Welcome to Metals and the Reactivity Series!
Have you ever wondered why gold jewellery stays shiny for thousands of years, but an iron nail left out in the rain turns reddish-brown and flaky within days? Or why magnesium burns with a blinding white light while copper barely changes when heated? It all comes down to chemical reactivity.
In this chapter of CCEA GCSE Science Single Award (Unit 2: Chemistry), you will learn how different metals behave, how to rank them in order of their reactivity, how we extract metals from the ground, and how to stop iron from rusting. Don't worry if science feels tough sometimes — we will break down every single idea step by step!
1. What is the Reactivity Series?
The reactivity series is a list that places metals in order from the most reactive at the top to the least reactive at the bottom. A metal's reactivity depends on how easily it loses electrons to form positive ions and react with other substances like water, air, or acids.
The Order of Metals
Here is the reactivity series from most reactive to least reactive that you need to know for your exam:
• Potassium (\(\text{K}\)) — Most reactive
• Sodium (\(\text{Na}\))
• Calcium (\(\text{Ca}\))
• Magnesium (\(\text{Mg}\))
• Aluminium (\(\text{Al}\))
• Zinc (\(\text{Zn}\))
• Iron (\(\text{Fe}\))
• Copper (\(\text{Cu}\)) — Least reactive (among common school laboratory metals)
• (Silver \(\text{Ag}\) and Gold \(\text{Au}\) are even lower down and unreactive)
Memory Trick to Remember the Order
Try memorising this catchy sentence to remember the first letters in order:
Please Send Callum Monkeys And Zebras In Cages
(Potassium, Sodium, Calcium, Magnesium, Aluminium, Zinc, Iron, Copper)
Key Takeaway: The higher up a metal is in the series, the more vigorously and quickly it reacts with other chemicals.
2. Reactions of Metals
A. Reactions of Metals with Cold Water
When reactive metals react with cold water, they produce a metal hydroxide (an alkaline solution) and hydrogen gas.
General Word Equation:
\( \text{metal} + \text{water} \rightarrow \text{metal hydroxide} + \text{hydrogen} \)
Observations to know:
• Potassium (\(\text{K}\)): Reacts violently and instantly. It melts into a sphere, darts around on the surface, catches fire with a distinctive lilac flame, fizzes (effervesces), and quickly disappears.
• Sodium (\(\text{Na}\)): Reacts very vigorously. It melts into a silvery ball, darts around the water surface, fizzes rapidly (giving off hydrogen), and produces a yellow/orange flame if ignited.
• Calcium (\(\text{Ca}\)): Reacts steadily with cold water. It sinks, bubbles rapidly (effervescence), releases heat, and forms a cloudy white alkaline suspension of calcium hydroxide (\(\text{Ca(OH)}_2\)).
• Magnesium (\(\text{Mg}\)): Reacts extremely slowly with cold water (only a few bubbles after several hours).
• Zinc, Iron, Copper: No reaction with cold water.
Example Equation:
\( 2\text{Na} + 2\text{H}_2\text{O} \rightarrow 2\text{NaOH} + \text{H}_2 \)
B. Reactions of Metals with Steam
Metals like magnesium, zinc, and iron react much faster with steam (hot gaseous water) than with cold water. When metals react with steam, they form a metal oxide and hydrogen gas (instead of a hydroxide).
General Word Equation:
\( \text{metal} + \text{steam} \rightarrow \text{metal oxide} + \text{hydrogen} \)
• Magnesium (\(\text{Mg}\)): Burns with a bright white light in steam to leave a white powder (magnesium oxide).
\( \text{Mg} + \text{H}_2\text{O} \rightarrow \text{MgO} + \text{H}_2 \)
• Copper (\(\text{Cu}\)): Does not react with water or steam at all.
C. Reactions of Metals with Dilute Acids (e.g. Hydrochloric Acid)
When metals react with dilute hydrochloric acid, they produce a metal salt (a metal chloride) and hydrogen gas.
General Word Equation:
\( \text{metal} + \text{hydrochloric acid} \rightarrow \text{metal chloride} + \text{hydrogen} \)
Observations:
• Potassium and Sodium: Dangerously explosive! Never added to acids in a school laboratory.
• Calcium: Reacts very vigorously, fizzes rapidly, temperature rises quickly.
• Magnesium: Reacts quickly, vigorous effervescence (bubbles of gas), tube gets warm, metal dissolves.
\( \text{Mg} + 2\text{HCl} \rightarrow \text{MgCl}_2 + \text{H}_2 \)
• Zinc: Steady stream of bubbles, slower reaction than magnesium.
\( \text{Zn} + 2\text{HCl} \rightarrow \text{ZnCl}_2 + \text{H}_2 \)
• Iron: Very slow reaction, tiny steady bubbles produced.
\( \text{Fe} + 2\text{HCl} \rightarrow \text{FeCl}_2 + \text{H}_2 \)
• Copper: No reaction (copper is below hydrogen in reactivity).
Testing for Hydrogen Gas
How do we prove that the gas made is hydrogen? Place a lighted wooden splint near the mouth of the test tube. If hydrogen is present, it will burn with a distinctive squeaky pop sound.
Key Takeaway: Potassium and sodium react violently with water and acid; magnesium, zinc, and iron react moderately with acid; copper shows no reaction at all.
3. Displacement Reactions
A displacement reaction occurs when a more reactive metal pushes out (displaces) a less reactive metal from a compound or solution.
An Everyday Analogy
Imagine a playground game: A stronger, more energetic player pushes past a less energetic player to grab a seat on the bench. The more reactive metal wants to be in a compound more than the less reactive metal does!
Rules for Displacement:
• A more reactive metal + a compound of a less reactive metal \(\rightarrow\) Reaction happens (displacement occurs).
• A less reactive metal + a compound of a more reactive metal \(\rightarrow\) No reaction.
Classic Example: Iron in Copper(II) Sulfate Solution
Iron (\(\text{Fe}\)) is higher in the reactivity series than Copper (\(\text{Cu}\)). When an iron nail is placed into blue copper(II) sulfate solution:
Word Equation:
\( \text{iron} + \text{copper(II) sulfate} \rightarrow \text{iron(II) sulfate} + \text{copper} \)
Chemical Equation:
\( \text{Fe} + \text{CuSO}_4 \rightarrow \text{FeSO}_4 + \text{Cu} \)
Visible Observations to write in exams:
1. The blue solution gradually fades and turns pale green (as copper sulfate turns into iron sulfate).
2. A reddish-brown coating forms on the surface of the iron nail (this is solid copper metal being pushed out).
3. The temperature increases slightly (it is an exothermic reaction).
What if we reverse it?
If you place copper metal into iron(II) sulfate solution:
\( \text{Copper} + \text{iron(II) sulfate} \rightarrow \text{No reaction} \)
Reason: Copper is lower down than iron and cannot displace it.
Key Takeaway: A metal can only displace another metal that sits below it in the reactivity series.
4. Extraction of Metals
Most metals are found in the Earth's crust chemically combined with other elements (like oxygen or sulfur) inside rocks called ores. To obtain the pure metal, we must extract it.
How Extraction Depends on Reactivity
The method used to extract a metal depends directly on its position in the reactivity series:
1. Very High Reactivity (\(\text{K}, \text{Na}, \text{Ca}, \text{Mg}, \text{Al}\)):
• These metals hold onto oxygen very tightly.
• They are extracted by electrolysis (breaking down molten compounds using large amounts of electricity).
• Disadvantage: Very expensive because it requires huge quantities of electrical energy.
2. Medium Reactivity (\(\text{Zn}, \text{Fe}, \text{Cu}\)):
• These metals are less reactive than carbon.
• They are extracted by heating with carbon (reduction) in a furnace (e.g. the Blast Furnace for iron).
• Carbon is cheaper than electricity, so this is economical.
• How it works: Carbon is more reactive than iron, so carbon removes (displaces) the oxygen from the iron oxide:
\( \text{iron(III) oxide} + \text{carbon} \rightarrow \text{iron} + \text{carbon dioxide} \)
\( 2\text{Fe}_2\text{O}_3 + 3\text{C} \rightarrow 4\text{Fe} + 3\text{CO}_2 \)
3. Very Low Reactivity (\(\text{Ag}, \text{Au}\)):
• Found uncombined (native) as pure elements in the Earth's crust because they are so unreactive (they do not easily react with oxygen).
Key Takeaway: High-reactivity metals require electrolysis; medium-reactivity metals are reduced with carbon; unreactive metals are found native.
5. Rusting of Iron
Rusting is a specific chemical reaction that applies only to iron and steel (all other metals simply "corrode"). Rusting destroys iron structures, bridges, cars, and ships over time.
What is Rust?
Rust is chemically known as hydrated iron(III) oxide. It is an orange-brown, flaky solid.
Word Equation for Rusting:
\( \text{iron} + \text{oxygen} + \text{water} \rightarrow \text{hydrated iron(III) oxide} \)
The Essential Conditions for Rusting
Both oxygen (from the air) AND water are strictly required for iron to rust. If either is missing, iron will NOT rust.
Classic Experiment: Three Test Tubes
• Tube 1 (Water + Air): Iron nail placed in tap water open to air \(\rightarrow\) RUSTS (both water and oxygen are present).
• Tube 2 (Dry Air only): Iron nail with anhydrous calcium chloride (a drying agent) and a stopper \(\rightarrow\) NO RUST (water is absent).
• Tube 3 (Water without Air): Iron nail in boiled water covered with a layer of oil \(\rightarrow\) NO RUST (boiling removes dissolved air, oil blocks oxygen from entering).
Methods of Rust Prevention
To prevent rusting, we must keep water and oxygen away from the iron surface:
1. Barrier Methods:
Coat the iron to create a physical barrier between the metal, oxygen, and moisture.
• Painting: Used for cars, gates, and bridges.
• Oiling / Greasing: Used on moving machinery parts and bicycle chains (where paint would rub off).
• Plastic coating: Used on dish racks and wire fencing.
2. Galvanising:
Coating iron with a layer of zinc.
• Zinc acts as a physical barrier.
• Even if the zinc is scratched, zinc is more reactive than iron, so it reacts with oxygen and water first, protecting the iron beneath it (sacrificial protection).
3. Sacrificial Protection:
Attaching blocks of a more reactive metal (such as magnesium or zinc) to the iron structure.
• Used on underground pipelines and large steel ship hulls.
• The more reactive metal oxidises and corrodes "sacrificially" instead of the iron. These blocks must be replaced periodically.
Key Takeaway: Rusting strictly needs both oxygen and water. Rust can be prevented by barriers (paint, oil, plastic), galvanising (zinc coating), or sacrificial protection.
6. Summary & Quick Review Checklist
Check off these essential points before your exam:
• Can you list the reactivity series from potassium down to copper?
• Do you remember that potassium makes a lilac flame with water, while sodium makes a yellow flame?
• Can you give the test for hydrogen gas (lighted splint gives a squeaky pop)?
• Can you explain displacement (a more reactive metal kicks out a less reactive metal from its solution)?
• Can you state why metals above carbon must be extracted using electrolysis?
• Do you know the two substances needed for iron to rust (water and oxygen)?
• Can you name three ways to prevent rusting (painting, oiling/greasing, galvanising/sacrificial protection)?
Well done! Keep practicing your word equations and observations, and you'll do brilliantly in your CCEA Chemistry exam!