Chemical Reactions

What Is a Chemical Reaction?

A chemical reaction is a process in which reactants are converted into products through the breaking and forming of chemical bonds. The atoms rearrange — no atoms are created or destroyed — but new substances with new properties are formed.

Signs that a chemical reaction has taken place:

• Colour change (e.g., copper turning green as it oxidises)
• Gas produced — fizzing or bubbling (e.g., acid on metal produces hydrogen)
• Temperature change (exothermic = heats up; endothermic = cools down)
• Precipitate forms — a solid appearing in a clear solution
• Light produced (e.g., burning)
• Irreversible change — the products are difficult to convert back to reactants

Word Equations and Symbol Equations

A word equation names reactants and products:

hydrogen + oxygen → water

A symbol equation uses chemical formulae:

2H2 + O2 → 2H2O

Symbol equations must be balanced — the same number of each type of atom on both sides.

Balancing Equations

The Law of Conservation of Mass states that atoms are neither created nor destroyed in a chemical reaction. Therefore, a balanced equation must have equal numbers of each atom on both sides.

Method:

1. Write the unbalanced equation with correct formulae
2. Count each type of atom on both sides
3. Add coefficients (numbers in front of formulae) to balance — never change the subscripts inside formulae
4. Check your answer by counting again

Example:

• Unbalanced: H2 + O2 → H2O (left: 2H, 2O; right: 2H, 1O — oxygen is unbalanced)
• Balanced: 2H2 + O2 → 2H2O (left: 4H, 2O; right: 4H, 2O — balanced)

Energy Changes in Reactions

All chemical reactions involve energy changes:

Exothermic Reactions

Energy is released to the surroundings (usually as heat). The temperature of the surroundings increases.

Examples:

• Combustion (burning fuels)
• Respiration (glucose + oxygen → CO2 + water + energy)
• Neutralisation (acid + alkali)
• Oxidation reactions (rusting)
• Many displacement reactions

Endothermic Reactions

Energy is absorbed from the surroundings. The temperature of the surroundings decreases.

Examples:

• Photosynthesis (CO2 + water + light → glucose + oxygen)
• Thermal decomposition (heating to break a compound apart)
• Dissolving some salts in water (e.g., ammonium nitrate — used in cold packs)

Types of Chemical Reactions

Reaction type	Description	Example
Combustion	Burning in oxygen; produces CO2 and H2O	Methane + oxygen → carbon dioxide + water
Decomposition	One compound breaks into simpler substances	Calcium carbonate → calcium oxide + CO2
Thermal decomposition	Decomposition caused by heat	CaCO3 heated → CaO + CO2
Oxidation	Gain of oxygen (or loss of electrons)	Iron + oxygen → iron oxide (rust)
Reduction	Loss of oxygen (or gain of electrons)	Iron oxide + carbon → iron + CO2
Neutralisation	Acid + base → salt + water	HCl + NaOH → NaCl + H2O
Precipitation	Two solutions produce an insoluble solid	Lead nitrate + potassium iodide → lead iodide (yellow precipitate)
Displacement	More reactive element replaces less reactive one	Zinc + copper sulfate → zinc sulfate + copper

Rates of Reaction

The rate of reaction measures how fast reactants are converted to products.

Factors that increase the rate:

Factor	Effect	Explanation
Increasing temperature	Faster	Particles move faster; more frequent and energetic collisions
Increasing concentration	Faster	More particles in the same volume; more frequent collisions
Increasing surface area	Faster	More particles exposed; more sites for collisions
Adding a catalyst	Faster	Provides an alternative reaction pathway with lower activation energy

A catalyst speeds up a reaction without being used up. It is not consumed and can be used again. Enzymes are biological catalysts.

Collision Theory

For a reaction to occur, particles must collide with enough energy (the activation energy). Factors that increase the rate of reaction work by increasing the frequency of collisions or the energy of collisions.

Key Terms

Term	Definition
Reactant	Starting material in a chemical reaction
Product	Substance formed in a chemical reaction
Exothermic	Reaction that releases heat energy
Endothermic	Reaction that absorbs heat energy
Catalyst	Substance that speeds up a reaction without being consumed
Activation energy	Minimum energy needed for a reaction to occur
Combustion	Burning in oxygen
Decomposition	Breaking down of one compound into simpler substances
Precipitation	Formation of an insoluble solid from two solutions

Worked Example

Q: Magnesium burns in oxygen. Write a balanced symbol equation for this reaction and state whether it is exothermic or endothermic.

Step 1: Write the word equation: magnesium + oxygen → magnesium oxide

Step 2: Write formulae: Mg + O2 → MgO

Step 3: Balance — left: 1 Mg, 2 O; right: 1 Mg, 1 O. Oxygen is unbalanced. Add coefficients: 2Mg + O2 → 2MgO Check: left: 2 Mg, 2 O; right: 2 Mg, 2 O. Balanced.

Step 4: Burning (combustion) is exothermic — energy is released as heat and light.

Common Mistakes

• Changing the subscript numbers inside a formula when balancing (e.g., writing H3O instead of H2O) — never do this, it changes the substance entirely.
• Confusing exothermic and endothermic — exothermic releases heat (surroundings warm up); endothermic absorbs heat (surroundings cool down).
• Forgetting that a catalyst speeds up reactions but is not used up — it remains unchanged at the end.
• Mixing up physical changes (reversible, no new substance) with chemical changes (usually irreversible, new substance formed).

Tips and Tricks

• Exo = exits, so heat exits the reaction to the surroundings (things get hot). Endo = enters, so heat enters the reaction from the surroundings (things get cold).
• When balancing equations, balance the most complex molecule first, then hydrogen, then oxygen last.
• OILRIG mnemonic: Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons).
• Thermal decomposition is always endothermic — you must supply heat to break the compound apart.