Upper secondary · 15 min

Acid rain and the pH of lakes

Trace $\ce{SO2}$ from a smokestack to an acidified lake, then see how limestone neutralises the acidity.

Goal

Explain why clean rain has pH ≈ 5.6, why polluted rain reaches pH ≈ 4, and why geology decides which lakes suffer.

Apparatus and reagents

Schematic diagram with two scenarios: acid formation and deposition, then a limestone-buffered or limed lake.

Procedure

Open the first scenario and follow the arrows from the stack to the lake: $\ce{SO2}$ is emitted, oxidised to $\ce{H2SO4}$ , and falls as acid rain.
Note the pH values on the rain node: compare polluted rain with clean rain.
Switch to the second scenario: follow how $\ce{CaCO3}$ in the lake bed or added $\ce{Ca(OH)2}$ neutralises the acid.

What to observe

Rain in contact with clean air is already slightly acidic (pH ≈ 5.6) because $\ce{CO2}$ dissolves to give carbonic acid.
Lakes on granite or sand acidify easily; lakes over limestone resist because $\ce{CaCO3 + H+ -> Ca^{2+} + HCO3^-}$ consumes the incoming acid.

Explanation

Burning coal and oil releases $\ce{SO2}$ and $\ce{NO_{x}}$ ; in clouds $\ce{SO2}$ is oxidised (by $\ce{O3}$ , $\ce{H2O2}$ , radicals) to $\ce{H2SO4}$ and $\ce{NO_{x}}$ to $\ce{HNO3}$ , pulling rain pH to 4 or lower. Because pH is logarithmic, a lake falling from 5.6 to 4.2 has become about 25 times more acidic — fish eggs fail and $\ce{Al^{3+}}$ leaches from soil. Limestone and liming restore the carbonate buffer: $\ce{HCO3^-}/\ce{CO2}$ holds pH near 6–7.

Chemists behind it

Svante Arrhenius