CCEA Chemistry Unit 2: Metals and Reactivity Series

Metals Reacting with Oxygen

When you heat metals in air, they react with oxygen to form metal oxides - and the metal actually gains mass because it's combining with oxygen atoms. You'll use a crucible setup with a tripod and pipeclay triangle, lifting the lid occasionally to let more air in.

The most reactive metals like potassium and sodium are too dangerous for school labs - they'd burn with coloured flames (lilac for potassium, yellow for sodium). Magnesium burns with that brilliant white light you've probably seen, whilst metals like copper just get covered in a black layer without actually burning.

Remember: The more reactive the metal, the more vigorous its reaction with oxygen - this pattern helps you work out the reactivity series!

Each metal has its own characteristic reaction, from aluminium only burning as a fine powder to iron filings creating orange sparks. These observations directly link to how reactive each metal is.

Metals Reacting with Water

Group 1 metals (lithium, sodium, potassium) react dramatically with cold water, producing metal hydroxides and hydrogen gas. You'll see them floating, fizzing, and moving about the surface - with sodium melting into a silvery ball and potassium burning with a lilac flame.

The general equation is: metal + water → metal hydroxide + hydrogen. These reactions are exothermic (they give out heat), which is why the metals get hot enough to melt or even burn.

For safety reasons, you'll always do these reactions behind a safety screen with just tiny pieces of metal. The fizzing you see is hydrogen gas being produced - quite exciting but potentially dangerous!

Safety tip: Never use large pieces of Group 1 metals with water - they can explode!

Less Reactive Metals and Water

Calcium and magnesium need a different setup because they react more slowly with water. You'll place them in a beaker with an inverted filter funnel and boiling tube to collect the hydrogen gas produced.

Calcium still fizzes and eventually disappears, leaving a cloudy solution of calcium hydroxide. Magnesium barely reacts with cold water - just a few slow bubbles if you're lucky.

To test for hydrogen gas, bring a lit splint near the test tube opening. You'll hear that characteristic "squeaky pop" sound that confirms hydrogen is present.

Quick test: The squeaky pop test for hydrogen is a classic exam question - make sure you can describe both the method and the result!

The cloudiness in calcium hydroxide solution and the slow reaction of magnesium show how reactivity decreases as you go down this part of the series.

Steam Reactions and Setup

When metals won't react with cold water, try steam instead! The setup uses damp mineral wool to generate steam, which then reacts with heated metals to produce metal oxides and hydrogen.

Magnesium glows with bright white light and forms a white solid (magnesium oxide). Zinc creates a yellow powder that turns white when it cools, whilst iron powder glows to form a black solid.

Aluminium has a quirky behaviour - foil won't react because of its protective oxide layer, but powdered aluminium reacts readily. This is why aluminium saucepans don't dissolve when you cook!

Watch out: Prevent "suck back" by removing the apparatus from water when you stop heating - otherwise cold water can rush back into the hot tube and crack it.

The general equation here is: metal + steam → metal oxide + hydrogen. Notice how steam reactions produce oxides rather than hydroxides.

Understanding Metal Reactivity

Metal reactivity comes down to how easily metals lose electrons to form positive ions. When metals react, they follow the pattern: M → M⁺ + e⁻ (where M represents any metal).

More reactive metals like potassium and sodium lose electrons very easily, making them highly reactive. Copper holds onto its electrons tightly, which explains why it's so unreactive and doesn't even burn properly in air.

Different metals lose different numbers of electrons - potassium loses one (K⁺), calcium loses two (Ca²⁺), and aluminium loses three (Al³⁺). This links directly to their position in the periodic table.

Key insight: The easier it is for a metal to lose electrons, the more reactive it is - this explains the entire reactivity series!

Understanding this electron-loss concept helps you predict and explain all the reaction patterns you've observed.

Displacement Reactions

Displacement reactions happen when a more reactive metal kicks out a less reactive one from its compound - think of it as chemical bullying! You'll see this with solid metals reacting with metal oxides or metal salt solutions.

A classic example is aluminium displacing iron: Fe₂O₃ + 2Al → 2Fe + Al₂O₃. This reaction is so energetic it's used in welding! When magnesium displaces copper from copper sulfate solution, you'll see the blue solution fade to colourless as brown copper forms.

The reactivity series (K, Na, Ca, Mg, Al, Zn, Fe, Cu from most to least reactive) lets you predict whether displacement will happen. A metal can only displace those below it in the series.

Exam tip: If asked whether a displacement will occur, check the reactivity series - the attacking metal must be higher up than the one being displaced!

These reactions are always exothermic, so you'll feel heat being released as the more reactive metal takes over.

Metal Extraction Methods

Extracting metals from their ores (rocks containing metal compounds) depends entirely on the metal's reactivity. The more reactive the metal, the harder it is to extract because it really doesn't want to give up being combined with other elements.

Highly reactive metals (potassium to aluminium) need electrolysis - using electricity to force them out of their compounds. This is expensive, which is why aluminium used to be more valuable than gold!

Less reactive metals (zinc to copper) can be extracted by heating their ores with carbon. Carbon is reactive enough to steal oxygen from these metal oxides, leaving the pure metal behind.

Cost matters: Electrolysis uses loads of electricity, making reactive metals expensive to extract - that's why recycling aluminium is so important!

The least reactive metals like gold are sometimes found naturally as pure elements, which is why they've been used for thousands of years. The extraction method tells you a lot about the metal's properties and uses.