Pollution

Environmental chemistry examines how human activities release pollutants into the atmosphere, hydrosphere, and lithosphere, and how those pollutants undergo chemical transformations that harm ecosystems and human health. The atmosphere is divided into layers; the troposphere (0–12 km) is where we live and where weather occurs, while the stratosphere (12–50 km) contains the ozone layer that shields Earth from ultraviolet radiation. Different pollutants cause distinct problems in each layer. Tropospheric pollution manifests as photochemical smog and acid rain. Smog forms when nitrogen oxides (NOₓ) and volatile organic compounds (VOCs) react under sunlight to produce ground-level ozone and peroxyacetyl nitrate (PAN), both powerful irritants. Acid rain results when sulfur dioxide (SO₂) and nitrogen dioxide (NO₂) dissolve in atmospheric water to form sulfurous, sulfuric, and nitric acids, lowering precipitation pH well below the natural value of 5.6 caused by dissolved CO₂. Stratospheric ozone depletion is driven by chlorofluorocarbons (CFCs) and other halogenated compounds that release chlorine radicals capable of catalytically destroying thousands of ozone molecules per radical pair. In response to these interconnected crises, the philosophy of Green Chemistry emerged, proposing twelve principles to design chemical processes that prevent pollution at the source rather than treating it after the fact. Understanding the chemistry behind each problem is essential for designing effective solutions.

Key takeaways
  • Formation of Photochemical Smog
  • Acid Rain Chemistry
  • Stratospheric Ozone Depletion — The Chapman and Rowland–Molina Mechanisms
  • Green Chemistry — Prevention Over Cure

Formation of Photochemical Smog

At high temperatures in internal combustion engines, atmospheric nitrogen and oxygen react: N₂ + O₂ → 2 NO. Nitric oxide (NO) is then oxidised to nitrogen dioxide: 2 NO + O₂ → 2 NO₂. Under UV sunlight, NO₂ photodissociates — NO₂ + hν → NO + O• — producing highly reactive oxygen radicals. These radicals combine with O₂ to form ground-level ozone: O• + O₂ → O₃. Ozone reacts with VOCs (hydrocarbons from unburned fuel) to produce aldehydes and peroxyacetyl nitrate (PAN, CH₃CO·O₂NO₂), the brownish haze characteristic of photochemical smog. PAN is a lachrymator (causes eye irritation) and phytotoxin.

Acid Rain Chemistry

Combustion of sulfur-containing coal and oil releases SO₂. In the atmosphere, SO₂ is oxidised: 2 SO₂ + O₂ → 2 SO₃, catalysed by particulate matter or in the presence of hydroxyl radicals (•OH + SO₂ → HOSO₂•, further oxidised to H₂SO₄). Sulfur trioxide dissolves readily in water droplets: SO₃ + H₂O → H₂SO₄. Similarly, NO₂ reacts with water: 3 NO₂ + H₂O → 2 HNO₃ + NO. The resulting acid precipitation has a pH as low as 2–4 in heavily industrialised regions. Acid rain leaches aluminium ions from soil, damages leaf cuticles, corrodes limestone (CaCO₃ + H₂SO₄ → CaSO₄ + H₂O + CO₂), and acidifies lakes.

Stratospheric Ozone Depletion — The Chapman and Rowland–Molina Mechanisms

Normally, ozone in the stratosphere is maintained in a steady state by the Chapman cycle: O₂ + hν → 2 O•; O• + O₂ → O₃; O₃ + hν → O₂ + O•; O• + O₃ → 2 O₂. CFCs (e.g., CCl₂F₂, Freon-12) are inert in the troposphere but in the stratosphere UV radiation cleaves the C–Cl bond: CCl₂F₂ + hν → •CClF₂ + Cl•. The chlorine radical destroys ozone catalytically: Cl• + O₃ → ClO• + O₂; ClO• + O• → Cl• + O₂. Net: O₃ + O• → 2 O₂. Crucially, Cl• is regenerated and can destroy ~100,000 ozone molecules before being deactivated. Bromine from halons is even more destructive per atom.

Green Chemistry — Prevention Over Cure

Green Chemistry, formalised by Paul Anastas and John Warner in 1998, provides 12 principles including: (1) Prevent waste rather than treat it; (2) Design atom-efficient reactions (atom economy = MW of desired product / sum of MW of all reactants × 100%); (3) Use renewable feedstocks; (4) Design for degradation; (5) Use catalysts instead of stoichiometric reagents to reduce by-products. A classic metric is atom economy: for the hydrogenation of propene (CH₂=CHCH₃ + H₂ → CH₃CH₂CH₃), atom economy = 44/46 × 100% ≈ 95.7%, showing nearly all atoms end up in the product. Contrast this with a Grignard reaction where MgX by-products are discarded (atom economy << 100%).

The Key Equation

\ceCl.+O3−>ClO.+O2then\ceClO.+O.−>Cl.+O2\ce{Cl^{.} + O3 -> ClO^{.} + O2}\quad\text{then}\quad\ce{ClO^{.} + O^{.} -> Cl^{.} + O2}
  • \ce{Cl^{.}} — Chlorine radical (dimensionless): Highly reactive species with an unpaired electron, released from CFCs by UV photolysis in the stratosphere.
  • \ce{O3} — Ozone (molecules cm⁻³): Triatomic oxygen forming the stratospheric shield that absorbs UV-B and UV-C radiation.
  • \ce{ClO^{.}} — Chlorine monoxide radical (dimensionless): Intermediate radical; reacts with atomic oxygen to regenerate Cl• and continue the catalytic cycle.
  • \ce{O^{.}} — Oxygen atom (atomic oxygen) (dimensionless): Produced by UV photolysis of O₂ and O₃ in the stratosphere; reacts with ClO• to regenerate Cl•.

Worked Example

Problem. A coal-fired power station burns coal containing 2.5% sulfur by mass. If 1000 kg of coal is burned completely, calculate: (a) the mass of SO₂ produced, and (b) the mass of H₂SO₄ that could form if all SO₂ is oxidised and hydrated. (Relative atomic masses: S = 32, O = 16, H = 1)

m(\ceSO2)=50 kg;m(\ceH2SO4)≈76.6 kgm(\ce{SO2}) = 50\,\text{kg};\quad m(\ce{H2SO4}) \approx 76.6\,\text{kg}

Answer. 50 kg of SO₂ and approximately 76.6 kg of H₂SO₄ could be produced.

Figure
Diagram
Tropospheric Smog vs Stratospheric Ozone DepletionPhotochemical Smog(Troposphere, 0–12 km)① N₂ + O₂ → 2 NO (engine heat)② 2 NO + O₂ → 2 NO₂③ NO₂ + hν → NO + O•④ O• + O₂ → O₃ (ground level!)⑤ O₃ + VOC → PAN + aldehydesNO₂brownhνO•radical+O₂O₃Effects: eye irritation,respiratory damage,crop damage, smog hazeOzone Depletion(Stratosphere, 12–50 km)① CCl₂F₂ + hν → •CClF₂ + Cl•② Cl• + O₃ → ClO• + O₂③ ClO• + O• → Cl• + O₂Net: O₃ + O• → 2O₂Cl• is recycled (catalyst!)Cl•radicalO₃→O₂×100,000Effects: UV-B increase,skin cancer, cataracts,phytoplankton damageMontreal Protocol (1987): CFCs banned
Origins · 1 minF. Sherwood Rowland and how this was discovered
FS
F. Sherwood Rowland
1927–2012 · American

Together with Mario Molina, Rowland proposed in 1974 that chlorofluorocarbons (CFCs) could catalytically destroy stratospheric ozone. They worked out the complete radical chain mechanism showing that Cl• released from CFCs under UV radiation destroys thousands of O₃ molecules per radical. Their predictions were confirmed when the Antarctic ozone hole was discovered in 1985. Rowland and Molina shared the 1995 Nobel Prize in Chemistry with Paul Crutzen for this work. When Rowland's wife asked him how his research was going shortly after the 1974 discovery, he replied, 'The work is going well — but it looks like it might be the end of the world.' He spent decades as a scientific activist pushing governments to ban CFCs.

  • Redox Reactions — The conversion of SO₂ to SO₃ and the formation of NO₂ from NO are oxidation reactions; understanding electron transfer in redox chemistry is essential for acid rain formation mechanisms.
  • Reaction Kinetics and Catalysis — The ozone depletion cycle is a classic example of catalytic radical chain reactions; kinetics explains why Cl• is so destructive despite being present in tiny concentrations.
  • Acids, Bases, and pH — Acid rain is characterised by its low pH; understanding the dissociation of H₂SO₄ and HNO₃ and the pH scale is required to quantify acid rain severity and its effects.
  • Photochemistry and Electromagnetic Radiation — Both smog formation and ozone depletion are initiated by photolysis — the absorption of UV photons to break chemical bonds — making photochemistry central to atmospheric pollution.