Carbon Dioxide and Carbonic Acid

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) + CO₃^2- (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;

CO₂ enters water through interface with the atmosphere and the biological processes of organic carbon digestion and photosynthesis.

Aqueous carbon dioxide, CO₂ (aq), reacts with water forming carbonic acid, H₂CO₃(aq).

Carbonic acid may loose protons to form bicarbonate, HCO₃^-, and carbonate, CO₃^2-. In this case the proton is liberated to the water, decreasing pH.

The complex chemical equilibria are described using two acid equilibrium equations.

The first acid equilibrium constant accounts for the CO₂ (aq) - H₂CO₃(aq) equilibrium. It concequently seems to have a high pK_a.

The 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 pK_a1 a (H₂CO₃) ~ 1

At pH = pK_a1, 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

K_CO2 =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