Introduction
Okay, Metals looks like a lot of random facts, but it's basically one idea wearing six different outfits. Here it is: a metal reacts by losing electrons to become a positive ion, and the easier it does that, the more reactive it is. That's the whole chapter's cheat code.
Line the metals up by how easily they lose electrons and you get the reactivity series, the one diagram that secretly answers everything: which metals fizz in acid, which displace which, which rust, and how each one gets pulled out of the ground. This is your fast refresh, not the full textbook: one key idea per topic, one worked example, and the exact move the marker wants. Let's lose some electrons and win some marks. ⚙️
Properties of MetalsE9.1
A metal is basically a grid of positive ions chilling in a "sea" of delocalised electrons that can roam around freely. Almost every metal property comes from that picture. Those free electrons drift when you connect a battery, so metals conduct electricity. And the layers of ions can slide over each other without the whole thing snapping, so metals are malleable (bendy) and ductile (pull into wire). Chemically, the one move metals make is losing electrons to form positive ions, e.g. $\text{Zn} \rightarrow \text{Zn}^{2+} + 2\text{e}^-$.
> 🔑 Key idea: Metal = positive ions + a sea of free electrons. Free electrons → conducts; sliding layers → bends; and metals react by losing electrons to make + ions. ⚡
Worked Example: Why Metals ConductE9.1
Problem In terms of metallic bonding, why does a metal conduct electricity?
Step 1 Structure: a metal is a lattice of positive ions surrounded by a sea of delocalised electrons.
Step 2 Link: connect a voltage and those electrons are free to move through the lattice. Moving electrons = current. So it conducts.
Just remember: it's the electrons that move, not the ions. ✅
Uses of MetalsE9.2
The two metals the syllabus loves are aluminium and copper, and they get picked for jobs that match their properties. Aluminium is light (low density) and doesn't corrode (it grows a protective oxide skin), so it's used for aircraft bodies, overhead power cables and drink cans. Copper is a brilliant electrical conductor and bends into wire easily, so it's the go-to for electrical wiring, plus it conducts heat well for saucepan bases. The trick in the exam is to link the property to why it helps the job, not just name it.
> 🔑 Key idea: Match property → use. Aluminium = light + corrosion-proof (planes, cables, cans). Copper = great conductor + ductile (wires, pans). Always say why the property helps. 🎯
Worked Example: Aluminium AircraftE9.2
Problem Why is aluminium chosen for aircraft bodies? Give two properties.
Step 1 Planes must be light → aluminium has a low density, so less fuel is needed.
Step 2 Planes sit out in the weather for years → aluminium resists corrosion thanks to its oxide layer.
So: low density and resistance to corrosion. Both, together. 🔥
AlloysE9.3
An alloy is just a mixture of a metal with other elements (usually other metals). Brass = copper + zinc, stainless steel = iron + chromium + nickel. Here's the magic: alloys are harder and stronger than the pure metal. Why? A pure metal has identical atoms in tidy layers that slide over each other (so it's soft). Mix in atoms of a different size and they jam up the layers so they can't slide. Boom, harder.
> 🔑 Key idea: An alloy is a mixture, and it's harder because different-sized atoms disrupt the neat layers so they can't slide. Not "stronger bonds", the sliding layers are the whole story. 💪
Worked Example: Brass vs CopperE9.3
Problem Brass is copper mixed with zinc. Why is brass stronger than pure copper?
Step 1 Pure copper: same-size atoms in neat layers that slide → soft.
Step 2 Brass: zinc atoms are a different size, so they stop the layers sliding → stronger and harder.
The marker wants "the layers can't slide", not "stronger bonds". ✅
The Reactivity SeriesE9.4
This is the spine of the chapter: metals lined up by how easily they lose electrons. Learn it: K, Na, Ca, Mg, Al, (C), Zn, Fe, (H), Cu, Ag, Au. The top three react with cold water (→ hydroxide + hydrogen). The middle ones react with steam (→ oxide + hydrogen) and with dilute acids (→ salt + hydrogen). The bottom three (Cu, Ag, Au) basically don't react. So a metal's behaviour with water, steam and acid tells you exactly where it sits.
> 🔑 Key idea: K Na Ca Mg Al [C] Zn Fe [H] Cu Ag Au. Cold water = top metals; steam/acid = middle; no reaction = bottom. Where a metal reacts pins down its place. 🪜
Worked Example: Name That MetalE9.4
Problem Metal Z reacts with dilute acid, reacts very slowly with cold water but fast with steam. Is it magnesium, gold, potassium or silver?
Step 1 Reacts with acid → above hydrogen. Cross off gold and silver.
Step 2 Slow with cold water but fast with steam → a middle metal. Cross off potassium (it's violent with cold water).
Left standing: magnesium. 🎯
DisplacementE9.4
If you drop a more reactive metal into a solution of a less reactive metal's salt, the reactive one displaces (kicks out) the other. Zinc is above copper, so zinc displaces copper: $\text{Zn} + \text{Cu}^{2+} \rightarrow \text{Zn}^{2+} + \text{Cu}$. The leftover ion (like sulfate) just watches, so it's a spectator and gets left out of the ionic equation. And the golden rule for ranking metals: if A displaces B, A is more reactive. The catch: ionic equations have to balance charge as well as atoms.
> 🔑 Key idea: A more reactive metal displaces a less reactive one from its salt. If A displaces B, A is more reactive. Balance ionic equations for charge and atoms. 🥊
Worked Example: Aluminium Displaces ZincE9.4
Problem Balance the ionic equation $\text{Al} + \text{Zn}^{2+} \rightarrow \text{Al}^{3+} + \text{Zn}$.
Step 1 Electrons lost = electrons gained. Al loses 3, $\text{Zn}^{2+}$ gains 2. Lowest common number = 6.
Step 2 Use 2 Al (lose 6) and 3 Zn (gain 6):
Charge check: left $6+$, right $6+$. Balanced. ✅
$$2\text{Al} + 3\text{Zn}^{2+} \rightarrow 2\text{Al}^{3+} + 3\text{Zn}$$
Corrosion and RustingE9.5
Rusting is iron corroding, and it needs both water and oxygen (no water or no air = no rust). The rust itself is hydrated iron(III) oxide. Two ways to stop it: barrier methods (paint, oil, plastic, plating) just block the water and air, but only while they're intact. Sacrificial protection is cleverer: bolt on a more reactive metal (zinc or magnesium) and it loses electrons instead of the iron, so it corrodes and the iron survives even if exposed. Galvanising (zinc coating) does both at once.
> 🔑 Key idea: Rust needs water + oxygen. Block it with a barrier, or use a more reactive metal that corrodes instead (sacrificial). Galvanising = barrier + sacrificial. 🛡️
Worked Example: GalvanisingE9.5
Problem An iron nail is coated in zinc. How does this stop it rusting?
Step 1 Barrier: while the zinc covers it, water and air can't reach the iron.
Step 2 Sacrificial: zinc is more reactive than iron, so even if it's scratched, the zinc corrodes first and the iron is safe.
So it's both, and it works because zinc is more reactive than iron. ✅
Extracting MetalsE9.6
Metals come from ores (usually joined to oxygen), and extracting them means removing the oxygen (reduction). The whole decision comes down to one comparison: is the metal above or below carbon? Below carbon (Zn, Fe, Cu) → just heat the ore with carbon (cheap). Above carbon (K, Na, Ca, Mg, Al) → carbon can't shift the oxygen, so you need electrolysis (pricey but necessary). Super-unreactive ones like gold are just found as the metal already.
> 🔑 Key idea: Below carbon → reduction with carbon (cheap). Above carbon → electrolysis. The comparison is metal vs carbon, not hydrogen. ⛏️
Worked Example: Zinc vs AluminiumE9.6
Problem Zinc is below carbon, aluminium is above carbon. How is each extracted, and why?
Step 1 Zinc is below carbon, so carbon can grab its oxygen → reduction with carbon: $\text{ZnO} + \text{C} \rightarrow \text{Zn} + \text{CO}$.
Step 2 Aluminium is above carbon, so carbon can't → electrolysis.
Compare to carbon, not hydrogen. 🎯
Iron in the Blast FurnaceE9.6
Iron comes from hematite (iron(III) oxide). You load iron ore, coke (carbon) and limestone in the top and blast hot air in the bottom. The coke burns to make heat ($\text{C} + \text{O}_2 \rightarrow \text{CO}_2$), that gas reacts with more coke to make the real reducing agent, carbon monoxide ($\text{CO}_2 + \text{C} \rightarrow 2\text{CO}$), and the CO strips the oxygen off the ore. The limestone mops up the sandy impurity as slag.
> 🔑 Key idea: Coke burns for heat, makes CO, and CO reduces the ore: $\text{Fe}_2\text{O}_3 + 3\text{CO} \rightarrow 2\text{Fe} + 3\text{CO}_2$. Limestone removes sand as slag. 🏭
Worked Example: Reducing the OreE9.6
Problem Carbon monoxide reduces iron(III) oxide. Write the balanced equation.
Step 1 $\text{Fe}_2\text{O}_3$ has 3 oxygens, so you need 3 CO to carry them off as 3 $\text{CO}_2$.
Step 2 Fe 2=2, C 3=3, O 6=6. Balanced. The reducer is CO, not carbon. ✅
$$\text{Fe}_2\text{O}_3 + 3\text{CO} \rightarrow 2\text{Fe} + 3\text{CO}_2$$
Aluminium by ElectrolysisE9.6
Aluminium is above carbon, so it's electrolysis time. Its ore bauxite is cleaned to aluminium oxide, which is then dissolved in molten cryolite to lower the melting point (saves energy). Electrolyse it with carbon electrodes: aluminium forms at the negative electrode (cathode) and oxygen at the positive electrode (anode). That oxygen burns the carbon anodes away, so they keep getting replaced.
> 🔑 Key idea: Cathode: $\text{Al}^{3+} + 3\text{e}^- \rightarrow \text{Al}$. Anode: $2\text{O}^{2-} \rightarrow \text{O}_2 + 4\text{e}^-$. Cryolite lowers the melting point; the oxygen burns the carbon anodes away. ⚡
Worked Example: At the CathodeE9.6
Problem Write the half-equation at the negative electrode in the electrolysis of aluminium oxide.
Step 1 $\text{Al}^{3+}$ ions are positive, so they go to the negative electrode (cathode).
Step 2 Each one gains 3 electrons to become an aluminium atom:
Gain = cathode (RED-CAT). ✅
$$\text{Al}^{3+} + 3\text{e}^- \rightarrow \text{Al}$$
Key Points to Remember
- A metal is positive ions in a sea of free electrons: free electrons → conducts, sliding layers → bends.
- Metals react by losing electrons to form + ions: acid → salt + hydrogen; cold water → hydroxide + hydrogen; steam → oxide + hydrogen; oxygen → basic oxide.
- Aluminium (light, corrosion-proof) → aircraft, cables, cans. Copper (great conductor, ductile) → wiring, pans.
- An alloy is a mixture; it's harder because different-sized atoms stop the layers sliding.
- Reactivity series: K Na Ca Mg Al [C] Zn Fe [H] Cu Ag Au. A higher metal displaces a lower one; ionic equations balance charge and atoms.
- Rusting needs water + oxygen → hydrated iron(III) oxide. Stop it with barriers or sacrificial protection (more reactive metal); galvanising does both.
- Extraction: below carbon → reduction with carbon (iron); above carbon → electrolysis (aluminium); gold is native.
- Iron: $\text{Fe}_2\text{O}_3 + 3\text{CO} \rightarrow 2\text{Fe} + 3\text{CO}_2$, limestone makes slag. Aluminium: cathode $\text{Al}^{3+} + 3\text{e}^- \rightarrow \text{Al}$, anode $2\text{O}^{2-} \rightarrow \text{O}_2 + 4\text{e}^-$.