Module 10 — Chemical Reactions in the Atmosphere: Ultraviolet Radiation and Ozone

• The atmosphere is not just a mixture of gases; it is a medium for many chemical reactions.
• Examples include the formation of ozone by the effect of ultraviolet rays and the oxidation of certain gases.
• These reactions affect air quality, climate change and the health of humans and other living organisms.
• This module follows the book's sequence:

1. Ultraviolet radiation: what it is, what absorbs it.
2. Formation of ozone in the stratosphere.
3. Depletion of the ozone layer, its effects on health, and how to reduce it.
4. Ozone in the troposphere (a pollutant) and smog.

You met the electromagnetic spectrum in Module 7: radiation emitted by the Sun travels through space as electromagnetic waves. Each region of the spectrum has a characteristic range of wavelengths (λ) and frequencies (ν), related by c = λν. One region is ultraviolet (UV), which lies just beyond the violet end of visible light.

• UV has shorter wavelengths and higher energy than visible light: photon energy E = hν = hc/λ. UV wavelengths range from about 10 nm to 400 nm.
• The Sun is a natural source; artificial sources are special lamps.

• Sterilisation: special UV lamps are used to sterilise surgical instruments and laboratory equipment and to disinfect water and air, avoiding microbial contamination, because UV-C damages the DNA of micro-organisms.
• Detecting forged banknotes: the security threads and marks on real banknotes contain fluorescent substances that glow visibly when illuminated with UV, which forged notes lack; shops use a UV lamp to verify the authenticity of money.
• Others: forensic science (finding body fluids), curing of dental fillings and paints, fluorescent lamps, and vitamin D synthesis in the skin.

• Nitrogen (N₂), oxygen (O₂) molecules and atomic oxygen absorb photons with wavelengths shorter than 240 nm (high-energy UV-C and shorter).
• The formation and destruction of ozone (O₃) absorb photons in the ultraviolet region with wavelengths from 240 nm to 310 nm (the rest of UV-C and most of UV-B).
• In this way the ozone present in the stratosphere protects us from these high-energy photons. Only part of UV-B and most of UV-A reach the surface.

Problem. Calculate the energy of one photon and of one mole of photons for wavelengths of 240 nm and 310 nm (h = 6.63 × 10⁻³⁴ J·s, c = 3.0 × 10⁸ m/s, N_A = 6.02 × 10²³). Compare with the bond energy of O₂ (498 kJ/mol).

Solution. E = hc/λ.

• λ = 240 nm: E = (6.63 × 10⁻³⁴ × 3.0 × 10⁸)/(240 × 10⁻⁹) = 8.29 × 10⁻¹⁹ J per photon; per mole = 8.29 × 10⁻¹⁹ × 6.02 × 10²³ = 4.99 × 10⁵ J = 499 kJ/mol.
• λ = 310 nm: E = 1.989 × 10⁻²⁵/3.10 × 10⁻⁷ = 6.42 × 10⁻¹⁹ J; per mole = 386 kJ/mol. The O=O bond needs 498 kJ/mol, so only photons with energy ≥ 498 kJ/mol, that is wavelength ≤ 240 nm, can break an O₂ molecule. A 310 nm photon (386 kJ/mol) is too weak to split O₂ but strong enough to split ozone (whose bond is much weaker, about 100–150 kJ/mol). This explains the book's two numbers, 240 nm and 310 nm.

Ozone

• is a molecule made of three oxygen atoms (O₃) bonded together.
• It is formed in the stratosphere by the effect of UV rays from the Sun on oxygen molecules (O₂), in two steps:

Step 1 — Photodissociation of oxygen.

• Photodissociation is the breaking of the chemical bonds of a molecule as a result of light absorption.
• UV with a wavelength shorter than 240 nm breaks the covalent bond in the oxygen molecule, giving two individual oxygen atoms:

O₂ + UV photon (λ < 240 nm) → O + O

Step 2 — Combination. An individual oxygen atom (in its atomic state, very reactive) combines with an oxygen molecule to form ozone:

O + O₂ → O₃ (in reality a third molecule M, usually N₂ or O₂, is needed to carry away the excess energy: O + O₂ + M → O₃ + M)

Overall: 3 O₂ → 2 O₃ (check: 6 oxygen atoms on each side).

Destruction and the ozone–oxygen cycle. Ozone itself absorbs UV in the 240–310 nm range and splits again:

O₃ + UV photon (240–310 nm) → O₂ + O

• and the free O atom can recombine with O₂ to make ozone again.
• In this way ozone is created and destroyed continuously, and each cycle converts UV energy into heat — that is why the temperature rises in the stratosphere (Module 9).
• The ozone layer is a steady state where formation and destruction balance; the concentration is highest at about 20–30 km (the "ozone layer", though even there ozone is only a few parts per million of the air).

Worked example 2 — mass and moles. Problem. How many moles and grams of ozone can be made from 96 g of oxygen? What volume (at 0 °C and 1 atm, molar volume 22.4 L) does the oxygen occupy and what volume the ozone formed? (O = 16)

Solution. 3O₂ → 2O₃. n(O₂) = 96/32 = 3.0 mol → n(O₃) = 2.0 mol → mass of O₃ = 2 × 48 = 96 g (mass is conserved). V(O₂) = 3.0 × 22.4 = 67.2 L; V(O₃) = 2.0 × 22.4 = 44.8 L: the volume decreases by one third when oxygen is converted into ozone (3 volumes → 2 volumes).

Ozone layer depletion

• is the gradual deterioration (thinning) of the protective ozone layer in the atmosphere.
• It occurs mainly because of the release of chemical compounds containing chlorine and bromine, such as chlorofluorocarbons (CFCs), which result from industrial activities.
• (CFCs were used as refrigerants in refrigerators and air conditioners, as propellants in aerosol sprays, in foam production and as cleaning solvents; halons, which contain bromine, were used in fire extinguishers.)

CFCs are very stable in the lower atmosphere, so they are not removed by rain or reactions; slowly they rise to the stratosphere, where strong UV radiation breaks them and releases chlorine atoms:

1. CFCl₃ + UV → CFCl₂ + Cl
2. Cl + O₃ → ClO + O₂
3. ClO + O → Cl + O₂ Net: O₃ + O → 2O₂

• The chlorine atom is regenerated in step 3 — it acts as a catalyst.
• A single chlorine atom can destroy tens of thousands (about 100 000) ozone molecules before it is removed from the stratosphere.
• Bromine atoms from halons are even more effective per atom (about 40–60 times).

The ozone hole over Antarctica develops each spring (September–October), because in the polar winter ice clouds in the stratosphere help to convert chlorine compounds into active forms; when the sun returns in spring the chlorine destroys ozone rapidly. The ozone column falls from about 300 Dobson units to about 100 or less (a Dobson unit measures the thickness of ozone if compressed to 0 °C and 1 atm: 100 DU = 1 mm).

• In 1985 a weather monitoring station in Antarctica, operated by the British scientists, discovered a large hole in the ozone layer.
• The discovery led to the Montreal Protocol (1987), an international agreement that banned the production and use of many ozone-depleting chemicals (CFCs and related compounds).
• Thanks to this cooperation, the ozone layer has begun to recover gradually (expected to return to the 1980 level by about the middle of the 21st century), showing humanity can solve major environmental problems through collective action.

Problem. If each chlorine atom destroys 100 000 ozone molecules, how many ozone molecules are destroyed by 1 mole of Cl atoms? How many grams of ozone is that? (N_A = 6.02 × 10²³; O₃ = 48 g/mol)

Solution. Number destroyed = 6.02 × 10²³ × 1.0 × 10⁵ = 6.02 × 10²⁸ molecules = 10⁵ mol of O₃ = 10⁵ × 48 g = 4.8 × 10⁶ g = 4800 kg of ozone. Only 35.5 g of chlorine (1 mol) can remove 4.8 tonnes of ozone — a good example of the power of a catalyst.

• With depletion, a greater amount of harmful UV (particularly UV-B) reaches the surface.
• (A rough rule: a 1 % decrease in ozone leads to about a 1.5–2 % increase in UV-B at the ground.)

1. On the skin. More UV causes damage to DNA in skin cells. The damage can cause genetic mutations that may result in skin cancers, such as melanoma (the most dangerous). The body attempts to repair the damage by activating special DNA repair enzymes, but repeated exposure can exceed the cells' ability to repair. Also sunburn, premature ageing and wrinkles.
2. On the immune system. UV affects certain immune cells, such as helper T cells, which leads to a weakened immune response, making the body more vulnerable to infections and diseases (and reducing the effect of vaccines).
3. On the eye. Absorption of UV by the eye's lens accelerates the development of cataract (clouding of the lens), which reduces the transmission of light to the retina. UV can also damage the cornea ("snow blindness") and retina.

Effects on other organisms (extension): damage to phytoplankton (the base of the marine food chain; Unit 1), to amphibian eggs and young fish, and to crop yield and plant growth.

Following simple actions helps reduce the emission of harmful chemical substances:

1. Choose ozone-friendly products. Many traditional cleaning products contain chlorine and bromine compounds that damage the ozone layer when emitted. Look for products labelled "eco-friendly" or use natural cleaners such as vinegar and baking soda.
2. Avoid ozone-depleting substances (ODS).
   • Refrigerants: avoid refrigerators and air conditioners that use CFCs. Choose modern appliances that use safer alternatives, such as refrigerant gas R-410A (which contains no chlorine and therefore does not harm ozone; the hydrofluorocarbons also warm the climate, so newer ones such as R-32 and natural refrigerants are replacing them).
   • Fire extinguishers: avoid extinguishers containing halon gas, which contains bromine, one of the most powerful ozone-depleting substances. Replace them with modern extinguishers free of it.
3. Proper disposal and recycling of old refrigerators and air conditioners so that the gas is captured, not released.
4. Support the laws (Montreal Protocol) and educate others.

• Ozone near the ground is formed in a different way — a "bad ozone" instead of the "good ozone" in the stratosphere.
• It forms in the troposphere as a result of the reaction between nitrogen oxides and hydrocarbons in the presence of sunlight:

1. Nitrogen dioxide absorbs sunlight (λ < 400 nm): NO₂ + light → NO + O
2. The oxygen atom reacts with O₂: O + O₂ → O₃
3. Ozone reacts back with NO (O₃ + NO → NO₂ + O₂) in a cycle that would balance itself; but volatile organic compounds/hydrocarbons convert NO back to NO₂ without using up ozone, so ozone accumulates to harmful levels, especially on hot, sunny afternoons in traffic-heavy cities.

Effects of ozone formed near the Earth's surface:

• Smog formation. Ozone combines with other pollutants in the air — sulfur dioxide, nitrogen dioxide and fine particles — and forms smog (a haze of pollutants that reduces visibility and irritates the eyes and lungs). The photochemical (ozone-type) smog of sunny cities differs from the older "London-type" smog (smoke + SO₂ + fog).
• Burning of plant leaves. Ozone enters the stomata and oxidises cell components, burning leaf tissue and reducing the plants' ability to carry out photosynthesis.
• Damage to plants and crops. Growth and quality are affected; yields of wheat, rice, cotton and soybean fall.
• Corrosion of materials such as plastic and rubber (tyres crack, plastics become brittle).
• Human health: irritation of the lungs, reduced respiratory capacity, asthma attacks (Module 12).

Comparison: stratospheric ozone protects, tropospheric ozone pollutes. The same molecule (O₃) has opposite effects depending on its location.

Problem. Write the balanced equation for the formation of ozone from NO₂ and show that it conserves atoms and charge.

Solution. NO₂ → NO + O (needs light); O + O₂ → O₃; add: NO₂ + O₂ → NO + O₃. Atoms: left N 1, O 4; right N 1, O 1 + 3 = 4 ✔. The reaction is driven by sunlight; the NO is later oxidised back to NO₂ by peroxy radicals formed from hydrocarbons, so NO₂ acts as a catalyst for net ozone production.

The book says the atmosphere is a medium for "the oxidation of certain gases". The main examples are:

1. Lightning and nitrogen. The temperature in a lightning channel is about 30 000 °C, enough to break the N≡N bond: N₂ + O₂ → 2NO (nitric oxide) 2NO + O₂ → 2NO₂ (nitrogen dioxide, brown gas) 3NO₂ + H₂O → 2HNO₃ + NO (nitric acid, carried to the soil by rain as nitrate fertiliser) (Atom check for the last equation: N 3 = 2 + 1; O 6 + 1 = 6 + 1; H 2 = 2.) This is the natural counterpart of the acid-rain reactions in Module 12. Industrial combustion produces the same oxides at high temperature.
2. The hydroxyl radical (OH) — the "detergent" of the atmosphere. Ozone absorbing UV can produce an excited oxygen atom, which reacts with water vapour to make OH radicals: O* + H₂O → 2OH. OH oxidises many pollutants and greenhouse gases (CO, CH₄, SO₂, NO₂) and so cleans the air, for example CO + OH → CO₂ + H, and NO₂ + OH → HNO₃. The lifetime of methane in the air (about 12 years) is controlled by this reaction.
3. Sulfur oxidation. SO₂ is oxidised to SO₃ and then dissolves to make sulfuric acid (Module 12), or forms fine sulfate particles that reflect sunlight and scatter light (sulfur from volcanoes cools the climate for a short time).
4. Combustion (burning) of fuels is oxidation: complete combustion gives CO₂ and water; incomplete gives CO and soot (Module 12).

• Ozone is measured by (i) ground spectrometers (the Dobson instrument) that compare sunlight at two UV wavelengths, one absorbed by ozone and one not; (ii) balloons carrying small ozone sensors (ozonesondes); and (iii) satellites measuring the UV scattered from the atmosphere.
• The column is expressed in Dobson units; the global average is about 300 DU, and a value below 220 DU is the conventional boundary of the "ozone hole".

• Egypt, being a sunny country near the Tropic of Cancer, has high to extreme UV index for much of the year, so protection (hat, sunglasses that block UV, light long clothing, sunscreen, shade) is important.
• Cloud does not block all UV and water, sand and snow reflect it.

• Confusing UV with infrared; UV has shorter wavelength than visible light.
• Writing that ozone is absorbed by UV in the formation step. In step 1 UV is absorbed by O₂; ozone absorbs UV in destruction.
• Saying that CFCs are destroyed in the troposphere. They are stable until they reach the stratosphere.
• Writing that all ozone is harmful. Stratospheric ozone is protective; ground-level ozone is a pollutant.
• Forgetting that the Montreal Protocol dealt with ozone-depleting substances, not with greenhouse gases in general.