Acid Rain and Its Effects

Introduction

All rain is slightly acidic. Carbon dioxide dissolves in raindrops to form weak carbonic acid, giving clean rain a pH of about 5.6. But in the twentieth century, scientists found rain in parts of Europe and North America with a pH of 4 or even lower. This acid rain killed fish, damaged forests and ate away at stone buildings. Its causes are acidic oxides from burning fuels, and its cures are neutralisation chemistry on an industrial scale.

Core explanation

Sulfur dioxide. Coal and crude oil contain small amounts of sulfur compounds. When these fuels burn, the sulfur is oxidised to sulfur dioxide:

S(s) + O₂(g) → SO₂(g)

Volcanoes also release SO₂, but in industrial regions burning fossil fuels was the main source. In the air, SO₂ is slowly oxidised further (helped by sunlight and catalysts such as dust particles) and dissolves in water droplets. The overall result is sulfuric acid, H₂SO₄, a strong acid:

2SO₂(g) + O₂(g) + 2H₂O(l) → 2H₂SO₄(aq)

SO₂ dissolving directly in water also gives the weaker sulfurous acid, H₂SO₃.

Nitrogen oxides. Air is about 78% nitrogen. Nitrogen is normally unreactive, but at the very high temperatures inside car engines and furnaces it combines with oxygen to form nitrogen monoxide, NO, which is quickly oxidised to nitrogen dioxide, NO₂. Together these are called NOₓ. Nitrogen dioxide reacts with water and oxygen to form nitric acid, HNO₃, another strong acid:

4NO₂(g) + O₂(g) + 2H₂O(l) → 4HNO₃(aq)

The fuel does not need to contain nitrogen: the nitrogen comes from the air.

Effects on lakes. Acid rain lowers the pH of lakes and streams, especially those on granite rock that cannot neutralise acid. Below about pH 5, many fish, snails and insect larvae cannot reproduce. Acid water also dissolves aluminium compounds from surrounding soil; aluminium ions damage fish gills.

Effects on forests. Acid rain leaches calcium and magnesium ions from soil and releases toxic aluminium, weakening trees. Acidic mist and fog damage leaves and needles directly. Weakened trees are more vulnerable to frost, drought and disease.

Effects on buildings and metals. Limestone and marble are both calcium carbonate, which reacts with acids. Statues lose their fine detail, and carvings crumble. Metal structures such as iron bridges and railings corrode faster in acidic conditions.

Reducing acid rain. Solutions attack the sources: removing sulfur from fuels before burning; flue gas desulfurisation in power stations, where limestone or lime slurry reacts with SO₂; catalytic converters in cars that reduce NOₓ to nitrogen; and switching to renewable energy. Lakes can be treated short-term by adding powdered limestone.

Formulae

Acid rain on limestone: CaCO₃(s) + H₂SO₄(aq) → CaSO₄(s) + H₂O(l) + CO₂(g)

Acid rain on iron: Fe(s) + H₂SO₄(aq) → FeSO₄(aq) + H₂(g)

Flue gas desulfurisation (simplified): CaCO₃(s) + SO₂(g) → CaSO₃(s) + CO₂(g)

Step-by-step reasoning

Tracing acid rain from power station to statue:

1. Coal containing sulfur burns, producing SO₂. 2. SO₂ travels with the wind, sometimes hundreds of kilometres. 3. SO₂ is oxidised and dissolves in cloud droplets, forming sulfuric acid. 4. Acid rain falls and reacts with calcium carbonate in a statue. 5. Calcium sulfate forms and is washed away, and the stone surface is lost.

Visual explanation

Picture a diagram: a chimney and a car exhaust on the left release SO₂ and NOₓ. Arrows carry the gases up into a cloud, labelled with H₂SO₄ and HNO₃. Rain falls on the right onto a lake with dead fish, a forest with bare trees and a marble statue whose face is worn smooth. Each arrow is labelled with the key equation.

Real-world analogy

Acid rain is like pollution that travels by post. A factory chimney "posts" its sulfur dioxide into the sky, and the wind delivers it as acid to a distant country that never burned the fuel. This is why acid rain became an international problem needing international agreements.

Real-world example

In the 1970s and 1980s, thousands of lakes in southern Norway and Sweden lost their fish populations, largely because of sulfur emissions carried from industrial areas of Britain and central Europe. After European emission controls, SO₂ emissions fell by more than 80%, and many lakes have partly recovered.

Why?

Why do some lakes suffer more than others from the same acid rain? Lakes on limestone or chalk are surrounded by calcium carbonate, which neutralises the acid as rain drains into them. Lakes on granite or sandstone have no natural base to neutralise the acid, so their pH falls.

Common misconception

"Acid rain is caused by carbon dioxide." Carbon dioxide makes all rain slightly acidic (about pH 5.6), but true acid rain is caused mainly by sulfur dioxide and nitrogen oxides, which form strong acids. Carbon dioxide is more important for climate change.

Worked example

Question: A lake had pH 6.5. After years of acid rain its pH is 4.5. By what factor has the hydrogen ion concentration increased?

Reasoning: Each fall of one pH unit multiplies [H⁺] by 10. The pH has fallen by 2 units, so [H⁺] has increased by 10 × 10 = 100.

Answer: The H⁺ concentration is 100 times greater.

Quick check

1. Which acid forms in the atmosphere from nitrogen dioxide? Answer: Nitric acid, HNO₃.

Exam focus

Examiners expect you to name both SO₂ and NOₓ, explain their sources (sulfur impurities in fuels; nitrogen from air at high temperature), and give effects on living things and buildings. Link each solution to the pollutant it removes: desulfurisation and low-sulfur fuel for SO₂, catalytic converters for NOₓ.

Advanced insight

Liming a lake treats the symptom, not the cause, and must be repeated because fresh acid keeps arriving. It can also harm plants adapted to soft water. The long-term fix is reducing emissions. Interestingly, the gypsum (CaSO₄) produced by power-station desulfurisation is sold to make plasterboard, turning a pollutant into a useful product.

Summary

Clean rain has pH about 5.6 due to dissolved CO₂. Acid rain forms when SO₂ from sulfur in fuels and NOₓ from high-temperature combustion become sulfuric and nitric acids. It acidifies lakes, harms fish and trees by releasing aluminium and leaching nutrients, and corrodes limestone, marble and metals. It is reduced by desulfurisation, low-sulfur fuels and catalytic converters.

Practice questions

1. Explain why burning a fuel that contains no nitrogen can still produce nitrogen oxides. Answer: The nitrogen comes from the air; at the high temperatures in engines and furnaces, nitrogen and oxygen from air combine to form NO and NO₂. 2. Write a balanced equation for sulfuric acid reacting with a marble statue. Answer: CaCO₃ + H₂SO₄ → CaSO₄ + H₂O + CO₂ 3. Suggest why acid rain can kill fish even when the water is not acidic enough to harm them directly. Answer: Acidic water dissolves aluminium from soils and rocks, and aluminium ions damage fish gills. 4. Name a method of reducing SO₂ emissions from a power station and state the type of reaction involved. Answer: Flue gas desulfurisation, in which a base such as calcium carbonate reacts with the acidic gas SO₂; it is a neutralisation reaction.