Chemical Equations and Reactions Revision Notes | GCSE Edexcel Higher Chemistry

1. From Words to Symbols

Chemical reactions happen everywhere - when you light a match, digest food, or charge your phone. But how do we write these reactions down clearly? You'll notice chemists have developed a shorthand system that's like learning a new language.

Think of it like translating from English to French. First you understand the English sentence, then you convert it word by word, then you check the grammar makes sense. Chemical equations work the same way - we start with words, convert to symbols, then balance everything.

Word equations are your starting point. They tell the story of what happens: reactants (starting materials) go in, products (what you get) come out. The arrow means 'reacts to form' or 'produces'.

Worked example: Writing a Word Equation

Magnesium burns in oxygen to produce magnesium oxide. Write the word equation.

1. magnesium + oxygen — These are the starting materials that react together
2. magnesium + oxygen → — The arrow means 'reacts to form' or 'produces'
3. magnesium + oxygen → magnesium oxide — This is what forms when the reactants combine

Answer: magnesium + oxygen → magnesium oxide

2. Converting to Symbol Equations

Word equations are clear but clunky. Symbol equations are the chemist's shorthand - they pack more information into less space. You need to know the chemical formulae for common compounds and elements.

Here's where students often stumble. Many elements exist as molecules, not single atoms. Hydrogen is $H_2$, not $H$. Oxygen is $O_2$, not $O$. These are called diatomic elements - they travel in pairs.

Worked example: Converting Word to Symbol Equation

Convert this word equation: magnesium + oxygen → magnesium oxide

1. Mg — Magnesium is a metal element, symbol Mg
2. Mg + $O_2$ — Oxygen is diatomic - it exists as $O_2$ molecules
3. Mg + $O_2$ → MgO — Magnesium forms Mg²⁺ ions, oxygen forms O²⁻ ions, so MgO

Answer: Mg + $O_2$ → MgO

But wait - this equation isn't finished yet. Count the oxygen atoms: 2 on the left, 1 on the right. Atoms can't just disappear! This is where balancing comes in.

3. The Balancing Act

Balanced equations follow the law of conservation of mass - atoms can't be created or destroyed, only rearranged. You must have the same number of each type of atom on both sides.

The key rule: you can only change the big numbers (coefficients) in front of formulae. Never change the small numbers (subscripts) in the formulae themselves - that would change the compound into something completely different.

Worked example: Balancing the Magnesium Equation

Balance: Mg + $O_2$ → MgO

1. Left: 1 Mg, 2 O Right: 1 Mg, 1 O — Oxygen is not balanced - 2 left, 1 right
2. Mg + $O_2$ → 2MgO — Put 2 in front of MgO to get 2 oxygen atoms on right
3. Left: 1 Mg Right: 2 Mg — Now magnesium is unbalanced - 1 left, 2 right
4. 2Mg + $O_2$ → 2MgO — Put 2 in front of Mg to balance both sides
5. Left: 2 Mg, 2 O Right: 2 Mg, 2 O — Both sides balanced - equation complete

Answer: 2Mg + $O_2$ → 2MgO

This looks intimidating, but there's a systematic approach. Balance the most complex molecule first, then work through the others. Always do a final atom count to check your work.

4. State Symbols - The Full Picture

Professional chemists always include state symbols to show whether substances are solid, liquid, gas, or dissolved in water. These matter because they affect how reactions happen and what you observe.

The complete magnesium equation becomes: 2Mg(s) + $O_2$(g) → 2MgO(s). This tells you magnesium metal burns in oxygen gas to form solid magnesium oxide - exactly what you'd see in the lab.

Worked example: Adding State Symbols

Add state symbols to: $HCl + NaOH → NaCl + H_2O$

1. Acid + alkali → salt + water — This is a neutralisation reaction in solution
2. Room temperature, in water — Neutralisation happens in aqueous solution
3. HCl(aq) + NaOH(aq) → NaCl(aq) + $H_2O$(l) — Acid, alkali and salt dissolved; water is liquid

Answer: HCl(aq) + NaOH(aq) → NaCl(aq) + $H_2O$(l)

5. Higher Tier - Ionic and Half Equations

For Higher Tier students, we can dig deeper. Many reactions involve ions - charged particles formed when atoms gain or lose electrons. Ionic equations show only the particles that actually react, cutting out the 'spectator ions' that just watch.

Take the neutralisation above. The full ionic equation shows all the ions: $H^+$(aq) + $Cl^-$(aq) + $Na^+$(aq) + $OH^-$(aq) → $Na^+$(aq) + $Cl^-$(aq) + $H_2O$(l). But $Na^+$ and $Cl^-$ appear unchanged on both sides - they're spectators.

Worked example: Writing an Ionic Equation

Write the ionic equation for: HCl(aq) + NaOH(aq) → NaCl(aq) + $H_2O$(l)

1. $H^+$ + $Cl^-$ + $Na^+$ + $OH^-$ → $Na^+$ + $Cl^-$ + $H_2O$ — Show all ions present in solution
2. $Na^+$ and $Cl^-$ appear on both sides — These ions don't actually react - they just watch
3. $H^+$(aq) + $OH^-$(aq) → $H_2O$(l) — This shows what actually happens in any neutralisation

Answer: $H^+$(aq) + $OH^-$(aq) → $H_2O$(l)

Half equations split reactions into two parts: oxidation (losing electrons) and reduction (gaining electrons). They're essential for understanding electrolysis and redox reactions.

6. Exam Success - Getting Full Marks

Examiners are looking for specific things when you write equations. Miss these details and you'll lose marks even if you understand the chemistry perfectly.

Common exam mistakes include: writing H instead of $H_2$ for hydrogen gas, forgetting to balance equations, mixing up reactants and products, and including spectator ions in ionic equations. Practice these skills with past paper questions.

Master these equation-writing skills and you'll have a powerful tool for understanding all of chemistry. Every reaction you meet - from photosynthesis to industrial processes - can be represented clearly using these methods.