The rate constant of the reaction O(g) + N₂(g) → NO(g) + N(g), which takes place in the stratosphere, is 9.7 * 10¹⁰ L · mol^–1 · s^–1 at 800. °C. The activation energy of the reaction is 315 kJ · mol^–1. What is the rate constant at 700. °C? (See Box 7E.1.)

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Box 7E.1 WHAT HAS THIS TO DO WITH... THE ENVIRONMENT? PROTECTING THE OZONE LAYER Every year our planet is bombarded with enough energy from the Sun to destroy all life. Only the ozone in the stratosphere protects us from that onslaught. The ozone, though, is threatened by mod- ern lifestyles. Chemicals used as refrigerants and propellants, such as chlorofluorocarbons (CFCs), and the nitrogen oxides in jet exhausts have been found to create holes in Earth's protective ozone layer, particularly over Antarctica. Because they act as catalysts, even small amounts of these chemicals can cause large changes in the vast reaches of the stratosphere. Ozone forms in the stratosphere in two steps. First, O₂ mol- ecules or other oxygen-containing compounds are broken apart into atoms by sunlight, a process called photodissociation: sunlight A<340 nm 0 +0 Then the O atoms, which are reactive radicals with two unpaired electrons, react with the more abundant O₂ molecules to form ozone. The ozone molecules are created in such a high-energy state that their vibrational motions would quickly tear them apart unless another molecule, such as O₂ or N₂, collides with them first. The other molecule, indicated as M, carries off some of the energy: 0 + 0₂ → 0₂* O₂ + MO₂ + M where * denotes a high-energy state. The net reaction, the sum of these two elementary reactions (after multiplying them by 2) and the reaction presented earlier, is 3 0₂ → 2 03. Some of the ozone is decomposed by ultraviolet radiation: UV,A<340 mm 0₂- 0+0₂ The oxygen atom produced in this step can react with oxygen molecules to produce more ozone, and so the ozone concentration Ozone hole area/(10 km²) 25 8 15 10 5 0 Seasonal variation in the mean area of the ozone hole over Antarctica for 2014. in the stratosphere normally remains constant, with seasonal variations. Because the decomposition of ozone absorbs ultraviolet radiation, ozone helps to shield the Earth from radiation damage. Warnings of the possibility that anthropogenic (human- generated) chemicals can threaten the ozone in the stratosphere began to appear in 1970 and 1971, when Paul Crutzen deter- mined experimentally that NO and NO₂ molecules catalyze the destruction of ozone. Nitrogen oxides are produced naturally in the troposphere by lightning and combustion in automobile and airplane engines. However, only N₂O is sufficiently unreactive to travel up to the stratosphere, where it is converted into NO₂. Mario Molina and Sherwood Rowland used Crutzen's work and other data in 1974 to build a model of the stratosphere that explained how chlorofluorocarbons could threaten the ozone layer.* In 1985, ozone levels over Antarctica were indeed found