| the photosynthesis reaction in green plants: | 6CO2 (g) + 6 H2O (l)  C6H12O6
(aq) + O2 (g) |
[1] |
| the solubility of CO2 in water: | CO2 (g) CO2 (aq) |
[2] |
Carbon dioxide makes up almost 95% of Mars' thin atmosphere. However on Earth it is only 0.035% of the atmosphere. Given the great importance paid to carbon dioxide emissions as a greenhouse gas, you may be surprised to learn that it makes up so little of the Earth's air. Why is there such a large proportion of CO2 in the Martian atmosphere as compared to Earth?
The reactions which regulate the amount of carbon dioxide in the Earth's atmosphere are very complex. However, they are effectively controlled by two processes:
| the photosynthesis reaction in green plants: | 6CO2 (g) + 6 H2O (l) C6H12O6
(aq) + O2 (g) |
[1] |
| the solubility of CO2 in water: | CO2 (g) CO2 (aq) |
[2] |
Reaction [2] is actually considerably more complex than shown for the aqueous CO2 will react with the water to form carbonic acid (H2CO3)
| formation of carbonic acid: | CO2 (aq) + H2O H2CO3 (aq) |
[3] |
Carbonic acid will dissociate in water to form carbonate and bicarbonate ions as shown below in Brönsted-Lowry format
| H2CO3 (aq) + H2O
H3O+ (aq) + HCO3- (aq) |
[4] |
| HCO3- (aq) + H2O
H3O+ (aq) + CO32- (aq) |
[5] |
Overall, equations [3], [4], and [5] mean that
| CO2 (g) + 3H2O (l) 2H3O+(aq) + CO32- (aq) |
[6] |
|
In the carbon cycle, atmospheric CO2 diffuses into lakes and oceans at the surface, where it may be used in photosynthesis by aquatic plants, or remain as dissolved carbonates and bicarbonates. Many oceanic (and also fresh water) calciferous animals such as crustaceans and corals convert carbonate and bicarbonate ions into fairly insoluble calcium carbonate sediments which settle to the ocean bottom. Over time these sediments become compressed into carbonate rocks such as limestone and dolomite. Terrestrial plants use CO2 from the atmosphere directly for photosynthesis. The remains of plants which did this eons ago have become preserved as fossil fuels. These processes form carbon dioxide "sinks" that remove it from the atmosphere.
Decomposition, and respiration return terrestrial and aquatic biomass back into the atmosphere. Man's major contribution to the carbon cycle is to increase the amount of CO2 by burning of fossil fuels.
Since Mars no longer has any liquid water (though it may have once), and no plant life,
the carbon dioxide in its atmosphere cannot be removed in a process equivalent to the
carbon cycle which occurs on earth.
| Carbon Dioxide Sinks | Approximate amount in Gigatonnes |
| Marine Sediments and Sedimentary Rocks | 100 000 000 |
| Ocean | 40 000 |
| Fossil Fuel Deposits | 4000 |
| Soil Organic Matter | 1500 |
| Atmosphere | 750 |
| Terrestrial Plants | 550 |
Both oceanic and atmospheric carbon dioxide amounts are increasing, but the atmospheric amount is increasing more rapidly, about 3.4 Gt per year.
| Warning: this experiment requires you to handle solid carbon dioxide (dry ice). The sublimation temperature of CO2 (s) is -78 ºC, so you cannot touch it with your bare skin. Always use insulated leather gloves, or tongs to handle dry ice. Direct contact with your skin will freeze it solid in just a few seconds. The result is that the cells in the area will die (dry ice doesn't really "burn" you, but the effect of severe frostbite is the same as a third degree burn). Also, never put dry ice into a sealed container. The container will explode, and could cause severe eye or limb damage. |
Part I
1. Add a few drops of methyl red acid-base indicator to about 500 mL of distilled water in a 600 mL beaker. Add bromothymol blue to a second 500 mL of water in another beaker. Note: CO2 from the air will almost always have dissolved into distilled water. The pH of the water will therefore always be slightly acid. To avoid this problem you should boil the water, and then allow it to cool in a sealed container.
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2. Adjust the acidity of the water by adding drop by drop 0.1 M NaOH solution, and stirring. The color of the methyl red solution should be adjusted until it is yellow. The bromothymol blue solution should be blue. Observe the color of the indicators in an (approximately) neutral solution.
3. Break off a small piece of solid dry ice (about 1 to 2 cm across).
Using tongs or gloves, lower it into one of the beakers. Repeat the procedure for
the other beaker. Observe the results. Note: since the water is far above the
temperature of solid dry ice, it immediately turns into CO2 (g). Thus the
reactions you are observing are due to the presence of carbon dioxide gas.
Part II
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4. Place about 50 mL of a saturated solution of calcium hydroxide (limewater) into a
100 mL beaker. Add a small pea-sized piece of dry ice. Observe the results.
1. Using an acid-base color chart, identify from the color changes what happens to the pH of the water when the dry ice dissolves.
2. Write out an answer which explain the indicator color change in terms of the equations given in the introduction to this lab.
3. Solution of calcium hydroxide, Ca(OH)2, are used as a simple test for the presence of carbon dioxide. The Ca2+ ions react with the CO32- ions that form from the CO2 (g) dissolving in the water. The following represents the net ionic equation for the formation of the precipitate of calcium carbonate:
Ca2+ (aq) + CO32- (aq) CaCO3 (s)
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Using le Châtelier's principle, and making use of the equations given in the introduction, write out an answer which explains how the addition of CO2 (g) to the water causes the formation of a solid precipitate. Note: the product is the same low solubility calcium carbonate, but the mechanism of its production is not that used by calciferous animals in the oceans.
