What Are Chemical Properties?

A chemical property describes how a substance behaves when it undergoes a chemical reaction.

When a chemical reaction occurs, atoms are rearranged and new substances are formed.

Important chemical properties of metals include how they react with:

• oxygen
• water
• acids
• other metal compounds

Different metals react at very different rates. Some metals are extremely reactive, while others are much less reactive.

These differences are summarized by the reactivity series.

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Why Do Metals React?

Metal atoms generally have relatively few electrons in their outer electron shell.

During many chemical reactions, metal atoms become more stable by losing electrons.

When an atom loses electrons, it forms a positive ion, or cation.

For example:

Na → Na⁺ + e⁻

A sodium atom loses one electron and becomes a sodium ion.

For magnesium:

Mg → Mg²⁺ + 2e⁻

The magnesium atom loses two electrons and forms a 2+ ion.

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Metals Form Positive Ions

The loss of electrons explains an important general property:

Metals tend to form positive ions during chemical reactions.

Common examples include:

• sodium → Na⁺
• potassium → K⁺
• magnesium → Mg²⁺
• calcium → Ca²⁺
• aluminium → Al³⁺
• zinc → Zn²⁺

Transition metals can sometimes form ions with different charges.

For example:

• Fe²⁺
• Fe³⁺
• Cu⁺
• Cu²⁺

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Why Do Metals Lose Electrons?

Metal atoms tend to have outer electrons that can be removed relatively easily compared with many non-metals.

When metals react, these electrons can be transferred to other atoms or species.

For example, when sodium reacts with chlorine:

2Na + Cl₂ → 2NaCl

Sodium atoms lose electrons.

Chlorine atoms gain electrons.

The resulting Na⁺ and Cl⁻ ions attract each other and form ionic sodium chloride.

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Oxidation of Metals

In terms of electron transfer:

oxidation = loss of electrons

Because metals commonly lose electrons during reactions, metals are often oxidized.

For example:

Mg → Mg²⁺ + 2e⁻

Magnesium has lost electrons.

Therefore:

magnesium has been oxidized.

A useful memory aid is:

OIL RIG

Oxidation Is Loss

Reduction Is Gain

of electrons.

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The Reactivity Series

Metals can be arranged according to how readily they react.

A simplified reactivity series is:

Potassium

Sodium

Lithium

Calcium

Magnesium

Aluminium

Carbon

Zinc

Iron

Hydrogen

Copper

Silver

Gold

Carbon and hydrogen are included as useful reference points even though they are not metals.

Metals near the top are generally more reactive.

Metals near the bottom are generally less reactive.

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What Does "More Reactive" Mean?

A more reactive metal loses electrons more readily during appropriate chemical reactions.

For example:

potassium is more reactive than sodium

magnesium is more reactive than zinc

zinc is more reactive than copper

copper is more reactive than silver

A highly reactive metal may react rapidly with water or oxygen.

A less reactive metal may react slowly or require stronger conditions.

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Metals Reacting with Oxygen

Many metals react with oxygen to form metal oxides.

General pattern:

metal + oxygen → metal oxide

For example:

2Mg + O₂ → 2MgO

Magnesium reacts with oxygen to form magnesium oxide.

When magnesium burns, observations include:

• intense white light
• heat released
• formation of a white solid

The white solid is magnesium oxide.

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Evidence of Chemical Reaction

How can we tell that a chemical reaction has occurred?

Possible evidence includes:

• colour change
• gas production
• temperature change
• light being produced
• formation of a new solid
• disappearance of a reactant
• formation of a new substance

Not every reaction shows all of these signs.

Also, an observation such as bubbling does not automatically prove a chemical reaction—boiling can also produce bubbles. Evidence should be interpreted in context.

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Metal + Oxygen Example: Copper

Copper also reacts with oxygen, although less dramatically than magnesium under ordinary laboratory conditions.

When copper is heated strongly in air, its surface can form black copper(II) oxide.

2Cu + O₂ → 2CuO

Observation:

reddish-brown copper → black surface

The colour change provides evidence that a new substance has formed.

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Metal + Oxygen Example: Iron

Iron reacts with oxygen in the presence of water to form rust.

Rusting is a form of corrosion.

Rusting requires:

• iron
• oxygen
• water

Salt can increase the rate of corrosion by helping charge move through the water.

This is why iron objects can corrode more rapidly in salty environments.

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Metals and Water

Some metals react with water.

The reaction depends strongly on the metal's position in the reactivity series.

Very reactive metals can react vigorously with cold water.

Less reactive metals may react slowly, react only with steam, or show little reaction under ordinary conditions.

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Group 1 Metals and Water

Group 1 metals are highly reactive.

For example, sodium reacts with water:

2Na + 2H₂O → 2NaOH + H₂

Products:

• sodium hydroxide
• hydrogen gas

Observations can include:

• movement across the water
• fizzing
• metal becoming smaller
• temperature increase
• hydrogen gas production

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Why Does Fizzing Occur?

Fizzing occurs because a gas is being produced.

In the reaction between sodium and water, the gas is:

hydrogen

Gas production is one piece of evidence that a chemical reaction is occurring.

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Calcium and Water

Calcium also reacts with cold water:

Ca + 2H₂O → Ca(OH)₂ + H₂

Possible observations include:

• bubbles of hydrogen
• calcium gradually disappearing
• cloudy or milky solution due to limited solubility of calcium hydroxide
• temperature change

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Magnesium and Water

Magnesium reacts much more slowly with cold water.

However, magnesium reacts more readily with steam.

Mg + H₂O(g) → MgO + H₂

This demonstrates that reaction conditions matter.

A substance may react slowly under one set of conditions but much faster when temperature is increased.

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Metals and Acids

Many metals react with dilute acids.

A common pattern is:

metal + acid → salt + hydrogen

For example:

Mg + 2HCl → MgCl₂ + H₂

Magnesium reacts with hydrochloric acid to produce:

• magnesium chloride
• hydrogen gas

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Evidence During a Metal-Acid Reaction

When a reactive metal is placed in dilute acid, observations may include:

• bubbles or fizzing
• metal becoming smaller
• temperature increase
• gas being produced

The gas can often be identified as hydrogen.

A common laboratory test for hydrogen uses a lit splint, producing a characteristic small squeaky pop when hydrogen burns.

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Example: Zinc and Hydrochloric Acid

Zn + 2HCl → ZnCl₂ + H₂

Zinc loses electrons:

Zn → Zn²⁺ + 2e⁻

Hydrogen ions gain electrons and eventually form hydrogen gas.

This demonstrates electron transfer during a metal-acid reaction.

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Comparing Metals Using Acid

One way to compare metal reactivity is to place similar pieces of different metals into the same concentration and volume of dilute acid.

Suppose we test:

• magnesium
• zinc
• iron
• copper

Typical observations might be:

Magnesium → vigorous reaction

Zinc → moderate reaction

Iron → slower reaction

Copper → little or no reaction with dilute hydrochloric acid

This supports the order:

Mg > Zn > Fe > Cu

for these conditions.

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How Can Reaction Rate Show Reactivity?

A more reactive metal often reacts more rapidly under the same conditions.

Possible comparisons include:

• amount of bubbling
• rate of gas production
• time required for the metal to disappear
• temperature change

However, a fair comparison requires controlled conditions.

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Designing a Fair Reactivity Test

To compare metals fairly, keep important variables constant.

For example:

• same mass or comparable amount of metal
• similar surface area
• same acid concentration
• same acid volume
• same temperature
• same apparatus

Then change only:

the type of metal

The dependent variable might be:

volume of hydrogen produced per minute

This gives quantitative evidence rather than relying only on visual descriptions such as "lots of bubbles."

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Metals Below Hydrogen

Metals below hydrogen in the reactivity series generally do not displace hydrogen from ordinary dilute non-oxidizing acids such as dilute hydrochloric acid.

Examples include:

• copper
• silver
• gold

Therefore, copper placed in dilute hydrochloric acid normally shows little or no reaction.

This behaviour can be predicted from the reactivity series.

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Displacement Reactions

A more reactive metal can displace a less reactive metal from a solution containing its ions.

For example:

Zn + CuSO₄ → ZnSO₄ + Cu

Zinc is more reactive than copper.

Therefore, zinc can displace copper from copper sulfate solution.

Possible observations include:

• reddish-brown copper forming
• blue copper sulfate solution becoming less intense
• zinc gradually dissolving

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Why Does Zinc Displace Copper?

Zinc loses electrons more readily than copper.

Zinc atoms become ions:

Zn → Zn²⁺ + 2e⁻

Copper ions gain electrons:

Cu²⁺ + 2e⁻ → Cu

Overall:

Zn + Cu²⁺ → Zn²⁺ + Cu

Zinc has been oxidized.

Copper ions have been reduced.

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Predicting Displacement Reactions

Use the reactivity series.

More reactive metal + ions of less reactive metal → reaction

For example:

Fe + CuSO₄ → FeSO₄ + Cu

Iron is more reactive than copper, so displacement occurs.

But:

Cu + FeSO₄ → no reaction

Copper is less reactive than iron and cannot displace it.

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Example 1: Predict the Reaction

Will magnesium react with zinc sulfate?

Compare:

Mg > Zn

Magnesium is more reactive.

Therefore:

Mg + ZnSO₄ → MgSO₄ + Zn

A displacement reaction occurs.

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Example 2: Predict the Reaction

Will silver react with copper sulfate?

Compare:

Cu > Ag

Silver is less reactive than copper.

Therefore:

Ag + CuSO₄ → no reaction

Silver cannot displace copper.

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Metal Reactivity and Electron Loss

The reactivity series can be understood partly in terms of how readily metal atoms lose electrons.

Highly reactive metals lose electrons relatively easily.

Less reactive metals hold their electrons more strongly.

Therefore:

greater tendency to lose electrons → generally greater metal reactivity

This connects metal reactions with atomic structure.

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Group 1 Reactivity

Within Group 1:

reactivity increases down the group

For example:

Li < Na < K

Why?

Going down the group:

• atoms have more occupied electron shells
• atomic radius increases
• shielding increases
• attraction between the nucleus and outer electron decreases
• the outer electron is lost more easily

Therefore, the metals become more reactive.

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Corrosion

Corrosion is the gradual destruction of a metal through chemical reactions with its environment.

Rusting is the corrosion of iron.

Other metals can corrode in different ways.

Corrosion is important because it can weaken:

• bridges
• vehicles
• buildings
• pipelines
• ships
• tools

Understanding metal reactivity helps engineers choose suitable materials and protection methods.

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Aluminium: Reactive but Resistant

Aluminium is relatively high in the reactivity series.

You might therefore expect aluminium objects to corrode very rapidly.

However, aluminium quickly develops a thin layer of aluminium oxide on its surface.

This oxide layer adheres strongly and helps prevent further reaction.

This process is called passivation.

Therefore, chemical behaviour depends not only on whether a metal reacts, but also on what happens to the products formed on its surface.

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Iron Is Different

Rust does not form a tightly protective layer in the same way.

Rust can crack and flake away, exposing fresh iron beneath.

The newly exposed iron can then continue reacting.

This is why rusting can progressively damage iron structures.

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Copper and Patina

Copper reacts relatively slowly with substances in the environment.

Over long periods, copper surfaces can develop a green coating called a patina.

The familiar green colour seen on old copper roofs and statues results from copper compounds formed through environmental reactions.

This is a visible everyday example of metal chemistry.

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Gold and Low Reactivity

Gold is very low in the reactivity series.

It resists many common chemical reactions and does not readily corrode under ordinary conditions.

This contributes to its use in:

• jewellery
• electrical contacts
• specialized electronics

Its usefulness is therefore connected partly to its low chemical reactivity.

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Everyday Example: Galvanising

Iron and steel can be protected by coating them with zinc.

This process is called galvanising.

Zinc provides a physical barrier between iron and the environment.

Zinc is also more reactive than iron, so it can provide additional protection if the coating is damaged.

This is an application of the reactivity series.

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Everyday Example: Sacrificial Protection

More reactive metals can be attached to less reactive metals to protect them from corrosion.

For example, blocks of magnesium or zinc may be used to help protect steel structures.

The more reactive metal oxidizes preferentially.

It is therefore sometimes called a sacrificial metal.

Applications can include:

• ship hulls
• pipelines
• storage tanks
• some water-heating systems

This is a practical use of differences in metal reactivity.

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Everyday Example: Stainless Steel

Another approach to corrosion is to change the material itself.

Stainless steels contain alloying elements, especially chromium, that help form a protective surface layer.

This gives stainless steel much greater corrosion resistance than ordinary carbon steel in many environments.

This makes stainless steel useful for:

• kitchen equipment
• medical instruments
• food-processing equipment
• buildings
• industrial systems

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Choosing Metals for Different Uses

The most reactive metal is not automatically the "best" metal.

Material selection depends on many properties.

For example:

A bridge needs:

• strength
• durability
• manageable cost
• corrosion protection

Electrical wiring needs:

• high conductivity
• ductility
• reasonable chemical stability

Jewellery may require:

• attractive appearance
• low reactivity
• corrosion resistance

Chemical properties must therefore be considered alongside physical and economic properties.

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Comparing Metal Behaviour

Consider four metals:

Magnesium

Reacts readily with acids and burns strongly in oxygen.

Zinc

Reacts with acids and can displace less reactive metals such as copper.

Iron

Reacts with acids but more slowly than magnesium; rusts in moist air.

Copper

Much less reactive; does not normally react with dilute hydrochloric acid.

These observations provide experimental evidence for differences in chemical behaviour.

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Example 3: Identifying an Unknown Metal

Three metals, X, Y, and Z, are tested with dilute hydrochloric acid.

X → vigorous fizzing

Y → slow fizzing

Z → no visible reaction

Based only on these observations:

X appears most reactive with the acid

Y shows intermediate reactivity

Z appears least reactive under these conditions

If the metals were magnesium, iron, and copper, a reasonable identification would be:

X → magnesium

Y → iron

Z → copper

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Example 4: Evidence from Displacement

Metal X displaces copper from copper sulfate.

Metal X does not displace magnesium from magnesium sulfate.

Therefore:

X is more reactive than copper

but:

X is less reactive than magnesium.

So its position must be somewhere between magnesium and copper in the relevant reactivity sequence.

This demonstrates how chemical reactions can be used to determine relative reactivity.

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Reaction Conditions Matter

When comparing metals, conditions must be considered.

Reaction rate can depend on:

• temperature
• concentration
• surface area
• purity
• oxide coatings
• physical form of the metal

For example, metal powder often reacts faster than a large solid piece because the powder has a greater surface area.

Therefore, comparisons should use similar conditions.

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Real-World Observation: Why Metals Are Stored Differently

Very reactive metals require special storage.

For example, some highly reactive metals must be kept away from water and moisture.

Less reactive metals can often be exposed to air with fewer problems.

This difference in storage requirements reflects their chemical properties.

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Real-World Observation: Coastal Corrosion

Iron and steel structures near the sea can experience significant corrosion.

Water and oxygen are necessary for rusting.

Dissolved salts can increase the rate of the electrochemical processes involved.

This is why corrosion protection is especially important for:

• ships
• offshore structures
• coastal bridges
• harbour equipment

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Investigating Metal Reactivity

A useful laboratory investigation could compare the reaction of several metals with dilute acid.

Possible metals:

• magnesium
• zinc
• iron
• copper

Measure:

volume of hydrogen gas produced over time

Rather than simply recording "fast" or "slow," students could plot:

hydrogen volume vs time

A steeper initial graph indicates faster gas production under the tested conditions.

This provides quantitative evidence for comparing reaction rates.

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A Useful Metal-Reaction Strategy

When analyzing a metal reaction:

1. Identify the metal.

2. Locate it in the reactivity series.

3. Identify the other reactant.

Is it:

• oxygen?
• water?
• acid?
• another metal compound?

4. Predict whether a reaction should occur.

5. Predict the products.

6. Look for evidence of chemical change.

7. Consider electron transfer.

Has the metal lost electrons?

8. Compare the reaction with other metals.

9. Connect the result to the reactivity series.

10. Relate the behaviour to practical uses or corrosion.

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Common Mistakes

Mistake 1: Saying metals gain electrons to form positive ions

Metals generally lose electrons to form positive ions.

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Mistake 2: Confusing electron loss with reduction

Loss of electrons is:

oxidation

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Mistake 3: Assuming all metals react equally

Different metals have very different chemical reactivities.

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Mistake 4: Assuming all metals react with cold water

Many do not.

Some react only slowly, with steam, or not appreciably under ordinary conditions.

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Mistake 5: Assuming every metal reacts with dilute hydrochloric acid

Metals below hydrogen in the reactivity series generally do not displace hydrogen from ordinary dilute non-oxidizing acids.

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Mistake 6: Predicting displacement backwards

The more reactive metal displaces the less reactive metal from its compound.

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Mistake 7: Assuming bubbling always proves a chemical reaction

Bubbles indicate gas formation, but boiling can also produce bubbles. Use several observations and the experimental context.

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Mistake 8: Assuming a reactive metal must always visibly corrode quickly

Protective oxide layers can slow further reaction, as with aluminium.

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Did You Know?

The chemical reactivity of metals affects everything from jewellery to bridges.

Gold's low reactivity helps it resist corrosion.

Zinc's greater reactivity allows it to protect steel.

Aluminium's protective oxide layer helps it resist further corrosion.

Iron's tendency to rust means steel structures often require coatings, alloys, or other forms of protection.

The reactivity series is therefore not just a list to memorize. It helps explain why metals behave differently in everyday life and why particular metals are chosen for particular applications.

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Key Terms

• Chemical property: A characteristic describing how a substance behaves during chemical reactions.
• Reactivity: How readily a substance undergoes chemical reactions.
• Reactivity series: Arrangement of metals according to their relative chemical reactivity.
• Cation: Positively charged ion.
• Oxidation: Loss of electrons.
• Reduction: Gain of electrons.
• Metal oxide: Compound formed between a metal and oxygen.
• Displacement reaction: Reaction in which a more reactive element replaces a less reactive element in a compound.
• Corrosion: Gradual chemical deterioration of a material through reactions with its environment.
• Rusting: Corrosion of iron involving oxygen and water.
• Passivation: Formation of a protective surface layer that reduces further reaction.
• Galvanising: Protecting iron or steel with a zinc coating.
• Sacrificial protection: Using a more reactive metal to protect a less reactive metal from corrosion.

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Key Reaction Patterns

Metal with oxygen:

metal + oxygen → metal oxide

Example:

2Mg + O₂ → 2MgO

Metal with dilute acid:

metal + acid → salt + hydrogen

Example:

Mg + 2HCl → MgCl₂ + H₂

Reactive metal with water:

metal + water → metal hydroxide + hydrogen

Example:

2Na + 2H₂O → 2NaOH + H₂

Some metals with steam:

metal + steam → metal oxide + hydrogen

Example:

Mg + H₂O(g) → MgO + H₂

Displacement:

more reactive metal + less reactive metal compound → more reactive metal compound + less reactive metal

Example:

Zn + CuSO₄ → ZnSO₄ + Cu

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Key Takeaways

• Chemical properties describe how metals behave during chemical reactions.
• Metals generally tend to lose electrons.
• Losing electrons produces positive ions called cations.
• Loss of electrons is oxidation.
• Metals differ greatly in chemical reactivity.
• The reactivity series allows us to compare and predict metal behaviour.
• More reactive metals generally lose electrons more readily.
• Many metals react with oxygen to form metal oxides.
• Some reactive metals react with water to produce hydrogen.
• Many metals above hydrogen react with dilute non-oxidizing acids to produce a salt and hydrogen.
• More reactive metals can displace less reactive metals from their compounds.
• Evidence of a metal reaction can include gas production, colour change, temperature change, light, and formation of new substances.
• Reaction conditions must be controlled when comparing metal reactivity.
• Iron corrosion is called rusting and requires oxygen and water.
• Aluminium can resist further corrosion because it forms a protective oxide layer.
• Zinc can protect iron through galvanising and sacrificial protection.
• Gold's low reactivity contributes to its resistance to corrosion.
• Metal reactivity helps explain everyday observations involving buildings, vehicles, jewellery, tools, ships, electronics, and other technologies.
• A useful reasoning sequence is:

identify the metal → consider its tendency to lose electrons → locate it in the reactivity series → identify the other reactant → predict the reaction → look for evidence → compare behaviour → connect the result to real-world use.