Lesson 15: Environmental Chemistry and Green Chemistry

Chemistry made the modern world and is now being asked to clean it up. Environmental chemistry tracks what our substances do after release — CO₂ warming the air, CFCs opening the ozone hole, SO₂ returning as acid rain — while green chemistry tries to redesign processes so the damage never happens. This lesson supplies the quantitative tools of both: ppm-scale accounting, the carbon arithmetic of fuels, and atom economy, the metric that grades a synthesis by how much of its reactants end up in the product rather than the waste bin.

Air: greenhouse gases, ozone, acid rain

Greenhouse gases (CO₂, CH₄, water vapour, N₂O) absorb infrared radiation re-emitted by Earth's surface; the enhanced greenhouse effect from burning fossil carbon is the driver of current climate change — atmospheric CO₂ has risen from about 280 ppm pre-industrially to about 420 ppm today. Stratospheric ozone does the opposite job, absorbing ultraviolet; CFCs delivered chlorine radicals that catalytically destroy O₃ (one Cl· can destroy ~100,000 O₃ molecules before being removed), which is why the Montreal Protocol phased CFCs out — environmental chemistry's clearest success story. Acid rain begins with SO₂ from sulfur-bearing fuels and NOx from hot engines, oxidising to H₂SO₄ and HNO₃ in cloud water; unpolluted rain is already mildly acidic (pH ≈ 5.6, dissolved CO₂ forming carbonic acid), and acid rain means pH below that, damaging limestone buildings, soils and aquatic life.

Worked example 1: the carbon arithmetic of a fuel

Octane, C₈H₁₈ (Mr = 114.23), is petrol's model molecule: C₈H₁₈ + 12½O₂ → 8CO₂ + 9H₂O. Burning 1.00 kg: n(octane) = 1000 ÷ 114.23 = 8.754 mol; n(CO₂) = 8 × 8.754 = 70.0 mol; mass = 70.0 × 44.01 = 3080 g ≈ 3.08 kg of CO₂ from 1.00 kg of fuel. The product outweighs the fuel because most of its mass is oxygen taken from the air — a result that surprises everyone the first time and that every carbon budget in the world is built on. Scale it: a car burning 40 L of petrol a week (density 0.75 kg/L → 30 kg) emits roughly 92 kg of CO₂ weekly, over 4.7 tonnes a year.

Atom economy

Atom economy = (Mr of desired product ÷ Σ Mr of all reactants) × 100. It is a design metric computed before any experiment: addition reactions and rearrangements score 100%; substitutions and eliminations always waste atoms. Crucially, atom economy is independent of yield — a reaction can run at 100% yield and still discard half its atoms as by-product.

Worked example 2: atom economy, two ways

(a) Limestone calcination: CaCO₃ → CaO + CO₂. If CaO is the product, atom economy = 56.08 ÷ 100.09 × 100 = 56.0% — the CO₂ is waste here (and, awkwardly, a greenhouse emission intrinsic to cement-making: even perfect fuel efficiency cannot remove it). (b) Making ethanol industrially by hydrating ethene: C₂H₄ + H₂O → C₂H₅OH. Atom economy = 46.07 ÷ (28.05 + 18.02) × 100 = 100% — every atom of every reactant is in the product. The older fermentation route, C₆H₁₂O₆ → 2C₂H₅OH + 2CO₂, scores 92.14 ÷ 180.16 = 51.1%. Green chemistry's twelve principles (prevent waste, maximise atom economy, use catalysis, design for degradation, prefer renewable feedstocks…) read like common sense precisely because each was learned from an expensive mistake.

In the lab: how acidic is the rain?

Measuring local rain or simulated acid deposition on marble chips (mass loss over a week) is a staple practical, and its uncertainty profile teaches restraint. A school pH probe, calibrated at pH 4 and 7, reads to ±0.05 — adequate — but rain is very weakly buffered, so the probe reading drifts as the sample absorbs or releases CO₂ while you measure it: the sample changes under observation. Marble-chip mass loss is worse: a week of exposure to pH 4.5 solution removes perhaps 0.02 g — the same size as the balance's ±0.01 g reading uncertainty doubled across two weighings, and meaningless if chips are handled with fingers (grease) or dried inconsistently. Honest design extends exposure, uses several chips averaged, controls chip surface area by sieving to a size fraction, and reports losses with their uncertainty rather than a single brittle number. Much real environmental data shares this character — small signals, patient methods — which is why monitoring networks average many stations over years.

Chemistry in the real world

Catalytic converters finish Lesson 12's story: platinum-group surfaces convert CO and unburned hydrocarbons to CO₂ and H₂O, and NOx back to N₂, within the gas's brief transit. Flue-gas desulfurisation sprays limestone slurry into power-station exhaust, turning SO₂ into saleable gypsum — acid rain chemistry run backwards as an industry.

Common mistakes

• Confusing the ozone layer with the greenhouse effect. Ozone absorbs UV high up; its depletion is a radiation-exposure problem. CO₂ warms by absorbing IR. Different gases, mechanisms, treaties.
• "Natural pH of rain is 7." Dissolved CO₂ makes clean rain pH ≈ 5.6; acid rain is defined against that baseline, not neutrality.
• Atom economy confused with percent yield. Yield measures how much product you got from the limiting reactant; atom economy measures how the reaction design divides atoms between product and by-product — independent numbers (Lessons 4 and this one).
• Treating catalytic ozone destruction as stoichiometric. The Cl· is regenerated; that is why trace CFCs did outsized damage.

Exam Experience

1. Which pair of gases are the principal acid-rain precursors (one of them is also the brown gas visible in polluted city air)? A) CO₂ and CH₄ B) SO₂ and NO₂ C) O₂ and N₂ D) O₃ and CFCs Answer: B. SO₂ oxidises in cloud water to sulfuric acid; NO₂ is itself brown and forms nitric acid. A names greenhouse gases; C names air's inert bulk; D belongs to the ozone-layer story.

2. A reaction has atom economy 40% and percent yield 90%. The fraction of reactant mass converted into desired product is: A) 90% B) 40% C) 36% D) 130% Answer: C. The metrics multiply: of every 100 g of reactants, the design places at most 40 g into product atoms, and the yield realises 90% of that possibility: 0.40 × 0.90 = 0.36 → 36%.

3. Which statement about catalytic converters is correct? A) They remove CO₂ from exhaust B) They use catalysts to convert CO and NOx into CO₂, H₂O and N₂ rapidly at exhaust temperatures C) They work only when the engine is cold D) They increase atom economy of the fuel Answer: B. Converters oxidise CO/hydrocarbons and reduce NOx simultaneously on Pt/Pd/Rh surfaces; they do not capture CO₂ (the carbon leaves as gas), need to be hot to work, and are kinetics devices, not atom-economy ones.

Constructed response (5 marks). A cement works calcines 1000 kg of pure CaCO₃ daily by CaCO₃ → CaO + CO₂. (a) Calculate the mass of CaO produced (Mr: CaCO₃ 100.09, CaO 56.08), assuming complete reaction. (b) Calculate the mass of CO₂ released. (c) The works' fuel adds a further 300 kg of CO₂ daily. Explain why the process CO₂ in (b) is described as "unavoidable" in a way the fuel CO₂ is not. Scoring guide: (a) 1 mark: n = 1,000,000 g ÷ 100.09 = 9991 mol ≈ 1.00 × 10⁴ mol; 1 mark: mass CaO = 9991 × 56.08 = 560 kg (accept 559–561). (b) 1 mark: mass CO₂ = 9991 × 44.01 = 440 kg. (c) 1 mark: the CO₂ is part of the reactant's own atoms — making CaO from CaCO₃ necessarily ejects it (atom economy 56%); 1 mark: fuel CO₂ could in principle be eliminated by switching to carbon-free energy (electric kilns, hydrogen), but the reaction CO₂ requires either a different raw material or carbon capture. Numbers that violate mass balance (a) + (b) ≠ 1000 kg forfeit the (b) mark.

Eutrophication: when nutrients become pollutants

Fertiliser runoff carries nitrate and phosphate — plant nutrients — into waterways, where algae bloom on the feast. The bloom itself is not the injury; its death is. Decomposers oxidising the dead algae consume dissolved oxygen faster than it dissolves (Lesson 7's gas solubility falling as water warms, compounding the crisis), and fish suffocate in water full of the products of growth. The chemistry is two of this course's threads tied together: quantitative nutrient accounting (farmers now match fertiliser moles to crop uptake, Lesson 4 stoichiometry with a catchment attached) and the oxygen equilibrium of Lesson 11. Buffer strips, phosphate-free detergents and precision dosing each attack a different step — input, transport, or consequence — and policy marks them accordingly.

Worked example 3: the CO₂ ledger of a journey

A petrol car emits from its fuel what Worked example 1 computed: 3.08 kg of CO₂ per kilogram of petrol burned. At a consumption of 7.0 L per 100 km (petrol density 0.75 kg/L), the fuel burned per 100 km is 7.0 × 0.75 = 5.25 kg, emitting 5.25 × 3.08 = 16.2 kg of CO₂ per 100 km = 162 g/km — the number on the vehicle's rating label, derived end-to-end from one combustion equation. An electric car on a grid averaging 200 g CO₂ per kWh and using 18 kWh per 100 km emits 3.6 kg per 100 km (36 g/km) — better by 4.5×, and improving automatically as the grid greens, while the petrol figure is frozen by stoichiometry forever. That asymmetry — a chemical constant versus a policy variable — is the deepest environmental chemistry in this lesson, and it needs only Lesson 4's mole ratios to prove.

Carbon capture chemistry: the ledger continues

Post-combustion capture scrubs CO₂ from flue gas into an amine solution: the amine (a base, Lesson 6) binds CO₂ as a carbamate at about 40 °C; heating the rich solvent to 120 °C reverses the binding (Lesson 11's temperature rule), releasing a pure CO₂ stream for compression and geological storage, and regenerating the amine for another cycle. The energy cost of the regeneration step — the "parasitic load", typically 20–30% of a power station's output — is the technology's central problem, and it is thermochemistry, not engineering laziness: the binding must be strong enough to grab CO₂ at 0.04–0.15 atm partial pressure yet weak enough to let go cheaply, a thermodynamic tightrope that new solvents, metal–organic frameworks and Lesson 20's membranes all walk. Mineral carbonation sidesteps the release step entirely by making carbonate rock (Worked example 2's limestone reaction run backwards) — permanent storage at the price of mining and grinding the rock.