Sulfur Chemistry and Acid Deposition

Introduction

Unpolluted rain is already slightly acidic because it dissolves carbon dioxide, giving a pH of about 5.6. In the second half of the twentieth century, rain in parts of Europe and North America regularly fell with a pH of 4.0–4.5, and fog readings were lower still. The extra acidity came overwhelmingly from sulfur dioxide and nitrogen oxides released by burning fuels. This page follows sulfur from emission to sulfuric acid, adds nitric acid, and examines why acid deposition damaged some landscapes badly while leaving others almost untouched.

Core explanation

Sources of SO₂. Coal and heavy fuel oil contain sulfur, typically 0.5–3% by mass in coal, in the form of pyrite (FeS₂) and organic sulfur compounds. Combustion converts it to SO₂. Smelting of sulfide ores (for example Cu₂S, ZnS) is another large source. Natural sources include volcanoes and the oxidation of dimethyl sulfide released by marine phytoplankton. SO₂ is only moderately soluble and is not itself a strong acid, so the key chemistry is its oxidation from sulfur(IV) to sulfur(VI).

Gas-phase oxidation. The hydroxyl radical, the main tropospheric oxidant, adds to SO₂:

OH + SO₂ + M → HOSO₂ + M HOSO₂ + O₂ → HO₂ + SO₃ SO₃ + H₂O → H₂SO₄

M is a third body that removes excess energy. The HO₂ produced can regenerate OH through reaction with NO, so the process is chain-like rather than consuming the oxidant permanently. Gas-phase oxidation is slow — roughly 1% of the SO₂ per hour in sunny conditions — and it depends on daylight because OH is photochemically produced.

Aqueous-phase oxidation. In cloud and fog droplets, SO₂ dissolves and establishes the equilibria SO₂(aq) + H₂O ⇌ HSO₃⁻ + H⁺ and HSO₃⁻ ⇌ SO₃²⁻ + H⁺. Dissolved S(IV) is then oxidised, chiefly by hydrogen peroxide and by ozone:

HSO₃⁻ + H₂O₂ → HSO₄⁻ + H₂O (overall)

The H₂O₂ route is especially important because its rate is almost independent of pH: the fall in S(IV) solubility as the droplet acidifies is compensated by acid catalysis. Ozone oxidation, by contrast, slows sharply as pH falls, so it is self-limiting. Transition metals such as Fe(III) and Mn(II) can also catalyse oxidation by O₂. Globally, more than half of sulfate forms in clouds, which is why sulfate production can be fast even at night or in winter.

Nitric acid. Nitrogen oxides from high-temperature combustion are oxidised in the daytime by OH + NO₂ + M → HNO₃ + M. At night, NO₃ radicals and N₂O₅ form, and N₂O₅ hydrolyses on wet particles to give HNO₃. Nitric acid is very soluble and is removed rapidly.

Deposition. Acids reach the surface by wet deposition (dissolved in rain, snow and cloud water) and dry deposition (SO₂, HNO₃ and particles absorbed directly by vegetation, soils and buildings). Because atmospheric lifetimes of SO₂ and sulfate are days, acid can be deposited hundreds to more than a thousand kilometres from its source, making acid rain a transboundary problem.

Effects. Where soils are thin and derived from granite or other slowly weathering rock, there is little buffering. Acid inputs leach base cations (Ca²⁺, Mg²⁺) from soil, mobilise aluminium as Al³⁺ and related species, and acidify lakes and streams. Dissolved aluminium damages fish gills, and many fish populations collapse below about pH 5. Forests suffer nutrient loss and aluminium stress on roots. Limestone and marble buildings dissolve as CaCO₃ + H₂SO₄ → CaSO₄ + H₂O + CO₂, and the gypsum crust formed spalls away.

Step-by-step reasoning

To predict whether a region will suffer from acid deposition:

1. Identify upwind sources of SO₂ and NOₓ and the prevailing wind direction. 2. Estimate transport distance from atmospheric lifetimes of several days. 3. Decide whether oxidation is mainly in the gas phase or in clouds, based on cloudiness and oxidant supply. 4. Assess the receiving soils and bedrock: carbonate-rich or deep soils neutralise acid; thin granite soils do not. 5. Combine loading with buffering capacity to judge the risk to lakes and forests.

Visual explanation

Sketch a power station on the left with a plume drifting right. Label SO₂ and NOₓ in the plume, then a sunny clear-air zone marked "OH oxidation, slow", and a cloud marked "H₂O₂ and O₃ oxidation, fast". Arrows descend as rain (wet deposition) and directly onto trees (dry deposition). On the ground, draw a limestone valley with a healthy lake and a granite upland with an acidified, fishless lake.

Real-world analogy

Acid deposition resembles a spillage onto two floors: one covered with an absorbent mat and one bare tile. The same amount of liquid lands on each, but the mat soaks it up, while the tile shows every drop. Carbonate soils act like the mat, neutralising acid, whereas thin granite soils behave like bare tile, passing acid straight into lakes and streams.

Real-world example

Emission controls under the UK Clean Air Acts, the European sulfur protocols and the US Clean Air Act Amendments of 1990 cut sulfur emissions by well over 80% in many countries since 1980. Flue-gas desulfurisation, low-sulfur fuels and the decline of coal were responsible. Rain has become markedly less acidic, and many Scandinavian and Scottish lakes show chemical recovery, although biological recovery is slower.

Why?

Why is sulfate formed in cloud droplets much faster than in clear air? The droplet concentrates dissolved S(IV) and oxidants such as H₂O₂ into a small volume, and aqueous ionic reactions have low barriers. The rate of the H₂O₂ reaction is almost insensitive to pH, so it continues even as the droplet acidifies, oxidising sulfur within minutes to hours.

Common misconception

"Acid rain is any rain with a pH below 7." Rain in equilibrium with atmospheric CO₂ naturally has a pH near 5.6, so it is not neutral. Acid rain is precipitation with pH significantly below about 5.6, caused by strong acids such as H₂SO₄ and HNO₃ rather than carbonic acid.

Worked example

Question: A rain sample contains 25 µmol L⁻¹ of sulfate from fully dissociated H₂SO₄ and no other significant acids or bases. Estimate its pH.

Reasoning: Each H₂SO₄ releases two H⁺ ions, so [H⁺] ≈ 2 × 25 × 10⁻⁶ = 5.0 × 10⁻⁵ mol L⁻¹. The contribution of dissolved CO₂ (about 2.5 × 10⁻⁶ mol L⁻¹) is small, and carbonic acid dissociation is further suppressed by the strong acid. pH = −log(5.0 × 10⁻⁵) = 4.30.

Answer: pH ≈ 4.3, well into the acid rain range.

Quick check

1. Why does SO₂ oxidation by ozone in cloud droplets slow down as the droplet becomes more acidic? Answer: Lower pH shifts S(IV) towards less-reactive dissolved SO₂ and reduces its solubility, so the ozone pathway is self-limiting.

Exam focus

Be ready to write the three-step OH oxidation of SO₂ and the overall aqueous H₂O₂ reaction, to distinguish wet from dry deposition, and to explain why lakes on granite are more sensitive than those on limestone. Quote the natural rain pH of about 5.6 with its reason.

Advanced insight

Reducing sulfate aerosol has a climate side effect. Sulfate particles scatter sunlight and brighten clouds, exerting a cooling influence. Cleaning up SO₂ therefore removes part of this "mask" on greenhouse warming, and studies attribute part of the recent acceleration in warming over Europe and East Asia to lower sulfate loadings. Air-quality and climate policy must be considered together.

Summary

SO₂ from fossil fuel and smelting is oxidised to H₂SO₄ slowly by OH in the gas phase and rapidly by H₂O₂ and O₃ in cloud droplets. NOₓ is oxidised to HNO₃. These strong acids lower precipitation pH below the natural value of about 5.6 and reach the surface by wet and dry deposition far from sources. Damage depends on buffering: poorly buffered soils lose base cations, release toxic aluminium and acidify lakes, while carbonate stone dissolves.

Practice questions

1. Write equations showing how OH converts SO₂ to H₂SO₄ in the gas phase. Answer: OH + SO₂ + M → HOSO₂ + M; HOSO₂ + O₂ → HO₂ + SO₃; SO₃ + H₂O → H₂SO₄. 2. Explain why unpolluted rain has a pH of about 5.6 rather than 7. Answer: Rain dissolves atmospheric CO₂, forming carbonic acid, which partly dissociates to give H⁺ and HCO₃⁻. 3. Why does acid deposition release aluminium into lake water, and why does this matter? Answer: Acid dissolves aluminium-bearing minerals once base cations are depleted, releasing Al³⁺ and hydroxo species; these are toxic to fish, damaging their gills. 4. Distinguish wet and dry deposition, giving one example of each. Answer: Wet deposition removes acid dissolved in rain, snow or fog, such as sulfate in raindrops; dry deposition is direct uptake of gases or particles by surfaces, such as SO₂ absorbed by leaves. 5. Explain why acid rain is described as a transboundary pollution problem. Answer: SO₂, NOₓ and their acid products survive for days in the atmosphere, so they are carried hundreds of kilometres and deposited in countries other than those that emitted them.