1. chemical reactions in atmosphere

Earth’s Atmosphere and the Ozone Layer

Earth’s atmosphere at sea level is an approximately 80:20 solution of nitrogen and oxygen gases, with small amounts of carbon dioxide, water vapor, and the noble gases, and trace amounts of a variety of other compounds (Table ‘b’ “The Composition of Earth’s Atmosphere at Sea Level*”). A key feature of the atmosphere is that its composition, temperature, and pressure vary dramatically with altitude. Consequently, scientists have divided the atmosphere into distinct layers, which interact differently with the continuous flux of solar radiation from the top and the land and ocean masses at the bottom. Some of the characteristic features of the layers of the atmosphere are illustrated in Figure ‘c’ “Variation of Temperature with Altitude in Earth’s Atmosphere”.

Table ‘b’ The Composition of Earth’s Atmosphere at Sea Level*

GasFormulaVolume (%)
carbon dioxideCO20.0314
nitrous oxideN2O0.00005
* In addition, air contains as much as 7% water vapor (H2O), 0.0001% sulfur dioxide (SO2), 0.00007% ozone (O3), 0.000002% carbon monoxide (CO), and 0.000002% nitrogen dioxide (NO2).
 Carbon dioxide levels are highly variable; the typical range is 0.01–0.1%.

Figure’c’ Variation of Temperature with Altitude in Earth’s Atmosphere

Note the important chemical species present in each layer. The yellow line indicates the temperature at various altitudes.

The troposphere is the lowest layer of the atmosphere, extending from Earth’s surface to an altitude of about 11–13 km (7–8 mi). Above the troposphere lies the stratosphere, which extends from 13 km (8 mi) to about 44 km (27 mi). As shown in Figure ‘c’ “Variation of Temperature with Altitude in Earth’s Atmosphere”, the temperature of the troposphere decreases steadily with increasing altitude. Because “hot air rises,” this temperature gradient leads to continuous mixing of the upper and lower regions within the layer. The thermally induced turbulence in the troposphere produces fluctuations in temperature and precipitation that we collectively refer to as “weather.” In contrast, mixing between the layers of the atmosphere occurs relatively slowly, so each layer has distinctive chemistry. We focus our attention on the stratosphere, which contains the highest concentration of ozone.

The sun’s radiation is the major source of energy that initiates chemical reactions in the atmosphere. The sun emits many kinds of radiation, including visible light, which is radiation that the human eye can detect, and ultraviolet light, which is higher energy radiation that cannot be detected by the human eye. This higher energy ultraviolet light can cause a wide variety of chemical reactions that are harmful to organisms. For example, ultraviolet light is used to sterilize items, and, as anyone who has ever suffered a severe sunburn knows, it can produce extensive tissue damage.

Light in the higher energy ultraviolet range is almost totally absorbed by oxygen molecules in the upper layers of the atmosphere, causing the O2 molecules to dissociate into two oxygen atoms in a cleavage reaction:

Equation 1

O2(g) −→light 2O(g)

In Equation 1, light is written above the arrow to indicate that light is required for the reaction to occur. The oxygen atoms produced in Equation 1 can undergo a condensation reaction with O2 molecules to form ozone:

Equation 2

O(g) + O2(g) → O3(g)

Ozone is responsible for the pungent smell we associate with lightning discharges and electric motors. It is also toxic and a significant air pollutant, particularly in cities.

In the stratosphere, the ozone produced via Equation 2 has a major beneficial effect. Ozone absorbs the less-energetic range of ultraviolet light, undergoing a cleavage reaction in the process to give O2 and O:

Equation 3.36

O3(g)−→lightO2(g) + O(g)

The formation of ozone (Equation 3) and its decomposition (Equation 3.36) are normally in balance, resulting in essentially constant levels of about 1015 ozone molecules per liter in the stratosphere. This so-called ozone layer acts as a protective screen that absorbs ultraviolet light that would otherwise reach Earth’s surface.

In 1974, F. Sherwood Rowland and Mario Molina published a paper claiming that commonly used chlorofluorocarbon (CFC) compounds were causing major damage to the ozone layer (Table 3.3 “Common CFCs and Related Compounds”). CFCs had been used as refrigerants and propellants in aerosol cans for many years, releasing millions of tons of CFC molecules into the atmosphere. Because CFCs are volatile compounds that do not readily undergo chemical reactions, they persist in the atmosphere long enough to be carried to the top of the troposphere, where they eventually enter the stratosphere. There they are exposed to intense ultraviolet light and undergo a cleavage reaction to produce a chlorine atom, which is shown for Freon-11:

Equation 3.37

CCl3F(g)−→lightCCl2F(g) + Cl(g)

The resulting chlorine atoms act as a homogeneous catalyst in two redox reactions (Equation 3.38 and Equation 3.39):

Equation 3.38

Cl(g) + O3(g) → ClO(g) + O2(g)

Equation 3.39

ClO(g) + O(g) → Cl(g) + O2(g)

Adding the two reactions in Equation 3.38 and Equation 3.39 gives

Equation 3.40

Cl(g) + O3(g) + ClO(g) + O(g) → ClO(g) + Cl(g) + 2O2(g)

Because chlorine and ClO (chlorine monoxide) appear on both sides of the equation, they can be canceled to give the following net reaction:

Equation 3.41

O3(g) + O(g) → 2O2(g)

In the presence of chlorine atoms, one O3 molecule and one oxygen atom react to give two O2 molecules. Although chlorine is necessary for the overall reaction to occur, it does not appear in the net equation. The chlorine atoms are a catalyst that increases the rate at which ozone is converted to oxygen.

Table 3.3 Common CFCs and Related Compounds

NameMolecular FormulaIndustrial Name
trichlorofluoromethaneCCl3FCFC-11 (Freon-11)
dichlorodifluoromethaneCCl2F2CFC-12 (Freon-12)
chlorotrifluoromethaneCClF3CFC-13 (Freon-13)
*Halons, compounds similar to CFCs that contain at least one bromine atom, are used as fire extinguishers in specific applications (e.g., the engine rooms of ships).

Because the stratosphere is relatively isolated from the layers of the atmosphere above and below it, once chlorine-containing species enter the stratosphere, they remain there for long periods of time. Each chlorine atom produced from a CFC molecule can lead to the destruction of large numbers of ozone molecules, thereby decreasing the concentration of ozone in the stratosphere. Eventually, however, the chlorine atom reacts with a water molecule to form hydrochloric acid, which is carried back into the troposphere and then washed out of the atmosphere in rainfall.

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