GCSE Chemistry Energy changes questions by topic

Classify each reaction as exothermic or endothermic.

1. (a)Burning natural gas in a boiler[1]
2. (b)A sports cold pack activated by squeezing[1]
3. (c)Magnesium dissolving in dilute acid[1]
4. (d)Thermal decomposition of limestone in a kiln[1]

A thermometer in a reaction mixture reads 19.8 °C initially and 23.1 °C at the end.

1. (a)Calculate the temperature change.[1]
2. (b)State whether the reaction is exothermic or endothermic.[1]
3. (c)State the direction of energy transfer and the sign of ΔH.[2]

Sketch or describe the reaction profile of an exothermic reaction.

1. (a)State whether products sit above or below reactants.[1]
2. (b)Name the hump on the graph and what it represents.[2]
3. (c)State the sign of ΔH.[1]

The temperature-change practical is repeated with a glass beaker instead of a polystyrene cup.

1. (a)State and explain the effect on the measured temperature rise.[2]
2. (b)State one improvement to reduce this problem.[1]
3. (c)Name the variable plotted on the x-axis of the class graph.[1]

A hand warmer contains iron filings that rust slowly when exposed to air.

1. (a)Name the type of reaction that warms the pouch.[1]
2. (b)Explain why the reaction transfers heat to the hands.[2]

Classify each everyday reaction as exothermic or endothermic, and give one piece of evidence for two of your answers.

1. (a)Burning natural gas in a hob.[1]
2. (b)Photosynthesis in a leaf.[1]
3. (c)Neutralising acid with alkali.[1]
4. (d)Thermal decomposition of limestone.[1]
5. (e)A sports injury pack that turns cold.[1]
6. (f)Give one observation that would support your answer to (a).[1]
7. (g)Give one observation that would support your answer to (e).[1]

A student dissolves 5 g of ammonium nitrate in 50 cm³ of water in a polystyrene cup. The temperature falls from 20.0 °C to 16.2 °C.

1. (a)State whether the dissolving is exothermic or endothermic, and give the evidence.[2]
2. (b)Calculate the temperature change of the water.[1]
3. (c)Sketch or describe the reaction profile for this change.[2]

A profile shows reactants at 250 kJ/mol, a peak at 390 kJ/mol, and products at 310 kJ/mol.

1. (a)Calculate ΔH and state its sign.[2]
2. (b)Calculate the activation energy.[2]
3. (c)Classify the reaction and explain what a catalyst would do to the profile.[2]

The temperature practical: 30 cm³ of acid at 18.5 °C receives 5 cm³ portions of alkali.

1. (a)Explain why the maximum temperature increases up to a certain volume and then decreases.[3]
2. (b)State what the peak of the graph tells you about the reacting volumes.[2]
3. (c)State one control variable.[1]

Explain, in terms of bond breaking and making, why reactions are exothermic or endothermic.

1. (a)State the energy change when bonds break and when bonds form.[2]
2. (b)Explain why making more (or stronger) bonds than are broken gives a negative ΔH.[2]
3. (c)State what must be true of a reaction with ΔH = +55 kJ/mol in bond terms.[2]

A chemical cell is made with magnesium and copper strips in sodium chloride solution.

1. (a)Explain which metal supplies the electrons to the circuit.[2]
2. (b)Predict whether a zinc–copper cell gives a higher or lower voltage than the magnesium–copper cell.[1]
3. (c)State two ways to increase the current the cell can deliver.[2]

Hydrogen bromide forms in the gas phase: H₂ + Br₂ → 2HBr. Bond energies: H–H = 436, Br–Br = 193, H–Br = 366 kJ/mol.

1. (a)Calculate the total energy needed to break the bonds in the reactants.[1]
2. (b)Calculate the total energy released forming the bonds in the products.[1]
3. (c)Calculate ΔH and state whether the reaction is exothermic or endothermic.[3]
4. (d)Explain why the answer is only approximate.[1]

A student adds 5 cm³ portions of sodium hydroxide solution to 25.0 cm³ of hydrochloric acid in a polystyrene cup, recording the temperature after each addition.

1. (a)Identify the volume of alkali at neutralisation, giving your evidence.[2]
2. (b)Calculate the maximum temperature rise.[1]
3. (c)Explain why the temperature change falls again beyond that volume.[2]
4. (d)State two improvements to the method that would make the maximum temperature more accurate.[1]

Bond energies: N≡N = 945, H–H = 436, N–H = 391 kJ/mol. Reaction: N₂ + 3H₂ → 2NH₃.

1. (a)Calculate the energy needed to break all bonds in the reactants.[2]
2. (b)Calculate the energy released forming the products.[2]
3. (c)Calculate ΔH and classify the reaction.[3]

The hydrogen–oxygen fuel cell. [Chemistry only, Higher]

1. (a)Write the overall equation.[1]
2. (b)Write both half equations in acidic conditions.[4]
3. (c)Give one advantage and one disadvantage of fuel cells over petrol engines.[2]

Two students measure the temperature rise when 25 cm³ of acid neutralises 25 cm³ of alkali. Student A uses a polystyrene cup with a lid and records a rise of 6.8 °C; Student B uses an open glass beaker and records 4.9 °C.

1. (a)Explain the difference between the results.[3]
2. (b)State which result is closer to the true ΔH of neutralisation and why.[2]
3. (c)Calculate the percentage difference between the two results, relative to A’s.[2]
4. (d)Suggest one further improvement to A’s method.[1]

Nitrogen monoxide forms in car engines: N₂ + O₂ → 2NO. Bond energies: N≡N = 945, O=O = 498, N=O = 607 kJ/mol.

1. (a)Explain why breaking the bond in N₂ needs so much energy.[2]
2. (b)Calculate ΔH for the reaction.[3]
3. (c)State what the sign of ΔH shows, and sketch or describe the reaction profile.[2]

A city bus operator must choose between hydrogen fuel cells and rechargeable lithium batteries for a new fleet. Hydrogen is made by electrolysis of water using surplus renewable electricity, but the city has few filling stations.

1. (a)Write the overall equation for the fuel cell reaction and name its only chemical product.[2]
2. (b)Give two advantages and two disadvantages of the hydrogen fuel cell choice.[4]
3. (c)Give one advantage of the battery choice and suggest a judgement about which the operator should pick.[2]