Welcome to Metals and the Reactivity Series!
Have you ever wondered why gold jewellery stays shiny forever, but an old iron bicycle left in the rain turns into a crumbly, orange mess? It all comes down to chemical reactivity. In this chapter, we will discover how different metals behave, how they compete with each other, and how we extract useful metals from the Earth's crust. Don't worry if chemistry sometimes feels overwhelming — we will break down every single reaction step-by-step!
---1. The Reactivity Series
The reactivity series is a list of metals arranged in order of their reactivity, from the most reactive at the top to the least reactive at the bottom.
When metals react, their atoms want to lose electrons to form positive ions (cations).
• More reactive metals lose electrons very easily and react vigorously.
• Less reactive metals hold onto their electrons much more tightly and react slowly or not at all.
The Order of Metals
Here is the reactivity series you need to know for your CCEA GCSE course, from most reactive to least reactive:
1. Potassium (\(\text{K}\)) — Most reactive
2. Sodium (\(\text{Na}\))
3. Calcium (\(\text{Ca}\))
4. Magnesium (\(\text{Mg}\))
5. Aluminium (\(\text{Al}\))
6. Zinc (\(\text{Zn}\))
7. Iron (\(\text{Fe}\))
8. Copper (\(\text{Cu}\)) — Least reactive
Memory Aid: Easy Mnemonic
To remember the order easily, use this popular rhyme:
Please (Potassium)
Send (Sodium)
Cats (Calcium)
Monkeys (Magnesium)
And (Aluminium)
Zebras (Zinc)
In (Iron)
Cages (Copper)
Key Takeaway: The higher a metal is in the series, the more easily it loses electrons and the more violently it reacts.
---2. Reactions of Metals
A. Reactions with Cold Water
Only the most reactive metals react with cold water. When they do, they produce a metal hydroxide and hydrogen gas.
General Word Equation:
\(\text{metal} + \text{water} \rightarrow \text{metal hydroxide} + \text{hydrogen}\)
1. Potassium (\(\text{K}\)):
• Reacts violently and instantly.
• Floats, melts into a ball, dashes across the surface.
• Produces a lilac flame.
• Equation: \(2\text{K}\text{ (s)} + 2\text{H}_2\text{O}\text{ (l)} \rightarrow 2\text{KOH}\text{ (aq)} + \text{H}_2\text{ (g)}\)
2. Sodium (\(\text{Na}\)):
• Reacts vigorously.
• Floats, melts into a silvery ball, fizzes rapidly (effervescence), moves across the water surface.
• May produce a yellow-orange flame if trapped.
• Equation: \(2\text{Na}\text{ (s)} + 2\text{H}_2\text{O}\text{ (l)} \rightarrow 2\text{NaOH}\text{ (aq)} + \text{H}_2\text{ (g)}\)
3. Calcium (\(\text{Ca}\)):
• Reacts steadily.
• Sinks, then rises as bubbles of hydrogen gas form on it.
• The water turns cloudy/milky because calcium hydroxide is only slightly soluble.
• Equation: \(\text{Ca}\text{ (s)} + 2\text{H}_2\text{O}\text{ (l)} \rightarrow \text{Ca(OH)}_2\text{ (aq/s)} + \text{H}_2\text{ (g)}\)
4. Magnesium (\(\text{Mg}\)):
• Reacts extremely slowly with cold water — only a few tiny bubbles over several days.
Note: Aluminium, Zinc, Iron, and Copper do not react with cold water at all.
B. Reactions with Steam
Metals in the middle of the reactivity series (like Magnesium, Zinc, and Iron) react much faster with steam (hot water vapour) than with liquid water. When a metal reacts with steam, it forms a metal oxide (NOT a hydroxide) and hydrogen gas.
General Word Equation:
\(\text{metal} + \text{steam} \rightarrow \text{metal oxide} + \text{hydrogen}\)
• Magnesium: Burns with a brilliant, blinding white light to form a white powder (\(\text{MgO}\)).
\(\text{Mg}\text{ (s)} + \text{H}_2\text{O}\text{ (g)} \rightarrow \text{MgO}\text{ (s)} + \text{H}_2\text{ (g)}\)
• Zinc: Reacts when heated strongly. Zinc oxide is yellow when hot and turns white when cold.
\(\text{Zn}\text{ (s)} + \text{H}_2\text{O}\text{ (g)} \rightarrow \text{ZnO}\text{ (s)} + \text{H}_2\text{ (g)}\)
• Iron: Must be heated glowing red hot to react slowly.
\(3\text{Fe}\text{ (s)} + 4\text{H}_2\text{O}\text{ (g)} \rightarrow \text{Fe}_3\text{O}_4\text{ (s)} + 4\text{H}_2\text{ (g)}\)
• Copper: No reaction with steam.
C. Reactions with Air (Oxygen)
When heated in air, metals react with oxygen to form metal oxides.
General Word Equation:
\(\text{metal} + \text{oxygen} \rightarrow \text{metal oxide}\)
• Potassium / Sodium: Burn readily with bright flames (lilac for \(\text{K}\), yellow for \(\text{Na}\)) to form white solid oxides.
• Calcium: Burns with a brick-red flame to produce white calcium oxide (\(\text{CaO}\)).
• Magnesium: Burns with a bright white flame to produce white magnesium oxide (\(\text{MgO}\)):
\(2\text{Mg}\text{ (s)} + \text{O}_2\text{ (g)} \rightarrow 2\text{MgO}\text{ (s)}\)
• Aluminium: Burns only when powdered to form white aluminium oxide (\(\text{Al}_2\text{O}_3\)).
• Zinc: Burns with a blue-green flame, forming \(\text{ZnO}\) (yellow when hot, white when cool).
• Iron: Iron filings glow with orange sparks to produce black iron oxide (\(\text{Fe}_3\text{O}_4\)).
• Copper: Does not burn; the surface slowly turns black as a layer of copper(II) oxide forms:
\(2\text{Cu}\text{ (s)} + \text{O}_2\text{ (g)} \rightarrow 2\text{CuO}\text{ (s)}\)
Key Takeaway: Metal + cold water \(\rightarrow\) metal hydroxide + hydrogen. Metal + steam \(\rightarrow\) metal oxide + hydrogen. Metal + oxygen \(\rightarrow\) metal oxide.
---3. Displacement Reactions
A displacement reaction is a reaction where a more reactive metal takes the place of a less reactive metal from its compound.
Everyday Analogy: Imagine a stronger person taking a seat from someone weaker. The stronger (more reactive) metal always wins the compound partner!
Displacement in Solutions
If you place a strip of magnesium into a blue solution of copper(II) sulfate:
1. Magnesium is more reactive than copper.
2. Magnesium displaces copper from the sulfate solution.
3. Word Equation:
\(\text{magnesium} + \text{copper(II) sulfate} \rightarrow \text{magnesium sulfate} + \text{copper}\)
4. Symbol Equation:
\(\text{Mg}\text{ (s)} + \text{CuSO}_4\text{ (aq)} \rightarrow \text{MgSO}_4\text{ (aq)} + \text{Cu}\text{ (s)}\)
What you would observe:
• The blue colour of the solution fades and becomes colourless (as copper sulfate turns into magnesium sulfate).
• Red-brown solid (copper metal) coats the magnesium or falls to the bottom.
• The beaker gets warm (this reaction is exothermic).
What if we put Copper into Magnesium Sulfate?
No reaction! Copper is less reactive than magnesium, so it cannot displace it.
Solid State Displacement (The Thermite Reaction)
Displacement also happens between solids when heated strongly.
Example: Aluminium powder heated with iron(III) oxide:
\(2\text{Al}\text{ (s)} + \text{Fe}_2\text{O}_3\text{ (s)} \rightarrow \text{Al}_2\text{O}_3\text{ (s)} + 2\text{Fe}\text{ (l)}\)
Because aluminium is more reactive than iron, it takes the oxygen. This reaction produces so much heat that the iron forms as a molten liquid! It is used to weld railway tracks together.
Key Takeaway: A more reactive metal will displace a less reactive metal from its oxide or salt solution.
---4. Extraction of Metals
Most metals are found in the Earth's crust chemically combined with oxygen or other elements in rocks called ores. Very unreactive metals (like gold and platinum) are found native (as uncombined pure elements).
The method used to extract a metal depends directly on its position in the reactivity series:
1. Above Carbon (Potassium, Sodium, Calcium, Magnesium, Aluminium):
• Extracted using electrolysis (splitting compounds using electricity).
• These metals form very strong bonds and require large amounts of electrical energy, making extraction expensive.
2. Below Carbon (Zinc, Iron, Copper):
• Extracted by reduction with carbon or carbon monoxide in a furnace.
• Carbon is cheaper than electricity, so heating with carbon is preferred whenever possible.
3. Least Reactive (Silver, Gold):
• Found native; no chemical extraction needed, just physical separation.
5. Extraction of Iron in the Blast Furnace
Iron is extracted from its ore, haematite (\(\text{Fe}_2\text{O}_3\)), inside a giant structure called a Blast Furnace.
Raw Materials Added at the Top:
1. Haematite — Iron ore containing \(\text{Fe}_2\text{O}_3\) (and sandy impurities like \(\text{SiO}_2\)).
2. Coke — Mostly carbon (\(\text{C}\)); acts as fuel and reducing agent.
3. Limestone — Calcium carbonate (\(\text{CaCO}_3\)); removes acidic impurities.
4. Hot Air — Blown in near the bottom to provide oxygen.
Step-by-Step Chemical Reactions:
Step 1: Combustion of Carbon
Coke burns in hot air to produce carbon dioxide. This reaction is highly exothermic and heats the furnace:
\(\text{C}\text{ (s)} + \text{O}_2\text{ (g)} \rightarrow \text{CO}_2\text{ (g)}\)
Step 2: Production of Carbon Monoxide
Carbon dioxide reacts with more hot coke to form carbon monoxide (the main reducing agent):
\(\text{CO}_2\text{ (g)} + \text{C}\text{ (s)} \rightarrow 2\text{CO}\text{ (g)}\)
Step 3: Reduction of Iron(III) Oxide
Carbon monoxide reduces the iron(III) oxide to molten iron, which sinks to the bottom:
\(\text{Fe}_2\text{O}_3\text{ (s)} + 3\text{CO}\text{ (g)} \rightarrow 2\text{Fe}\text{ (l)} + 3\text{CO}_2\text{ (g)}\)
(Reduction means loss of oxygen).
Step 4: Removing Acidic Impurities (Slag Formation)
Haematite contains sandy impurities of silicon dioxide (\(\text{SiO}_2\)).
• Limestone thermally decomposes into calcium oxide:
\(\text{CaCO}_3\text{ (s)} \rightarrow \text{CaO}\text{ (s)} + \text{CO}_2\text{ (g)}\)
• Calcium oxide (a basic oxide) reacts with silicon dioxide (an acidic oxide) to form liquid slag (calcium silicate):
\(\text{CaO}\text{ (s)} + \text{SiO}_2\text{ (s)} \rightarrow \text{CaSiO}_3\text{ (l)}\)
• Molten slag floats on top of molten iron and is tapped off separately.
Key Takeaway: Carbon monoxide reduces iron(III) oxide to molten iron. Limestone is added to remove sandy impurities as slag.
---6. Rusting of Iron and Corrosion
Corrosion is the destruction of a metal by reaction with substances in its environment. Rusting is the specific term used ONLY for the corrosion of iron and steel.
Chemical Name of Rust: Hydrated iron(III) oxide (\(\text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O}\)).
Conditions Required for Rusting
Iron will only rust if BOTH of the following are present:
1. Oxygen (from air)
2. Water
Did you know? Salt (sodium chloride) does not cause rust by itself, but it speeds up the rusting process significantly!
Methods of Rust Prevention
1. Barrier Methods:
Keep oxygen and water away from the iron surface.
• Painting: Used for car bodies, gates, bridges.
• Oiling / Greasing: Used for moving engine parts and bicycle chains.
• Plastic Coating: Used on dish drainers and wire fences.
Drawback: If the barrier is scratched, the iron underneath will rust.
2. Sacrificial Protection:
Attaching blocks of a more reactive metal (like zinc or magnesium) to iron.
• The more reactive metal oxidises (loses electrons) preferentially instead of iron.
• Used on underground steel pipes and ship hulls.
3. Galvanising:
Coating iron or steel with a layer of zinc.
• Zinc acts as both a barrier and provides sacrificial protection if the coating is scratched.
Key Takeaway: Rusting needs both water and oxygen. We prevent it using barriers, sacrificial protection, or galvanising.
---7. Quick Review: Common Exam Traps to Avoid
• Don't say "rusting" for other metals: Copper tarnishes, aluminium corrodes, but only iron and steel rust.
• Watch the products: Metal + cold water makes a hydroxide; metal + steam makes an oxide.
• Remember the slag equation: \(\text{CaO} + \text{SiO}_2 \rightarrow \text{CaSiO}_3\). Calcium oxide is basic; silicon dioxide is acidic — this is a neutralisation reaction!