Energy Changes (Exothermic, Endothermic, Profiles)

Quick-Fire Definitions

Exothermic: A reaction that transfers energy to the surroundings (temperature increases).
Endothermic: A reaction that takes in energy from the surroundings (temperature decreases).
Activation energy (Eₐ): The minimum energy needed for a reaction to start - shown by the "hump" on a reaction profile.
Bond energy: The energy needed to break one mole of a particular covalent bond (in kJ/mol).
Catalyst: A substance that speeds up a reaction by providing an alternative pathway with lower activation energy. Not used up.
Fuel cell: An electrochemical cell that converts the chemical energy of a fuel (e.g. Hydrogen) directly into electrical energy.

Exothermic & Endothermic Reactions

Exothermic Reactions

An exothermic reaction transfers energy to the surroundings, usually by heating. The temperature of the surroundings increases.

Combustion: burning fuels (CH₄ + 2O₂ → CO₂ + 2H₂O)
Neutralisation: acid + alkali → salt + water
Oxidation: metals reacting with oxygen

Everyday uses: self-heating cans, hand warmers.

Endothermic Reactions

An endothermic reaction takes in energy from the surroundings, usually by heating. The temperature of the surroundings decreases.

Thermal decomposition: CaCO₃ → CaO + CO₂
Dissolving ammonium nitrate: caused by breaking ionic bonds.
Citric acid + sodium hydrogencarbonate

Everyday uses: instant cold packs for sports injuries.

All reactions involve both bond breaking (endothermic - requires energy) and bond making (exothermic - releases energy). Whether the overall reaction is exothermic or endothermic depends on which process transfers more energy.

Required Practical: Temperature Changes Required Practical

Investigate the variables that affect temperature changes in reacting solutions, e.g. Acid + alkali neutralisation.

Measure a fixed volume of dilute acid (e.g. 25 cm³ of HCl) into a polystyrene cup using a measuring cylinder.
Record the starting temperature of the acid using a thermometer.
Add a measured volume of alkali (e.g. 25 cm³ of NaOH) and stir.
Record the highest temperature reached.
Calculate the temperature change: ΔT = final temperature − initial temperature.
Repeat with different concentrations of alkali to investigate how concentration affects the temperature change.

The polystyrene cup minimises heat loss to the surroundings, providing a more precise measurement of the maximum temperature reached.

A polystyrene cup is used because it is a good insulator - it minimises heat loss to the surroundings, making the temperature measurement more accurate.

Interpreting temperature data

A student adds NaOH to HCl in a polystyrene cup. The temperature rises from 21°C to 28°C. Is this exothermic or endothermic?

Step 1: ΔT = 28 − 21 = +7°C.

Step 2: The temperature of the surroundings (solution) increased.

Step 3: Energy was transferred to the surroundings → this is an exothermic reaction.

A common mistake is confusing the direction of energy transfer. If the temperature goes up, the reaction is exothermic (energy is released into the solution). If the temperature goes down, the reaction is endothermic (energy is taken from the solution).

Reaction Profiles

Reaction profiles are energy diagrams showing the energy of reactants and products.

Reaction profiles showing the energy changes. Notice how adding a catalyst (purple dotted line) lowers the activation energy (Eₐ) without changing the overall energy change (ΔH).

Exothermic Profile

Products are at a lower energy level than reactants. The energy difference is released to the surroundings. The "hump" represents the activation energy (Eₐ) - the minimum energy needed to start the reaction.

Endothermic Profile

Products are at a higher energy level than reactants. The energy difference is taken in from the surroundings.

Effect of Catalysts on Profiles

A catalyst provides an alternative reaction pathway with a lower activation energy. On the profile, the "hump" is smaller, but the overall energy change (ΔH) remains the same.

When drawing reaction profiles, always label: reactants, products, activation energy (Eₐ), and the overall energy change. The x-axis is "Progress of reaction" and the y-axis is "Energy".

Bond Energy Calculations (HT)

Chemical reactions involve breaking old bonds (endothermic) and making new bonds (exothermic).

Overall energy = Energy to break bonds − Energy released making bonds

If more energy is released making bonds than is needed to break bonds → exothermic (negative value).
If more energy is needed to break bonds than is released making bonds → endothermic (positive value).

Worked Example 1: H₂ + Cl₂ → 2HCl (exothermic)

Bond energies: H–H = 436 kJ/mol, Cl–Cl = 242 kJ/mol, H–Cl = 431 kJ/mol

Breaking: 1 × H–H + 1 × Cl–Cl = 436 + 242 = 678 kJ

Making: 2 × H–Cl = 2 × 431 = 862 kJ

Overall: 678 − 862 = −184 kJ/mol (exothermic)

Worked Example 2: N₂ + O₂ → 2NO (endothermic)

Bond energies: N≡N = 941 kJ/mol, O=O = 498 kJ/mol, N=O = 587 kJ/mol

Breaking: 1 × N≡N + 1 × O=O = 941 + 498 = 1439 kJ

Making: 2 × N=O = 2 × 587 = 1174 kJ

Overall: 1439 − 1174 = +265 kJ/mol (endothermic - more energy needed to break bonds than is released making them)

Worked Example 3: Combustion of methane CH₄ + 2O₂ → CO₂ + 2H₂O

Bond energies: C–H = 413, O=O = 498, C=O = 805, O–H = 464 kJ/mol

Breaking: 4 × C–H + 2 × O=O = (4 × 413) + (2 × 498) = 1652 + 996 = 2648 kJ

Making: 2 × C=O + 4 × O–H = (2 × 805) + (4 × 464) = 1610 + 1856 = 3466 kJ

Overall: 2648 − 3466 = −818 kJ/mol (exothermic - combustion always releases energy)

A negative answer = exothermic (energy released). A positive answer = endothermic (energy absorbed). Always show your working clearly in the exam. The methane combustion calculation is one of the most commonly set exam questions.

Chemical & Fuel Cells Chemistry Only

Chemical Cells (Batteries)

A chemical cell produces a voltage when two different metals are dipped into an electrolyte and connected by a wire. Electrons flow from the more reactive metal to the less reactive metal through the external circuit.

The greater the difference in reactivity between the two metals, the greater the voltage.

Predicting voltage from reactivity

Three cells are set up using these metal pairs in the same electrolyte: (A) Zn and Cu, (B) Mg and Cu, (C) Fe and Cu. Which gives the highest voltage?

Step 1: The reactivity order is Mg > Zn > Fe > Cu.

Step 2: The biggest gap in reactivity gives the highest voltage.

Answer: Cell B (Mg and Cu) gives the highest voltage because magnesium and copper are furthest apart in the reactivity series.

Non-rechargeable cells go flat when the reactants are used up. Rechargeable cells can be recharged because the chemical reactions are reversible - applying an external electrical current reverses the reaction.

Hydrogen Fuel Cells

Hydrogen fuel cells react hydrogen with oxygen to produce water and electricity - with no other waste products.

A hydrogen fuel cell. Hydrogen enters the anode, is oxidised, and its electrons travel through the external circuit to the cathode, generating electricity. The hydrogen ions travel through the electrolyte to react with oxygen, forming water.

2H₂ + O₂ → 2H₂O + electrical energy

Fuel Cell Half-Equations Higher Tier

At the anode (negative electrode):

2H₂(g) → 4H⁺(aq) + 4e⁻

Hydrogen is oxidised - it loses electrons.

At the cathode (positive electrode):

O₂(g) + 4H⁺(aq) + 4e⁻ → 2H₂O(l)

Oxygen is reduced - it gains electrons.

In a hydrogen fuel cell, the overall reaction is simply 2H₂ + O₂ → 2H₂O. The cell converts chemical energy directly into electrical energy, which is more efficient than burning the hydrogen as a fuel.

Advantages

Only product is water - no CO₂ greenhouse gas emissions at point of use.
Do not need recharging - continuous supply of fuel.
More efficient than combustion engines at converting chemical energy to useful work.

Disadvantages

Hydrogen is difficult and expensive to store and transport (must be compressed or liquefied).
Hydrogen is most commonly produced from natural gas (fossil fuel) - so not truly zero-carbon unless produced by electrolysis using renewable electricity.
Platinum catalysts are expensive.
Hydrogen fuel infrastructure (refuelling stations) is limited.

A common exam question compares hydrogen fuel cells with rechargeable batteries. Key differences: fuel cells need a continuous fuel supply but don't go flat; batteries store a fixed amount of energy and eventually need recharging or replacing.

Students Also Studied

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Topic 4: Chemical Changes

Reactivity series, extraction of metals, acids, bases and electrolysis.

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Topic 6: Rates of Reaction

Factors affecting rate, collision theory, catalysts and reversible reactions.

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