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Chapter 20: Problem 16

How do CFCs and nitrogen oxides destroy ozone in the stratosphere?

Short Answer

Expert verified
CFCs and nitrogen oxides destroy ozone in the stratosphere through several chemical reactions. CFCs release chlorine atoms which convert ozone into molecular oxygen. Nitrogen oxides react with ozone to also produce oxygen. Both reactions thus reduce the amount of ozone, allowing more harmful UV rays to reach the Earth's surface.

Step by step solution

01

Understanding the role of CFCs

CFCs are volatile compounds that are used in refrigeration systems, aerosol propellants, and other industrial applications. When these are released into the atmosphere, they slowly rise to the stratosphere. There, they are broken apart by solar radiation and release chlorine atoms.
02

Explaining the destruction process

The chlorine atoms can react with ozone (O3) to form chlorine monoxide (ClO) and molecular oxygen (O2). The ClO can then react with another molecule of ozone to create more molecular oxygen and regenerate the chlorine atom. So, one chlorine atom can destroy many molecules of ozone before it is removed from the stratosphere.
03

Role of Nitrogen Oxides

Nitrogen oxides also play a role in ozone depletion. They are produced by various natural and human activities. In the stratosphere, these nitrogen oxides can react with and consequently destroy ozone. The reaction produces nitrogen dioxide and oxygen.
04

Effect of these reactions

The net result of these reactions is the conversion of ozone, which blocks deadly ultraviolet (UV) radiation, into molecular oxygen, which does not provide the same protection against UV rays. This leads to what is commonly known as 'ozone hole'.

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Most popular questions from this chapter

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Explain why, for maximum performance, supersonic airplanes need to fly at a high altitude (in the stratosphere).

What causes the polar ozone holes?

The equilibrium constant \(\left(K_{P}\right)\) for the reaction $$ \mathrm{N}_{2}(g)+\mathrm{O}_{2}(g) \rightleftharpoons 2 \mathrm{NO}(g) $$ is \(4.0 \times 10^{-31}\) at \(25^{\circ} \mathrm{C}\) and \(2.6 \times 10^{-6}\) at \(1100^{\circ} \mathrm{C},\) the temperature of a running car's engine. Is this an endothermic or exothermic reaction?

A 14-m by \(10-\mathrm{m}\) by \(3.0-\mathrm{m}\) basement had a high radon content. On the day the basement was sealed off from its surroundings so that no exchange of air could take place, the partial pressure of \({ }^{222} \mathrm{Rn}\) was \(1.2 \times 10^{-6} \mathrm{mmHg} .\) Calculate the number of \({ }^{222} \mathrm{Rn}\) isotopes \(\left(t_{1}=3.8 \mathrm{~d}\right)\) at the beginning and end of 31 days. Assume STP conditions.
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