C15: Reactivity of Metals
The reactivity series orders metals by how easily they lose electrons. Understanding reactivity helps predict displacement reactions, explain methods of metal extraction, and describe how to prevent rusting and corrosion.
The Reactivity Series
The reactivity series lists metals in order of their reactivity, from most reactive to least reactive. A more reactive metal loses electrons more easily and forms positive ions more readily.
Metals at the top of the series are more reactive because they lose outer electrons more easily. This is related to their atomic structure: metals with fewer outer electrons and larger atomic radii tend to be more reactive, as the outer electrons are further from the nucleus and more shielded, so they are lost more easily.
| Reactivity | Metal | Reaction with Cold Water | Reaction with Dilute Acid | Reaction with Oxygen |
|---|---|---|---|---|
| Most reactive | Potassium (K) | Violent, catches fire | Explosive | Burns brightly |
| Sodium (Na) | Rapid, fizzes vigorously | Explosive | Burns with yellow flame | |
| Lithium (Li) | Fizzes steadily | Very rapid | Burns with red flame | |
| Calcium (Ca) | Fizzes steadily | Rapid | Burns with red flame | |
| Magnesium (Mg) | Very slow reaction | Rapid, bright flame | Bright white flame | |
| Aluminium (Al) | No reaction (oxide layer) | Moderate (if oxide removed) | Bright white flame | |
| Zinc (Zn) | No reaction | Moderate | Burns with yellow flame | |
| Iron (Fe) | No reaction | Slow | Burns with yellow sparks | |
| Tin (Sn) | No reaction | Very slow | No reaction | |
| Lead (Pb) | No reaction | Very slow | No reaction | |
| Copper (Cu) | No reaction | No reaction | No reaction | |
| Silver (Ag) | No reaction | No reaction | No reaction | |
| Least reactive | Gold (Au) | No reaction | No reaction | No reaction |
Remember the reactivity series using this mnemonic: Please Send Lions Cats Monkeys And Cute Zebras Into Lovely Ponds Covered With Snow (Potassium, Sodium, Lithium, Calcium, Magnesium, Aluminium, Carbon, Zinc, Iron, Lead, Copper, Silver).
Reactions with Water
Metals above hydrogen in the reactivity series react with water to produce a metal hydroxide (or oxide for some) and hydrogen gas. The more reactive the metal, the more vigorous the reaction.
Potassium + Water
2K(s) + 2H₂O(l) → 2KOH(aq) + H₂(g)
The reaction is extremely vigorous. The hydrogen gas catches fire and burns with a lilac flame. The potassium melts into a ball and darts around on the surface of the water.
Sodium + Water
2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)
Sodium melts into a ball and fizzes rapidly on the water surface. The hydrogen produced may ignite with an orange flame from the sodium vapour.
Calcium + Water
Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g)
Calcium reacts steadily with cold water, fizzing as hydrogen gas is produced. The calcium hydroxide formed is only slightly soluble, making the water turn cloudy.
Magnesium + Water
Mg(s) + 2H₂O(l) → Mg(OH)₂(aq) + H₂(g)
Magnesium reacts very slowly with cold water but reacts more readily with steam to form magnesium oxide and hydrogen:
Mg(s) + H₂O(g) → MgO(s) + H₂(g)
When a metal reacts with cold water, the product is a metal hydroxide. When a metal reacts with steam, the product is a metal oxide. This is because the metal hydroxide decomposes at the higher temperature of steam.
Reactions with Dilute Acids
Metals above hydrogen in the reactivity series react with dilute acids to produce a salt and hydrogen gas. The more reactive the metal, the faster the reaction.
The general equation for the reaction of a metal with a dilute acid is:
Metal + Acid → Salt + Hydrogen
The speed of the reaction is indicated by the rate of bubbling (effervescence) and the temperature change. More reactive metals produce hydrogen faster and release more heat energy.
Magnesium + Hydrochloric Acid
Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)
Rapid fizzing is observed. The magnesium ribbon dissolves and the test tube becomes warm. Collecting the gas and testing it with a lit splint produces a squeaky pop, confirming hydrogen.
Zinc + Sulfuric Acid
Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g)
Moderate fizzing. The zinc gradually dissolves. Hydrogen gas is confirmed by the squeaky pop test.
Iron + Hydrochloric Acid
Fe(s) + 2HCl(aq) → FeCl₂(aq) + H₂(g)
Slow fizzing. The iron takes a long time to dissolve. The solution turns pale green due to iron(II) ions.
Metals below hydrogen in the reactivity series (copper, silver, gold) do not react with dilute acids because they cannot displace hydrogen from the acid.
Displacement Reactions
A more reactive metal will displace a less reactive metal from its compound in solution. This is because the more reactive metal forms ions more readily, transferring electrons to the ions of the less reactive metal.
Displacement reactions are a way of establishing the order of reactivity of metals. If a piece of metal is placed in a solution of a salt of a different metal, a reaction will only occur if the metal is more reactive than the metal in the salt.
Iron + Copper Sulfate
Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)
Iron is more reactive than copper, so it displaces copper from copper sulfate solution. A brown-red coating of copper forms on the iron nail, and the blue solution fades to pale green as iron(II) sulfate is formed.
Copper + Magnesium Sulfate
No reaction. Copper is less reactive than magnesium, so it cannot displace magnesium from magnesium sulfate solution.
Zinc + Iron(II) Sulfate
Zn(s) + FeSO₄(aq) → ZnSO₄(aq) + Fe(s)
Zinc is more reactive than iron and displaces it from solution. The pale green solution turns colourless and a dark grey deposit of iron forms on the zinc.
In a displacement reaction, look at the reactivity series. If the free metal is higher (more reactive) than the metal in the compound, a displacement reaction occurs. If it is lower, no reaction happens.
Extracting Metals from Ores
How a metal is extracted from its ore depends on its position in the reactivity series. Metals below carbon can be extracted by reduction with carbon. Metals above carbon must be extracted by electrolysis.
Most metals are found in the Earth's crust combined with other elements in compounds called ores. The method used to extract a metal from its ore depends on how reactive the metal is.
| Reactivity | Extraction Method | Examples | Reason |
|---|---|---|---|
| Potassium, Sodium, Lithium, Calcium, Magnesium, Aluminium | Electrolysis | Aluminium from bauxite | Too reactive for carbon reduction; carbon cannot displace these metals from their oxides |
| Zinc, Iron, Tin, Lead, Copper | Reduction with carbon (or carbon monoxide) | Iron from haematite (Fe₂O₃) | Carbon is more reactive and can displace these metals from their oxides |
| Silver, Gold | Found native (uncombined) | Gold nuggets | So unreactive they exist as elements in the Earth's crust |
Extracting Iron from Haematite
Iron is extracted from its ore haematite (Fe₂O₃) in a blast furnace using carbon monoxide as the reducing agent:
Fe₂O₃(s) + 3CO(g) → 2Fe(l) + 3CO₂(g)
The carbon monoxide is produced by reacting coke (carbon) with hot air:
C(s) + O₂(g) → CO₂(g)
CO₂(g) + C(s) → 2CO(g)
Carbon can be used because it is more reactive than iron, so it can displace iron from its oxide.
Extracting metals by electrolysis is very expensive because it requires large amounts of electricity. Reduction with carbon is much cheaper, which is why iron is far less expensive than aluminium despite aluminium being more abundant in the Earth's crust.
Oxidation and Reduction
Oxidation is the gain of oxygen (or loss of electrons). Reduction is the loss of oxygen (or gain of electrons). If oxidation and reduction happen at the same time, it is called a redox reaction.
In the context of metal extraction and reactivity:
- When a metal reacts with oxygen to form a metal oxide, the metal is oxidised (it gains oxygen).
- When a metal oxide loses oxygen to form the metal, the oxide is reduced (it loses oxygen).
Redox in the Blast Furnace
Fe₂O₃(s) + 3CO(g) → 2Fe(l) + 3CO₂(g)
Iron(III) oxide is reduced (loses oxygen) to form iron. Carbon monoxide is oxidised (gains oxygen) to form carbon dioxide. This is a redox reaction because oxidation and reduction happen simultaneously.
Oxidation of Magnesium
2Mg(s) + O₂(g) → 2MgO(s)
Magnesium is oxidised because it gains oxygen to form magnesium oxide. Oxygen is reduced because it gains electrons from magnesium.
Rusting of Iron
Iron corrodes in the presence of water and oxygen to form hydrated iron(III) oxide, which is commonly called rust. The word equation is: iron + water + oxygen → hydrated iron(III) oxide.
Rusting is a specific type of corrosion that affects iron and steel. For rusting to occur, both water (or water vapour) and oxygen must be present. If either one is absent, rusting does not occur.
4Fe(s) + 3O₂(g) + 6H₂O(l) → 4Fe(OH)₃(s) → 2Fe₂O₃·3H₂O(s) (rust)
Unlike aluminium oxide, which forms a protective layer, rust is porous and flaky. This means it falls off the surface, exposing fresh iron to further rusting. Eventually the iron will corrode completely through.
| Condition | Does Rusting Occur? | Reason |
|---|---|---|
| Iron in dry air (no water) | No | Water is required for rusting |
| Iron in boiled water with oil layer (no oxygen) | No | Oxygen is required for rusting |
| Iron in air and water | Yes | Both water and oxygen are present |
| Iron in salt water and air | Yes (faster) | Ions in salt water speed up the electrochemical process |
Rusting requires BOTH water and oxygen. Removing either one prevents rusting. Salt accelerates rusting because dissolved ions increase the conductivity of the water, speeding up the electrochemical process.
Preventing Rusting
Rusting can be prevented by creating a barrier between the iron and water/oxygen, or by using sacrificial protection where a more reactive metal is attached to the iron.
There are several methods of preventing rusting:
- Barrier methods: Painting, oiling/greasing, coating with plastic, or using a metal coating. These all work by keeping water and oxygen away from the iron surface.
- Galvanising: Coating iron or steel with a layer of zinc. The zinc acts as a barrier and also provides sacrificial protection if the coating is scratched.
- Sacrificial protection: Attaching a more reactive metal (such as zinc or magnesium) to the iron. The more reactive metal corrodes instead of the iron, protecting it.
| Method | How It Works | Advantages | Disadvantages |
|---|---|---|---|
| Painting | Barrier to water and oxygen | Cheap, decorative, easy to apply | Can chip or scratch, needs reapplying |
| Oiling/Greasing | Barrier to water and oxygen | Good for moving parts | Needs regular reapplication, messy |
| Plastic coating | Barrier to water and oxygen | Durable, comes in many colours | Can crack or peel over time |
| Galvanising | Zinc barrier + sacrificial protection | Long-lasting, protects even if scratched | More expensive, limited appearance options |
| Sacrificial protection | More reactive metal corrodes instead | Protects even when coating is damaged | Blocks need replacing as they corrode away |
Galvanising and Sacrificial Protection in Detail
Galvanising involves coating iron or steel with zinc. If the zinc layer is scratched, the zinc still protects the iron because zinc is more reactive and will corrode preferentially. This is sacrificial protection.
Sacrificial protection is used in situations where rusting would be dangerous or difficult to repair:
- Ship hulls have blocks of zinc attached to them, which corrode instead of the steel hull.
- Underground iron and steel pipes are connected to bars of magnesium, which corrode sacrificially.
- Galvanised buckets, fencing, and roofing are coated with zinc, providing both a barrier and sacrificial protection.
Sacrificial Protection of a Ship's Hull
Blocks of zinc are bolted to the steel hull of a ship. Because zinc is more reactive than iron, it oxidises more readily. The zinc loses electrons and forms zinc ions, while the iron remains unreacted. The zinc blocks gradually corrode away and must be replaced periodically, but they protect the steel hull from rusting.
Aluminium does not corrode in the same way as iron because it forms a thin, hard, transparent layer of aluminium oxide on its surface. This layer is impermeable to water and oxygen, so it protects the aluminium underneath from further corrosion. This is why aluminium is used for window frames and aircraft bodies despite being a reactive metal.
Alloys and Their Uses
An alloy is a mixture of two or more elements where at least one is a metal. Alloys are often harder and more useful than pure metals because the different-sized atoms disrupt the regular arrangement, preventing layers from sliding over each other.
| Alloy | Metals in Alloy | Properties | Uses |
|---|---|---|---|
| Bronze | Copper + Tin | Hard, resistant to corrosion | Statues, coins, decorative items |
| Brass | Copper + Zinc | Hard, golden appearance | Door fittings, musical instruments |
| Stainless steel | Iron + Chromium + Nickel | Hard, does not rust | Cutlery, surgical instruments, chemical equipment |
| Gold jewellery | Gold + Copper (or silver) | Harder than pure gold | Jewellery (pure gold is too soft) |
Practice Questions
1. Write the word equation for the reaction of sodium with water and describe what you would observe.
Sodium + water → sodium hydroxide + hydrogen. The sodium melts into a ball, fizzes rapidly on the water surface, and moves around. The hydrogen may ignite with an orange flame. The solution becomes alkaline (turns universal indicator purple/blue).
2. Explain why gold is found native in the Earth's crust but aluminium is not.
Gold is very unreactive and is found at the bottom of the reactivity series, so it does not readily form compounds with other elements. Aluminium is very reactive and is near the top of the reactivity series, so it readily combines with oxygen and other elements to form compounds such as aluminium oxide (bauxite).
3. A student places a copper coin into a solution of magnesium sulfate. Explain why no reaction occurs.
Copper is less reactive than magnesium. For a displacement reaction to occur, the free metal must be more reactive than the metal in the compound. Since copper is below magnesium in the reactivity series, it cannot displace magnesium from its salt solution.
4. Explain why galvanising provides better protection than painting, even if the zinc layer is scratched.
Galvanising provides both a barrier and sacrificial protection. If the zinc coating is scratched, the exposed iron is still protected because zinc is more reactive than iron. The zinc corrodes preferentially, losing electrons and forming zinc ions, while the iron remains unreacted. Painting only provides a barrier; if the paint is scratched, water and oxygen can reach the iron and it will rust.
5. Write a balanced symbol equation for the displacement reaction between zinc and copper(II) sulfate solution.
Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)
6. Explain why aluminium is extracted by electrolysis but iron is extracted by reduction with carbon.
Aluminium is more reactive than carbon, so carbon cannot displace aluminium from its oxide. Electrolysis is needed to decompose aluminium oxide using electricity. Iron is less reactive than carbon, so carbon can displace iron from its oxide in a blast furnace. Reduction with carbon is much cheaper than electrolysis, so it is used whenever possible.
Required Practical
Investigating the Reactivity Series (Displacement Reactions)
This required practical involves using displacement reactions to establish the order of reactivity of metals. Place pieces of different metals into solutions of different metal salts and observe whether a reaction occurs. If the free metal is more reactive than the metal in the salt, a displacement reaction will happen.
Method:
- Prepare small pieces of four or five different metals (e.g. magnesium, zinc, iron, copper).
- Prepare solutions of the corresponding metal sulfates (e.g. magnesium sulfate, zinc sulfate, iron(II) sulfate, copper(II) sulfate).
- Place each metal into each solution in a dimple tray or test tubes.
- Observe whether a reaction occurs: look for a metal coating forming on the added metal, a colour change in the solution, or gas bubbles.
- Record your observations in a results table, noting which combinations produce a reaction and which do not.
- Use the results to place the metals in order of reactivity.
For example, placing iron nails into copper(II) sulfate solution produces a brown-red coating of copper on the nail and the blue solution fades to pale green. This confirms iron is more reactive than copper. No reaction when copper is placed in magnesium sulfate confirms copper is less reactive than magnesium.
Variables: The independent variable is the type of metal. The dependent variable is whether a displacement reaction occurs (observed as a colour change or metal deposit). Control variables include the concentration of the salt solutions, the surface area of the metal pieces, the temperature, and the time the metal is left in the solution.
Safety: Wear eye protection. Some metal salts are harmful — avoid skin contact and wash hands after the practical. Dispose of solutions as instructed, not down the sink.
Maths Skills
Writing Ionic Equations and Half Equations
When given experimental observations (e.g. whether a metal reacts with water, acid, or oxygen), use the results to deduce the order of reactivity. A metal that reacts vigorously with cold water is more reactive than one that only reacts with steam, which is more reactive than one that only reacts with acid.
You may need to interpret data presented in tables or graphs. Look for patterns: faster reaction rates, more vigorous fizzing, or larger temperature increases all indicate higher reactivity. Use comparative language such as "metal X reacts more vigorously than metal Y, therefore X is more reactive."
Writing ionic equations for displacement reactions: Mg(s) + Cu2+(aq) → Mg2+(aq) + Cu(s). The sulfate ion is a spectator and is cancelled. Writing half equations shows the electron transfer: Mg → Mg2+ + 2e− (oxidation); Cu2+ + 2e− → Cu (reduction).
Common Misconceptions
Reactivity and Reaction Vigour
Wrong: Reactivity is the same as how vigorous the reaction looks Correct: Reactivity is about the tendency of a metal to lose electrons and form positive ions. While more reactive metals do tend to react more vigorously, the visual appearance of a reaction can be misleading. For example, aluminium appears unreactive because its oxide layer prevents reaction, but aluminium itself is very reactive. Reactivity is a fundamental chemical property, not just about what you can see
Unreactive Metals
Wrong: Unreactive metals like gold do not react at all Correct: Even very unreactive metals can react under the right conditions. Gold does not react with most acids or oxygen, but it will dissolve in aqua regia (a mixture of concentrated nitric and hydrochloric acids). "Unreactive" means the metal does not react under normal conditions, not that it is chemically inert
Wrong: Gold is unreactive because it is rare Correct: Gold is unreactive because its atoms hold onto their outer electrons very tightly due to their electronic structure and nuclear attraction. Its rarity is unrelated to its reactivity
Wrong: Rusting only needs water Correct: Rusting requires both water AND oxygen. Removing either one prevents rusting. This is why iron in boiled water with an oil layer (no oxygen) does not rust
6-Mark Extended Question
Native Metals and the Reactivity Series
Explain why gold is found native in the Earth's crust but iron is not. Use the reactivity series in your answer.
Gold is at the bottom of the reactivity series and is one of the least reactive metals. It has a very low tendency to lose electrons and form positive ions, so it does not readily react with oxygen, water, or other elements in the Earth's crust. As a result, gold is found as the pure elemental metal (native gold) rather than as a compound. [2 marks]
Iron is much higher in the reactivity series and is a moderately reactive metal. It readily loses electrons to form iron ions, which combine with oxygen and water in the environment to form iron oxides (such as haematite, Fe2O3). This means iron is always found combined with other elements as an ore in the Earth's crust, never as native iron. [2 marks]
The position of a metal in the reactivity series determines how it is found and extracted. Metals above carbon (like aluminium) are so reactive that they form very stable compounds and must be extracted by electrolysis. Metals below carbon but above hydrogen (like iron) are found as oxides and can be extracted by reduction with carbon. Only the very least reactive metals (copper, silver, gold) can be found native because they do not readily form compounds. [2 marks]
AO3: Analyse and Evaluate
Placing Metals in Order of Reactivity
A student tests four metals (P, Q, R, S) by adding each to dilute hydrochloric acid and to solutions of the other metals' sulfates. The results are shown below: P fizzes rapidly with acid and displaces Q and S from their salts. Q fizzes slowly with acid and displaces R and S from their salts. R fizzes very slowly with acid and displaces S only. S does not fizz with acid and does not displace any metal. Deduce the order of reactivity from most to least reactive and justify your answer.
Most reactive: P (fizzes rapidly with acid, displaces two metals) → Q (fizzes slowly, displaces two metals, but not P, so Q is less reactive than P) → R (fizzes very slowly, displaces only S, so less reactive than P and Q) → S (no fizzing, displaces none, so least reactive). Order: P > Q > R > S. Justification: P displaces Q so P > Q. Q displaces R so Q > R. R displaces S so R > S. The rate of fizzing with acid also confirms this order: faster fizzing means more reactive. No contradictions exist in the data, confirming this order is consistent.
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