, the pressure is higher and the temperature is lower; the K_sp is greater under these conditions and calcium carbonate is therefore more soluble.
• So as the calcium carbonate moves deeper down the oceans, it begins to dissolve, forming a marine snow (lots of remains of dead marine creatures).
• Any organic material that reaches the lower regions of the ocean is then consumed by other organism, which release CO₂.
• The extra CO₂ promotes the dissolving of calcium carbonate even further.
• There are no shells on the ocean floor; they have all dissolved by the time they reach the low temperatures and high pressures.
• This means that the calcium carbonate cliffs must have formed in shallow seas; it is likely that the water was also warm and tropical.

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