Reactivity Series & Corrosion of Metals
Revise Reactivity Series & Corrosion of Metals for Chemistry 0620 (O Level) — revision notes and instant AI marking.
The Big Idea
Metals are reactive because they lose electrons easily. The more readily a metal loses electrons, the higher it sits on the reactivity series. This explains how metals react with water and acid, why iron rusts (needing both oxygen AND water), and how we can protect metals from corrosion using barrier methods or sacrificial protection.
- The reactivity series orders metals by how easily they lose electrons
- Rusting requires both oxygen and water — never just one
- More reactive metals displace less reactive metals from salt solutions
- Sacrificial protection & galvanising are two distinct anti-corrosion methods
The Reactivity Series
What is the Reactivity Series?
The reactivity series is a list of metals (and two non-metals) arranged in order of how readily they lose electrons when they react with other substances. It's essentially a ranking of "how keen each metal is to react."
Potassium → Sodium → Calcium → Magnesium → Aluminium → Carbon → Zinc → Iron → Hydrogen → Copper → Silver → Gold
Notice that carbon and hydrogen are included (they're non-metals), because they're used to extract metals from their oxides in industry. The mnemonic is:
"Please Send Cats, Monkeys And Cute Zebras Into Hot Countries Signed Gordon"
P-S-C-M-A-C-Z-I-H-C-S-G = each first letter = each metal/non-metal in order.
Why This Order?
Metals at the top of the series lose electrons very easily, so they react vigorously. Metals at the bottom lose electrons reluctantly, so they barely react at all. The "strength" of a metal's desire to lose electrons determines its position on this list.
Metals Reacting with Cold Water
Only the most reactive metals (potassium, sodium, and calcium) react with cold water. All three form a metal hydroxide and hydrogen gas:
Potassium (K)
Reacts violently with cold water — the reaction is so vigorous it often ignites the hydrogen gas.
Sodium (Na)
Reacts quickly with cold water — slower than potassium but still energetic.
Calcium (Ca)
Reacts less strongly with cold water — noticeably slower than sodium.
When describing these reactions, use words like "vigorous," "rapidly," or "slowly" to show you understand the difference in reactivity. Don't just say "they react" — show you can observe and interpret the difference in reaction rates.
Metals Reacting with Steam
Magnesium doesn't react much with cold water, but it does react with steam. It forms a metal oxide (not a hydroxide) and hydrogen gas:
This is an important distinction: with cold water, you get hydroxides; with steam, you get oxides.
Metals Reacting with Dilute Hydrochloric Acid
All metals above hydrogen in the reactivity series will react with dilute HCl to produce a salt and hydrogen gas. Those below hydrogen (copper, silver, gold) don't react at all.
Metals That React (Above Hydrogen)
- Magnesium: Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)
- Zinc: Zn(s) + 2HCl(aq) → ZnCl₂(aq) + H₂(g)
- Iron: Fe(s) + 2HCl(aq) → FeCl₂(aq) + H₂(g)
Metals That Don't React (Below Hydrogen)
- Copper: Cu + dilute HCl → no reaction
- Silver: Ag + dilute HCl → no reaction
- Gold: Au + dilute HCl → no reaction
Students often think copper reacts with HCl. It doesn't — copper is too unreactive. Copper is found in nature as the metal (not always as an ore), which is a clue to its low reactivity. Only the very reactive metals can "kick out" hydrogen from acids.
Deducing Reactivity from Observations
In the lab, you can figure out the reactivity order by making observations:
- The more vigorous the reaction, the more reactive the metal. If potassium and sodium both react with water, the one that reacts faster and more energetically is more reactive.
- If metals don't react with water, use their reactions with dilute HCl to compare them. The one that reacts faster with acid is more reactive.
- Temperature change is a clue. The greater the temperature increase during a reaction, the more reactive the metal (more energy is released).
Practice Question
Two unknown metals, X and Y, are added to dilute hydrochloric acid. Metal X produces vigorous bubbling with a temperature increase of 18°C. Metal Y produces slow bubbling with a temperature increase of 5°C. Which metal is more reactive? Explain your reasoning.
Explaining Reactivity (Extended Tier)
This section covers the "why" behind reactivity — electron transfer, reducing agents, and displacement reactions. If you're sitting Core tier, you can skim this, but understanding it will give you much deeper exam confidence.
Why Are Some Metals More Reactive Than Others?
Metal reactivity is all about electron loss. When a metal atom reacts, it loses electrons to form a positive ion:
A metal's reactivity is a measure of how easily it loses electrons.
Metals high on the series lose electrons very readily → very reactive.
Metals low on the series lose electrons reluctantly → less reactive.
Displacement Reactions Between Metals and Metal Salts
Here's a powerful rule: Any metal will displace another metal that is below it in the reactivity series from a solution of one of its salts.
Why? Because a more reactive metal loses electrons more easily than a less reactive one. So when a reactive metal meets a solution containing a less reactive metal's ions, the reactive metal "wins" the electrons from the less reactive metal, pushing it out of the solution.
Imagine electrons are like seats at a concert. More reactive metals are "desperate" to give away their electrons (get rid of the "seats"). Less reactive metals hold onto their electrons ("seats") more tightly. When a desperate metal meets a reluctant metal, the desperate one pushes the reluctant one out and takes the seats.
The Magnesium & Copper Sulfate Reaction
A classic example is magnesium metal added to a solution of copper sulfate (CuSO₄):
What you observe: The blue color of the copper sulfate solution fades to colorless (as CuSO₄ is replaced by MgSO₄, which is colorless). A reddish-brown copper powder coats the magnesium metal and settles at the bottom.
Why this happens: Magnesium is more reactive than copper, so it loses electrons more easily. It "displaces" the copper from the solution:
- Mg loses 2 electrons → becomes Mg²⁺
- Cu²⁺ gains 2 electrons → becomes Cu (pure copper metal)
When you describe a displacement reaction, always explain it in terms of electron transfer and reactivity. Say things like: "Magnesium is more reactive than copper, so it loses electrons more readily. The copper ions gain those electrons and become copper metal, which precipitates out."
Using Displacement Reactions to Deduce Reactivity Order
You can design experiments with different metal-salt pairs and observe which reactions happen. A reaction occurs only if the metal is more reactive than the metal in the salt.
| Mixture | Reaction? | What It Tells Us |
|---|---|---|
| Mg + FeSO₄ | ✓ Yes | Mg is more reactive than Fe |
| Zn + FeSO₄ | ✓ Yes | Zn is more reactive than Fe |
| Fe + CuSO₄ | ✓ Yes | Fe is more reactive than Cu |
| Cu + FeSO₄ | ✗ No | Cu is less reactive than Fe |
| Ag + CuSO₄ | ✗ No | Ag is less reactive than Cu |
By combining these observations, you can build up the order: Mg > Zn > Fe > Cu > Ag — which matches the reactivity series!
Why Aluminium Looks Unreactive
Here's a common exam trick question: Aluminium is high on the reactivity series, yet it doesn't seem to react with water or acids. Why?
The answer: Aluminium rapidly oxidises in air, forming a protective layer of aluminium oxide (Al₂O₃). This layer is so thin it's invisible, but it completely blocks water and acids from reaching the metal beneath. So aluminium behaves as if it's unreactive, even though it's actually very reactive underneath that protective coating.
Examiners love asking "Why doesn't aluminium react with water if it's high on the reactivity series?" Show them you understand that a thin layer of aluminium oxide protects it, and you'll pick up marks others miss.
Practice Question
Zinc powder is added to a solution of copper(II) sulfate. A displacement reaction occurs. Explain why this reaction happens in terms of electron transfer and reactivity.
Rusting of Iron
What is Rusting?
Rusting is a chemical reaction between iron, water, and oxygen in the air. The product is a reddish-brown compound called hydrated iron(III) oxide.
iron + water + oxygen → hydrated iron(III) oxide (rust)
Note: Steel (which contains iron) rusts under the same conditions.
Many students think iron rusts with just water or just oxygen. This is wrong. For rusting to occur, iron must be exposed to BOTH water and oxygen simultaneously. Miss one, and no rusting happens.
Investigating Rusting: The Experiment
The classic school experiment uses three test tubes, each with an iron nail and different conditions:
| Tube 1 | Tube 2 | Tube 3 |
|---|---|---|
| Conditions: Air + Water | Conditions: Water only (no air) | Conditions: Air only (no water) |
| Boiled water with layer of oil on top to seal air out | Actually wait — read carefully. Oil layer keeps out air, and water is boiled to remove dissolved oxygen | Calcium chloride inside to absorb any water vapor |
| Result: Nail DOES rust | Result: Nail does NOT rust | Result: Nail does NOT rust |
Conclusion: Only the nail in Tube 1 (exposed to both air and water) rusts. This proves both are essential.
The oil layer in Tube 2 seals the water and prevents oxygen (air) from dissolving into it. The boiling removes dissolved oxygen that was already in the water. Together, they ensure Tube 2 has water but no oxygen. Tube 3 has anhydrous (water-free) air — no moisture at all.
Rust Prevention: Barrier Methods
If you coat iron or steel with a barrier that prevents water and oxygen from reaching the metal, it won't rust. Common barrier methods include:
- Paint — creates a physical seal
- Oil or grease — coats the surface and excludes water
- Plastic coating — prevents contact with air and water
The catch: If the coating gets scratched or damaged, the iron underneath is exposed to air and water again — and it will rust.
When asked to describe rust prevention, always mention both what the barrier does (excludes water and oxygen) AND the limitation (it's only effective as long as the coating remains intact). This shows deeper understanding.
Practice Question
A car body is painted to prevent rusting. Explain why painting prevents rusting, and explain what happens if the paint is scratched.
Galvanising & Sacrificial Protection (Extended Tier)
This section is required for Extended tier but optional for Core. However, understanding sacrificial protection is brilliant for deepening your grasp of the reactivity series, so read it even if you're Core.
Sacrificial Protection: The Concept
Instead of a barrier, you can protect iron using sacrificial protection. The idea: attach a more reactive metal to your iron object. This metal will oxidise (corrode) first, leaving the iron safe.
A more reactive metal attached to iron will lose electrons (oxidise) before iron does, protecting iron from corrosion. The more reactive metal is "sacrificed" so the iron is saved.
Why it works: More reactive metals lose electrons more easily. When both metals are present and exposed to water, the more reactive one loses its electrons first and corrodes. The less reactive metal (iron) stays protected because the sacrificial metal is "taking the hit."
Real-World Example: Zinc Bars on Ships
Steel ships are exposed to seawater constantly. Rust would damage the hull dangerously. The solution: attach blocks of zinc to the ship's hull.
Why zinc? Zinc is more reactive than iron, so it oxidises instead of iron:
How much zinc is needed? Surprisingly, the zinc blocks don't need to cover the entire hull. As long as they're electrically connected to the steel, they protect nearby areas too. They're typically placed on edges and joints where corrosion is most likely.
Students often ask: "But if you need to protect the whole ship, don't you need zinc covering everything?" The answer is no — electrical connection allows the protective effect to spread. This is a great detail to mention in an exam answer to show you really understand the mechanism.
The maintenance catch: The zinc blocks gradually corrode away. They have to be replaced regularly to maintain protection.
Galvanising: Protection by Coating
Galvanising is a process where iron is coated with a layer of zinc. This can be done by:
- Electroplating — using electric current to deposit zinc onto iron
- Hot dipping — dipping iron into molten zinc
When zinc reacts with oxygen and carbon dioxide in the air, it forms zinc carbonate (ZnCO₃), which protects the iron by the barrier method.
Galvanising provides TWO types of protection:
1. Barrier protection: The zinc coating stops water and oxygen reaching the iron.
2. Sacrificial protection: If the coating is damaged and the iron is exposed, zinc is still present nearby and more reactive, so it corrodes instead of iron.
This makes galvanising very robust. Even if the surface is scratched, the iron is still protected by the sacrificial effect of zinc.
Choosing a Sacrificial Metal
For sacrificial protection to work, the metal must be more reactive than iron. If you used a less reactive metal (like copper), the iron would corrode instead — the opposite of what you want!
You might be asked: "Why would copper not be suitable as a sacrificial metal to protect steel?" Answer: "Copper is less reactive than iron, so iron would corrode in preference to copper. For sacrificial protection to work, the protecting metal must be more reactive than iron."
Practice Question
A steel offshore platform uses zinc plates attached to its steel frame to prevent corrosion. Explain how this method of protection works, and explain why the zinc plates must be replaced periodically.
What to Memorise
These are the core facts and definitions you need to remember for your exam.
The Reactivity Series (in order)
K → Na → Ca → Mg → Al → C → Zn → Fe → H → Cu → Ag → Au
Remember with: "Please Send Cats, Monkeys And Cute Zebras Into Hot Countries Signed Gordon"
What Rusting Requires
BOTH iron AND water AND oxygen
Rust is hydrated iron(III) oxide (Fe₂O₃·xH₂O). It's reddish-brown.
Metals + Cold Water
Only K, Na, Ca react with cold water.
Product: metal hydroxide + H₂ gas
Metals + Dilute HCl
Only metals above hydrogen in the series react.
Product: salt + H₂ gas
Mg + Steam
Mg(s) + H₂O(g) → MgO(s) + H₂(g)
Forms oxide (not hydroxide), unlike reaction with cold water.
Displacement Reactions
A more reactive metal displaces a less reactive metal from a salt solution.
Example: Mg + CuSO₄ → MgSO₄ + Cu
Barrier Methods
Paint, oil, grease, plastic prevent water and oxygen reaching metal.
Weakness: if coating is scratched, corrosion starts.
Sacrificial Protection
A more reactive metal corrodes instead of iron.
Example: Zinc bars on ships protect steel.
Galvanising
Iron coated with zinc (by electroplating or hot dipping).
Provides both barrier AND sacrificial protection.
Key Fact: Aluminium
High on reactivity series, but appears unreactive because an invisible layer of Al₂O₃ protects it.
Concepts Checklist
Check off each concept as you master it. Aim for 100%.
Exam Tips & Common Mistakes
Avoid these traps and master the techniques examiners love.
The "Rusting Trap"
Students often say rusting requires "water and oxygen" but forget to emphasize that BOTH are essential. Never say "rusting needs moisture or oxygen" — say "rusting requires water AND oxygen together." Many marks are lost this way.
Copper with HCl
A common misconception: copper reacts with dilute HCl. It doesn't. Copper is below hydrogen in the series, so it's too unreactive. Learn this rule: only metals above hydrogen react with dilute HCl.
"Rusting" vs. "Corrosion"
Only iron and steel "rust." Other metals "corrode." If you say "copper rusts," you've made a terminology error that costs marks. Examiners are strict about this.
Explaining Displacement
Always explain displacement reactions using the language of electron transfer and reactivity. Say "Mg is more reactive, so it loses electrons more readily, and Cu²⁺ gains these electrons" — not just "Mg displaces Cu."
The Aluminium Question
If asked why aluminium doesn't react with water despite being high on the reactivity series, mention the thin, invisible layer of Al₂O₃ that forms and protects it. Extended tier especially loves this answer.
Barrier vs. Sacrificial
Know the difference: Barrier methods exclude water and oxygen. Sacrificial methods use a more reactive metal to corrode instead. Galvanising does both, which is why it's so effective.
Extended Tier Strength
If you're Extended tier, emphasize electron transfer and reactivity in every answer. Show you understand the "why" behind observations, not just the "what." This separates high marks from average marks.
Observation & Deduction
Examiners ask: "A metal reacts faster with acid than another metal — which is more reactive?" Always link faster reaction → more vigorous → more reactive. Temperature changes are also clues: larger temperature rise = more reactive.
Final Practice Questions
Question 1: Reactivity Series Ranking
Iron, copper, and zinc are three metals. Iron reacts with dilute hydrochloric acid, copper does not, and zinc reacts more vigorously than iron. Use this information to arrange the three metals in order of reactivity, starting with the most reactive.
Question 2: Rusting Prevention
A steel window frame is protected from rusting by painting it. Explain why painting prevents rusting. In areas where the paint is chipped, rust forms. Explain why rust forms at these locations.
Question 3: Displacement Reaction (Extended Tier)
When magnesium is added to a solution of iron(II) sulfate, a displacement reaction occurs. The solution changes from pale green to colorless. Explain this reaction in terms of reactivity and electron transfer.
Question 4: Sacrificial Protection (Extended Tier)
Suggest why magnesium is not used as a sacrificial metal to protect iron in seawater, but zinc is. (Hint: think about the reactivity series and practicality.)
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