Atmospheric Chemistry

01 / UNDERSTAND & APPLY

1. Follow a pollutant from source to chemical fate

Atmospheric chemistry connects emissions, chemical transformation, transport and removal. A primary pollutant is emitted directly, such as carbon monoxide from incomplete combustion or sulfur dioxide from sulfur-containing fuels. A secondary pollutant is produced in air from precursor reactions, such as much ground-level ozone or sulfate aerosol. The classification concerns origin, not severity.

The lower atmosphere, the troposphere, contains weather and the air we breathe. Above it is the stratosphere, where much of the atmospheric ozone absorbs ultraviolet radiation. Layer boundaries vary with location and season; a diagram should be read as a conceptual organization rather than a fixed wall at one altitude.

A chemical can be beneficial in one context and harmful in another. Stratospheric ozone reduces exposure to damaging ultraviolet radiation, whereas elevated ozone near the surface is an oxidizing pollutant. These are the same O₃ molecules; location, concentration and exposure determine the different consequences.

For any exam statement, ask four questions: what is emitted, what chemistry converts it, what environmental effect follows, and what removes it? This avoids mixing acid deposition, photochemical smog, greenhouse warming and ozone-layer depletion into one vague phenomenon.

02 / UNDERSTAND & APPLY

2. Concentration units and atmospheric lifetime

A gas mixing ratio expresses the amount of a species relative to the air mixture. One ppm by mole means one mole per million moles of mixture; one ppb means one per billion. For ideal gases at common temperature and pressure, mole fraction equals volume fraction. Mass concentration, such as μg m⁻³, is a different quantity and depends on pressure, temperature and molar mass.

To convert an ideal-gas mole fraction x to mass per volume, use c_mass=xPM/(RT), with consistent units. If M is in g mol⁻¹ and P/(RT) in mol m⁻³, the result is g m⁻³. Multiply by 10⁶ for μg m⁻³. Never convert ppm to mass concentration using only a decimal-place shift.

Atmospheric lifetime describes removal relative to burden under a specified model. For a well-mixed box with first-order removal, dC/dt=−kC and C=C₀e^(−kt). Real lifetimes vary with sunlight, oxidants, deposition and transport. A long-lived compound can travel far from its emission source; short-lived reactive products may be important locally.

03 / UNDERSTAND & APPLY

3. Combustion products and reactive precursors

Complete oxidation of carbon gives CO₂, while limited oxygen or incomplete mixing can produce CO and soot. Carbon monoxide is not simply another name for carbon dioxide: it has different bonding, reactivity and effects. Both emission chemistry and exposure matter when comparing pollutants.

At high combustion temperatures, nitrogen and oxygen can form NO even if the fuel contains no nitrogen. NO can be converted to NO₂ in air; together these are commonly represented as NOx in pollution discussions. Fuel-bound nitrogen can also contribute. Do not confuse NOx with nitrous oxide N₂O, which has different atmospheric behaviour.

Sulfur-containing fuels can yield SO₂. Further oxidation through gas-phase and cloud-water reactions produces sulfuric acid and sulfate. A balanced overall equation summarizes atom conservation but is not necessarily an elementary collision mechanism. Atmospheric water droplets provide reaction environments as well as simple dissolution.

Particulate matter is suspended solid or liquid material, not one chemical compound. Soot, dust and sea salt are examples of primary particles; sulfate, nitrate and organic material can form secondarily. Particle size affects transport and deposition, while composition affects chemistry. PM₂.₅ and PM₁₀ denote size-based categories, not molecular masses.

04 / UNDERSTAND & APPLY

4. Acid deposition and logarithmic acidity

Dissolved CO₂ makes otherwise clean rain mildly acidic through carbonic-acid equilibria. A commonly quoted idealized background value is near pH 5.6, but actual natural rain varies because of dust, sea salt and other gases. Therefore pH below seven alone does not prove anthropogenic acid deposition.

Oxidation of SO₂ and NOx produces sulfuric and nitric acids that can be transported and deposited. Wet deposition carries acidic material in rain, snow or fog; dry deposition transfers acidic gases and particles to surfaces without precipitation. “Acid rain” is a common shorthand, while acid deposition covers both paths.

A representative net sulfur balance is 2SO₂+O₂+2H₂O → 2H₂SO₄. Nitric acid can form through NO₂+OH+M → HNO₃+M in the gas phase; M is a third-body collision partner carrying excess energy away. These routes do not mean that nitrogen or sulfur atoms disappear: trace them through precursor, acid and eventual salts.

Acidity can consume carbonate buffering and affect soils, waters and carbonate stone. CaCO₃+2H⁺ → Ca²⁺+CO₂+H₂O represents acid consumption by carbonate. The consequences depend on buffering capacity and exposure, not only the pH of one rain sample. pH measures hydrogen-ion activity approximately; it does not directly give total acid-neutralization capacity.

05 / UNDERSTAND & APPLY

5. Why sunlight and precursor chemistry matter

Photochemical smog is associated with sunlight-driven reactions among NOx, volatile organic compounds and oxidants. Ground-level ozone is a key secondary product; aldehydes, peroxyacyl nitrates and secondary particles may also be involved. Classical sulfurous smog is associated more with smoke, SO₂ and cool humid conditions. These broad categories can overlap in real air.

NO₂ absorbs suitable sunlight and dissociates to NO+O. Atomic oxygen then combines with O₂ in the presence of a third body to form O₃. But NO can consume ozone: NO+O₃ → NO₂+O₂. These three reactions alone largely cycle NOx and ozone rather than explaining sustained net ozone accumulation.

Oxidation of organic precursors produces peroxy radicals that can convert NO to NO₂ without consuming ozone. For example, RO₂+NO → RO+NO₂ is a representative propagation step. Subsequent NO₂ photolysis can then support net ozone production. This is why the roles of VOCs, radicals and NOx must be considered together.

Reducing one precursor does not necessarily give a locally linear ozone response, because chemical regimes and transport differ. A school-level explanation should identify the causal chemistry without presenting a universal numerical control rule. Ozone may peak away from fresh emission sources after sunlight, reaction time and wind transport.

06 / UNDERSTAND & APPLY

6. Protective ozone and catalytic depletion

In the stratosphere, ultraviolet photons can split O₂ into oxygen atoms. An atom can combine with O₂ and a third body to form ozone. Ozone itself absorbs radiation and can dissociate, while other reactions remove odd oxygen. The ozone layer is a dynamic balance of formation and destruction, not a permanent sheet of molecules.

Long-lived chlorine- and bromine-containing source gases can reach the stratosphere, where photochemistry releases reactive halogen species. In a simple chlorine cycle, Cl+O₃ → ClO+O₂ and ClO+O → Cl+O₂. Chlorine is regenerated, so one chlorine atom can participate in many turnovers before entering a reservoir or being removed.

Adding these steps gives O₃+O → 2O₂. This particular cycle requires atomic oxygen and is not the complete explanation of polar lower-stratospheric ozone loss, where other catalytic cycles and heterogeneous reactions on polar stratospheric cloud particles are important. A schematic cycle should state its scope rather than claim it explains every region identically.

The ozone hole refers to severe seasonal depletion, not an actual empty opening through the atmosphere. Ultraviolet shielding and greenhouse infrared absorption are distinct mechanisms. Some ozone-depleting substances also act as greenhouse gases, but that overlap does not make ozone depletion the main cause of greenhouse warming.

07 / UNDERSTAND & APPLY

7. Infrared-active gases and the energy balance

Earth absorbs incoming solar energy and emits thermal infrared radiation. Greenhouse gases absorb and emit in parts of that infrared spectrum, changing how efficiently energy escapes to space and warming the lower atmosphere and surface relative to a comparable atmosphere without those absorbers. Energy is redistributed and ultimately emitted; it is not created by the gases.

CO₂, CH₄, N₂O, water vapour and ozone are examples of greenhouse gases. Infrared absorption depends on molecular vibrational and rotational transitions, not simply on whether a gas molecule is heavy. N₂ and O₂, though abundant, lack the ordinary strong dipole-changing vibrational absorption of many greenhouse molecules.

The natural greenhouse effect is essential to Earth’s climate; increased concentrations of long-lived greenhouse gases alter the radiative balance. Water vapour is strongly temperature-dependent and acts importantly as a feedback. Aerosols can scatter or absorb radiation and influence clouds, so “all pollution warms equally” is false.

Greenhouse action concerns mainly infrared energy exchange, whereas ozone-layer protection concerns absorption of ultraviolet radiation. Keep the wavelength regions and mechanisms separate. This lesson teaches stable chemical concepts, not current legal limits, emissions inventories or a numerical climate forecast.

08 / UNDERSTAND & APPLY

8. Weather changes concentration without creating atoms

Wind transports pollutants and turbulence dilutes them. A temperature inversion can suppress vertical mixing when warmer air overlies cooler near-surface air, allowing pollutants to accumulate. The inversion changes dispersion; it is not itself a chemical source of sulfur or nitrogen oxides.

Control choices should match the source and chemical pathway. More complete combustion can reduce CO and soot; sulfur removal targets sulfur precursors; particle capture targets particulate emissions; appropriate catalytic systems reduce selected exhaust pollutants. No single control removes every gaseous and particulate contaminant equally.

Secondary pollutants require attention to precursors and atmospheric transformation. Transport can carry both precursors and products across large distances, so a clean-looking local source inventory does not prove the absence of imported pollution. Distinguish a measured concentration, an emitted mass rate and an exposure duration before comparing results.

09 / UNDERSTAND & APPLY

9. Revision by mechanism

For acid deposition, follow SO₂/NOx oxidation to acids and then wet or dry deposition, with buffering controlling the response. For photochemical smog, follow sunlight, NOx, VOC radicals and secondary ozone. For stratospheric ozone, follow ultraviolet chemistry and catalytic loss with explicit regeneration.

For greenhouse warming, identify infrared-active gases and energy balance. For concentration problems, preserve the distinction between mixing ratio and mass per volume; state temperature, pressure and molar mass when converting. For pH problems, use powers of ten rather than linear differences.

All worked cases use original hypothetical numbers or explicit elementary equations. They are chemical reasoning exercises, not claimed past-paper examples or statements of current air-quality standards.