Carbon Dioxide and Carbonic Acid

The most common source of acidity in water is dissolved carbon dioxide.

Carbon dioxide enters the water through equilibrium with the atmosphere

CO2 (aq) CO2 (g)

and biological degradation/photosynthesis involving organic carbon, {CH2O}

{CH2O} + O2(aq) CO2 (aq) + H2O

Aqueous CO2 (aq) also undergoes a number of important inorganic equilibrium reactions. First, it can dissolve limestone

CaCO3 + CO2 (aq) + H2O Ca2+(aq) + 2 HCO3- (aq)

Second, it can react with the water to form carbonic acid

CO2 (aq) + H2O H2CO3 (aq)

Only a small fraction exists as the acid

and the kinetics to form H2CO3 are relatively slow (on the time scale of seconds).

Carbon Dioxide and Carbonic Acid-Base Equilibria

Dissolved CO2 in the form of H2CO3 may loose up to two protons through the acid equilibria

H2CO3 (aq) H+ (aq) + HCO3- (aq)

HCO3- (aq) H+ (aq) + CO32- (aq)

The equilibrium equations for these are labeled as "1" and "2" hence

To account for the fact that CO2 (aq) is in equilibrium with H2CO3 (aq), the first acid equilibrium is normally given by

The acid equilibrium equations can be solved to give the fraction of carbonates in a particular form.

Relative H2CO3 concentration is really CO2 (aq) in equilibrium with water.

In summary;
bulletCO2 enters water through interface with the atmosphere and the biological processes of organic carbon digestion and photosynthesis.
bulletAqueous carbon dioxide, CO2 (aq), reacts with water forming carbonic acid, H2CO3 (aq).
bulletCarbonic acid may loose protons to form bicarbonate, HCO3- , and carbonate, CO32-. In this case the proton is liberated to the water, decreasing pH.
bulletThe complex chemical equilibria are described using two acid equilibrium equations.
bulletThe first acid equilibrium constant accounts for the CO2 (aq) - H2CO3 (aq) equilibrium. It concequently seems to have a high pKa.
bulletThe fraction of the inorganic carbon in a particular form is call the "alpha" and there are simple equation to describe this alpha.

Graphical Results

Consequences of the fractional amount or "alpha" equations may be understood by examination of the graphical results. Shown to the left is a plot of the various alpha as a function of pH.

Some important points to observe are:

For pH well below pKa1 a (H2CO3) ~ 1

At pH = pKa1, a (H2CO3)= a (HCO3-)

For 7<pH<10, HCO3- is the predominant species

The pH of water (no lime)

We now have enough information to calculate the pH of water.

First, we calculate the amount of CO2 dissolved in water under an atmosphere of pressure from Henry’s Law

Since CO2 makes up 0.0355% of the atmosphere (on the average) and

KCO2 =2x10-3

Since is in equilibrium with H2CO3 (aq), the first acid equilibrium is normally given by

is predominant. Also since

CO2 (aq) + H2O H+ (aq) + HCO3- (aq)

The proton and bicarbonate concentrations are equal. Thus

When we substitute the carbon dioxide concentration, and solve for pH, we get

pH = 5.65

Since rain is in equilibrium with the atmosphere, this is the pH expected for natural rain. It is also the pH expected if the body of water is in equilibrium with the atmosphere, and does not contact limestone (e.g., CaCO3).

This page edited Thursday, December 21, 2006

This page was last edited Thursday, December 21, 2006