Preparation of Oxygen and Its Properties

Lessons

Senior 2 · Chemistry

Preparation of Oxygen and Its Properties

4 lessons

Lesson 1 of 4

Preparing Oxygen in the Laboratory and in Industry

25 min

Why this matters

About 21% of the air you breathe, roughly one-fifth by volume, is oxygen. That share keeps every animal alive, makes charcoal glow in an imbabura and lets a welder cut through steel. In hospitals, patients with pneumonia and premature babies often need air that is richer in oxygen than the air around us, so oxygen has to be produced, stored and delivered safely.

You cannot scoop pure oxygen out of the air in a school laboratory. Instead, chemists release it from compounds that contain a lot of oxygen. In this lesson you will learn the standard laboratory methods, the chemistry behind each one, and how oxygen is produced on a large scale in industry.

The idea behind every laboratory method

In the laboratory, oxygen is prepared by breaking down an oxygen-rich compound. A reaction in which one substance breaks down into two or more simpler substances is a decomposition reaction. Some compounds decompose only when heated (thermal decomposition); others decompose at room temperature once a catalyst is added.

A catalyst is a substance that speeds up a chemical reaction without being used up. At the end of the reaction it has the same mass and the same chemical composition as at the start, so it can be filtered off, dried and used again.

Method 1: hydrogen peroxide and manganese(IV) oxide (the usual school method)

Hydrogen peroxide, H₂O₂, is a colourless liquid, usually sold as a dilute solution in water; pharmacies sell a weak solution as an antiseptic. On its own it decomposes very slowly. Add a little black manganese(IV) oxide powder, MnO₂ (often called manganese dioxide), and the solution fizzes steadily as oxygen is given off:

2H₂O₂(aq) → 2H₂O(l) + O₂(g) (catalyst: MnO₂)

Apparatus. About 1 g of manganese(IV) oxide is placed in a flat-bottomed flask closed by a two-holed stopper. A tap funnel (dropping funnel) containing the hydrogen peroxide solution passes through one hole. A delivery tube passes through the other and leads under the water in a trough, up into a gas jar that was filled with water and turned upside down. When the tap is opened, peroxide drips onto the catalyst, and bubbles of oxygen travel along the tube and push the water out of the jar. You will study this way of collecting gases in the next lesson.

Schools prefer this method because it needs no heating, the rate is easy to control with the tap, and the only by-product is water.

Worked example: proving that MnO₂ is a catalyst

Uwase weighs out 1.00 g of manganese(IV) oxide and uses it to decompose 50 cm³ of hydrogen peroxide solution. Then she filters the mixture, washes and dries the black solid, and weighs it again.

Result 1: the dried solid still has a mass of 1.00 g and is still a black powder.
Result 2: when she adds the recovered solid to fresh hydrogen peroxide, it fizzes again just as fast.
Conclusion: the solid speeded up the reaction but was neither used up nor changed chemically, so it is a catalyst. All of the oxygen came from the hydrogen peroxide, none from the MnO₂.

Method 2: heating potassium chlorate with a catalyst

Potassium chlorate, KClO₃, is a white solid. When it is heated with a little manganese(IV) oxide as a catalyst, it gives off oxygen and leaves potassium chloride:

2KClO₃(s) → 2KCl(s) + 3O₂(g) (heat, MnO₂ catalyst)

Check the balance: there are 2 K, 2 Cl and 6 O atoms on each side. The catalyst lowers the temperature at which the chlorate releases its oxygen. This was the classic laboratory method for many years. It is used less today because potassium chlorate can explode if it is contaminated with carbon, sulfur, sugar or other combustible matter.

Method 3: heating potassium permanganate

Potassium permanganate (systematic name potassium manganate(VII)), KMnO₄, forms purple-black crystals. Heating it strongly releases oxygen:

2KMnO₄(s) → K₂MnO₄(s) + MnO₂(s) + O₂(g) (heat)

The solid left behind is a mixture of green potassium manganate(VI) and black manganese(IV) oxide. The method is simple and fairly safe, but each gram of solid gives only a small volume of oxygen.

Other ways of obtaining oxygen

Heating the nitrate of a very reactive metal: 2KNO₃(s) → 2KNO₂(s) + O₂(g). The white solid left is potassium nitrite.
Heating mercury(II) oxide: 2HgO(s) → 2Hg(l) + O₂(g). In 1774 Joseph Priestley obtained oxygen by focusing sunlight onto this red powder. Schools no longer use the method because mercury vapour is highly poisonous.
Electrolysis of water: passing a direct current through water made slightly acidic gives hydrogen at one electrode and oxygen at the other: 2H₂O(l) → 2H₂(g) + O₂(g).

Oxygen from air: the industrial method

Most of the oxygen sold in cylinders is separated from air by the fractional distillation of liquid air:

Air is filtered to remove dust. Water vapour and carbon dioxide are removed, because they would freeze and block the pipes.
The air is compressed and cooled repeatedly until it becomes a liquid at about −200 °C.
The liquid air is allowed to warm up slowly in a tall fractionating column. Nitrogen has the lower boiling point (−196 °C), so it boils off first and leaves from the top. Oxygen boils at −183 °C, so it stays behind as a liquid and is drawn off.

Many hospitals also run small on-site units called oxygen concentrators and oxygen plants, which pass compressed air through a material that holds back nitrogen and lets oxygen-rich air through.

Oxygen in nature

Green plants make oxygen by photosynthesis, using the energy of sunlight: 6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂. This process keeps the oxygen in the air topped up.

Warning

Safety in oxygen preparations

Wear safety goggles. Hydrogen peroxide stronger than about 6% burns the skin and eyes, so use only the dilute solution your teacher provides and wash off any splash at once with plenty of water.
Hydrogen peroxide, potassium chlorate and potassium permanganate are oxidisers: they make other things burn fiercely. Their containers carry the "flame over a circle" hazard symbol. Keep them away from fuels, paper and oil.
Never grind or mix potassium chlorate with sulfur, charcoal, sugar or any other combustible substance, because the mixture can explode.
When you heat a solid and collect the gas over water, take the delivery tube out of the water before you remove the flame. Otherwise cold water is sucked back into the hot tube (suck-back) and the glass can crack.
Clamp a tube of solid being heated with its mouth slightly lower than its closed end, so that water which condenses cannot run back onto the hot glass. Point it away from everyone and heat gently at first.

Tip

Exam technique: name the catalyst and count the atoms

In an answer about the peroxide method, always name the catalyst (manganese(IV) oxide) and write it beside the arrow; a catalyst never appears among the products. For KClO₃ and KMnO₄, write "heat". Before you move on, count every element on both sides: in 2KClO₃ → 2KCl + 3O₂ there are 6 oxygen atoms on each side, and that is what earns the mark.

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Practice

Try each question before you read the solution under it.

Write a balanced equation, with state symbols, for the preparation of oxygen from hydrogen peroxide. Name the catalyst.

Show solutionHide solution

Solution: 2H₂O₂(aq) → 2H₂O(l) + O₂(g). Check: 4 H and 4 O atoms on each side. The catalyst is manganese(IV) oxide, MnO₂.

Give two reasons why schools prepare oxygen from hydrogen peroxide rather than by heating potassium chlorate.

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Solution: Any two of: no heating is needed; the rate can be controlled with the tap funnel; the only by-product is water; potassium chlorate carries a risk of explosion if contaminated.

68 g of hydrogen peroxide decomposes completely. Calculate the mass of oxygen formed and its volume at STP. (H = 1, O = 16; molar gas volume at STP = 22.4 dm³)

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Solution: M(H₂O₂) = 2(1) + 2(16) = 34 g/mol, so moles of H₂O₂ = 68 ÷ 34 = 2.0 mol. From the equation, 2 mol H₂O₂ give 1 mol O₂, so 1.0 mol O₂ forms. Mass = 1.0 × 32 = 32 g. Volume at STP = 1.0 × 22.4 = 22.4 dm³.

Balance: __KMnO₄ → K₂MnO₄ + MnO₂ + O₂.

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Solution: The right-hand side has 2 K and 2 Mn, so put 2 in front of KMnO₄: 2KMnO₄ → K₂MnO₄ + MnO₂ + O₂. Oxygen check: left 8; right 4 + 2 + 2 = 8.

In the fractional distillation of liquid air, which gas leaves the top of the column first, and why?

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Solution: Nitrogen. Its boiling point (−196 °C) is lower than that of oxygen (−183 °C), so it turns to gas first as the liquid air warms.

Summary

In the laboratory, oxygen is made by decomposing an oxygen-rich compound. The usual school method is 2H₂O₂ → 2H₂O + O₂, with manganese(IV) oxide as the catalyst and no heating.
Heating KClO₃ (with MnO₂) or KMnO₄ also gives oxygen. Guard against contamination of the chlorate and against suck-back.
A catalyst speeds up a reaction and is recovered unchanged in mass and composition.
Industry separates oxygen from air by fractional distillation of liquid air: nitrogen (boiling point −196 °C) boils off before oxygen (boiling point −183 °C).

Sources and credits

Chemistry 2e, section 18.9: Occurrence, Preparation, and Compounds of Oxygen (separating oxygen from liquid air; electrolysis of water; photosynthesis) by Paul Flowers, Klaus Theopold, Richard Langley and William R. Robinson (OpenStax), hosted by LibreTexts, licensed under CC BY 4.0. Adapted by Ganzaa.
Oxygen (history of its discovery; the chlorate method; liquefaction of air), Wikipedia contributors, Wikipedia, licensed under CC BY-SA 4.0. Adapted by Ganzaa.
Image: GHS-pictogram-rondflam.svg, GHS pictogram (UNECE), redrawn by Wikimedia contributors, via Wikimedia Commons, public domain.
Everything else in this lesson, including the worked examples, the safety notes and the Practice questions, was written by Ganzaa.

This lesson adapts CC BY-SA material and is shared under CC BY-SA 4.0.

1Why this matters

2The idea behind every laboratory method

3Method 1: hydrogen peroxide and manganese(IV) oxide (the usual school method)

Worked example: proving that MnO2 is a catalyst

4Method 2: heating potassium chlorate with a catalyst

5Method 3: heating potassium permanganate

6Other ways of obtaining oxygen

7Oxygen from air: the industrial method

8Oxygen in nature

Practice

Why this matters

The idea behind every laboratory method

Method 1: hydrogen peroxide and manganese(IV) oxide (the usual school method)

Worked example: proving that MnO₂ is a catalyst

Method 2: heating potassium chlorate with a catalyst

Method 3: heating potassium permanganate

Other ways of obtaining oxygen

Oxygen from air: the industrial method

Oxygen in nature