R1.1 Measuring enthalpy changes
Practice exam-style IB Chemistry questions for Measuring enthalpy changes, aligned with the syllabus and grouped by topic.
Question 1
A reacting mixture in a beaker is defined as the system. What are the surroundings?
A.
Only the air above the reacting mixture
B.
Only the thermometer placed in the beaker
C.
Everything outside the reacting mixture that can exchange energy with it
D.
The reactants before mixing and the products after mixing
Question 2
A salt dissolves in water and the temperature of the solution decreases. What is the sign of the enthalpy change for the dissolving process?
A.
Positive, because energy is transferred from the surroundings to the system
B.
Negative, because the solution loses thermal energy
C.
Negative, because all dissolving processes release heat
D.
Positive, because energy is transferred from the system to the surroundings
Question 3
What is the best description of heat?
A.
Energy transferred because of a temperature difference
B.
The average kinetic energy of particles in a sample
C.
A state function measured in degrees Celsius
D.
The total chemical potential energy stored in bonds
Question 4
In an exothermic reaction, what is correct about the relative potential energies of reactants and products?
A.
Products have lower potential energy than reactants.
B.
Products have no potential energy after the reaction.
C.
Reactants and products must have equal potential energy.
D.
Reactants have lower potential energy than products.
Question 5
Distinguish between heat and temperature.
[1]
Question 6
A student adds ammonium nitrate to water and observes that the beaker becomes cold.
State the type of process occurring.
[1]
Outline the direction of energy transfer between system and surroundings.
[1]
Question 7
50.0 g of water is heated from 21.0 °C to 27.0 °C. What is the heat transferred to the water? Use c = 4.18 J g⁻¹ K⁻¹.
A.
12.5 kJ
B.
1.25 kJ
C.
5640 kJ
D.
5.64 kJ
Question 8
What labels should be used on the axes of an energy profile for a chemical reaction?
A.
x-axis: temperature; y-axis: enthalpy change
B.
x-axis: reaction coordinate; y-axis: potential energy
C.
x-axis: potential energy; y-axis: reaction coordinate
D.
x-axis: time; y-axis: temperature
Question 9
A reaction releases 2.40 kJ of heat to the solution when 0.0200 mol of limiting reactant reacts. What is ΔH for the reaction?
A.
−120 kJ mol⁻¹
B.
+120 kJ mol⁻¹
C.
−48.0 kJ mol⁻¹
D.
+48.0 kJ mol⁻¹
Question 10
Two solutions, each of volume 25.0 cm³ and density 1.00 g cm⁻³, are mixed in a polystyrene cup. The temperature rises by 8.0 K. What value of Q should be used for the heat gained by the solution? Use c = 4.18 J g⁻¹ K⁻¹.
A.
1670 kJ
B.
3.34 kJ
C.
1.67 kJ
D.
0.836 kJ
Question 11
What does the symbol ⦵ in ΔH⦵ indicate?
A.
Substances are in standard states under standard conditions.
B.
Only one mole of every reactant must be used experimentally.
C.
The reaction is carried out at constant volume.
D.
The reaction has zero enthalpy change.
Question 12
50.0 cm³ of an aqueous solution of density 1.00 g cm⁻³ increases in temperature by 5.2 K during a reaction. Assume c = 4.18 J g⁻¹ K⁻¹.
Calculate the heat gained by the solution, in kJ.
[1]
State the sign of ΔH for the reaction.
[1]
Question 13
The combustion of methanol is exothermic.
State which is more stable: the reactants or the products.
[1]
Explain your answer using potential energy.
[1]
Question 14
Sketch an energy profile for an endothermic reaction. Include labels for axes, reactants, products and ΔH.
[1]
Question 15
A student investigated the reaction between magnesium and excess hydrochloric acid using a polystyrene cup calorimeter. The temperature was recorded before and after adding magnesium.
Determine the temperature change, ΔT, from the graph.
[1]
Calculate the heat gained by 50.0 g of solution using c = 4.18 J g⁻¹ K⁻¹.
[1]
The amount of magnesium reacted was 0.00206 mol. Calculate ΔH, in kJ mol⁻¹, for the reaction.
[1]
Question 16
The figure shows an energy profile for a reaction.
State whether the reaction is exothermic or endothermic.
[1]
Identify the arrow that represents ΔH.
[1]
Identify the arrow that represents the activation energy for the forward reaction.
[1]
State which is more stable, reactants or products.
[1]
Question 17
A spirit burner experiment gives an enthalpy of combustion for ethanol that is less exothermic than the data-book value. What is the most likely reason?
A.
Some energy heats the air and apparatus instead of the water.
B.
The specific heat capacity of water is zero.
C.
All heat released by combustion is transferred to the water.
D.
Combustion reactions are normally endothermic.
Question 18
A hot reaction mixture cools noticeably while temperature readings are being taken. How does this affect the calculated ΔH for an exothermic reaction if the highest recorded temperature is used without extrapolation?
A.
The magnitude of ΔH is too large because ΔT is underestimated.
B.
The magnitude of ΔH is too small because ΔT is underestimated.
C.
The sign of ΔH changes from negative to positive.
D.
There is no effect because enthalpy is a state function.
Question 19
25.0 cm³ of 1.00 mol dm⁻³ HCl is mixed with 25.0 cm³ of 2.00 mol dm⁻³ NaOH. Which amount should be used as n in ΔH = −Q/n for the neutralization reaction H⁺(aq) + OH⁻(aq) → H₂O
A.
0.0500 mol, the amount of OH⁻ initially present
B.
50.0 mol, the total volume of solution in cm³
C.
0.0250 mol, the amount of H⁺ reacted
D.
0.0750 mol, the total amount of ions mixed
Question 20
0.460 g of propan-1-ol, C₃H₇OH, is burned to heat water. The water absorbs 13.5 kJ. What is the experimental enthalpy change of combustion of propan-1-ol? M(C₃H₇OH) = 60.1 g mol⁻¹.
A.
−1760 kJ mol⁻¹
B.
−0.104 kJ mol⁻¹
C.
+1760 kJ mol⁻¹
D.
−104 kJ mol⁻¹
Question 21
Why is a bomb calorimeter generally more precise than a simple spirit burner and beaker arrangement for measuring combustion energy?
A.
It uses a controlled oxygen atmosphere and a well-insulated heat-absorbing system.
B.
It eliminates the need to know the temperature change.
C.
It prevents any chemical reaction from occurring outside water.
D.
It makes combustion endothermic so temperature changes are easier to measure.
Question 22
An energy profile has reactants at 120 kJ mol⁻¹, products at 45 kJ mol⁻¹ and a peak at 180 kJ mol⁻¹. What are ΔH and the activation energy for the forward reaction?
A.
ΔH = +75 kJ mol⁻¹; Eₐ = 60 kJ mol⁻¹
B.
ΔH = +60 kJ mol⁻¹; Eₐ = −75 kJ mol⁻¹
C.
ΔH = −75 kJ mol⁻¹; Eₐ = 135 kJ mol⁻¹
D.
ΔH = −75 kJ mol⁻¹; Eₐ = 60 kJ mol⁻¹
Question 23
A student calculates ΔH for an acid–base reaction using the volume of acid only to estimate the mass of solution, although equal volumes of acid and alkali were mixed. What is the likely effect on the calculated magnitude of ΔH?
A.
It changes sign because mass is used in the calculation.
B.
It is unaffected because the same temperature change is used.
C.
It is too large because the mass of solution is overestimated.
D.
It is too small because the heat transferred to the solution is underestimated.
Question 24
25.0 cm³ of 1.00 mol dm⁻³ HCl is mixed with excess sodium hydroxide solution in a polystyrene cup. The solution gains 1.42 kJ of heat.
Calculate the amount, in mol, of HCl reacted.
[1]
Calculate ΔH, in kJ mol⁻¹, for the neutralization per mole of HCl.
[1]
State one assumption made in this calculation.
[1]
Question 25
Explain, in terms of bond breaking and bond formation, why a reaction may be exothermic overall.
[1]
Question 26
A simple combustion experiment uses a spirit burner to heat water in a beaker.
Suggest two reasons why the experimental enthalpy of combustion is less exothermic than the data-book value.
[1]
Suggest one safety precaution for this experiment.
[1]
Question 27
A student burns 0.315 g of ethanol, C₂H₅OH, to heat 200.0 g of water. The water temperature increases by 18.6 K. Use c = 4.18 J g⁻¹ K⁻¹ and M(C₂H₅OH) = 46.1 g mol⁻¹.
Calculate the heat gained by the water, in kJ.
[1]
Calculate the amount of ethanol burned.
[1]
Calculate the experimental enthalpy change of combustion, in kJ mol⁻¹.
[1]
Question 28
40.0 cm³ of 0.500 mol dm⁻³ sulfuric acid, H₂SO₄(aq), is mixed with 50.0 cm³ of 1.00 mol dm⁻³ sodium hydroxide. The reaction is:
H₂SO₄(aq) + 2NaOH(aq) → Na₂SO₄(aq) + 2H₂O
Determine the limiting reactant.
[1]
State the amount, in mol, to use for calculating ΔH per mole of H₂SO₄ reacted.
[1]
State why using the total moles of acid and base would be incorrect.
[1]
Question 29
Nitrogen and oxygen can react to form nitrogen monoxide in an endothermic reaction.
State the sign of ΔH.
[1]
Explain why the reaction is endothermic in terms of the bonding in nitrogen and energy changes.
[1]
Question 30
A bomb calorimeter is used to determine the energy released by a food sample.
Describe two features of a bomb calorimeter that improve reliability compared with a simple open flame calorimeter.
[1]
State the measurement that is converted into energy released.
[1]
Question 31
A thermometric titration was carried out by adding sodium hydroxide solution to 25.0 cm³ of hydrochloric acid. The temperature after each addition was recorded.
Identify the volume of sodium hydroxide at which neutralization is complete.
[1]
State how the graph shows that neutralization is exothermic.
[1]
Explain why the temperature decreases after the equivalence volume.
[1]
Suggest one source of systematic error in this experiment.
[1]
Question 32
Four trials were carried out for the same exothermic reaction in a polystyrene cup. The table gives the temperature changes obtained.
| Trial | Temperature change / K |
|---|---|
| 1 | 12.4 |
| 2 | 12.6 |
| 3 | 8.0 |
| 4 | 12.5 |
Identify the anomalous trial.
[1]
Calculate the mean temperature change excluding the anomalous trial.
[1]
Explain why repeated trials improve the reliability of the result.
[1]
State why the experimental value of |ΔH| may still be smaller than the true value.
[1]
Question 33
A student burned equal masses of two alcohols separately to heat equal masses of water. The graph compares the temperature rise of the water with time.
Identify which alcohol transfers energy to the water at the greater average rate during the first minute.
[1]
State which trial gives the larger value of Q for the water.
[1]
Suggest one reason why equal masses of different alcohols may not correspond to equal amounts in mol.
[1]
Suggest one experimental factor, other than fuel identity, that could affect the temperature rise.
[1]
Question 34
A temperature probe records data before and after two solutions are mixed in a cup calorimeter. The highest measured temperature is lower than the temperature obtained by extrapolating the cooling curve back to the mixing time.
Explain why extrapolation gives a more reliable value of ΔT.
[1]
State the effect on the magnitude of ΔH for an exothermic reaction if extrapolation is not used.
[1]
Identify whether this error is mainly random or systematic.
[1]
Question 35
A solid dissolves in 100.0 g of water and the temperature decreases from 22.4 °C to 18.1 °C. Assume the solution has the same specific heat capacity as water, c = 4.18 J g⁻¹ K⁻¹. The amount of solid dissolved is 0.0500 mol.
Calculate Q for the water, in kJ, including its sign.
[1]
Calculate ΔH for the dissolving process, in kJ mol⁻¹.
[1]
Question 36
An energy profile has reactants at 80 kJ mol⁻¹, products at 140 kJ mol⁻¹ and the transition peak at 210 kJ mol⁻¹.
Determine ΔH for the forward reaction.
[1]
Determine the activation energy for the forward reaction.
[1]
State whether the forward reaction is endothermic or exothermic and justify your answer.
[1]
Question 37
A student calculates an enthalpy change by assuming the density and specific heat capacity of a salt solution are the same as those of pure water.
State why this assumption is made in school calorimetry.
[1]
Suggest one way this assumption could affect the calculated value of Q.
[1]
Suggest two improvements to reduce uncertainty or error in the experiment.
[1]
Question 38
A data logger was used in a cup calorimetry experiment. The graph shows the temperature before mixing, after mixing, and an extrapolated cooling line.
Read the corrected maximum temperature from the extrapolation.
[1]
Calculate the corrected ΔT using the initial temperature.
[1]
Explain why using the recorded maximum instead of the extrapolated value affects ΔH.
[1]
State whether the correction makes ΔH for an exothermic reaction more negative or less negative.
[1]
Question 39
A hydrocarbon fuel was burned in a simple calorimeter. The table shows the mass of burner before and after heating, the mass of water and the water temperature change.
| Initial burner mass / g | Final burner mass / g | Mass of water / g | Water ΔT / K | Fuel molar mass / g mol⁻¹ |
|---|---|---|---|---|
| 52.48 | 51.76 | 150.0 | 22.4 | 86.0 |
Calculate the mass of fuel burned.
[1]
Calculate the heat gained by the water using c = 4.18 J g⁻¹ K⁻¹.
[1]
Calculate the amount of fuel burned using the molar mass given in the table.
[1]
Calculate the experimental enthalpy change of combustion.
[1]
Question 40
A student mixed different volumes of 1.00 mol dm⁻³ hydrochloric acid with 25.0 cm³ of 1.00 mol dm⁻³ sodium hydroxide and recorded the maximum temperature change.
Identify the volume of acid that gives the largest temperature change.
[1]
Explain why the temperature change increases before this volume.
[1]
Explain why the temperature change decreases after this volume.
[1]
Determine which reactant is limiting when the acid volume is much less than the volume identified in (a).
[1]
Question 41
A student determines the enthalpy change for the reaction between zinc and copper(II) sulfate solution using a polystyrene cup calorimeter.
Outline how the temperature change should be measured to obtain a reliable value of ΔT.
[1]
Explain how the data are used to calculate ΔH and why the experimental value may differ from the accepted value.
[1]
Question 42
Define exothermic and endothermic reactions in terms of energy transfer between system and surroundings.
[1]
Compare exothermic and endothermic reactions in terms of temperature observations, sign of ΔH, bond energy changes and relative stability of reactants and products.
[1]
Question 43
Describe the essential features of an energy profile for an exothermic reaction.
[1]
Explain how the energy profile shows activation energy, enthalpy change and relative stability.
[1]
Question 44
Two methods were used to measure the enthalpy change of the same exothermic reaction: a polystyrene cup calorimeter and a well-insulated calorimeter with a known heat capacity. The table compares the results.
| Method | ΔT / K | Solution mass / g | Calorimeter C / J K⁻¹ | Energy q / kJ | ΔH / kJ mol⁻¹ |
|---|---|---|---|---|---|
| Polystyrene cup | 5.2 | 100.0 | not included | 2.17 | -43.4 |
| Insulated calorimeter | 5.6 | 100.0 | 85 | 2.82 | -56.4 |
Identify which method gives the larger magnitude of ΔH.
[1]
Suggest why this method gives a larger magnitude.
[1]
Explain why including the calorimeter heat capacity changes the calculated energy transferred.
[1]
State one limitation that may remain even with the improved calorimeter.
[1]
Question 45
The graph shows two possible energy profiles, X and Y, for reactions starting from the same reactants.
Identify which profile represents the more exothermic reaction.
[1]
Explain your answer to (a).
[1]
Identify which profile has the greater activation energy.
[1]
State which products, X or Y, are more stable.
[1]
Question 46
A simple spirit burner is used to determine the enthalpy change of combustion of butan-1-ol by heating water in a metal can.
State the measurements needed to calculate the experimental enthalpy change of combustion.
[1]
Discuss the main limitations of this method and how the design could be improved.
[1]
Question 47
A student measures the enthalpy change of neutralization by mixing 50.0 cm³ of acid with 50.0 cm³ of alkali in a polystyrene cup. The solutions are initially at slightly different temperatures.
Outline how an appropriate initial temperature could be chosen for calculating ΔT.
[1]
Evaluate the assumptions and sources of error in using the temperature change to determine ΔH.
[1]
Question 48
A research laboratory uses a bomb calorimeter to measure the energy released by a food sample, while a school laboratory uses a burning food sample below a boiling tube of water.
State two measurements common to both methods.
[1]
Discuss why the bomb calorimeter gives a more accurate value and identify limitations that may remain.
[1]
Question 49
Two reactions have energy profiles with the same activation energy but different enthalpy changes. Reaction A is exothermic and reaction B is endothermic.
Describe how the positions of reactants, products and peak compare for the two profiles.
[1]
Compare and contrast what can be deduced from the profiles about ΔH, activation energy and stability.
[1]
Question 50
A student obtains the following information for a dissolving process: 6.00 g of solid X, molar mass 120.0 g mol⁻¹, is added to 100.0 g of water. The temperature changes from 20.0 °C to 16.5 °C. Assume c = 4.18 J g⁻¹ K⁻¹ for the solution.
Calculate the enthalpy change of dissolving, in kJ mol⁻¹.
[1]
Evaluate two assumptions in the calculation and explain how one experimental improvement could increase confidence in the result.
[1]
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