68 The Chemistry of Acid Rain
Saylor Academy
Learning Objectives
- To understand the chemistry of acid rain.
Acid–base reactions can have a strong environmental impact. For example, a dramatic increase in the acidity of rain and snow over the past 150 years is dissolving marble and limestone surfaces, accelerating the corrosion of metal objects, and decreasing the pH of natural waters. This environmental problem is called acid rain and has significant consequences for all living organisms. To understand acid rain requires an understanding of acid–base reactions in aqueous solution.
The term acid rain is actually somewhat misleading because even pure rainwater collected in areas remote from civilization is slightly acidic (pH ≈ 5.6) due to dissolved carbon dioxide, which reacts with water to give carbonic acid, a weak acid:
Equation 4.44
[latex]\text{CO}_2 (g) + \text{H}_2\text{O} (l) \rightleftharpoons \text{H}_2\text{CO}_3 (aq) \rightleftharpoons \text{H} ^+ (aq) + \text{HCO}_3^- (aq)[/latex]
The English chemist Robert Angus Smith is generally credited with coining the phrase acid rain in 1872 to describe the increased acidity of the rain in British industrial centers (such as Manchester), which was apparently caused by the unbridled excesses of the early Industrial Revolution, although the connection was not yet understood. At that time, there was no good way to measure hydrogen ion concentrations, so it is difficult to know the actual pH of the rain observed by Smith. Typical pH values for rain in the continental United States now range from 4 to 4.5, with values as low as 2.0 reported for areas such as Los Angeles. Recall that rain with a pH of 2 is comparable in acidity to lemon juice, and even “normal” rain is now as acidic as tomato juice or black coffee.
What is the source of the increased acidity in rain and snow? Chemical analysis shows the presence of large quantities of sulfate ([latex]\text{SO}_4^{2-}[/latex]) and nitrate ([latex]\text{NO}_3^-[/latex]) ions, and a wide variety of evidence indicates that a significant fraction of these species come from nitrogen and sulfur oxides produced during the combustion of fossil fuels. At the high temperatures found in both internal combustion engines and lightning discharges, molecular nitrogen and molecular oxygen react to give nitric oxide:
Equation 4.45
[latex]\text{N}_2 (g) + \text{O}_2 (g) \rightarrow \text{2NO} (g)[/latex]
Nitric oxide then reacts rapidly with excess oxygen to give nitrogen dioxide, the compound responsible for the brown color of smog:
Equation 4.46
[latex]\text{2NO} (g) + \text{O}_2 (g) \rightarrow \text{2NO}_2 (g)[/latex]
When nitrogen dioxide dissolves in water, it forms a 1:1 mixture of nitrous acid and nitric acid:
Equation 4.47
[latex]\text{2NO}_2 (g) + \text{H}_2\text{O} (l) \rightarrow \text{HNO}_2 (aq) + \text{HNO}_3 (aq)[/latex]
Because molecular oxygen eventually oxidizes nitrous acid to nitric acid, the overall reaction is
Equation 4.48
[latex]\text{2N}_2 (g