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Module 3 · L7 of 12 35 min ⚡ +50 XP in Learn · +25 to complete

Metal Activity Series & Reactions of Metals

Steel ships, pipelines and offshore structures are protected from rusting by bolting on blocks of a more reactive metal, such as zinc or magnesium. These "sacrificial anodes" corrode in place of the steel because they sit higher in the activity series, so they lose electrons more readily. The activity series predicts which metal corrodes first, and engineers use it deliberately to choose which metal to sacrifice.

Today's hook, Steel ships, pipelines and offshore structures are protected from rusting by bolting on blocks of a more reactive metal, such as zinc or magnesium. These "sacrificial anodes" corrode in place of the steel because they sit higher in the activity series, so they lose electrons more readily. The activity series predicts which metal corrodes first, and engineers use it deliberately to choose which metal to sacrifice.

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Galvanised steel (zinc-coated steel) contains both steel (iron alloy) and zinc. When both metals are exposed to oxygen and moisture, the zinc corrodes preferentially, greatly reducing corrosion of the iron while the zinc is gradually consumed (the protection lasts while zinc and electrical contact remain).

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What you'll master

Know

Key facts

The standard HSC activity series from K (most reactive) to Au (least reactive)
Metals above hydrogen in the series react with dilute acids; metals below do not
Displacement reactions are also redox reactions: the more reactive metal is oxidised

Understand

Concepts

Reactivity reflects ease of electron loss (ionisation energy); more reactive metals have lower ionisation energy and are found as ores rather than native
The activity series was derived from experimental observations of water, acid, and displacement reactions
Aluminium appears anomalously unreactive due to a protective Al₂O₃ passivation layer, its true reactivity is higher than its behaviour suggests

Can do

Skills

Use the activity series to predict whether a displacement reaction will occur and write the balanced equation
Explain why metals above H dissolve in dilute acid while metals below H do not
Identify the oxidised and reduced species in a displacement reaction

Key terms

Activity series

A ranking of metals from most to least reactive based on their tendency to lose electrons (be oxidised).

Single displacement reaction

A more reactive metal displaces a less reactive metal ion from its salt solution: A + BC → AC + B.

Passivation

Formation of a protective oxide layer on a metal surface that slows further corrosion; e.g., Al₂O₃ on aluminium.

Metal + water reaction

Very reactive metals (Na, K) react vigorously with cold water; Mg reacts with hot water; Fe reacts with steam.

Metal + acid reaction

Active metals above hydrogen in the activity series react with dilute acids: Metal + HCl → salt + H₂(g).

Thermite reaction

A highly exothermic single displacement: Al + Fe₂O₃ → Al₂O₃ + Fe; used in welding and incendiary devices.

What Is Metal Reactivity and Why Does It Vary?

core concept

Drop sodium into water: violent fizzing, hydrogen gas, the metal dissolves in seconds. Drop copper into the same water: nothing. Leave iron in water overnight: slight orange staining, then nothing dramatic. All three are metals, all three react differently with the same substance. The difference is measurable, predictable, and can be listed in order: the metal activity series.

Three atomic properties determine how easily electrons are lost:

Trend for Increasing Reactivity

Increases (larger)

Decreases (lower)

Why

Outermost electrons further from nucleus, less nuclear attraction, easier to remove

Less energy needed to remove the outermost electron

Less tendency to attract or hold electrons

Down Group 1 of the periodic table, atomic radius increases, ionisation energy decreases, and electronegativity decreases, reactivity increases. This is why caesium is more reactive than lithium, and why potassium reacts more violently with water than sodium. These periodic trends explain reactivity within a group, but the full activity series (which compares metals from different groups) is established from experimental reactions and E° data, not from atomic trends alone.

Must Know: In HSC answers explaining metal reactivity, you must connect the periodic trend to the electron-losing ability explicitly. “Potassium is more reactive than sodium because potassium has a larger atomic radius, so its outermost electron is held less tightly by the nucleus and is lost more easily” is a full-mark answer. “Potassium is lower in Group 1” is not.

Common Error: Students often say “more reactive metals have more electrons.” This is wrong, reactivity is about how easily electrons are lost, not how many there are. Gold has more electrons than potassium but is far less reactive.

Metal reactivity reflects how readily a metal loses electrons (M → Mⁿ⁺ + ne⁻). Periodic trends (larger atomic radius and lower ionisation energy down a group) partly explain this, but the activity series is ranked from observed reactions and standard electrode potentials (E°), not derived from atomic trends alone, because aqueous reactivity also depends on atomisation, hydration and conditions.

Pause, copy the highlighted definition into your book before moving on.

Odd one out: Which of these properties increases as metal reactivity increases?

Reactions of Metals, Constructing the Activity Series

core concept

We just saw that metal reactivity varies due to several factors, including ionisation energy and atomic radius. That raises a question: how do chemists experimentally measure and rank this reactivity to build the activity series? This card answers it → by comparing reactions with oxygen, cold water, and dilute acids, then combining displacement observations to rank metals in order.

Safety, metal reactivity, teacher demonstration only. The most vigorous reactions in this lesson are demonstrations or secondary data, not student practicals. Alkali metals (Na, K) reacting with water, and the thermite reaction (Al + Fe₂O₃), are teacher demonstrations behind a safety screen or shown by video/data, because they are violently exothermic and produce hydrogen or molten metal. For the safe student tests (metals in dilute acid), use small quantities of dilute acid, wear safety glasses, work in a well-ventilated room, and treat any hydrogen produced as flammable, an ignition (squeaky-pop) test is done by the teacher on a small sample. Never heat or ignite these mixtures yourself.

The activity series is constructed experimentally by comparing how vigorously different metals react with the same reagents under the same conditions. Three standard investigations are used:

Metal	Reaction with O₂	Reaction with Cold Water	Reaction with Dilute HCl
Potassium (K)	Burns vigorously	Explosive	Explosive (not used)
Sodium (Na)	Burns vigorously	Very vigorous	Explosive (not used)
Calcium (Ca)	Burns	Vigorous, bubbles	Vigorous
Magnesium (Mg)	Burns brightly	Very slow (reacts with steam)	Vigorous
Aluminium (Al)	Burns	Slow or none at first (passivation; oxide can be removed)	Moderate (slowed by oxide layer)
Zinc (Zn)	Burns	No reaction	Moderate
Iron (Fe)	Burns slowly (as powder)	No reaction	Slow
Lead (Pb)	Tarnishes	No reaction	Very slow
Copper (Cu)	Surface oxide only	No reaction	No reaction
Silver (Ag)	No reaction	No reaction	No reaction
Gold (Au)	No reaction	No reaction	No reaction

Must Know: Aluminium appears to be less reactive than its position in the series suggests because it forms a tough, impermeable aluminium oxide (Al₂O₃) layer on its surface that prevents further reaction. This is passivation not low reactivity. In HSC questions about aluminium, always acknowledge this distinction.

Common Error: The NESA activity series lists hydrogen (H) as a reference point, metals above H displace it from acids; metals below H do not. Students often omit hydrogen from their series. It must appear between lead and copper in the standard series.

The activity series is built from three experimental tests: vigour of reaction with O₂, cold water, and dilute acid. Aluminium appears less reactive than expected due to passivation (protective Al₂O₃ layer). Hydrogen sits between Pb and Cu as a reference point for acid reactions.

Add the highlighted point to your notes before the check below.

Fill the gap: Aluminium appears less reactive than expected because it rapidly forms a protective [___] layer on its surface that prevents further reaction with water or dilute acids.

The Standard Activity Series

core concept

We just saw how experimental observations with water, oxygen, and acids are used to rank metal reactivity. That raises a question: what is the specific NESA-approved order that you need to memorise for HSC exams? This card answers it → K, Na, Ca, Mg, Al, Zn, Fe, Pb, H, Cu, Ag, Au, with hydrogen as the reference point between Pb and Cu.

The standard activity series used in HSC Chemistry (confirm against the current NESA data sheet), from most reactive to least reactive:

Mnemonic: Please Stop Calling Me A Zombie, I Like Having Copper Silverware Guaranteed. (K, Na, Ca, Mg, Al, Zn, Fe, Pb, H, Cu, Ag, Au)

Must Know: The NESA activity series and the standard reduction potential table (L09) encode the same information in different forms. The activity series ranks by relative reactivity; the reduction potential table gives quantitative values. Begin building the connection now.

Insight: Hydrogen is included in the activity series as a reference point, not because it is a metal. Any metal above hydrogen will reduce H⁺ ions in solution to H₂ gas. Any metal below hydrogen cannot do this. This is why copper, silver, and gold do not dissolve in dilute hydrochloric or sulfuric acid.

Standard HSC activity series (most → least reactive; check the current NESA data sheet): K, Na, Ca, Mg, Al, Zn, Fe, Pb, H, Cu, Ag, Au. Mnemonic: "Please Stop Calling Me A Zombie, I Like Having Copper Silverware Guaranteed." Metals above H react with dilute acids to produce H₂; metals below H (Cu, Ag, Au) do not.

Pause, write the highlighted series and mnemonic into your book.

Quick check: According to the NESA activity series, which metal would NOT react when added to dilute hydrochloric acid?

Displacement Reactions, Predicting from the Activity Series

core concept

We just saw the NESA activity series in order. That raises a question: how do we use the series to predict which displacement reactions will occur, and write the correct balanced equations? This card answers it → if the metal added is ABOVE the ion in solution, displacement occurs; if BELOW, no reaction.

A more reactive metal will always displace a less reactive metal from its salt solution, the more reactive metal has a greater tendency to form ions, so it “takes” the electron-release role away from the less reactive metal’s ions.

Prediction rule: Is the metal being added higher (more reactive) than the metal ion in solution? If yes → reaction occurs. If no → no reaction.

Metal Added	Solution	Reaction?	Reason
Zn(s)	CuSO₄(aq)	Yes ✓	Zn more reactive than Cu
Fe(s)	CuSO₄(aq)	Yes ✓	Fe more reactive than Cu
Cu(s)	ZnSO₄(aq)	No ✗	Cu less reactive than Zn
Mg(s)	FeSO₄(aq)	Yes ✓	Mg more reactive than Fe
Ag(s)	HCl(aq)	No ✗	Ag below H in series
Zn(s)	HCl(aq)	Yes ✓	Zn above H in series

For Zn(s) + CuSO₄(aq): The blue colour of CuSO₄ solution fades (Cu²⁺ ions removed); reddish-brown solid copper deposits on the zinc surface; the zinc gradually dissolves.
Equation: Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s). Atom check: 1Zn, 1Cu, 1S, 4O each side. ✓

Must Know: When writing displacement equations, the sulfate (or nitrate) anion acts as a spectator, it does not change. Only the cations exchange. Always verify the equation by checking that the more reactive metal has become the ion on the right and the less reactive metal has become the solid.

Common Error: Students sometimes predict that a reaction will occur simply because two substances are mixed. The activity series is the decision tool. If the metal being added is BELOW the metal ion in solution, no reaction occurs. Copper does not react with iron sulfate solution; iron reacts with copper sulfate solution.

Displacement rule: if the added metal is ABOVE the ion in solution in the activity series, displacement occurs. E.g. Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s) ✓; Cu(s) + ZnSO₄(aq) → no reaction ✗. The sulfate anion is a spectator.

Add the highlighted rule to your notes before the check below.

Odd one out: Which of these combinations would result in a displacement reaction occurring?

Galvanised Steel, Why Zinc Protects Steel

core concept

We just saw that more reactive metals preferentially displace less reactive metals. That raises a question: how does this principle apply to real corrosion protection, such as galvanised steel and sacrificial anodes? This card answers it → zinc (higher in the activity series than iron) preferentially corrodes, sacrificially protecting the steel underneath.

Steel (iron alloy) and zinc are both susceptible to oxidation in the presence of oxygen and moisture. When they are in electrical contact, as they are when zinc bolts fasten steel panels, zinc preferentially loses electrons because it is higher in the activity series than iron. The zinc corrodes while the iron is protected. This is called sacrificial protection or cathodic protection.

The same principle is used in galvanised steel (zinc-coated steel), where the zinc coating sacrificially corrodes even if scratched, continuing to protect the underlying iron.

If copper bolts were used instead, the situation reverses, copper is below iron in the activity series, so iron would corrode preferentially. This is a galvanic corrosion problem.

Must Know: The concept of sacrificial protection introduced here is the conceptual foundation for galvanic cells in L09 and cathodic protection in L10. Note it now, it will return in quantitative form using standard reduction potentials.

Common Error: Students confuse sacrificial protection with galvanising. Galvanising applies a zinc coating that acts as both a physical barrier AND a sacrificial anode if the coating is broken. Sacrificial protection (a separate zinc anode connected to the structure) has no physical barrier component, it works purely through the electrochemical activity series principle.

Sacrificial protection: because Zn is more reactive than Fe, zinc preferentially oxidises (Zn → Zn²⁺ + 2e⁻), protecting steel from corrosion even when scratched. If copper bolts were used, iron would corrode instead (Fe is more reactive than Cu).

Pause, write the highlighted point into your book.

Fill the gap: Zinc protects iron in galvanised steel because zinc is [___] in the activity series than iron, so it preferentially loses electrons and corrodes first.

Cross-lesson links: In L01–L06 (IQ1) you built the reaction toolkit, synthesis, combustion, precipitation, acid-base. The activity series in this lesson predicts the outcome of metal displacement reactions using the same observable-evidence approach. In L08, you will formalise this as redox chemistry, with oxidation numbers giving a precise language for what the activity series describes qualitatively.

Worked examples · reveal as you go

Worked example +5 XP on full reveal

Predict whether a reaction will occur in each case. Where a reaction occurs, write the balanced equation with state symbols and describe one observable change. (a) Iron nail placed in copper(II) sulfate solution. (b) Copper wire placed in silver nitrate solution. (c) Silver wire placed in copper(II) nitrate solution.

(a) Fe in CuSO₄, Activity series check: Fe is above Cu → Fe will displace Cu²⁺. Reaction occurs.

Iron is more reactive than copper, so it can donate electrons to copper ions in solution.

Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s) Atom check: 1Fe, 1Cu, 1S, 4O each side. ✓

The sulfate ions are spectators. Only the metal cations exchange their roles.

Observable: blue colour of solution fades (Cu²⁺ removed); reddish-brown copper deposits on iron nail surface.

The displacement of Cu²⁺ removes the blue colour. Metallic copper forms a visible red coating.

(b) Cu in AgNO₃, Activity series check: Cu is above Ag → Cu will displace Ag⁺. Reaction occurs. Cu(s) + 2AgNO₃(aq) → Cu(NO₃)₂(aq) + 2Ag(s) Atom check: 1Cu, 2Ag, 2N, 6O each side. ✓ Observable: solution gradually becomes blue (Cu²⁺ ions form); silver crystals deposit on copper wire surface. (c) Ag in Cu(NO₃)₂, Activity series check: Ag is below Cu → Ag cannot displace Cu²⁺. No reaction occurs. Observable: no change in solution colour; no solid deposits on the silver wire.

Use the NESA activity series (K, Na, Ca, Mg, Al, Zn, Fe, Pb, H, Cu, Ag, Au) to predict whether the added metal is higher or lower than the metal ion in solution.

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Explain why potassium reacts more vigorously with cold water than magnesium does, using the concepts of atomic radius, ionisation energy, and electronegativity.

Compare positions in periodic table: K is in Period 4, Group 1. Mg is in Period 3, Group 2.

Potassium has more electron shells than magnesium, placing it lower in the periodic table.

Atomic radius: K has a larger atomic radius than Mg. The valence electron (4s) is further from the nucleus than Mg's 3s electrons.

More electron shells means the outermost electrons are further from the nucleus, experiencing weaker nuclear attraction.

Ionisation energy: K has a lower first ionisation energy (419 kJ/mol) compared to Mg (738 kJ/mol). K loses its valence electron more readily under the same conditions.

The larger atomic radius in K means less nuclear attraction on the outermost electron, requiring less energy to remove.

Electronegativity: K has a lower electronegativity (0.82) than Mg (1.31). K has less tendency to attract electrons back toward the nucleus, reinforcing its greater tendency to exist as K⁺ rather than neutral K. Conclusion: Potassium reacts more vigorously because it has a larger atomic radius, lower ionisation energy, and lower electronegativity than magnesium. These properties mean K loses its outermost electron more readily, making it more reactive toward water. (This is a qualitative periodic comparison, not a derivation of the full activity series, which also depends on electron stoichiometry and on hydration/solution energetics and is established from observed reactions and E° data.)

All three trends work together: bigger atoms with lower ionisation energies and lower electronegativities lose electrons more easily, producing more vigorous reactions.

Formula reference · this lesson

core formula

📐

Key Patterns, This Lesson

$\text{M}_1(s) + \text{M}_2^{n+}(aq) \rightarrow \text{M}_1^{n+}(aq) + \text{M}_2(s)$ if $\text{M}_1$ more reactive

Only occurs if the added metal is HIGHER in the activity series, otherwise no reaction

Activity series (most → least reactive): K, Na, Ca, Mg, Al, Zn, Fe, Pb, H, Cu, Ag, Au

Mnemonic: Please Stop Calling Me A Zombie, I Like Having Copper Silverware Guaranteed

$r\uparrow \;\Rightarrow\; \text{IE}\downarrow \;\Rightarrow\; \chi\downarrow \;\Rightarrow\; \text{easier to lose } e^- \;\Rightarrow\; \text{more reactive}$

$r$ = atomic radius, IE = ionisation energy, $\chi$ = electronegativity, all trend together down Group 1

Common errors · the 3 traps that cost marks

Common misconception

A metal will displace any other metal from a compound if it is higher in the activity series.

Fix: Ordinary metal-ion displacement occurs only if the metal is higher in the activity series AND the reaction occurs in aqueous solution, where ions are free to move. (High-temperature reactions between solids, such as the thermite reaction Al + Fe₂O₃, are a separate case driven by strong heating, not this aqueous rule.)

Aluminium is less reactive than iron because it doesn't react with cold water

Students observe that aluminium doesn't react with dilute acid or cold water and rank it below iron or copper in the activity series.

Fix: Aluminium is more reactive than iron, it sits between Mg and Zn in the NESA activity series. However, Al rapidly forms a dense Al₂O₃ passivation layer that prevents further reaction. This protective oxide layer masks aluminium's true reactivity; if the oxide is removed (for example by acid over time, or by amalgamation), the aluminium underneath does react. In HSC answers about Al reactivity, mention the passivation layer as the reason for its apparent low reactivity.

A more reactive metal always displaces a less reactive one, even from solid compounds

Students apply the displacement rule to solid compounds, predicting for example that sodium metal dropped into solid CuSO₄ would displace copper.

Fix: Ordinary metal-ion displacement needs aqueous conditions so that ions are mobile and can migrate to the metal surface to be reduced. Sodium dropped onto solid CuSO₄ would react with any moisture present rather than performing a clean displacement. The activity series predicts displacement from aqueous solutions; high-temperature solid-state reactions such as thermite follow a different pathway.

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Quick-fire practice · 5 reps +2 XP per reveal

Q1 (4 marks): Explain, using the concepts of atomic radius, ionisation energy, and electronegativity, why sodium reacts more vigorously with water than lithium, even though both are in Group 1.

Q2 (4 marks): A student places a piece of iron metal into a solution of copper(II) sulfate. (a) Predict whether a reaction will occur, with reference to the activity series. (b) Write the balanced equation with state symbols. (c) Describe two observable changes the student would see. (1 + 2 + 1 marks)

Q3 (5 marks): A coastal steel structure is protected with sacrificial zinc anodes bolted to the steel. (a) Explain, using the activity series, why the zinc corrodes preferentially while the iron is protected. (b) Calculate the mass loss if 0.050 mol of Zn corrodes (Zn = 65.4 g/mol). (c) A maintenance engineer proposes replacing the zinc anodes with stainless steel (iron-based) to reduce replacement frequency. Evaluate this proposal with reference to the activity series and the consequences for the structure. (2 + 1 + 2 marks)

Arrange the following metals in order of increasing reactivity: Ag, Fe, Na, Cu, Zn. For the Zn/Cu pair, write the balanced equation for the displacement reaction and identify which metal is oxidised.

A student places a copper strip in silver nitrate solution. (a) Predict whether a reaction will occur using the activity series. (b) Write the balanced ionic equation. (c) Describe two observable changes that would confirm the reaction is occurring.

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Revisit your thinking

Sacrificial protection is used wherever steel must survive a corrosive environment: ships' hulls, underground pipelines and offshore structures carry blocks of a more reactive metal (zinc or magnesium). Because that metal is higher in the activity series, it loses electrons more readily (e.g. Zn → Zn²⁺ + 2e⁻ occurs more readily than Fe → Fe²⁺ + 2e⁻), so it corrodes as the anode while the steel is protected as the cathode. Engineers choose the sacrificial metal, and pair metals carefully, using the activity series; if two metals far apart in reactivity are joined in a wet, salty environment, the more reactive one can corrode quickly.

Now revisit your initial response. What did you get right? What has changed in your thinking?

Look back at your initial response in your book. Annotate it with what you now understand differently.

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Interactive Tool, Redox Reactions Lab Open fullscreen ↗

The Redox tool shows that oxidation is defined as…

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Short answer

UnderstandBand 3

Q1. (4 marks) Explain, using the concepts of atomic radius, ionisation energy, and electronegativity, why sodium reacts more vigorously with water than lithium, even though both are in Group 1. (4 marks)

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ApplyBand 4

Q2. (4 marks) A student places a piece of iron metal into a solution of copper(II) sulfate. (a) Predict whether a reaction will occur, with reference to the activity series. (b) Write the balanced equation with state symbols. (c) Describe two observable changes the student would see. (1 + 2 + 1 marks)

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EvaluateBand 5

Q3. (5 marks) A coastal steel structure is protected with sacrificial zinc anodes bolted to the steel. (a) Explain, using the activity series, why the zinc corrodes preferentially while the iron is protected. (b) Calculate the mass loss if 0.050 mol of Zn corrodes (Zn = 65.4 g/mol). (c) A maintenance engineer proposes replacing the zinc anodes with stainless steel (iron-based) to reduce replacement frequency. Evaluate this proposal with reference to the activity series and the consequences for the structure. (2 + 1 + 2 marks)

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