Ozone depletion

Formation. Ozone ($O_3$) is an unstable, three-atom form of oxygen built in the stratosphere when energetic ultraviolet radiation strikes ordinary two-atom oxygen ($O_2$). The UV photon supplies enough energy to split some $O_2$ molecules into free oxygen atoms, and each free atom then combines with an intact $O_2$ molecule to give $O_3$. The overall change is written as one balanced equation:

where: $O_2$ is molecular (diatomic) oxygen, $O_3$ is ozone, and the label above the arrow shows that ultraviolet radiation is the energy source driving the reaction — without it, the conversion does not happen at ordinary temperatures.

Function. The same ozone layer that UV radiation builds is also what absorbs UV radiation, so it acts as a continuous filter: most of the sun's harmful ultraviolet rays are soaked up in the stratosphere and never reach the Earth's surface. This shielding protects living cells from UV-induced damage — sunburn, skin cancer, and eye cataracts in humans, and harm to plant tissue and ocean plankton, which sit at the base of many food chains.

Depletion. Ozone depletion is what happens when the rate of ozone destruction overtakes the rate of ozone formation, so the layer thins over time instead of holding steady. The syllabus asks you to name, not explain the chemistry of, the substances responsible: chlorofluorocarbons (CFCs) are the main culprit, once used widely as refrigerants in refrigerators and air conditioners, as aerosol-can propellants, and as foam-blowing agents. Halons (used in fire extinguishers) and certain stratospheric nitrogen oxides are additional contributors worth knowing by name.

The goal here is to see why ultraviolet radiation specifically, rather than ordinary sunlight or heat, is needed to build ozone, and why the very chemical stability that makes CFCs excellent refrigerants is exactly what turns them into a stratospheric hazard once they leave the appliance.

Step 1. The oxygen-oxygen double bond in $O_2$ is strong, so ordinary warmth or visible light cannot break it apart. Ultraviolet radiation carries far more energy per photon than visible light, which is why only the UV portion of sunlight is energetic enough to split some $O_2$ molecules into separate oxygen atoms — the essential first step before $O_3$ can form.

Step 2. This UV-splitting happens mainly in the stratosphere because that is where enough unfiltered, high-energy UV radiation is still available; at ground level, most of that same UV has already been absorbed on the way down, so ozone formation is essentially a stratosphere-only process.

Step 3. Once formed, ozone itself is very good at absorbing UV radiation — the molecule captures the photon's energy and converts it to heat, warming the stratosphere and preventing that UV from ever completing its journey to the surface. This absorbing behaviour is precisely the protective "shielding" function described in the syllabus.

Step 4. CFCs became a threat for an ironic reason: chemical stability. Manufacturers valued CFCs precisely because they do not react with anything at ground level — they neither burn nor corrode nor decompose, which is why they made safe refrigerants and propellants. That same unreactivity means CFCs released at the surface survive the slow, multi-year drift upward without breaking down, arriving intact in the stratosphere where they can disturb the ozone balance. Being asked only to name CFCs, and not to write their destruction reactions, follows directly from this: the syllabus wants you to know which chemicals reach the stratosphere and why their stability lets them get there, not the detailed chemistry of what happens once they arrive.

Question. If $9.6\ \text{g}$ of oxygen gas ($O_2$) is completely converted into ozone according to $3O_2 \xrightarrow{\text{UV}} 2O_3$, calculate the mass of ozone formed and the number of ozone molecules produced. (Molar mass: $O_2 = 32\ \text{g/mol}$, $O_3 = 48\ \text{g/mol}$; Avogadro's number $= 6.022\times10^{23}\ \text{per mol}$.)

Step 1. Convert the given mass of $O_2$ to moles: $\text{moles of } O_2 = \dfrac{9.6\ \text{g}}{32\ \text{g/mol}} = 0.3\ \text{mol}$.

Step 2. Use the equation's mole ratio. The balanced equation shows $3$ mol of $O_2$ produces $2$ mol of $O_3$, so: $\text{moles of } O_3 = 0.3\ \text{mol} \times \dfrac{2}{3} = 0.2\ \text{mol}$.

Step 3. Convert moles of ozone to mass: $\text{mass of } O_3 = 0.2\ \text{mol} \times 48\ \text{g/mol} = 9.6\ \text{g}$. Notice this equals the starting mass of $O_2$ — expected, since the same oxygen atoms are simply rearranged and no mass is lost.

Step 4. Convert moles of ozone to number of molecules: $\text{molecules of } O_3 = 0.2\ \text{mol} \times 6.022\times10^{23}\ \text{per mol} = 1.2044\times10^{23}$ molecules.

Final answer: $9.6\ \text{g}$ of ozone is formed, containing about $1.2\times10^{23}$ ozone molecules.

• Confusing ozone depletion with global warming — CFCs destroying stratospheric ozone and greenhouse gases trapping heat are different problems, even though both sit in the same "atmospheric chemistry" chapter.
• Trying to write the reaction mechanism for how CFCs destroy ozone — the syllabus explicitly says this is not required; only naming CFCs as the responsible chemicals is expected.
• Writing $3O_2 \rightarrow 2O_3$ without the UV label — the equation is incomplete without noting that ultraviolet radiation drives the reaction.
• Mixing up stratospheric ozone with ground-level ozone — stratospheric ozone protects by absorbing UV, while ground-level ozone is a harmful pollutant component of smog; this topic concerns only the protective stratospheric layer.

Ask the tutor: Want to see how CFCs replaced by newer refrigerants changed things, or practice another mole-ratio calculation with the ozone equation? Ask the AI tutor.