Pearson Edexcel International GCSE (9–1) · CHEMISTRY · Unit
Physical Chemistry
442 practice questions
Topics covered
Practice questions
1–10 of 71 · Hard
Question 1
2 marks · Hard
Answer and see the solution →Hydrogen peroxide can be decomposed to form water and oxygen gas. When manganese(IV) oxide is added to hydrogen peroxide solution in a test tube, you notice that the sides of the test tube become noticeably warmer.
Explain, in terms of energy transfer, why this reaction is classified as exothermic.
Question 2
5 marks · Hard
Answer and see the solution →The decomposition of a certain substance is shown by the energy profile diagram below.
(a)[0 marks]Energy profile diagram for a chemical reaction; energies in kJ/mol (b)[2 marks]State whether this reaction is exothermic or endothermic, and explain your answer using the energy values from the diagram.
(c)[2 marks]Calculate the value of ΔH (in kJ/mol) for this reaction using the energy profile diagram.
(d)[1 mark]The activation energy for the reaction is the energy difference between the reactants and the highest point on the diagram. Use the diagram to state the value of the activation energy for this reaction.
Question 3
2 marks · Hard
Answer and see the solution →A student says that 'all reactions which produce heat and raise the temperature are exothermic.' Give one example of a reaction that is exothermic, and explain how the energy changes in this type of reaction.
Question 4
2 marks · Hard
Answer and see the solution →When 3.5 mol of substance X react with an acid, the reaction absorbs 96 kJ of heat energy from the surroundings.
Calculate the enthalpy change, , for this reaction in kJ/mol. State whether it is endothermic or exothermic.
Question 5
5 marks · Hard
Answer and see the solution →A student mixed 6.0 g of potassium nitrate (KNO₃, molar mass = 101 g/mol) in 50.0 g of water in a calorimeter. The temperature of the water fell by 3.2 °C. The specific heat capacity of water is 4.18 J g⁻¹ °C⁻¹.
(a)[2 marks]Calculate the number of moles of KNO₃ dissolved. (2 marks)
(b)[2 marks]Calculate the amount of heat (q, in J) absorbed by the water during the dissolving process. (2 marks)
(c)[1 mark]Calculate the enthalpy change, ΔH, for dissolving KNO₃ in kJ/mol. (1 mark)
Question 6
2 marks · Hard
Answer and see the solution →A reaction releases 212 kJ of heat when 7.0 mol of a reactant are completely used up. Calculate the enthalpy change, , in kJ/mol for this reaction. State the sign of and what it shows about the direction of energy transfer.
Question 7
2 marks · Hard
Answer and see the solution →A student wants to measure the enthalpy change for the reaction between hydrochloric acid and sodium carbonate solution using a polystyrene cup calorimeter.
State two steps the student should follow to ensure the experiment produces reliable results.
Question 8
5 marks · Hard
Answer and see the solution →Copper(II) sulfate crystals dissolve in water in a polystyrene cup calorimeter. A student wants to determine the enthalpy change of the process.
(a)[2 marks]State two measurements the student must take to calculate the enthalpy change. (2 marks)
(b)[2 marks]The student dissolves 10.0 g of copper(II) sulfate ( ) in 50.0 g of water. The temperature rises by 7.4 °C. Calculate the heat absorbed or released by the water. (Specific heat capacity of water = 4.2 J/g°C) Give your answer to three significant figures. (2 marks)
(c)[1 mark]Explain whether the dissolving of copper(II) sulfate in water is exothermic or endothermic, based on your calculation. (1 mark)
Question 9
2 marks · Hard
Answer and see the solution →In a calorimetry experiment, a student measures a temperature decrease during the reaction.
Explain what this temperature change tells you about the type of reaction and the direction of energy transfer.
Question 10
2 marks · Hard
Answer and see the solution →A student carries out a simple calorimetry experiment to measure the enthalpy change of combustion of a fuel. State two specific sources of error or limitations in this type of calorimetry experiment that could cause the measured value to be less exothermic (less negative) than the accepted value.