chemistry

Ozone–oxygen cycle

Ozone–oxygen cycle is a chemistry topic covered in the lgStudy science library. This page brings together a partial reference excerpt, illustrations, worked examples, real-world applications and a short study plan, so you can understand Ozone–oxygen cycle rather than just read about it. In short: The ozone–oxygen cycle is the process by which ozone is continually regenerated in Earth's stratosphere, converting ultraviolet radiation (UV) into heat. In 1930 Sydney Chapman resolved the chemistry involved.

Key takeaways

Ozone–oxygen cycle belongs to chemistry; place it in that map before memorising details.
Learn the definition first, then one example that makes the definition concrete.
Connect Ozone–oxygen cycle to a quantity you can measure, compute or draw — that is where exam questions come from.
Reproduce the core statement of Ozone–oxygen cycle from memory before moving on to harder problems.

Reference excerpt

The ozone–oxygen cycle is the process by which ozone is continually regenerated in Earth's stratosphere, converting ultraviolet radiation (UV) into heat. In 1930 Sydney Chapman resolved the chemistry involved. The process is commonly called the Chapman cycle by atmospheric scientists. Most of the ozone production occurs in the tropical upper stratosphere and mesosphere. The total mass of ozone produced per day over the globe is about 400 million metric tons. The global mass of ozone is relatively constant at about 3 billion metric tons, meaning the Sun produces about 12% of the ozone layer each day.

Photochemistry

The Chapman cycle describes the main reactions that naturally determine, to first approximation, the concentration of ozone in the stratosphere. It includes four processes - and a fifth, less important one - all involving oxygen atoms and molecules, and UV radiation:

Creation An oxygen molecule is split (photolyzed) by higher frequency UV light (top end of UV-B, UV-C and above) into two oxygen atoms (see figure):

1. oxygen photodissociation: O2 + ℎν(<242 nm) → 2 O Each oxygen atom may then combine with an oxygen molecule to form an ozone molecule:

2. ozone creation: O + O2 + A → O3 + A where A denotes an additional molecule or atom, such as N2 or O2, required to maintain the conservation of energy and momentum in the reaction. Any excess energy is produced as kinetic energy.

The ozone–oxygen cycle The ozone molecules formed by the reaction (above) absorb radiation with an appropriate wavelength between UV-C and UV-B. The triatomic ozone molecule becomes diatomic molecular oxygen, plus a free oxygen atom (see figure):

3. ozone photodissociation: O3 + ℎν(240–310 nm) → O2 + O The atomic oxygen produced may react with another oxygen molecule to reform ozone via the ozone creation reaction (reaction 2 above). These two reactions thus form the ozone–oxygen cycle, wherein the chemical energy released by ozone creation becomes molecular kinetic energy. The net result of the cycle is the conversion of penetrating UV-B light into heat, without any net loss of ozone. While keeping the ozone layer in stable balance, and protecting the lower atmosphere from harmful UV radiation, the cycle also provides one of two major heat sources in the stratosphere (the other being kinetic energy, released when O2 is photolyzed into individual O atoms).

Removal If an oxygen atom and an ozone molecule meet, they recombine to form two oxygen molecules:

4. ozone conversion: O3 + O → 2 O2 Two oxygen atoms may react to form one oxygen molecule:

5. oxygen recombination: 2O + A → O2 + A as in reaction 2 (above), A denotes another molecule or atom, like N2 or O2 required for the conservation of energy and momentum. Note that reaction 5 is of the least importance in the stratosphere, since, under normal conditions, the concentration of oxygen atoms is much lower than that of diatomic oxygen molecules. This reaction is therefore less common than ozone creation (reaction 2). The overall amount of ozone in the stratosphere is determined by the balance between production from solar radiation and its removal. The removal rate is slow, since the concentration of free O atoms is very low.

Additional reactions In addition to these five reactions, certain free radicals - the most important being hydroxyl (OH), nitric oxide (NO), and atomic chlorine (Cl) and bromine (Br) - catalyze the recombination reaction, leading to an ozone layer that is thinner than it would be if the catalysts were not present. Most OH and NO are naturally present in the stratosphere, but human activity - especially emissions of chlorofluorocarbons (CFCs) and halons - has greatly increased the concentration of Cl and Br, leading to ozone depletion. Each Cl or Br atom can catalyze tens of thousands of decomposition reactions before it is removed from the stratosphere.

Main reactions in different atmospheric layers

Thermosphere For given relative reactants concentrations, The rates of ozone creation and oxygen recombination (reactions 2 and 5) are proportional to the air density cubed, while the rate of ozone conversion (reaction 4) is proportional to the air density squared, and the photodissociation reactions (reactions 1 and 3) have a linear dependence on air density. Thus, at the upper thermosphere, where air density is very low and photon flux is high, oxygen photodissociation is fast while ozone creation is low, thus its concentration is low. Thus the most important reactions are oxygen photodissociation and oxygen recombination, with most of the oxygen molecules dissociated to oxygen atoms. As we go to the lower thermosphere (e.g. 100 km height and below), the photon flux in the <170 nm wavelengths drops sharply due to absorption by oxygen in the oxygen photodissociation reaction (reaction 1). This wavelength regime has the highest cross section for this reaction (10−17 cm2 per oxygen molecule), and thus the rate of oxygen photodissociation per oxygen molecule decreases significantly at these altitudes, from more than 10−7 per second (about once a month) at 100 km to 10−8 per second (about once every few years) at 80 km . As a result, the atomic oxygen concentration (both relative and absolute) decreases sharply, and ozone creation (reaction 2) is ongoing, leading to a small but non-negligible ozone presence. Note that temperatures also drop as altitude decreases, because lower photon photodissociation rates mean lower heat production per air molecule.

Below thermosphere: Reaction rates at steady state Odd oxygen species (atomic oxygen and ozone) have net creation rate only by oxygen dissociation (reaction 1), and net destruction by either ozone conversion or oxygen recombination (reactions 4 and 5). At steady state these processes are balanced, so the rates of these reactions obey:

(rate of reaction 1) = (rate of reaction 4) + (rate of reaction 5). At steady state, ozone creation is also balanced with its removal. so:

(rate of reaction 2) = (rate of reaction 3) + (rate of reaction 4). It thus follows that:

… excerpt ends here. Continue reading the full article.

Illustrations

Worked examples

Example 1 — a first encounter with Ozone–oxygen cycle

Start with the simplest possible case. Write down what Ozone–oxygen cycle claims or describes in one sentence, then invent the smallest concrete situation in which that sentence is true. In chemistry, the smallest case is usually a single object, a single equation or a single measurement. Check that every symbol or term in your sentence has a meaning in that case.

Example 2 — changing one variable

Take the situation from Example 1 and change exactly one quantity: double it, halve it, or set it to zero. Predict what should happen to Ozone–oxygen cycle before you calculate. Comparing your prediction with the result is the fastest way to find out whether you understand the idea or only the words.

Example 3 — an exam-style question

Typical questions about Ozone–oxygen cycle ask you to (a) state it precisely, (b) apply it to given data, and (c) explain a limitation. Practise writing all three answers in under five minutes; the third part is what separates a full-mark answer from an average one.

Applications of Ozone–oxygen cycle

In research: Ozone–oxygen cycle appears in chemistry research whenever the underlying quantities have to be modelled precisely. Papers usually cite it as a starting assumption and then explore where it breaks down.
In technology and industry: Engineering practice reuses Ozone–oxygen cycle in design rules, simulations and safety margins. Knowing the idea lets you read a specification sheet and understand why the numbers look the way they do.
In the classroom: Ozone–oxygen cycle is common in secondary-school and first-year university syllabi. It links to neighbouring topics Atmospheric chemistry, Biogeochemical cycle, Oxygen, so understanding it makes those chapters shorter.
In everyday life: Look for Ozone–oxygen cycle outside the textbook — in sport, cooking, traffic, electronics or the sky above you. An example you found yourself is remembered far longer than one you were given.

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How to study Ozone–oxygen cycle in 20 minutes

Read the reference excerpt below once, without taking notes.
Close the page and write down what Ozone–oxygen cycle means in your own words.
Compare your version with the excerpt and mark what you missed.
Work through the three examples above with pen and paper.
Explain Ozone–oxygen cycle out loud to somebody else — or to Teacher Smith in the lgStudy chat.

Frequently asked questions

What is Ozone–oxygen cycle in simple terms?

Why does Ozone–oxygen cycle matter?

Because it connects several chemistry ideas at once: it gives you a definition you can apply, a quantity you can calculate, and a way to check whether a result is plausible.

How should I study Ozone–oxygen cycle?

Read the excerpt, restate it from memory, then work through the examples and applications listed on this page. The five-step study plan above takes about twenty minutes.

What does this page cover?

It gives you a compact reference excerpt plus original lgStudy explanations, examples, applications and study material on Ozone–oxygen cycle.