Chemistry • Year 11 • Module 3 • Lesson 7
Metal Activity Series & Reactions of Metals
Apply the activity series to experimental data, real corrosion scenarios, and a graph of metal reactivity — moving from recall to reasoning.
1. Interpret the graph — reaction rate of metals with dilute HCl
The bar graph below shows the approximate volume of hydrogen gas produced per minute when equal masses of five metals are each placed in excess 1.0 mol L−1 HCl at 25 °C. 8 marks
Figure 1.1. Approximate rate of H2 gas production when 0.10 g of each metal reacts with excess 1.0 mol L−1 HCl at 25 °C. Data adapted from standard laboratory investigations.
1.1 Describe the overall trend shown in the graph. 2 marks
1.2 Explain why copper produces no hydrogen gas in this experiment. Use the terms “activity series” and “H+” in your answer. 2 marks
1.3 Predict the approximate volume of H2 per minute you would expect if aluminium (Al) were added under the same conditions (without its oxide layer). Justify your estimate by reference to aluminium's position in the NESA activity series. 2 marks
1.4 Write the balanced equation (with state symbols) for the reaction of zinc with dilute HCl. 2 marks
2. Interpret displacement reaction data
A student placed small pieces of five metals into five different salt solutions. After 10 minutes the student recorded whether a reaction occurred and any observable changes. 7 marks
| Metal added | Salt solution | Reaction? (Y/N) | Observable change |
|---|---|---|---|
| Zn(s) | CuSO4(aq) | Y | Blue colour fades; reddish-brown solid on zinc surface |
| Cu(s) | ZnSO4(aq) | N | No change |
| Fe(s) | CuSO4(aq) | Y | Blue colour fades; copper deposits on iron |
| Ag(s) | CuSO4(aq) | N | No change |
| Mg(s) | ZnSO4(aq) | Y | Grey solid deposits on Mg surface; Mg slowly dissolves |
2.1 Use the data to deduce the relative positions of Cu, Zn, Fe, Mg and Ag in the activity series. List them from most to least reactive and justify with one piece of evidence from the table. 3 marks
2.2 Write the balanced ionic equation for the reaction of zinc with copper sulfate solution. Include state symbols and identify the spectator ion. 2 marks
2.3 Predict and explain whether a reaction would occur if platinum (Pt) metal were placed in silver nitrate solution. (Platinum is below gold in the activity series.) 2 marks
3. Case study — galvanised steel on Australian farms
Read the scenario and answer the questions below. 6 marks
Scenario. Hot-dip galvanising is widely used on Australian farm infrastructure — fencing, water tanks, and shed roofing — to protect mild steel from corrosion. In the galvanising process, clean steel is dipped into a bath of molten zinc (at around 450 °C), producing a zinc coating typically 45–85 µm thick. The coating provides two layers of protection: (1) a physical barrier while the zinc layer is intact, and (2) sacrificial (cathodic) protection even when the coating is scratched or breached. A contractor on a rural property near Port Macquarie notices that after five years, some galvanised roofing sheets have small rust spots near copper roof-fixing screws — even though the zinc coating is otherwise undamaged.
3.1 Explain why normal galvanised steel protects iron even when the coating is scratched. Use the terms “activity series”, “sacrificial anode”, and “oxidised” in your answer. 3 marks
3.2 Account for the rust spots appearing near the copper screws, even though the zinc coating is intact elsewhere. Use your knowledge of the activity series and galvanic corrosion to explain what is happening at the contact point between the copper screw and the zinc-coated steel. 3 marks
4. Predict and justify — iron pipe in copper water system
5 marks
Scenario. A plumber connects a section of iron pipe to an existing copper pipe in a hot-water system in a coastal Australian home. The pipes are joined with a standard threaded fitting — placing iron and copper in direct metal contact. Within two years, a pinhole leak develops at the iron-copper junction but the copper pipe itself is undamaged.
4.1 Using your knowledge of the activity series and galvanic corrosion, predict which pipe corrodes and explain why. Refer to electron transfer in your answer. 3 marks
4.2 A plumber recommends installing a “dielectric union” fitting between the iron and copper pipes — a non-conducting plastic insert that electrically isolates the two metals. Justify why this would prevent the galvanic corrosion, using what you know about the conditions needed for a displacement-type corrosion reaction to occur. 2 marks
5. Compare and contrast — galvanising vs sacrificial anodes
Complete the comparison table below. Where a cell is already filled in, it is provided as a guide. 6 marks
| Feature | Galvanised steel | Sacrificial anode (zinc blocks) |
|---|---|---|
| Physical barrier? | No — anodes are separate from the structure | |
| Which metal corrodes preferentially? | Zinc coating | |
| Protection when coating/anode is fully consumed? | No — anodes must be periodically replaced | |
| Activity series principle used | ||
| Typical Australian application | Farm fencing, shed roofing | |
| One limitation |
Q1.1 — Trend description
The volume of H2 produced per minute decreases across the metals in the order Mg > Zn > Fe > Pb > Cu. This matches the NESA activity series ranking, confirming that more reactive metals (higher in the series) produce hydrogen at a faster rate when reacting with dilute HCl. Copper produces no hydrogen because it is below hydrogen in the series.
Q1.2 — Why copper produces no hydrogen
Copper is below hydrogen in the NESA activity series, meaning it has less tendency to be oxidised (lose electrons) than H. Therefore copper cannot reduce H+ ions to H2 gas. No electron transfer from Cu to H+ occurs, so no reaction takes place and no hydrogen is produced.
Q1.3 — Prediction for aluminium without oxide layer
In the NESA activity series taught in this lesson, the order is Mg > Al > Zn. Aluminium sits between magnesium and zinc — it is less reactive than magnesium but more reactive than zinc. Without passivation, Al's true reactivity would produce a rate of H2 between that of Mg (52 mL min−1) and Zn (28 mL min−1) — accept estimates in the range of approximately 30–50 mL min−1 with appropriate justification. Justification: Al sits below Mg but above Zn in the NESA series, meaning Al loses electrons more readily than Zn but less readily than Mg under the same conditions.
Q1.4 — Balanced equation for Zn + HCl
Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g) (1 mark for equation correct, 1 mark for state symbols)
Q2.1 — Deduce activity series order
From the data alone, the deducible relationships are: Mg > Zn (Mg displaces Zn2+); Zn > Cu and Fe > Cu (both displace Cu2+); Ag < Cu (Ag cannot displace Cu2+). The data does not contain a direct Fe vs Zn comparison, so both Mg > Fe > Zn and Mg > Zn > Fe are consistent with the experimental data. Accept either order for Fe and Zn provided the student notes the data is insufficient to separate them. Most to least reactive (consistent with the NESA series taught in this lesson): Mg > Zn > Fe > Cu > Ag. Supporting evidence (accept any one): Mg displaces Zn2+ from ZnSO4 (grey zinc solid deposits) confirming Mg > Zn [1]; Zn and Fe both displace Cu2+ from CuSO4 confirming both are above Cu [1]; Ag does not displace Cu2+, confirming Ag < Cu [1]. (3 marks for correct order with valid evidence; 2 marks for mostly correct order with evidence; 1 mark for partially correct.)
Q2.2 — Ionic equation for Zn + CuSO4
Full equation: Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s)
Net ionic equation: Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
Spectator ion: SO42− (it does not appear in the net ionic equation). [1 mark for correct equation; 1 mark for identifying the spectator ion]
Q2.3 — Platinum in silver nitrate
No reaction would occur. Platinum is below gold in the activity series (least reactive group), which means Pt is below Ag. A metal will only displace a less reactive metal ion from solution. Since Pt is less reactive than Ag, Pt cannot displace Ag+ from solution. No electron transfer occurs. [1 mark for correct prediction; 1 mark for correct reasoning using relative activity series positions]
Q3.1 — Why galvanised steel protects even when scratched
Zinc is higher than iron in the activity series, meaning zinc has a greater tendency to lose electrons (be oxidised) [1]. When the coating is scratched and both zinc and iron are exposed to moisture (an electrolyte), zinc acts as the sacrificial anode — it loses electrons preferentially (Zn → Zn2+ + 2e−) [1]. Iron acts as the cathode and receives electrons, so iron is not oxidised. The zinc “sacrifices” itself to protect the underlying steel [1].
Q3.2 — Rust near copper screws
Copper is below zinc in the activity series. When the copper screw is in direct metal contact with the zinc-coated steel (and moisture is present as an electrolyte), a galvanic cell forms [1]. In this pair, zinc is more reactive than copper, so zinc near the screw corrodes preferentially (acting as the anode) — this is normal sacrificial protection [1]. However, the underlying iron of the steel sheet is now exposed at the screw hole edge, and iron is above copper in the series. Iron near the copper screw therefore becomes the anode and is preferentially oxidised (corrodes = rust) relative to the copper screw, explaining the rust spots at the junction [1].
Q4.1 — Which pipe corrodes
Iron (Fe) is higher than copper (Cu) in the activity series, meaning Fe has a greater tendency to lose electrons [1]. When in direct contact with an electrolyte (the hot-water solution), iron acts as the anode and copper acts as the cathode [1]. Electrons flow from the iron pipe to the copper pipe: Fe(s) → Fe2+(aq) + 2e−. The iron corrodes; the copper is protected. This explains why the pinhole develops in the iron section at the junction [1].
Q4.2 — Why the dielectric union prevents galvanic corrosion
Galvanic corrosion requires three conditions: two metals of different reactivity, an electrolyte (the water), AND an electrical (metallic) connection between the two metals for electrons to flow [1]. The dielectric union breaks the electrical connection between iron and copper, so electrons cannot flow from iron to copper. Without electron flow, no oxidation of iron occurs — no galvanic corrosion [1].
Q5 — Compare and contrast table
Physical barrier? (Galvanised): Yes — the zinc layer physically covers and seals the steel surface.
Which metal corrodes preferentially? (Sacrificial anode): The zinc block (anode) corrodes preferentially, protecting the iron structure (cathode).
Protection when consumed? (Galvanised): No — once the zinc coating is fully corroded through, the iron is unprotected and rusts.
Activity series principle (both): Zinc is higher than iron in the activity series (more reactive), so zinc is preferentially oxidised in both cases.
Typical Australian application (Sacrificial anode): Steel jetties, ship hulls, underground pipes, oil rigs (e.g. coastal infrastructure in Sydney Harbour, Port Hedland offshore platforms).
One limitation (Galvanised): Once the zinc layer is fully consumed, iron is exposed and corrosion accelerates. / One limitation (Sacrificial anode): Zinc blocks must be inspected and replaced regularly (typically every 1–5 years), adding maintenance costs.