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.

 

 

 

6

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.

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.

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.

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

 

 

 

6

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.

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.

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.

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.

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.

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.

 

 

 

6

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.

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.

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.

 

 

 

6

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.

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.

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.

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.

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

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.

 

 

 

5

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.

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.

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.

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

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.

Worked Example: Zinc Oxide

Word equation:

zinc + oxygen → zinc oxide

Formula equation:

Zn + O₂ → ZnO

Balanced:

2Zn + O₂ → 2ZnO

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

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.

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.

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.

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.

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.

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.

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.

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.

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.

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

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.

Reactions with Water

Metals show very different behaviours when they come into contact with water. Some metals react extremely vigorously with cold water, some react slowly, some react only with steam, and others show almost no reaction at all.

These differences provide useful evidence about the reactivity of metals.

For many reactive metals, the general pattern with cold water is:

metal + water → metal hydroxide + hydrogen

However, metals that react with steam commonly form a metal oxide and hydrogen instead.

Studying these reactions allows us to use the reactivity series to predict how an unfamiliar metal may behave.

 

 

 

7

Why Metals React Differently with Water

Metals differ in how readily their atoms lose electrons and form positive ions.

More reactive metals generally lose electrons more readily.

For example:

Na → Na⁺ + e⁻

Sodium loses electrons readily, so it reacts strongly with water.

Copper is much less reactive and does not normally react with cold water or steam under ordinary classroom conditions.

This gives us a general pattern:

more reactive metal → greater tendency to react with water

But the exact behaviour depends on the metal and the conditions.

The Reactivity Series and Water

A useful simplified section of the reactivity series is:

Most reactive

Potassium
Sodium
Lithium
Calcium
Magnesium
Aluminium
Zinc
Iron
Hydrogen
Copper
Silver
Gold

Least reactive

The metals near the top react much more readily with water than metals near the bottom.

A simplified comparison is:

Metals Reacting with Cold Water

Very reactive metals can react directly with cold water.

For several important examples, the products are:

• a metal hydroxide
• hydrogen gas

The general word equation is:

metal + water → metal hydroxide + hydrogen

For example:

sodium + water → sodium hydroxide + hydrogen

The reaction can release a considerable amount of energy.

Potassium and Water

Potassium is one of the most reactive metals commonly encountered when learning about the reactivity series.

It reacts extremely vigorously with cold water.

Word equation:

potassium + water → potassium hydroxide + hydrogen

Balanced symbol equation:

2K + 2H₂O → 2KOH + H₂

The reaction is strongly exothermic.

 

 

 

5

Typical Observations for Potassium

Depending on the conditions, observations can include:

• rapid movement across the water
• vigorous fizzing
• rapid disappearance of the metal
• considerable heat
• ignition of the hydrogen or metal
• a lilac-coloured flame

These observations indicate that potassium is extremely reactive.

Because of the hazards involved, reactions of highly reactive alkali metals with water are normally performed only as controlled demonstrations.

Sodium and Water

Sodium also reacts vigorously with cold water.

Word equation:

sodium + water → sodium hydroxide + hydrogen

Balanced symbol equation:

2Na + 2H₂O → 2NaOH + H₂

Sodium is less reactive than potassium but still reacts rapidly.

 

 

 

5

Typical Observations for Sodium

Sodium may:

• float on the water
• move rapidly across the surface
• fizz
• melt into a small ball
• gradually disappear
• release hydrogen gas

Why might it melt?

The reaction releases heat.

Sodium has a relatively low melting point, so the heat generated can melt the metal.

The resulting sodium hydroxide dissolves in the water, producing an alkaline solution.

Lithium and Water

Lithium also reacts with cold water.

Word equation:

lithium + water → lithium hydroxide + hydrogen

Balanced symbol equation:

2Li + 2H₂O → 2LiOH + H₂

Lithium reacts less vigorously than sodium or potassium.

Typical observations include:

• floating
• fizzing
• movement
• gradual disappearance

This provides a useful comparison within Group 1.

Comparing Lithium, Sodium and Potassium

The order of reactivity is:

potassium > sodium > lithium

As we move down this part of Group 1, reactions with water become more vigorous.

A simplified comparison is:

Lithium

Fizzes and moves relatively slowly.

Sodium

Fizzes more rapidly and often melts.

Potassium

Reacts extremely rapidly and may ignite.

These observations provide experimental evidence for their relative positions in the reactivity series.

 

 

 

5

Calcium and Water

Calcium also reacts with cold water.

Word equation:

calcium + water → calcium hydroxide + hydrogen

Balanced symbol equation:

Ca + 2H₂O → Ca(OH)₂ + H₂

Calcium is less reactive than the alkali metals discussed above, so its reaction is generally less violent.

Observing Calcium and Water

Typical observations can include:

• bubbles forming on the calcium
• gradual disappearance of the metal
• hydrogen gas being produced
• the water becoming cloudy

The cloudiness can result from calcium hydroxide, which is only slightly soluble in water.

The solution becomes alkaline because calcium hydroxide is a base.

 

 

 

4

Testing the Hydrogen

When a suitable metal reacts with water, hydrogen gas can be produced.

The familiar laboratory test for hydrogen is the squeaky pop test.

A small sample of gas is exposed to a flame.

Hydrogen can burn rapidly, producing a characteristic pop.

The reaction is:

hydrogen + oxygen → water

Balanced:

2H₂ + O₂ → 2H₂O

Hydrogen testing should only be carried out using appropriate laboratory procedures and supervision.

Magnesium and Cold Water

Magnesium is less reactive than calcium.

It reacts only very slowly with cold water under ordinary conditions.

A simplified reaction can be represented as:

magnesium + water → magnesium hydroxide + hydrogen

Balanced:

Mg + 2H₂O → Mg(OH)₂ + H₂

However, the reaction can become slow partly because magnesium hydroxide can form on the surface and reduce further contact between magnesium and water.

Magnesium and Steam

Magnesium reacts much more readily with steam.

Importantly, the product changes.

With steam:

magnesium + steam → magnesium oxide + hydrogen

Balanced symbol equation:

Mg + H₂O(g) → MgO + H₂

Notice the difference:

Cold water → metal hydroxide + hydrogen

Steam → metal oxide + hydrogen

for the relevant metals and conditions.

 

 

 

4

Why Does Steam Make a Difference?

Steam involves much higher temperatures than cold water.

At higher temperatures:

• particles have greater kinetic energy
• collisions occur with greater energy
• a larger proportion of collisions can lead to reaction

Therefore, metals that react too slowly with cold water to be easily observed may react more readily with steam.

This is an important reminder that:

reactivity and reaction conditions both influence what we observe.

Zinc and Water

Zinc does not react significantly with cold water under ordinary conditions.

However, hot zinc can react with steam.

Word equation:

zinc + steam → zinc oxide + hydrogen

Symbol equation:

Zn + H₂O(g) → ZnO + H₂

This reaction provides evidence that zinc is reactive enough to remove hydrogen from water under sufficiently energetic conditions, but not reactive enough to react readily with cold water.

Iron and Water

Iron does not normally react significantly with cold water in the same way as calcium or sodium.

However, sufficiently hot iron can react with steam.

A simplified word equation is:

iron + steam → iron oxide + hydrogen

One commonly encountered equation is:

3Fe + 4H₂O(g) → Fe₃O₄ + 4H₂

The exact iron oxide produced depends on reaction conditions.

This is why word equations may sometimes simply refer to iron oxide unless a particular oxide is specified.

Copper and Water

Copper is much less reactive than magnesium, zinc or iron.

Copper does not normally react with:

• cold water
• hot water
• steam

under the conditions normally considered in introductory chemistry.

Therefore:

copper + water → no reaction

This behaviour is consistent with copper's low position in the reactivity series.

Silver and Gold

Silver and gold are even less reactive.

They do not react significantly with water under ordinary conditions.

This contributes to their usefulness in applications where chemical stability is important.

Gold, in particular, can remain relatively unchanged for extremely long periods.

Cold Water vs Steam

This distinction is very important.

Very Reactive Metals

Can react with cold water.

Examples:

• potassium
• sodium
• lithium
• calcium

General pattern:

metal + water → metal hydroxide + hydrogen

Moderately Reactive Metals

May react little or not at all with cold water but can react with steam.

Examples include:

• magnesium
• zinc
• iron

General pattern with steam:

metal + steam → metal oxide + hydrogen

Low-Reactivity Metals

Show little or no reaction even with steam under ordinary classroom conditions.

Examples:

• copper
• silver
• gold

Predicting Reactions Using the Reactivity Series

The reactivity series allows us to make predictions.

Suppose you are asked:

Will calcium react with cold water?

Calcium is relatively high in the reactivity series.

Prediction:

Yes.

Products:

calcium hydroxide + hydrogen

Worked Prediction: Copper

Question:

Will copper react with cold water?

Copper is relatively low in the reactivity series.

Prediction:

No significant reaction.

Copper does not readily displace hydrogen from water under these conditions.

Worked Prediction: Zinc

Question:

Will zinc react vigorously with cold water?

No.

However, zinc can react with steam when heated.

Products:

zinc oxide + hydrogen

Word equation:

zinc + steam → zinc oxide + hydrogen

Comparing Metal Reactivity Experimentally

Suppose four metals are tested with cold water.

Based only on these observations, a reasonable preliminary order is:

A > B > C > D

However, if D reacts strongly with steam, we gain additional information.

Therefore, several experimental conditions may be needed to place metals accurately in a reactivity series.

What Does Vigour Tell Us?

A more vigorous reaction can provide evidence of greater metal reactivity when the experimental conditions are comparable.

For a fair comparison, we should control variables such as:

• mass of metal
• surface area
• water volume
• water temperature
• purity of metal
• observation time

Otherwise, we may incorrectly attribute differences in reaction rate to differences in reactivity.

Surface Area

Imagine comparing:

• a large magnesium block
• finely divided magnesium

The finely divided magnesium has a much larger surface area exposed to the surroundings.

It may therefore react more rapidly.

This does not mean the powdered magnesium is a different position in the reactivity series.

It is still magnesium.

The difference is reaction rate, not fundamental chemical reactivity.

Protective Surface Layers

Surface coatings can also affect metal-water reactions.

Aluminium is an important example.

Aluminium is relatively high in the reactivity series but normally shows little obvious reaction with water because its surface is protected by a thin layer of aluminium oxide.

This oxide layer prevents water from readily reaching the aluminium underneath.

 

 

 

5

Aluminium and the Reactivity Series

If we judged aluminium only by placing aluminium foil in water, we might conclude that it is not very reactive.

That conclusion would be incorrect.

Other chemical evidence shows that aluminium is more reactive than metals such as:

• zinc
• iron
• copper

The protective oxide coating hides its underlying reactivity.

This demonstrates why chemists use multiple experiments when establishing a reactivity series.

Why Is Hydrogen Produced?

During many metal-water reactions, the metal loses electrons.

For example, sodium forms sodium ions:

Na → Na⁺ + e⁻

Species derived from water gain electrons, ultimately producing hydrogen gas.

Therefore, metal-water reactions involve electron transfer.

The metal is oxidized because it loses electrons.

Hydrogen-containing species are reduced as hydrogen gas is formed.

These are therefore redox reactions.

Oxidation and Reduction

Remember:

Oxidation = loss of electrons

Reduction = gain of electrons

A useful memory aid is:

OIL RIG

Oxidation Is Loss

Reduction Is Gain

During a metal-water reaction:

metal → positive metal ions

The metal loses electrons and is oxidized.

Why More Reactive Metals React More Vigorously

More reactive metals have a stronger tendency to undergo oxidation under suitable conditions.

Their atoms can lose electrons more readily.

This means electron transfer can occur more readily when they encounter water.

As a result, metals higher in the reactivity series generally react with water more easily than metals lower in the series.

However, the observed reaction can also depend on:

• temperature
• surface area
• oxide coatings
• concentration or availability of reactants

So the reactivity series describes an underlying chemical tendency, while the visible rate depends on conditions too.

Energy Changes

Many reactions between reactive metals and water are exothermic.

This means they transfer energy to the surroundings.

The temperature can increase.

For highly reactive metals, enough energy may be released to ignite the hydrogen produced.

This helps explain why reactions involving metals such as potassium can become particularly vigorous.

Metal Hydroxides

When some reactive metals react with cold water, they form metal hydroxides.

Examples include:

sodium hydroxide — NaOH

potassium hydroxide — KOH

lithium hydroxide — LiOH

calcium hydroxide — Ca(OH)₂

Metal hydroxides are generally basic, and many produce alkaline solutions in water.

For example, sodium reacting with water produces sodium hydroxide, causing the solution to become strongly alkaline.

Metal Oxides from Steam

When metals such as magnesium and zinc react with steam, they form metal oxides rather than metal hydroxides under the usual conditions studied.

Examples:

Mg + H₂O(g) → MgO + H₂

Zn + H₂O(g) → ZnO + H₂

This distinction is important when predicting reaction products.

Writing Word Equations

Sodium and Water

sodium + water → sodium hydroxide + hydrogen

Calcium and Water

calcium + water → calcium hydroxide + hydrogen

Magnesium and Steam

magnesium + steam → magnesium oxide + hydrogen

Zinc and Steam

zinc + steam → zinc oxide + hydrogen

Iron and Steam

iron + steam → iron oxide + hydrogen

Balancing Metal-Water Equations

Consider:

Na + H₂O → NaOH + H₂

This is not balanced.

Balanced:

2Na + 2H₂O → 2NaOH + H₂

Check:

Left:

Na = 2
H = 4
O = 2

Right:

Na = 2
H = 4
O = 2

The equation is balanced.

Worked Example: Calcium

Start with:

Ca + H₂O → Ca(OH)₂ + H₂

Calcium is already balanced.

To provide the required oxygen atoms, place 2 before H₂O:

Ca + 2H₂O → Ca(OH)₂ + H₂

Check:

Ca = 1 on both sides

O = 2 on both sides

H = 4 on both sides

Balanced equation:

Ca + 2H₂O → Ca(OH)₂ + H₂

Worked Example: Magnesium and Steam

Word equation:

magnesium + steam → magnesium oxide + hydrogen

Formula equation:

Mg + H₂O → MgO + H₂

Count the atoms.

Left:

Mg = 1
H = 2
O = 1

Right:

Mg = 1
H = 2
O = 1

The equation is already balanced:

Mg + H₂O(g) → MgO + H₂

Connecting Water Reactions to the Reactivity Series

The pattern can be summarized approximately as:

very high reactivity → reaction with cold water

moderate reactivity → reaction may require steam

low reactivity → little or no reaction with water

This is one of several ways chemists establish and apply the reactivity series.

 

 

 

4

Real-World Importance

Metal-water reactions help explain why metals must be stored, handled and used differently.

Sodium and Potassium

Because they react strongly with water and moisture, they require special storage conditions.

Aluminium

Its protective oxide layer allows aluminium to be used widely despite its underlying reactivity.

Iron

Iron can be used in many structures even though it can undergo corrosion under suitable environmental conditions.

Copper

Its relatively low reactivity contributes to its usefulness in water pipes and other applications.

Understanding reactivity therefore helps engineers choose appropriate materials.

Common Mistakes

Assuming All Metals React with Cold Water

They do not.

Some require steam, while others do not react significantly with water.

Assuming Metal + Water Always Produces an Oxide

For many highly reactive metals with cold water:

metal hydroxide + hydrogen

is produced.

With steam, certain metals produce:

metal oxide + hydrogen

Forgetting Hydrogen

Hydrogen gas is an important product in many metal-water reactions.

Confusing Hydrogen with Oxygen

The gas commonly released is:

H₂

not O₂.

Thinking No Visible Reaction Means a Metal Is Completely Unreactive

It may require higher temperature or steam, or it may have a protective surface layer.

Thinking Aluminium Is Low in the Reactivity Series

Aluminium is reactive but protected by an oxide coating.

Comparing Reactions Under Different Conditions

A fair comparison requires similar mass, surface area, temperature and other conditions.

Confusing Reactivity with Reaction Rate

Reactivity is an underlying chemical tendency.

Reaction rate also depends on experimental conditions.

Forgetting That Steam Is Water

Steam is gaseous water:

H₂O(g)

Forgetting the Product Difference

Remember:

cold water → often hydroxide + hydrogen

steam → often oxide + hydrogen

for the metals typically studied.

Changing Chemical Formulas When Balancing

Use coefficients.

Do not alter the formulas of the compounds.

Key Terms

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

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

Cold water reaction — A reaction between a metal and liquid water without requiring strong heating.

Steam — Water in the gas phase, H₂O(g).

Metal hydroxide — A compound containing a metal ion and hydroxide ions.

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

Hydrogen gas — H₂, a gas produced in many reactions between reactive metals and water.

Alkali — A soluble base that produces an alkaline solution in water.

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

Oxidation — Loss of electrons.

Reduction — Gain of electrons.

Redox reaction — A reaction in which oxidation and reduction occur together.

Electron transfer — Movement of electrons from one species to another.

Metal ion — A positively charged particle formed when a metal atom loses electrons.

Surface area — The amount of a material exposed to its surroundings.

Reaction rate — How quickly reactants are converted into products.

Passivation — Protection of a metal by a surface layer that slows further reaction.

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

Word equation — A chemical equation written using substance names.

Symbol equation — A chemical equation written using chemical formulas.

Balanced equation — An equation with equal numbers of each type of atom on both sides.

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

Squeaky pop test — A common laboratory test used to identify a small sample of hydrogen gas.

Key Takeaways

• Different metals react with water in very different ways.
• The differences provide evidence about their relative reactivity.
• Metals near the top of the reactivity series generally react more readily with water.
• Potassium reacts extremely vigorously with cold water.
• Sodium reacts very vigorously with cold water.
• Lithium reacts with cold water but less vigorously than sodium and potassium.
• Calcium reacts readily with cold water.
• Magnesium reacts only slowly with cold water under ordinary conditions.
• Magnesium reacts more readily with steam.
• Zinc and iron can react with steam under suitable conditions.
• Copper, silver and gold do not normally react significantly with water.
• For many reactive metals:

metal + cold water → metal hydroxide + hydrogen

• For several moderately reactive metals:

metal + steam → metal oxide + hydrogen

• Hydrogen gas can be identified using the appropriate laboratory test.
• Metal hydroxides are basic, and soluble hydroxides form alkaline solutions.
• More reactive metals generally lose electrons more readily.
• Metal-water reactions involve electron transfer and are redox reactions.
• The metal is oxidized during the reaction.
• Many metal-water reactions are exothermic.
• More vigorous reactions generally indicate greater reactivity when experimental conditions are comparable.
• Temperature can strongly affect the observed reaction rate.
• Surface area can strongly affect reaction rate.
• Protective oxide layers can prevent an apparently reactive metal from reacting visibly.
• Aluminium is an important example of passivation.
• A lack of reaction with cold water does not necessarily mean that a metal cannot react with steam.
• Experimental evidence from water reactions can help establish the reactivity series.
• The reactivity series can then be used to predict how metals will behave with water.

The main patterns to remember are:

More reactive metal → generally reacts more readily with water

Reactive metal + cold water → metal hydroxide + hydrogen

Some moderately reactive metals + steam → metal oxide + hydrogen

Check Your Understanding

1. Why do different metals react differently with water?

2. Write the general word equation for a reactive metal reacting with cold water.

3. What gas is commonly produced during metal-water reactions?

4. Describe two observations when sodium reacts with water.

5. Write the word equation for sodium reacting with water.

6. Write the balanced symbol equation for sodium reacting with water.

7. Why can sodium melt during its reaction with water?

8. Arrange lithium, sodium and potassium from least reactive to most reactive.

9. How does the reaction with water change as you move from lithium to potassium?

10. Write the word equation for calcium reacting with water.

11. Write the balanced symbol equation for calcium reacting with water.

12. Why might water become cloudy when calcium reacts with it?

13. How can hydrogen gas be identified in a laboratory?

14. Describe magnesium's reaction with cold water.

15. How does magnesium's reaction with steam differ from its reaction with cold water?

16. Write the word equation for magnesium reacting with steam.

17. Write the balanced symbol equation for magnesium reacting with steam.

18. What are the products when zinc reacts with steam?

19. Why can iron react with steam even though it does not react significantly with cold water?

20. Predict whether copper will react with cold water.

21. Predict whether gold will react significantly with steam under ordinary classroom conditions.

22. Explain why aluminium's behaviour with water can be misleading when judging its reactivity.

23. What is passivation?

24. Explain why a powdered metal might react faster than a large piece of the same metal.

25. Explain the difference between reactivity and reaction rate.

26. Why must experimental conditions be controlled when comparing metals?

27. What happens to metal atoms in terms of electrons during a metal-water reaction?

28. Why can metal-water reactions be described as redox reactions?

29. Explain how reactions with water can help establish a reactivity series.

30. Challenge: Four unknown metals, A, B, C and D, are tested.

• A reacts violently with cold water.
• B reacts steadily with cold water.
• C does not react with cold water but reacts with steam.
• D does not react with cold water or steam.

a. Arrange the metals from most reactive to least reactive based on the evidence.
b. Which metal is likely to be highest in the reactivity series?
c. Which is likely to be lowest?
d. Which metal might behave similarly to zinc or iron? Explain.
e. Predict the general products when A reacts with cold water.
f. Predict the general products when C reacts with steam.
g. Explain why hydrogen may be produced in both reactions.
h. Describe how you could confirm that the gas produced is hydrogen.
i. Identify three variables that should be controlled if the reactions are compared experimentally.
j. Explain why surface area could affect the observations.
k. Explain why temperature could affect the observations.
l. Explain why a protective oxide coating might make one of the metals appear less reactive than it really is.
m. Suggest another type of chemical reaction that could provide additional evidence for the relative reactivity of A, B, C and D.

Comparing Metal Reactivities

Different metals have different tendencies to take part in chemical reactions. Some react rapidly with water, acids or oxygen, while others react slowly or show little visible reaction.

Scientists can compare metals by collecting experimental evidence and looking for consistent patterns.

Useful evidence can come from:

• reactions with water
• reactions with acids
• reactions with oxygen
• displacement reactions
• measurements of reaction rate

By combining evidence from several investigations, metals can be arranged into a reactivity series and their behaviour in unfamiliar reactions can be predicted.

 

 

 

5

What Does Metal Reactivity Mean?

Reactivity describes how readily a metal undergoes chemical reactions.

At the particle level, metals commonly react by losing electrons and forming positive ions.

For example:

Mg → Mg²⁺ + 2e⁻

A metal that loses electrons more readily is generally more reactive.

This helps explain why magnesium reacts readily with dilute acids while copper does not under the same conditions.

The Reactivity Series

A simplified reactivity series is:

Most reactive

Potassium
Sodium
Lithium
Calcium
Magnesium
Aluminium
Carbon
Zinc
Iron
Hydrogen
Copper
Silver
Gold

Least reactive

Carbon and hydrogen are not metals but are included as useful reference points.

The reactivity series is not simply something that must be memorized. It is supported by experimental evidence.

Using Experimental Evidence

Suppose four metals are placed separately into identical samples of dilute hydrochloric acid.

Observations:

The evidence suggests:

magnesium > zinc > iron > copper

The faster reactions indicate greater reactivity, provided the experimental conditions were the same.

 

 

 

5

Observations vs Measurements

Scientists distinguish between qualitative and quantitative evidence.

Qualitative Evidence

Descriptive observations such as:

• vigorous fizzing
• slow bubbling
• bright flame
• colour change
• no visible reaction

Quantitative Evidence

Numerical measurements such as:

• 42 cm³ hydrogen produced in 30 seconds
• temperature increased by 8°C
• metal disappeared in 45 seconds
• mass decreased by 0.12 g
• 25 cm³ hydrogen produced per minute

Quantitative evidence usually allows more precise comparisons.

Comparing Reaction Rates

Reaction rate describes how quickly reactants are converted into products.

Suppose equal amounts of magnesium, zinc and iron react with identical samples of hydrochloric acid.

After 20 seconds:

Under these controlled conditions:

magnesium reacted fastest

followed by:

zinc

then:

iron

This supports:

Mg > Zn > Fe

in reactivity.

Rate Is Not the Same as Amount

Suppose two metals eventually produce the same volume of hydrogen:

Metal A: 50 cm³ in 25 seconds

Metal B: 50 cm³ in 90 seconds

Both produce the same amount of gas.

But Metal A has the greater reaction rate.

Therefore:

reaction rate = how quickly something happens

while:

amount of product = how much is eventually produced

These should not be confused.

Measuring Reaction Rate with a Gas Syringe

For metal-acid reactions, hydrogen can be collected using a gas syringe.

The experiment can measure:

hydrogen volume against time

For example:

The magnesium reaction has the greatest initial rate.

 

 

 

4

Interpreting Reaction-Rate Graphs

The gas-volume data can be plotted on a graph.

Horizontal axis:

time

Vertical axis:

volume of hydrogen

The slope tells us about reaction rate.

steeper slope → faster reaction

shallower slope → slower reaction

horizontal line → reaction has stopped producing gas

Therefore, reaction graphs provide quantitative evidence for comparing metal behaviour.

Comparing Initial Rates

The beginning of a reaction is often particularly useful.

Why?

Because conditions are most comparable near the start.

If the initial slope for magnesium is steeper than zinc, and zinc is steeper than iron:

Mg reacts faster than Zn

and:

Zn reacts faster than Fe

This supports the ranking:

Mg > Zn > Fe

 

 

 

Fair Tests Are Essential

Reaction rate is affected by many variables besides the identity of the metal.

If we want to compare metal reactivity, other variables must be controlled.

For a metal-acid investigation, control:

• amount of metal
• metal surface area
• acid concentration
• acid volume
• temperature
• apparatus
• measurement method

The independent variable should be:

type of metal

A possible dependent variable is:

volume of hydrogen produced in a fixed time

Why Surface Area Matters

Imagine comparing:

• 1.0 g of powdered zinc
• 1.0 g of magnesium ribbon

The zinc may have much more exposed surface area.

More particles can interact with the acid at the same time.

Therefore, zinc might appear to react unusually quickly.

This could produce a misleading comparison.

Surface area should therefore be kept as similar as possible.

Why Temperature Matters

At higher temperatures, particles:

• move faster
• collide more frequently
• collide with greater average energy

Therefore, increasing temperature generally increases reaction rate.

If magnesium is tested at 20°C and iron at 60°C, their rates cannot be fairly compared as evidence of intrinsic reactivity.

Why Concentration Matters

A more concentrated acid contains more acid particles in a given volume.

This can increase collision frequency at the metal surface.

Therefore:

higher concentration → generally faster reaction

If different acid concentrations are used for different metals, the results may not accurately represent differences in metal reactivity.

Comparing Metals Using Water

Reactions with water provide another source of evidence.

A simplified pattern is:

Potassium

Extremely vigorous reaction with cold water.

Sodium

Very vigorous reaction with cold water.

Lithium

Vigorous reaction with cold water.

Calcium

Readily reacts with cold water.

Magnesium

Very slow with cold water; reacts more readily with steam.

Zinc and Iron

Can react with steam under suitable conditions.

Copper, Silver and Gold

No significant reaction with water under ordinary conditions.

These observations help establish relative reactivity.

 

 

 

6

Comparing Metals Using Oxygen

Metals can also be compared by observing their reactions with oxygen.

For example:

Magnesium

Burns vigorously when ignited.

Zinc

Reacts when strongly heated.

Iron

Fine iron or iron wool can burn when strongly heated.

Copper

Oxidizes more slowly when heated.

Gold

Shows very little reaction with oxygen under ordinary conditions.

This provides further evidence about relative reactivity.

Why One Experiment Is Not Always Enough

Suppose aluminium is placed in cold water.

You might observe:

no obvious reaction

Does that mean aluminium is unreactive?

No.

Aluminium has a protective surface layer of aluminium oxide.

This layer prevents water from readily reaching the metal underneath.

Other experiments show that aluminium is actually relatively high in the reactivity series.

Therefore:

observations must be interpreted carefully.

 

 

 

4

Combining Evidence

Imagine three metals A, B and C.

Acid Experiment

A reacts rapidly.

B reacts slowly.

C shows no reaction.

This suggests:

A > B > C

Water Experiment

A reacts with cold water.

B reacts only with steam.

C does not react.

Again:

A > B > C

Displacement Experiment

A displaces B.

A displaces C.

B displaces C.

Again:

A > B > C

Because several independent experiments support the same order, we can have greater confidence in the ranking.

Displacement Reactions

Displacement reactions are particularly useful for comparing metal reactivities.

The rule is:

a more reactive metal displaces a less reactive metal from its compound

For example:

zinc + copper sulfate → zinc sulfate + copper

Symbol equation:

Zn + CuSO₄ → ZnSO₄ + Cu

Because zinc displaces copper:

Zn > Cu

 

 

 

6

Using Displacement Results

Suppose:

• A displaces B.
• B displaces C.
• C displaces D.

Then:

A > B

B > C

C > D

Therefore:

A > B > C > D

This method is powerful because it compares two metals directly.

When No Displacement Occurs

Suppose copper is placed into zinc sulfate solution.

No reaction occurs.

We know:

Cu cannot displace Zn

Therefore:

Zn > Cu

A "no reaction" result can therefore provide useful evidence.

Absence of displacement tells us that the solid metal is less reactive than the metal represented by the ions in solution.

Building a Reactivity Series from Unknown Metals

Suppose four unknown metals W, X, Y and Z give these results:

• W displaces X.
• W displaces Y.
• Z displaces W.
• X displaces Y.
• Y cannot displace X.

We can reason:

Z > W

W > X

X > Y

Therefore:

Z > W > X > Y

This is an example of using experimental evidence to construct a reactivity series rather than memorizing one.

Explaining Trends Using Electron Loss

Why is one metal more reactive than another?

Metal atoms generally react by losing electrons.

For example:

Mg → Mg²⁺ + 2e⁻

and:

Zn → Zn²⁺ + 2e⁻

If magnesium forms positive ions more readily under the relevant conditions, it is more reactive.

Therefore:

greater tendency to lose electrons → generally greater metal reactivity

This connects the reactivity series to oxidation.

Reactivity and Oxidation

Oxidation is the loss of electrons.

A reactive metal tends to be oxidized readily.

For example:

Zn → Zn²⁺ + 2e⁻

In a displacement reaction:

Zn + Cu²⁺ → Zn²⁺ + Cu

zinc loses electrons.

Therefore, zinc is oxidized.

Copper ions gain electrons:

Cu²⁺ + 2e⁻ → Cu

Therefore, copper ions are reduced.

Reactivity and Position in the Periodic Table

There are also periodic trends in metal reactivity.

For example, in Group 1, reactivity increases down the group:

lithium < sodium < potassium

Their reactions with water become increasingly vigorous.

This trend can be explained by changes in atomic structure.

As we move down Group 1:

• atoms have more occupied electron shells
• the outer electron is farther from the nucleus
• shielding increases
• nuclear attraction for the outer electron becomes weaker overall
• the outer electron is lost more easily

Therefore, reactivity increases.

Group 2 Metals

A similar general trend occurs down Group 2.

For example:

magnesium < calcium

in their reactions with water.

Calcium reacts readily with cold water, whereas magnesium reacts only very slowly with cold water.

Again, easier loss of outer electrons contributes to increased reactivity down the group.

Predicting an Acid Reaction

Suppose you know:

magnesium > zinc > iron > hydrogen > copper

Question:

Which metals will react with dilute hydrochloric acid to produce hydrogen?

Metals above hydrogen:

• magnesium
• zinc
• iron

will generally react.

Copper is below hydrogen.

Therefore, copper will not normally release hydrogen from dilute hydrochloric acid.

Predicting a Displacement Reaction

Question:

Will magnesium react with copper sulfate?

Compare the metals:

Mg > Cu

Therefore, magnesium can displace copper.

Prediction:

magnesium + copper sulfate → magnesium sulfate + copper

Predicting the Reverse Reaction

Question:

Will copper react with magnesium sulfate?

Compare:

Mg > Cu

Copper is less reactive.

Therefore, copper cannot displace magnesium.

Prediction:

no reaction

Predicting Reactions with Water

Suppose a metal is known to be between calcium and zinc in the reactivity series.

We might predict that it will be:

• less reactive with water than calcium
• more reactive than zinc overall

However, specific reaction conditions and protective coatings must still be considered.

The reactivity series helps make predictions, but chemical knowledge of the particular metal remains important.

Using Evidence to Identify an Unknown Metal

Suppose an unknown metal X:

• reacts slowly with dilute hydrochloric acid
• does not react with cold water
• reacts with steam
• displaces copper from copper sulfate

These observations suggest that X is:

• above hydrogen
• above copper
• not among the extremely reactive metals
• capable of reacting with steam

Metals such as zinc or iron would be reasonable possibilities.

Additional tests could distinguish between them.

Anomalous Results

Sometimes an experimental result does not fit the overall pattern.

Suppose hydrogen volumes after 30 seconds are:

Magnesium:

42 cm³, 41 cm³, 18 cm³, 43 cm³

The 18 cm³ result is very different from the others.

It may be an anomalous result.

Possible causes include:

• gas leakage
• inaccurate timing
• oxide coating
• incorrect metal mass
• equipment problems

Scientists should investigate unusual results rather than simply ignoring them.

Repeating Experiments

Repeated trials improve reliability.

Suppose zinc produces:

22 cm³
24 cm³
23 cm³

Mean:

mean = (22 + 24 + 23) / 3

mean = 23 cm³

Repeated measurements reduce the influence of random variation and make comparisons more convincing.

Reliability vs Validity

These terms describe different aspects of experimental quality.

Reliability refers to whether repeated measurements give consistent results.

Validity refers to whether the experiment actually tests what it is intended to test.

For example, repeating an unfair comparison many times may produce reliable measurements but still not provide a valid comparison of metal reactivity.

A good investigation should aim for both.

Evaluating Evidence

Before ranking metals, ask:

• Were the same conditions used?
• Were measurements quantitative?
• Were trials repeated?
• Were there anomalous results?
• Was surface area controlled?
• Was temperature controlled?
• Was concentration controlled?
• Could an oxide layer have affected the reaction?
• Do different experiments support the same conclusion?

Scientific conclusions should be based on the overall pattern of evidence.

Real-World Importance of Comparing Reactivity

Understanding relative metal reactivity has many practical applications.

Metal Extraction

A metal's position in the reactivity series helps determine how it can be extracted from its ore.

Corrosion Prevention

More reactive metals can sometimes protect less reactive metals.

Material Selection

Engineers choose metals partly according to their chemical stability.

Batteries

Differences in the tendency of metals to undergo oxidation are important in electrochemical cells.

Construction

Materials must be selected according to the environment in which they will be used.

Jewellery

Low-reactivity metals such as gold are useful because they resist chemical change.

Sacrificial Protection

Iron can be protected using a more reactive metal such as zinc.

Because:

Zn > Fe

zinc can oxidize preferentially.

This protects the iron.

The zinc is effectively "sacrificed."

This is an excellent example of using relative metal reactivity to make a practical prediction.

 

 

 

5

Common Mistakes

Assuming the Fastest Visible Reaction Always Means the Most Reactive Metal

This is only valid when conditions are comparable.

Ignoring Surface Area

Powders can react much faster than large pieces of the same metal.

Ignoring Temperature

Hotter reactions generally proceed faster.

Ignoring Concentration

Higher reactant concentrations can increase reaction rate.

Confusing Reaction Rate with Total Product

A reaction can be faster without producing more final product.

Assuming No Visible Reaction Provides No Information

A no-reaction result can be very useful, especially in displacement experiments.

Assuming Aluminium Is Unreactive

Its protective oxide layer can hide its underlying reactivity.

Using Only One Experiment

Several types of evidence provide a stronger conclusion.

Ignoring Anomalous Results

Unexpected data should be investigated.

Confusing Reliability and Validity

Repeating an unfair experiment does not make the comparison valid.

Forgetting the Displacement Rule

More reactive metal displaces less reactive metal.

Reversing the Reactivity Series

Metals near the top are generally more reactive.

Key Terms

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

Reactivity series — An ordering of metals according to their relative reactivity.

Reaction rate — How quickly reactants are converted into products.

Experimental evidence — Observations and measurements used to support a scientific conclusion.

Qualitative data — Descriptive information rather than numerical measurements.

Quantitative data — Numerical measurements.

Observation — Information directly recorded during an investigation.

Inference — An interpretation made from observations or evidence.

Independent variable — The variable deliberately changed.

Dependent variable — The variable measured.

Control variable — A factor kept constant during a fair test.

Fair test — An investigation in which relevant variables other than the independent variable are controlled.

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

Concentration — The amount of dissolved substance in a given volume.

Initial rate — The reaction rate near the beginning of a reaction.

Displacement reaction — A reaction in which a more reactive element replaces a less reactive element in a compound.

Oxidation — Loss of electrons.

Reduction — Gain of electrons.

Redox reaction — A reaction involving both oxidation and reduction.

Metal ion — A positively charged ion formed when a metal atom loses electrons.

Anomalous result — A result that differs substantially from the overall pattern.

Repeat — Performing an experiment again under the same conditions.

Reliability — The extent to which repeated measurements are consistent.

Validity — The extent to which an investigation actually tests the intended question.

Mean — The sum of values divided by the number of values.

Prediction — A statement about an expected result based on scientific evidence or understanding.

Trend — A general pattern shown by data.

Passivation — Formation of a protective surface layer that slows further reaction.

Sacrificial protection — Protection of a metal using a more reactive metal that oxidizes preferentially.

Key Takeaways

• Metals can be compared using experimental evidence.
• Useful evidence comes from reactions with water, acids, oxygen and metal compounds.
• Reaction rates can provide evidence about relative reactivity.
• Faster reaction under identical conditions generally indicates greater reactivity.
• Reaction rate and total amount of product are not the same thing.
• Quantitative measurements usually allow stronger comparisons than simple descriptions.
• Hydrogen volume can be measured to compare metal-acid reaction rates.
• A steeper gas-volume graph represents a faster reaction.
• Experimental conditions must be controlled.
• Surface area affects reaction rate.
• Temperature affects reaction rate.
• Concentration affects reaction rate.
• The type of metal should be the independent variable when comparing metal reactivity.
• Repeated measurements improve reliability.
• Anomalous results should be investigated.
• Reliable results are not automatically valid results.
• Several independent experiments provide stronger evidence than one experiment alone.
• Reactions with water help distinguish highly reactive metals.
• Reactions with dilute acids help compare metals above hydrogen.
• Reactions with oxygen provide additional evidence about reactivity.
• Displacement reactions directly compare the relative reactivity of two metals.
• A more reactive metal can displace a less reactive metal from its compound.
• A less reactive metal cannot displace a more reactive metal.
• More reactive metals generally lose electrons more readily.
• Metal oxidation involves loss of electrons.
• Protective oxide layers can affect observations.
• Aluminium is an important example of a reactive metal protected by an oxide layer.
• Reactivity trends can sometimes be explained using atomic structure.
• Group 1 metals become more reactive down the group.
• Experimental evidence can be used to construct a reactivity series.
• Once a reactivity order is established, it can be used to predict unfamiliar reactions.
• Reactivity comparisons have applications in extraction, corrosion prevention, batteries and material selection.

The most important reasoning pattern is:

collect evidence → compare results → establish reactivity order → use the order to make predictions

And for displacement reactions:

more reactive metal → can displace a less reactive metal

Check Your Understanding

1. What does metal reactivity mean?

2. Why can reactions with acids be used to compare metal reactivity?

3. Explain the difference between qualitative and quantitative evidence.

4. Give two examples of quantitative measurements that could be collected during a metal reactivity investigation.

5. Three metals produce 40 cm³, 23 cm³ and 9 cm³ of hydrogen in 20 seconds under identical conditions. Rank their reaction rates.

6. Why must experimental conditions be identical when comparing the metals?

7. Explain why surface area must be controlled.

8. Explain why acid concentration must be controlled.

9. Explain why temperature must be controlled.

10. What does the slope of a hydrogen-volume-against-time graph represent?

11. What does a steeper slope indicate?

12. Two metals eventually produce the same volume of hydrogen, but one reaches that volume sooner. What can you conclude?

13. Explain how reactions with water provide evidence about reactivity.

14. Explain how reactions with oxygen provide evidence about reactivity.

15. Why might aluminium produce misleading results in some reactivity experiments?

16. State the rule for metal displacement reactions.

17. Zinc displaces copper from copper sulfate. What does this tell you about zinc and copper?

18. Copper does not displace magnesium from magnesium sulfate. What can you conclude?

19. Metal A displaces B, and B displaces C. Rank the three metals.

20. Explain why displacement reactions are useful when constructing a reactivity series.

21. What happens to metal atoms in terms of electrons when they are oxidized?

22. Why are more reactive metals generally more easily oxidized?

23. Describe the trend in reactivity down Group 1.

24. Explain this trend using atomic structure.

25. What is an anomalous result?

26. Why should experiments be repeated?

27. Explain the difference between reliability and validity.

28. Why is evidence from several different experiments more convincing than evidence from only one?

29. Give two real-world situations in which comparing metal reactivity is useful.

30. Challenge: Four unknown metals, P, Q, R and S, are investigated.

Acid experiment:

• P produces 38 cm³ H₂ in 30 seconds.
• Q produces 15 cm³ H₂ in 30 seconds.
• R produces no hydrogen.
• S produces 51 cm³ H₂ in 30 seconds.

Displacement experiments:

• P displaces Q.
• S displaces P.
• Q displaces R.
• R cannot displace Q.

a. Rank P, Q, R and S from most reactive to least reactive.
b. Which result indicates that S is more reactive than P?
c. Which result indicates that P is more reactive than Q?
d. Which result indicates that Q is more reactive than R?
e. Explain why the acid results alone might not provide enough evidence to place R accurately.
f. Explain why the displacement evidence helps.
g. Predict whether S will displace Q.
h. Predict whether R will displace P.
i. If P, Q and S eventually produce the same total volume of hydrogen, explain why their initial rates can still differ.
j. Identify the independent variable in the acid experiment.
k. Identify a suitable dependent variable.
l. Give four control variables.
m. Explain how unequal surface areas could affect the ranking.
n. Suggest how the reliability of the investigation could be improved.
o. Suggest another experiment that could provide additional evidence.
p. Explain why combining all the evidence produces a stronger scientific conclusion than relying on a single observation.