Reactions with Oxygen

Most metals can react with oxygen to form compounds called metal oxides.

The general pattern is:

metal + oxygen → metal oxide

However, metals do not all react with oxygen at the same rate or with the same intensity. Some metals react extremely rapidly, while others require strong heating, react only slowly, or show very little reaction under ordinary conditions.

These differences provide important evidence about a metal's position in the reactivity series.

Metals and Oxygen

Oxygen makes up about 21% of Earth's atmosphere, so metals are frequently exposed to oxygen.

When a metal reacts with oxygen, the metal is oxidized.

For many common metals:

metal + oxygen → metal oxide

Examples include:

magnesium + oxygen → magnesium oxide

copper + oxygen → copper oxide

iron + oxygen → iron oxide

calcium + oxygen → calcium oxide

The product contains both the metal and oxygen.

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What Is a Metal Oxide?

A metal oxide is a compound containing a metal chemically combined with oxygen.

Examples include:

Metal oxides have different properties from the metals from which they formed.

For example, magnesium is a shiny metallic solid, while magnesium oxide is a white solid.

This is evidence that a chemical reaction has occurred.

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Oxidation

When a substance reacts with oxygen, we can describe the process as oxidation.

For example:

magnesium + oxygen → magnesium oxide

Magnesium has been oxidized.

Similarly:

copper + oxygen → copper oxide

Copper has been oxidized.

At a more advanced level, oxidation can also be defined as the loss of electrons.

For metals, these two ideas are closely connected.

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Magnesium and Oxygen

Magnesium provides one of the clearest examples of a metal reacting with oxygen.

When magnesium ribbon is heated strongly, it begins to burn with an extremely bright white light.

A white powder forms.

This powder is mainly magnesium oxide.

Word equation:

magnesium + oxygen → magnesium oxide

Balanced symbol equation:

2Mg + O₂ → 2MgO

Observations When Magnesium Burns

You might observe:

• an intense white light
• considerable heat
• magnesium ribbon disappearing
• formation of a white solid
• a rapid reaction once ignition occurs

These observations show that magnesium reacts vigorously with oxygen.

The reaction releases energy to the surroundings and is therefore exothermic.

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Why Does the Mass Increase?

An interesting feature of metal oxidation is that the final product can have a greater mass than the original metal.

Suppose magnesium is heated in oxygen.

Before:

magnesium

After:

magnesium oxide

The magnesium has combined with oxygen from the surrounding air.

Therefore:

mass of magnesium oxide = mass of magnesium + mass of oxygen that reacted

So the mass increases because oxygen atoms have become part of the solid product.

This does not violate conservation of mass.

If the oxygen and metal are both included, total mass is conserved.

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Worked Example: Mass Increase

Suppose 2.4 g of magnesium reacts completely with oxygen and produces 4.0 g of magnesium oxide.

Mass of oxygen gained:

mass of oxygen = mass of magnesium oxide − mass of magnesium

mass of oxygen = 4.0 − 2.4

mass of oxygen = 1.6 g

Therefore, 1.6 g of oxygen combined with the magnesium.

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Copper and Oxygen

Copper also reacts with oxygen, but it is less reactive than magnesium.

When copper is heated strongly in air, its surface changes.

Copper metal has a characteristic reddish-brown colour.

When heated in oxygen, a black coating of copper(II) oxide can form.

Word equation:

copper + oxygen → copper oxide

A commonly encountered balanced symbol equation is:

2Cu + O₂ → 2CuO

Comparing Magnesium and Copper

If magnesium and copper are heated in air:

Magnesium

• reacts rapidly once ignited
• burns with an intense white light
• releases substantial heat
• forms magnesium oxide

Copper

• reacts more slowly
• generally requires sustained heating
• forms a dark oxide coating
• does not burn as dramatically as magnesium

Therefore, the observations support the conclusion that:

magnesium is more reactive than copper

This agrees with the reactivity series.

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Calcium and Oxygen

Calcium is a reactive metal and reacts readily with oxygen.

Word equation:

calcium + oxygen → calcium oxide

Balanced symbol equation:

2Ca + O₂ → 2CaO

Calcium oxide is a white solid.

Because calcium is relatively high in the reactivity series, it reacts with oxygen more readily than many metals lower in the series.

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Aluminium and Oxygen

Aluminium presents an interesting case.

Aluminium is relatively reactive, so we might expect it to react rapidly with oxygen.

It does react with oxygen:

aluminium + oxygen → aluminium oxide

Balanced equation:

4Al + 3O₂ → 2Al₂O₃

However, aluminium objects often seem surprisingly resistant to further reaction.

Why?

Because aluminium quickly forms a thin, strongly attached layer of aluminium oxide.

Protective Oxide Layers

The aluminium oxide layer prevents oxygen from easily reaching the aluminium underneath.

This greatly slows further oxidation.

This process is called passivation.

The important point is:

Aluminium is reactive, but its oxide layer protects it.

Therefore, visible reaction rate does not always directly reveal a metal's underlying chemical reactivity.

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Zinc and Oxygen

Zinc reacts with oxygen when heated.

Word equation:

zinc + oxygen → zinc oxide

Balanced symbol equation:

2Zn + O₂ → 2ZnO

Zinc oxide is an interesting substance because it can show a colour change when strongly heated.

In typical classroom demonstrations, zinc oxide may appear:

• yellow when hot
• white when cool

The important chemical change is the formation of zinc oxide from zinc and oxygen.

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Iron and Oxygen

Iron reacts with oxygen, but its behaviour depends strongly on the conditions.

When iron is strongly heated, it can react with oxygen to form iron oxides.

Very fine iron particles or iron wool have a large surface area and can react much more rapidly than a solid iron nail.

Surface Area Matters

Why does iron wool react much more readily than an iron nail?

Iron wool contains many thin strands.

This gives it a much larger surface area exposed to oxygen.

Greater surface area allows more metal particles to interact with oxygen at the same time.

Therefore:

greater surface area → potentially faster reaction

This is an important distinction:

reactivity describes the chemical tendency to react.

reaction rate describes how quickly the reaction occurs under particular conditions.

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Iron Wool and Mass

When iron wool is heated in oxygen, its mass can increase.

Again, this is because oxygen becomes incorporated into the solid product.

Before:

iron

After:

iron oxide

Therefore:

mass after reaction > original mass of iron

if only the original iron was weighed initially.

This experiment provides useful evidence that oxygen has combined chemically with the metal.

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Gold and Oxygen

Gold is near the bottom of the reactivity series.

Under ordinary conditions, gold shows very little tendency to react with oxygen.

This is one reason gold:

• remains shiny
• resists corrosion
• can survive for extremely long periods
• has been widely used for jewellery and decorative objects

Gold's resistance to oxidation is a consequence of its very low reactivity.

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Silver and Oxygen

Silver is also relatively unreactive.

It does not readily oxidize in ordinary air in the same way as highly reactive metals.

Silver objects can tarnish, but common silver tarnish involves reactions with sulfur-containing substances, not simply the formation of silver oxide from oxygen.

This is an important example of why visible surface changes should not automatically be described as oxidation by oxygen without considering the chemistry involved.

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Comparing Reactions with Oxygen

A simplified comparison might look like this:

The exact behaviour depends on conditions such as:

• temperature
• surface area
• oxygen concentration
• physical form of the metal
• presence of an oxide coating

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Oxygen and the Reactivity Series

Reactions with oxygen provide evidence for the reactivity series.

In general, metals higher in the series tend to react more readily with oxygen than metals lower in the series.

For example:

magnesium > copper

Magnesium burns vigorously when ignited.

Copper requires heating and reacts much less dramatically.

Similarly:

copper > gold

Copper can be oxidized relatively readily when heated.

Gold resists oxidation.

Using Observations to Rank Metals

Suppose three unknown metals are heated equally in oxygen.

Metal A

Burns rapidly with a bright flame.

Metal B

Slowly develops an oxide coating.

Metal C

Shows no obvious reaction.

Based on these observations alone, a reasonable preliminary ranking would be:

A > B > C

However, scientists should be cautious.

Factors such as protective oxide layers and differences in surface area can affect what we observe.

Additional experiments would strengthen the conclusion.

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Writing Word Equations

A word equation shows the names of the reactants and products.

The general pattern is:

metal + oxygen → metal oxide

Examples:

magnesium + oxygen → magnesium oxide

calcium + oxygen → calcium oxide

zinc + oxygen → zinc oxide

copper + oxygen → copper oxide

Word equations are useful because they clearly show what substances are reacting and what substance is produced.

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Writing Symbol Equations

Symbol equations use chemical formulas.

For magnesium:

Mg + O₂ → MgO

However, this is not balanced.

There are two oxygen atoms on the left but only one on the right.

Balanced:

2Mg + O₂ → 2MgO

Now there are:

• 2 Mg atoms on each side
• 2 O atoms on each side

The equation obeys conservation of atoms.

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Worked Example: Calcium Oxide

Start with the word equation:

calcium + oxygen → calcium oxide

Write formulas:

Ca + O₂ → CaO

Count oxygen atoms.

Left: 2 O

Right: 1 O

Place 2 before CaO:

Ca + O₂ → 2CaO

Now there are two calcium atoms on the right, so place 2 before Ca:

2Ca + O₂ → 2CaO

Balanced equation:

2Ca + O₂ → 2CaO

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Worked Example: Aluminium Oxide

Word equation:

aluminium + oxygen → aluminium oxide

Formulas:

Al + O₂ → Al₂O₃

We need to balance both elements.

Balanced equation:

4Al + 3O₂ → 2Al₂O₃

Check:

Left:

Al = 4
O = 6

Right:

Al = 4
O = 6

The equation is balanced.

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Worked Example: Zinc Oxide

Word equation:

zinc + oxygen → zinc oxide

Formula equation:

Zn + O₂ → ZnO

Balanced:

2Zn + O₂ → 2ZnO

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Oxidation as Electron Loss

At a more detailed level, metal oxidation involves electrons.

When magnesium forms magnesium ions:

Mg → Mg²⁺ + 2e⁻

Magnesium loses electrons.

Therefore, magnesium is oxidized.

This leads to the more general definition:

Oxidation = loss of electrons

A useful memory aid is:

OIL RIG

Oxidation Is Loss

Reduction Is Gain

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Formation of Magnesium Oxide

During magnesium oxidation, magnesium atoms lose electrons:

Mg → Mg²⁺ + 2e⁻

Oxygen gains electrons and forms oxide ions.

The resulting Mg²⁺ and O²⁻ ions attract one another to form ionic magnesium oxide.

So the reaction:

2Mg + O₂ → 2MgO

can also be understood as an electron-transfer reaction.

This connects oxidation reactions to the study of redox chemistry.

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Oxidation and Reduction Occur Together

If magnesium loses electrons, something else must gain them.

In magnesium oxide formation:

• magnesium is oxidized
• oxygen is reduced

Therefore, the overall reaction is a redox reaction.

The term redox comes from:

REDuction

•  

OXidation

These processes occur together because electrons lost by one species must be gained by another.

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Oxidation vs Corrosion

Oxidation and corrosion are related but are not identical terms.

Oxidation describes a chemical process involving electron loss and, in many introductory examples, reaction with oxygen.

Corrosion refers to the gradual chemical deterioration of a material through reactions with its surroundings.

For example, iron corrosion produces rust under suitable conditions.

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Rusting Is More Complicated

Rusting requires both:

• oxygen
• water

Iron reacts through a series of processes that eventually produce hydrated iron oxides commonly called rust.

Therefore, rusting is not simply the same as rapidly burning iron in oxygen.

Environmental conditions are important.

Why Oxidation Matters in Everyday Life

Metal oxidation has many practical consequences.

Corrosion

Oxidation can damage structures, vehicles, tools and machinery.

Protective Coatings

Paint and other coatings can reduce contact between metals and oxygen.

Galvanizing

Zinc coatings can protect iron and steel.

Aluminium Products

Aluminium's oxide layer helps protect objects from further corrosion.

Jewellery

Low-reactivity metals such as gold resist oxidation and remain attractive for long periods.

Metal Extraction

Understanding metal oxides is important when extracting metals from ores.

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Oxide Layers Can Behave Differently

Not every oxide layer has the same effect.

Some oxide layers are:

• thin
• strongly attached
• relatively impermeable

These can protect the metal beneath.

Aluminium oxide is a good example.

Other corrosion products may be:

• porous
• flaky
• poorly attached

These may expose fresh metal underneath.

Rust on iron does not protect iron as effectively as aluminium oxide protects aluminium.

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Designing an Experiment

Suppose you wanted to compare the reactions of magnesium, zinc, iron and copper with oxygen.

You could heat similar samples under controlled conditions.

To make the comparison fair, try to control:

• mass of metal
• surface area
• heating time
• heating intensity
• oxygen availability
• apparatus

Possible observations include:

• whether ignition occurs
• flame or glow
• colour change
• mass change
• speed of oxide formation

Safety precautions would also be essential because some metal oxidation reactions can be extremely vigorous.

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Interpreting Experimental Evidence

Imagine the following observations:

A reasonable interpretation is:

A is probably the most reactive

followed by:

B

then:

C

with:

D probably the least reactive

However, these conclusions should ideally be checked against other experiments because reaction rate can be affected by more than reactivity alone.

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Common Mistakes

Thinking Oxygen Is Always Written as O

Elemental oxygen normally exists as:

O₂

not O.

Forgetting the Product

For the basic reaction:

metal + oxygen → metal oxide

Thinking Oxidation Always Produces a Flame

Many oxidation reactions occur without visible flames.

Thinking All Metals React Equally with Oxygen

Different metals have very different reactivities.

Assuming a Faster Reaction Always Means Greater Reactivity

Temperature, surface area and oxide coatings can also affect reaction rate.

Thinking Aluminium Is Unreactive

Aluminium is reactive but protected by a surface oxide layer.

Thinking an Increase in Mass Violates Conservation of Mass

The additional mass comes from oxygen entering the solid product.

Forgetting to Balance Oxygen

For example:

Mg + O₂ → MgO

is not balanced.

Correct:

2Mg + O₂ → 2MgO

Changing Chemical Formulas to Balance an Equation

Never change MgO into another formula just to make the atom numbers match.

Change the coefficients, not the chemical formulas.

Confusing Oxidation with Rusting

Rusting is a particular corrosion process involving iron, oxygen and water.

Thinking Gold Cannot React with Anything

Gold is very unreactive compared with most metals, but saying it can never undergo any chemical reaction is too strong.

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Key Terms

Oxygen — A reactive non-metal element that normally exists as O₂ molecules.

Metal oxide — A compound containing a metal chemically combined with oxygen.

Oxidation — A process involving loss of electrons; often introduced as reaction with oxygen.

Reactivity — The tendency of a substance to undergo chemical reactions.

Reactivity series — A ranking of metals according to their relative reactivity.

Combustion — A rapid reaction involving oxidation that releases energy.

Exothermic reaction — A reaction that transfers energy to the surroundings.

Reactant — A substance present at the beginning of a chemical reaction.

Product — A substance formed during a chemical reaction.

Word equation — An equation using substance names to represent a chemical reaction.

Symbol equation — An equation using chemical formulas.

Balanced equation — A chemical equation containing equal numbers of each type of atom on both sides.

Coefficient — A number placed before a chemical formula when balancing an equation.

Conservation of mass — The principle that total mass is conserved in a closed chemical reaction.

Oxide ion — An O²⁻ ion.

Electron transfer — Movement of electrons between particles during a redox reaction.

Reduction — Gain of electrons.

Redox reaction — A reaction involving oxidation and reduction.

Surface area — The amount of material exposed to the surroundings.

Oxide layer — A layer of oxide formed on the surface of a metal.

Passivation — Protection caused by formation of a stable surface layer that slows further reaction.

Corrosion — Gradual chemical deterioration of a material through reactions with its environment.

Rusting — Corrosion of iron involving oxygen and water.

Rust — Hydrated iron oxide products formed during corrosion of iron.

Galvanizing — Coating iron or steel with zinc to protect against corrosion.

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Key Takeaways

• Most metals can react with oxygen to form metal oxides.
• The general word equation is:

metal + oxygen → metal oxide

• Reaction with oxygen is an example of oxidation.
• At a more advanced level, oxidation means loss of electrons.
• Metals differ greatly in how vigorously they react with oxygen.
• More reactive metals generally react more readily with oxygen.
• Reactions with oxygen therefore provide evidence about the reactivity series.
• Magnesium burns vigorously in oxygen and forms magnesium oxide.
• Copper reacts more slowly with oxygen when heated.
• Gold is very resistant to oxidation under ordinary conditions.
• Aluminium is reactive but develops a protective aluminium oxide layer.
• Protective oxide layers can make some reactive metals appear less reactive.
• Surface area can strongly affect reaction rate.
• Fine metal wool or powder can react much faster than a large piece of the same metal.
• When a metal forms an oxide, the solid product can have greater mass than the original metal because oxygen has been added.
• Conservation of mass still applies when all reactants and products are considered.
• Metal-oxygen equations must be balanced by changing coefficients, not formulas.
• Elemental oxygen is normally represented as O₂.
• Magnesium oxidation can be represented as:

2Mg + O₂ → 2MgO

• Calcium oxidation can be represented as:

2Ca + O₂ → 2CaO

• Zinc oxidation can be represented as:

2Zn + O₂ → 2ZnO

• Aluminium oxidation can be represented as:

4Al + 3O₂ → 2Al₂O₃

• During metal oxidation, metal atoms lose electrons.
• Oxygen gains electrons, so oxidation and reduction occur together.
• Oxidation reactions are therefore examples of redox reactions.
• Oxidation is important in combustion, corrosion, metal extraction and corrosion prevention.
• Experimental observations must be interpreted carefully because temperature, surface area and protective coatings also affect reaction rate.

The central relationship is:

more reactive metal → generally greater tendency to react with oxygen

And the fundamental reaction pattern is:

metal + oxygen → metal oxide

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Check Your Understanding

1. What is a metal oxide?

2. Write the general word equation for a metal reacting with oxygen.

3. What is meant by oxidation?

4. Name the product when magnesium reacts with oxygen.

5. Write the word equation for magnesium reacting with oxygen.

6. Write the balanced symbol equation for magnesium reacting with oxygen.

7. Describe two observations you might make when magnesium burns.

8. Why does magnesium oxide have a greater mass than the original magnesium?

9. A sample of magnesium has a mass of 3.0 g. After complete oxidation, the magnesium oxide has a mass of 5.0 g. What mass of oxygen reacted?

10. What happens when copper is strongly heated in oxygen?

11. Which reacts more vigorously with oxygen: magnesium or copper?

12. What does this suggest about their relative positions in the reactivity series?

13. Write the word equation for zinc reacting with oxygen.

14. Balance:

Zn + O₂ → ZnO

15. Balance:

Al + O₂ → Al₂O₃

16. Why can aluminium appear surprisingly resistant to oxidation?

17. What is passivation?

18. Why does iron wool react more rapidly with oxygen than a large iron nail?

19. Explain the difference between reactivity and reaction rate.

20. Why does gold show very little reaction with oxygen under ordinary conditions?

21. What happens to metal atoms, in terms of electrons, during oxidation?

22. What happens to oxygen during formation of an ionic metal oxide?

23. Explain why metal oxidation can be described as a redox reaction.

24. Explain the difference between oxidation and rusting.

25. Why are oxygen and water both important in the rusting of iron?

26. Why does aluminium oxide protect aluminium more effectively than rust protects iron?

27. Name three variables that should be controlled when experimentally comparing metal reactions with oxygen.

28. Explain how reactions with oxygen can provide evidence for the reactivity series.

29. Why should observations from oxygen reactions ideally be combined with evidence from other types of reactions?

30. Challenge: Four metals, A, B, C and D, are heated under similar conditions.

• A burns rapidly with an intense flame.
• B glows strongly and quickly forms an oxide.
• C slowly develops an oxide coating.
• D shows no obvious reaction.

a. Suggest an order of reactivity from greatest to least.
b. Which observation provides evidence for the most reactive metal?
c. Which metal appears least reactive?
d. Explain why this ranking should be treated as preliminary rather than absolute.
e. Describe how surface area could affect the observations.
f. Describe how temperature could affect the observations.
g. Explain how a protective oxide layer could produce misleading results.
h. Suggest a second type of experiment that could be used to compare these metals.
i. Explain how the second experiment could strengthen the proposed reactivity order.
j. If the mass of B increases from 5.2 g to 7.0 g during oxidation, calculate the mass of oxygen that combined with B.
k. Explain why this mass increase does not violate conservation of mass.
l. Explain how the formation of the metal oxide involves electron transfer.