Ch 18: Chemistry of the Environment: Earth's Atmosphere and Atmospheric Chemistry

Earth's Atmosphere

Atmospheric Layers and Boundaries

The Earth's atmosphere is divided into distinct layers, each characterized by unique temperature and pressure profiles. The boundaries between these layers are denoted by the suffix -pause (e.g., tropopause, stratopause).

• Troposphere: Extends from the Earth's surface up to about 10 km. Temperature ranges from approximately 290 K at the surface and decreases with altitude.

• Stratosphere: Ranges from 10 km to 50 km altitude. Temperature increases from about 215 K to 275 K.

• Mesosphere: Spans 50 km to 85 km altitude. Temperature decreases from 275 K to 190 K.

• Thermosphere: Above 85 km, temperature increases with altitude.

Example: Commercial airplanes typically fly in the lower stratosphere to avoid weather disturbances found in the troposphere.

Composition of the Atmosphere

Major and Minor Components

The atmosphere is primarily composed of nitrogen and oxygen, with several trace gases present in much smaller amounts. The composition affects both physical and chemical processes in the environment.

Component	Content (mole fraction)	Molar Mass (g/mol)
Nitrogen (N2)	0.78084	28.013
Oxygen (O2)	0.20948	31.998
Argon (Ar)	0.00934	39.948
Carbon dioxide (CO2)	0.000382	44.0099
Neon (Ne)	0.00001818	20.183
Helium (He)	0.00000524	4.003
Methane (CH4)	0.000002	16.043
Krypton (Kr)	0.00000114	83.80
Hydrogen (H2)	0.0000005	2.0159
Nitrous oxide (N2O)	0.0000005	44.0128
Xenon (Xe)	0.000000087	131.30

Key Point: Trace gases such as ozone, sulfur dioxide, nitrogen dioxide, ammonia, and carbon monoxide play significant roles in atmospheric chemistry despite their low concentrations.

Photochemistry of the Atmosphere

Solar Spectrum and Energy Relationships

Photochemistry involves chemical reactions initiated by solar radiation. The energy of electromagnetic radiation is related to its frequency and wavelength:

• Energy of a photon: $E = h\nu$

• Relationship between frequency and wavelength: $c = \nu\lambda$

• Energy in terms of wavelength: $E = \frac{hc}{\lambda}$

Key Point: Higher frequency (shorter wavelength) radiation carries more energy, which is crucial for driving photochemical reactions in the atmosphere.

Example: Ultraviolet light has enough energy to break chemical bonds, leading to photodissociation and photoionization processes.

Photodissociation

Mechanism and Atmospheric Importance

Photodissociation is the process by which a chemical bond is broken due to absorption of a photon. For this to occur, the photon must have sufficient energy to break the bond, and the molecule must absorb the photon.

• In the upper atmosphere, photodissociation of oxygen molecules produces atomic oxygen:

$O_2(g) + h\nu \rightarrow 2O(g)$

Example: The formation of atomic oxygen is a key step in the creation of ozone in the stratosphere.

Bond Enthalpies and Photochemical Processes

Average Bond Enthalpies

Bond enthalpy is the energy required to break one mole of a specific bond in a molecule. It is crucial for understanding which bonds can be broken by photodissociation.

Bond	Enthalpy (kJ/mol)
C–H	413
O–H	463
O=O	495
C–Cl	339
Cl–Cl	242
N≡N	941

Key Point: The lower the bond enthalpy, the easier it is for photodissociation to occur under solar radiation.

Example: The C–Cl bond in chlorofluorocarbons (CFCs) is susceptible to photodissociation, leading to ozone depletion.

Photoionization

Ionization by Radiation

Photoionization is the process by which a photon removes an electron from an atom or molecule, creating an ion. This process requires photons of even higher energy than photodissociation.

• Occurs in the upper atmosphere, contributing to the formation of the ionosphere.

• Ionization energy determines the minimum photon energy required for ionization.

Example: Oxygen and nitrogen are ionized by high-energy ultraviolet radiation, affecting radio communication and atmospheric chemistry.

Ozone Formation and Depletion

Ozone Formation Reactions

Ozone (O3) is formed in the stratosphere through a series of photochemical reactions involving oxygen molecules and atoms.

• Step 1: Photodissociation of O2 produces atomic oxygen:

• $O_2(g) + h\nu \rightarrow 2O(g)$

• Step 2: Atomic oxygen reacts with molecular oxygen to form ozone:

• $O(g) + O_2(g) \rightarrow O_3^*(g)$

• Step 3: Excited ozone can lose energy by decomposing or transferring energy to other molecules (M, usually N2 or O2):

• $O_3^*(g) + M(g) \rightarrow O_3(g) + M^*(g)$

Key Point: Maximum ozone formation occurs in the stratosphere, where sufficient UV radiation is present.

Ozone Depletion: CFCs and Nitrogen Oxides

Ozone depletion is primarily caused by catalytic reactions involving chlorofluorocarbons (CFCs) and nitrogen oxides (NOx).

• CFCs undergo photodissociation, releasing chlorine atoms:

• $CF_2Cl_2(g) + h\nu \rightarrow CF_2Cl(g) + Cl(g)$

• Chlorine atoms catalyze ozone destruction:

• $Cl(g) + O_3(g) \rightarrow ClO(g) + O_2(g)$

• $ClO(g) + h\nu \rightarrow Cl(g) + O(g)$

• Net reaction: $2O_3(g) \rightarrow 3O_2(g)$

• Nitrogen oxides also catalyze ozone destruction via similar mechanisms.

Example: The Antarctic ozone hole is a result of increased CFC concentrations and unique polar stratospheric cloud chemistry.

Chemistry of the Troposphere

Acid Rain and pH of Rainwater

Acid rain is caused by the presence of sulfur dioxide (SO2) and nitrogen oxides (NOx) in the atmosphere, which are oxidized to form sulfuric and nitric acids.

• Sulfur dioxide is produced by combustion of fossil fuels and volcanic activity.

• SO2 is oxidized to SO3, which reacts with water:

• $SO_3(g) + H_2O(l) \rightarrow H_2SO_4(aq)$

• Normal rainwater has a pH of about 5 due to dissolved CO2; acid rain can have a pH below 4, which is harmful to aquatic life.

Example: Acid rain damages forests, aquatic ecosystems, and buildings.

Removal of Sulfur Dioxide

SO2 can be removed from flue gases during combustion using powdered limestone (CaCO3), which reacts to form calcium sulfate (CaSO4).

• CaCO3 decomposes to CaO, which reacts with SO2:

• $CaO(s) + SO_2(g) \rightarrow CaSO_3(s)$

• Further oxidation produces CaSO4 (gypsum), which can be removed as a wet slurry.

Example: Flue-gas desulfurization is widely used in power plants to reduce SO2 emissions.