Chemical changes, for the exam.

Part 1Ordering the metals

Metals react by losing electrons to form positive ions. The more easily a metal loses its electrons, the more reactive it is. Putting the metals in order of how readily they react gives the reactivity series.

You also need two non-metals — carbon and hydrogen — placed in the series so you can compare them with the metals. Their positions decide which metals can be extracted with carbon, and which react with acids.

Part 2Reactions with water and acid

The reactive metals at the top fizz with cold water, giving a metal hydroxide and hydrogen. Less reactive metals don't react with water but still react with dilute acid, giving a salt and hydrogen. The more reactive the metal, the more vigorous the fizzing.

A metal below hydrogen in the series (copper, silver, gold) will not react with dilute acids at all — that's why copper coins don't dissolve in your hand.

The general patterns to remember:

metal + water → metal hydroxide + hydrogen (for the reactive metals). metal + acid → salt + hydrogen (for metals above hydrogen).

Part 1Ores and native metals

Most metals are found as compounds — usually oxides — in rocks called ores. To get the metal out, you have to remove the oxygen. Removing oxygen from a compound is called reduction.

A few very unreactive metals, like gold, are found in the Earth as the pure metal — we say they are found native. They're so unreactive they never combined with anything in the first place, so no chemistry is needed to extract them.

Part 2Extraction with carbon

A metal can be extracted from its oxide by heating with carbon — but only if the metal is less reactive than carbon. The carbon takes the oxygen, so the metal is reduced. This is how iron and zinc are extracted cheaply.

Metals more reactive than carbon (like aluminium) can't be displaced by carbon — they hold onto their oxygen too tightly — so they are extracted by electrolysis instead (Topic 7), which is far more expensive.

Part 1In terms of oxygen

The simplest definition: oxidation is gaining oxygen; reduction is losing oxygen. When one substance is oxidised, another is reduced in the same reaction — together these are called redox reactions.

When magnesium burns, magnesium gains oxygen (it's oxidised) to form magnesium oxide. When iron oxide is heated with carbon, the iron oxide loses oxygen (it's reduced).

Part 2In terms of electrons

There's a deeper definition based on electrons, which you'll use most when metals form ions and in electrolysis.

Part 1Acids make salts

An acid reacts with a base to produce a salt and water — this is called neutralisation. A base is a substance that neutralises an acid; a soluble base is called an alkali. Metal oxides and metal hydroxides are bases; metal carbonates also neutralise acids.

Which salt you get depends on the acid: hydrochloric acid makes chlorides, sulfuric acid makes sulfates, and nitric acid makes nitrates.

Part 2Naming the salt

To name the salt, take the metal from the base and the ending from the acid. So copper oxide + sulfuric acid gives copper sulfate; sodium hydroxide + hydrochloric acid gives sodium chloride.

Part 1Acids, alkalis and the ions

The pH scale runs from 0 to 14 and measures how acidic or alkaline a solution is. pH 7 is neutral; below 7 is acidic; above 7 is alkaline. You can measure pH with universal indicator (which gives a colour) or, more precisely, a pH probe.

When acids dissolve in water they release hydrogen ions, H⁺. Alkalis dissolve to release hydroxide ions, OH⁻. In neutralisation these combine to make water:

H⁺ + OH⁻ → H₂O

Part 2Strong vs weak acids

A strong acid is one that is fully ionised in water — every molecule splits up to release H⁺ ions. Hydrochloric, nitric and sulfuric acids are strong. A weak acid is only partially ionised — most stays as whole molecules. Ethanoic, citric and carbonic acids are weak.

The more H⁺ ions in solution, the lower the pH. So at the same concentration, a strong acid has a lower pH than a weak one.

Part 1The plan

To make a pure, dry sample of a soluble salt, react a warm acid with an insoluble base (a metal oxide or carbonate). Using an insoluble base is the clever bit: add it until no more dissolves, and you know all the acid has reacted — the leftover excess base just sits at the bottom and can be filtered off. That guarantees a pure salt with no leftover acid.

Part 2From solution to crystals

Part 1What electrolysis is

Electrolysis uses electricity to break down an ionic compound into its elements. The liquid that's broken down is the electrolyte; it must be molten or dissolved so the ions are free to move. The electrodes are usually inert (carbon/graphite).

Ions move to the electrode of opposite charge: positive ions (cations) go to the negative cathode; negative ions (anions) go to the positive anode. There they gain or lose electrons and become neutral atoms or molecules.

Part 2Predicting the products

Molten compounds are simplest: the metal forms at the cathode and the non-metal at the anode. So molten lead bromide gives lead at the cathode and bromine at the anode.

Aqueous solutions are trickier because the water also provides H⁺ and OH⁻ ions. Use these rules:

At the cathode: you get hydrogen — unless the metal is less reactive than hydrogen (e.g. copper, silver), in which case the metal is deposited.

At the anode: you get oxygen — unless a halide (Cl⁻, Br⁻, I⁻) is present, in which case the halogen (chlorine, bromine, iodine) is produced.