Energy Changes

Chemistry obeys the law of conservation of energy — the total amount of energy in the universe stays the same. In a chemical reaction, energy is transferred between two places:

1. The chemicals (the reactants and products themselves).
2. The surroundings (everything else — the solution, the apparatus, the lab air).

You can detect the transfer because the temperature of the surroundings changes. Specifically, if the chemicals lose energy, the surroundings gain it (and warm up), and vice versa.

Key idea. A thermometer in the reaction mixture tells you whether energy is being released by, or absorbed by, the reaction.

Direction of energy flow	Effect on surroundings	Type of reaction
From chemicals → to surroundings	Temperature rises	Exothermic
From surroundings → to chemicals	Temperature falls	Endothermic

This is purely about energy, not mass. The mass of reactants always equals the mass of products (the law of conservation of mass — covered in topic 5.3). Energy and mass are separate accountings.

Exothermic reaction. The reaction releases energy to its surroundings. The mixture (and any nearby thermometer) gets hotter. Common types:

• Combustion — burning of any fuel (methane, propane, candles, petrol, wood). Releases a lot of energy as heat and light.
• Neutralisation — acid + alkali → salt + water. A small but reliable temperature rise (typically 5–10 °C in a school lab).
• Many oxidation reactions — including rusting of iron (slow exothermic), the thermite reaction (very exothermic).
• Displacement reactions between a reactive metal and a metal-salt solution (e.g. Zn + CuSO₄ warms the test tube).
• Respiration in living cells (glucose + oxygen → CO₂ + water — this is the biology-equivalent of combustion).

Endothermic reaction. The reaction absorbs energy from its surroundings. The mixture (and the thermometer) gets colder. Common types:

• Thermal decomposition — heating a solid until it breaks down (e.g. CaCO₃ → CaO + CO₂; NaHCO₃ → Na₂CO₃ + H₂O + CO₂). The energy 'taken in' is the heat you supply, but in any environment the decomposition continues to cool until the input matches the demand.
• The reaction of citric acid with sodium hydrogencarbonate — this is the AQA-specified endothermic example you should learn for the lab. The mixture cools by 1–4 °C.
• Photosynthesis — plants absorb light energy to convert CO₂ + H₂O → glucose + O₂. Endothermic with respect to chemical-energy storage.

Memory trick. Exothermic = exits the chemicals (heat exits → surroundings warm). Endothermic = enters the chemicals (heat enters → surroundings cool).

British exam boards love asking about consumer products that exploit chemical energy changes. Here are the standard examples for AQA:

Hand warmers (exothermic). Most disposable hand warmers contain iron filings, salt, water and air-permeable pouch. When activated, the iron is oxidised:

4Fe + 3O₂ → 2Fe₂O₃

The reaction releases enough heat to keep your hands warm for hours. Once the iron is fully rusted, the warmer is finished (cannot be recharged).

Reusable hand warmers (also exothermic, but reversible). A pouch of saturated sodium ethanoate solution. Click the metal disc — crystals start to form, releasing energy. Boil the pouch in water and the crystals redissolve, ready to use again.

Self-heating cans (exothermic). A separate compartment under the drink contains calcium oxide and water:

CaO + H₂O → Ca(OH)₂

This very exothermic reaction warms the drink without a heat source.

Sports cold packs (endothermic). The pack contains a sealed inner pouch of water inside ammonium nitrate crystals. Punching the pack breaks the inner pouch, dissolving the salt:

NH₄NO₃ + H₂O → NH₄⁺(aq) + NO₃⁻(aq) (endothermic dissolution)

The temperature drops to around 0 °C, useful for reducing swelling after a sports injury.

Instant cold packs (alternative). Some use ammonium chloride or urea — same idea, endothermic dissolution lowers the temperature.

Exam tip. The mark scheme rewards you for naming a specific product and saying whether the reaction is exothermic or endothermic and what the user gets out of it (warmth, cooling, etc.).

RP4 asks you to investigate the variables that affect the temperature change of a reacting solution. The textbook version uses hydrochloric acid + sodium hydroxide (a strongly exothermic neutralisation).

Apparatus:

• Polystyrene cup (insulator — reduces heat loss to surroundings).
• Lid with a small hole for the thermometer.
• Beaker to hold the cup upright.
• 0.5 mol/dm³ HCl(aq) and 0.5 mol/dm³ NaOH(aq).
• 25-cm³ measuring cylinders.
• Thermometer (0.5 °C precision or better).
• Stopclock and a way to stir.

Method (independent variable = volume of NaOH).

1. Measure 30 cm³ of 0.5 mol/dm³ HCl into the polystyrene cup. Place the cup in a beaker for support.
2. Measure the initial temperature of the acid. Record.
3. Measure 5 cm³ of 0.5 mol/dm³ NaOH. Add to the acid, stir gently with the thermometer, and record the maximum temperature reached.
4. Repeat the experiment, but each time add a different volume of NaOH (10, 15, 20, 25, 30, 35, 40 cm³). Keep the volume of HCl constant.
5. Calculate the temperature change ΔT = T_max − T_initial for each run.
6. Plot a graph of ΔT (y-axis) vs volume of NaOH added (x-axis).

Expected result. ΔT rises as more NaOH is added (more neutralisation happens) up to a maximum at the stoichiometric point (equal moles of acid and alkali). Beyond that, ΔT falls because excess NaOH cools the mixture without giving extra heat.

Why use a polystyrene cup?

• Polystyrene is a good thermal insulator → minimises heat loss to the surroundings.
• A lid further reduces heat loss by convection and evaporation.
• A beaker around it gives extra stability.

Sources of error:

• Heat loss to the lid/thermometer/air despite the insulation.
• Imprecise thermometer reading (±0.5 °C is typical).
• The stoichiometric maximum may sit between two measured volumes, so the peak ΔT is slightly under-recorded.

Examiners often give you an equation or scenario and ask you to classify the energy change. Here's how to think about each:

Example 1 — Combustion of methane.

CH₄ + 2O₂ → CO₂ + 2H₂O ; ΔT ↑ in the lab. Classification: exothermic (combustion always releases energy to the surroundings).

Example 2 — Heating limestone.

CaCO₃ → CaO + CO₂ ; requires constant heating from a Bunsen burner. Classification: endothermic (thermal decomposition needs energy from the surroundings; once you stop heating, the reaction stops).

Example 3 — Magnesium ribbon + dilute HCl.

Mg + 2HCl → MgCl₂ + H₂ ; temperature of solution rises by ~10 °C. Classification: exothermic (displacement of hydrogen by a more reactive metal).

Example 4 — Citric acid + sodium hydrogencarbonate (sherbet).

Citric acid + 3NaHCO₃ → Sodium citrate + 3CO₂ + 3H₂O ; temperature falls. Classification: endothermic (the AQA-named example of an endothermic reaction).

Example 5 — Photosynthesis.

6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂ ; absorbs sunlight. Classification: endothermic (light energy stored as chemical energy in glucose).

Exam wording. AQA mark schemes give marks for:

1. Stating exothermic or endothermic (1 mark).
2. Describing the energy transfer ("energy is given out to / taken in from the surroundings") (1 mark).
3. Naming the observation in the lab (temperature rises / falls) (1 mark).

Three marks for a confidently written one-sentence answer.

A reaction profile (sometimes called an energy-level diagram) shows how the energy of the chemicals changes as a reaction proceeds. The axes are:

• x-axis: 'reaction progress' or 'reaction coordinate' (not a number — just left-to-right as the reaction happens).
• y-axis: energy of the chemicals.

You always see three features:

1. A flat line on the LEFT at the energy of the reactants.
2. A peak in the middle — the activated complex at the top of the activation-energy hump.
3. A flat line on the RIGHT at the energy of the products.

A curve connects the three.

Two simple shapes to memorise:

• Exothermic — products LOWER than reactants. The downhill slope to the right of the peak corresponds to energy released to the surroundings.
• Endothermic — products HIGHER than reactants. The uphill rise to the right of the peak corresponds to energy absorbed from the surroundings.

Activation energy (E_a) is shown as a vertical arrow from the reactants level up to the peak. Even an exothermic reaction needs some E_a — particles must collide with enough energy to break the old bonds.

Overall energy change (ΔE) is shown as a vertical arrow from the reactants level to the products level. Negative for exothermic (products lower); positive for endothermic (products higher).

For an exothermic reaction:

• Reactants sit at a higher energy level.
• Products sit at a lower energy level.
• A peak in between marks the activation energy.
• The drop from reactants to products represents the energy released to the surroundings.

Worked example — combustion of methane. CH₄ + 2O₂ → CO₂ + 2H₂O is strongly exothermic. The reactants (methane + oxygen molecules) sit high up. They must first absorb enough energy to break the C–H and O=O bonds (the activation energy hump). At the peak the activated complex exists very briefly. Then new bonds form in CO₂ and H₂O, releasing energy and dropping the chemicals to a lower energy level. The net release of energy is what warms the surroundings.

Where does the released energy come from? The new bonds in the products are STRONGER (release more energy when they form) than the old bonds in the reactants (which absorbed energy when they broke). The surplus is what escapes to the surroundings as heat.

Equation for the overall energy change:

ΔE = E(products) − E(reactants)

For exothermic: products < reactants → ΔE is negative.

Labels you must include in your exam diagram.

1. Reactants line (left).
2. Products line (right) — LOWER than reactants.
3. The curve connecting them, going up to a peak.
4. Vertical arrow showing activation energy (from reactants up to peak).
5. Vertical arrow showing overall energy change (from reactants down to products) — labelled as energy 'given out' or 'released'.

If any of those five labels are missing, AQA mark schemes penalise you per item.

For an endothermic reaction:

• Reactants sit at a lower energy level.
• Products sit at a higher energy level.
• The peak (activation energy) is still above both reactants and products.
• The rise from reactants to products represents the energy absorbed from the surroundings.

Worked example — thermal decomposition of calcium carbonate. CaCO₃ → CaO + CO₂ requires constant heating. The reactant (CaCO₃) starts low. To form CaO + CO₂ the chemicals must first reach the peak (E_a) by absorbing energy from the surroundings, then settle at a HIGHER energy level (products). The net absorption of energy is what cools the surroundings (or, in the lab, requires constant heat input).

Where does the absorbed energy go? The new bonds in the products are WEAKER (release less energy when they form) than the old bonds in the reactants (which required more energy to break). The deficit is supplied by the surroundings.

ΔE = E(products) − E(reactants) is positive for endothermic.

Labels for the diagram (same as exothermic):

1. Reactants line (LOW).
2. Products line (HIGH).
3. Peak above both.
4. E_a arrow (reactants up to peak).
5. ΔE arrow (reactants up to products) — labelled 'energy taken in' or 'absorbed'.

Activation energy (E_a) is the minimum energy that colliding particles must possess for the collision to be 'successful' — i.e. lead to a reaction.

Why is there an activation energy at all? Before new bonds can form in the products, the OLD bonds in the reactants must be broken. Breaking bonds requires energy. So the chemicals have to climb a hill before they can roll down it.

Practical evidence.

• A mixture of hydrogen and oxygen sits stable at room temperature — even though the reaction is hugely exothermic. A spark provides the energy needed to reach E_a, then the released energy keeps the reaction going (and you get an explosion).
• Magnesium ribbon does not catch fire just lying in air. You need to heat it with a Bunsen first.
• Petrol does not spontaneously combust in your fuel tank — it needs the spark from a spark plug.

Why higher temperatures speed up reactions (link to rates topic 5.6). Heating gives more particles enough energy to clear E_a. So more collisions are 'successful' per second → faster rate.

Why some reactions are slow even though exothermic. A LARGE E_a means few particles have enough energy at room temperature → slow rate. Rusting is exothermic but takes weeks because its E_a is high enough that only a tiny fraction of collisions succeed at any moment.

A catalyst's job (next section). A catalyst lowers E_a — more particles can now clear the hump, so the reaction goes faster.

A catalyst is a substance that speeds up a reaction without being used up. It does this by providing an alternative reaction pathway with a lower activation energy.

On the profile diagram:

• The reactants line and the products line are in the same position.
• The peak is lower when the catalyst is present.
• The ΔE (vertical gap between reactants and products) is identical to the uncatalysed reaction.

Examiner-ready way of saying it:

"The catalyst provides an alternative reaction pathway with a lower activation energy. More particles now have enough energy to react successfully when they collide. The catalyst is not used up and the overall energy change is unchanged."

Why this matters. Catalysts let industry run reactions:

• At lower temperatures (saving fuel costs).
• Faster (higher production rate).
• Without changing the thermodynamics — the same products, same energy balance.

Examples on the AQA spec:

• Iron in the Haber process (N₂ + 3H₂ ⇌ 2NH₃).
• Platinum in catalytic converters (oxidising CO to CO₂).
• Vanadium(V) oxide in the Contact process (SO₂ → SO₃).
• Enzymes in living cells (biological catalysts).

Diagram-drawing tip. When asked to draw a profile WITH and WITHOUT a catalyst, draw the same reactants and products and TWO curves over the same start/end levels — one with a tall peak (uncatalysed) and a dashed lower peak (catalysed). Label both peaks 'E_a' with subscripts 'no catalyst' and 'with catalyst'. The ΔE arrow only needs to be drawn once.

In any chemical reaction, the bonds in the reactants are broken and new bonds in the products are formed.

Breaking a bond requires energy — you have to pull the two atoms apart against the electrostatic attraction holding them together. So:

Bond breaking is ENDOTHERMIC (energy goes IN, taken from the surroundings).

Making a bond releases energy — two atoms snap together, releasing the energy that was the potential between them. So:

Bond making is EXOTHERMIC (energy comes OUT, given to the surroundings).

The amount of energy needed/released for each specific bond is called its bond energy (sometimes "bond dissociation energy"). For example:

Bond	Energy (kJ/mol)
H–H	436
Cl–Cl	243
H–Cl	432
C–H	413
C=O	805
O=O	498
O–H	463

These numbers are positive (energy to break). When the same bond is FORMED in the products, the same number of kilojoules is RELEASED.

Why a reaction is overall exothermic or endothermic: it depends on whether the energy released in making the new bonds is MORE or LESS than the energy used to break the old bonds.

The overall energy change of a reaction is calculated by comparing the two totals.

ΔE = (Total energy to break reactant bonds) − (Total energy to make product bonds)

The sign of ΔE tells you the nature of the reaction:

Sign of ΔE	Meaning
Negative (−)	Exothermic — products are lower in energy than reactants
Positive (+)	Endothermic — products are higher in energy than reactants

Step-by-step method:

1. Write a balanced equation with displayed (full structural) formulae if possible — so you can see and count every bond.
2. List every bond broken in the reactants. Multiply each bond energy by the number of that bond. Sum to get total IN.
3. List every bond made in the products. Multiply each bond energy by the number of that bond. Sum to get total OUT.
4. Subtract: ΔE = IN − OUT.

Worked example: H₂ + Cl₂ → 2HCl

Bond energies (kJ/mol): H–H = 436, Cl–Cl = 243, H–Cl = 432.

Bonds broken (reactants):

1 × H–H + 1 × Cl–Cl = 436 + 243 = 679 kJ/mol (in)

Bonds made (products):

2 × H–Cl = 2 × 432 = 864 kJ/mol (out)

ΔE = 679 − 864 = −185 kJ/mol → exothermic.

The reaction releases 185 kJ per mol of H₂ reacted.

Example 1 — Combustion of methane:

CH₄ + 2O₂ → CO₂ + 2H₂O

Use displayed formulae:

H–C(H)(H)–H + 2 O=O → O=C=O + 2 H–O–H

Bonds broken (reactants):

• 4 × C–H = 4 × 413 = 1652
• 2 × O=O = 2 × 498 = 996
• Total in = 2648 kJ/mol

Bonds made (products):

• 2 × C=O = 2 × 805 = 1610
• 4 × O–H = 4 × 463 = 1852
• Total out = 3462 kJ/mol

ΔE = 2648 − 3462 = −814 kJ/mol → strongly exothermic (which is why methane is a fuel!).

Example 2 — Hydrogen iodide decomposition:

H₂ + I₂ → 2HI

Bond energies: H–H = 436, I–I = 151, H–I = 298.

Broken: 436 + 151 = 587. Made: 2 × 298 = 596.

ΔE = 587 − 596 = −9 kJ/mol → very slightly exothermic.

The small value reflects that this reaction is nearly thermo-neutral (and reversible — which is why HI dissociates back to H₂ + I₂ on heating).

Example 3 — Decomposition of ammonia (endothermic):

2NH₃ → N₂ + 3H₂

Bond energies: N–H = 391, N≡N = 945, H–H = 436.

Broken: 6 × N–H = 2346. Made: N≡N + 3 × H–H = 945 + 1308 = 2253.

ΔE = 2346 − 2253 = +93 kJ/mol → endothermic.

(In practice ammonia decomposition needs heat or a catalyst because energy in > energy out — consistent with our positive ΔE.)

Once you have ΔE, the interpretation is mechanical:

ΔE	Reaction type	Energy diagram
Negative	Exothermic	Products lower than reactants
Positive	Endothermic	Products higher than reactants
≈ 0	Thermo-neutral	Products at almost the same level

Magnitude:

• |ΔE| < 50 kJ/mol → small energy change; reaction may be slow or reversible.
• |ΔE| between 50 and 200 kJ/mol → typical reaction (e.g. neutralisation, displacement).
• |ΔE| > 500 kJ/mol → large (e.g. combustion of fuels).

Why our calculations are approximate: Bond energies are average values taken across many compounds. The actual energy of, say, the C–H bond in CH₄ is slightly different from the C–H bond in C₂H₆. So calculated ΔE is usually within ±5–10% of the experimental value.

Linking back to reaction profiles (5.5.1.2):

• Exothermic profile: reactants high → activation energy peak → products low. ΔE is the drop.
• Endothermic profile: reactants low → activation energy peak (taller) → products high. ΔE is the rise.

You may be asked to label ΔE on a reaction profile in an exam — it's always the vertical gap between reactants and products, NOT the activation energy.

A chemical cell (often called a battery) converts the chemical energy of a redox reaction into electrical energy. The simplest setup is:

1. Two different metals (the electrodes) dipping into an electrolyte (a salt solution that conducts ions).
2. A wire connecting the two electrodes externally.
3. A voltmeter or load (lamp, motor) in the circuit.

Because the metals have different reactivities, one of them loses electrons more readily than the other. Those electrons flow through the external wire from the more reactive metal (the negative electrode) to the less reactive one (the positive electrode), producing a current. Ions move through the electrolyte to balance the charge.

Classic example — the Daniell cell: A zinc strip in zinc sulfate solution and a copper strip in copper sulfate solution, joined by a salt bridge.

• Zinc is more reactive than copper → Zn dissolves to make Zn²⁺, releasing 2 electrons.
• Electrons flow through the wire to the copper electrode.
• Cu²⁺ ions in solution accept the electrons → Cu metal deposits.

Half-equations:

At Zn electrode: Zn → Zn²⁺ + 2e⁻ (oxidation, negative electrode) At Cu electrode: Cu²⁺ + 2e⁻ → Cu (reduction, positive electrode)

Voltage is read on the voltmeter — typically about 1.1 V for the Daniell cell.

Two key factors control the voltage produced by a chemical cell:

1. The reactivity DIFFERENCE between the two metals.

• A bigger difference in reactivity → a bigger voltage.
• For example: Mg (very reactive) paired with Cu (less reactive) → high voltage (~2.7 V).
• Cu paired with Ag (both unreactive, close in the series) → tiny voltage (~0.5 V).

2. The electrolyte used.

• Different electrolytes give different voltages even with the same metals.
• The electrolyte's concentration and the ions it contains both matter.

Metal pair	Approximate voltage (in dilute sulfate)
Mg / Cu	~2.7 V
Zn / Cu	~1.1 V
Fe / Cu	~0.78 V
Cu / Ag	~0.46 V

Important experiment (Required Practical 12 in 8462): Investigate how the voltage produced by a cell depends on the metals used and the electrolyte concentration. Plot voltage vs. position of the metal in the reactivity series — a clear linear trend emerges.

Why this matters: This is the principle behind every battery — from a 1.5 V AA cell to the 12 V car battery. Choose the right metals (or chemistry) and you can engineer the voltage you need.

Non-rechargeable cells (like an AA dry cell) work until the reactants are used up — then they're thrown away (or recycled).

• One of the electrode reactions cannot be reversed by simply running the current backwards.
• Example: zinc–carbon dry cells (1.5 V) used in remote controls and torches.

Rechargeable cells can be 'recharged' by passing an external current backwards through them.

• The current forces the cell reactions in REVERSE — products are converted back to reactants.
• The cell then has stored chemical energy again, ready to discharge.
• Examples: car lead-acid battery (12 V), phone Li-ion battery (3.7 V), laptop NiMH batteries.

Comparison:

Feature	Non-rechargeable	Rechargeable
Cost per use	Higher	Lower (after many cycles)
Lifetime	Single use	Many cycles
Environmental impact	Higher (waste)	Lower per use
Use cases	Smoke alarms, remotes	Phones, cars, EVs

Why some cells can be recharged and others can't: It depends on whether the chemistry inside is REVERSIBLE. In a lead-acid battery, both reactions (Pb + PbO₂ + H₂SO₄ ⇌ 2PbSO₄ + 2H₂O) reverse cleanly when current is applied. In a zinc-carbon cell, the products are physically scattered (gas evaporates) and cannot reform the original chemistry.

A fuel cell uses an external supply of a fuel (such as hydrogen) and oxygen to produce electricity. Unlike a battery, the reactants are continuously supplied from outside — so the cell never needs recharging, only refuelling.

The hydrogen–oxygen fuel cell:

Overall: 2H₂(g) + O₂(g) → 2H₂O(l) ΔE is large and negative → exothermic → electrical energy released.

Construction:

• Hydrogen gas is fed to the negative electrode (anode).
• Oxygen (or air) is fed to the positive electrode (cathode).
• An electrolyte (often a proton-conducting membrane) separates them.
• Electrons flow through the external circuit from anode to cathode, doing useful work.

Half-equations (in acidic electrolyte):

Anode (negative, oxidation): 2H₂ → 4H⁺ + 4e⁻ Cathode (positive, reduction): O₂ + 4H⁺ + 4e⁻ → 2H₂O

Add them together — electrons cancel, H⁺ cancels — and you recover the overall equation 2H₂ + O₂ → 2H₂O.

Advantages:

• Only product is water — no CO₂, no NOₓ, no soot. Pollution-free at the point of use.
• Higher efficiency than burning H₂ in an engine (~60% vs ~25%).
• No noise, no moving parts.
• Continuous operation — refuel in minutes, don't wait hours for a recharge.

Disadvantages:

• Hydrogen is expensive to produce — usually from natural gas (releasing CO₂!) or by electrolysis of water (energy-intensive).
• Hydrogen is explosive and hard to store safely (high-pressure tanks).
• Fuel cells contain expensive platinum catalysts.
• No widespread refuelling infrastructure in the UK.

Compared with rechargeable batteries (e.g. for cars):

Feature	Hydrogen fuel cell	Rechargeable Li-ion battery
Energy density (Wh/kg)	Higher	Lower
Refuel time	Minutes	Hours
Emissions at point of use	Water only	None
Infrastructure (UK 2026)	Sparse	Widespread
Cost	Expensive	Falling fast
Source of energy upstream	Depends on H₂ production	Depends on grid mix

Both are areas of intense research as the UK transitions away from fossil fuels.