This check needs JavaScript to record answers. The reviewed questions are listed below for study.
Calorimetry to enthalpy change.
A reaction warms 100 g of solution by $3.0\ ^\circ\text{C}$. Use $c=4.18\ \text{J g}^{-1}\text{K}^{-1}$. The amount of limiting reagent reacting is 0.0500 mol. Assume negligible heat loss and negligible heat absorbed by the vessel.
- Part a: Calculate $q_{\text{solution}}$. Report to two significant figures. (1 mark)
- Include the unit J. Scientific notation makes the precision of trailing zeros clear.
- Part b: Calculate $\Delta H$ in kJ mol$^{-1}$ for the reaction. Report to two significant figures. (1 mark) Enter the final answer here and show your method in the following field.
- Include the unit kJ/mol.
- Working for part b (1 mark)
Hess' Law with formation enthalpies.
For $\ce{CH4 + 2O2 -> CO2 + 2H2O(l)}$, use:
$\Delta H_f^\ominus(\ce{CO2})=-394$, $\Delta H_f^\ominus(\ce{H2O(l)})=-286$, $\Delta H_f^\ominus(\ce{CH4})=-75$ kJ mol$^{-1}$.
- Part a: Calculate $\Delta H^\ominus_{\text{reaction}}$. (1 mark) Enter the final answer here and show your method in the following field.
- Include the unit kJ/mol.
- Working for part a (1 mark)
- Part b: State the sign change rule for $\Delta H$ if you reverse a thermochemical equation. (1 mark)
Bond enthalpy estimate.
Estimate $\Delta H$ for $\ce{H2 + Cl2 -> 2HCl}$ using mean bond enthalpies:
$E(\ce{H-H})=436$, $E(\ce{Cl-Cl})=243$, $E(\ce{H-Cl})=431$ kJ mol$^{-1}$.
- Part a: Calculate the estimated $\Delta H$. (1 mark) Enter the final answer here and show your method in the following field.
- Include the unit kJ/mol.
- Working for part a (1 mark)
- Part b: State one reason this is an estimate rather than an exact value. (1 mark)
Gibbs free energy decision.
At 298 K, a reaction has $\Delta H=+35.0$ kJ mol$^{-1}$ and $\Delta S=+150$ J mol$^{-1}$ K$^{-1}$.
- Part a: Calculate $\Delta G$ at 298 K in kJ mol$^{-1}$. (1 mark) Enter the final answer here and show your method in the following field.
- Include the unit kJ/mol.
- Working for part a (1 mark)
- Part b: State whether the reaction is feasible under these conditions. (1 mark)
Lattice energy and Born-Haber logic.
- Part a: State two factors that increase the magnitude of lattice enthalpy. (2 marks)
- Part b: In a Born-Haber cycle, why must all intermediate steps use consistent sign conventions? (1 mark)
Energy-profile interpretation and enthalpy sign.
A reaction profile has reactants at 120 kJ mol$^{-1}$ and products at 65 kJ mol$^{-1}$, with a peak at 185 kJ mol$^{-1}$.
- Part a: Calculate $\Delta H$ for the reaction. (1 mark)
- Include the unit kJ/mol.
- Part b: State whether the reaction is exothermic or endothermic. (1 mark)
- Part c: Calculate the activation energy for the forward reaction. (1 mark)
- Include the unit kJ/mol.
Born–Haber cycle for magnesium chloride.
Data in kJ mol$^{-1}$: standard enthalpy change of formation of $\ce{MgCl2(s)}$ $-641$; enthalpy change of atomisation of Mg $+148$; first and second ionisation energies of Mg $+738$ and $+1451$; Cl–Cl bond energy $+242$; first electron affinity of Cl $-349$.
In this question, lattice energy means the lattice formation enthalpy: one mole of solid forms from its separated gaseous ions, so the value is negative. Build the cycle for one mole of $\ce{MgCl2(s)}$.
Complete the alternative directed route; each row changes only the listed particles| Route step | Directed change | Signed enthalpy / kJ mol⁻¹ |
|---|
| Direct formation | Mg(s) + Cl₂(g) → MgCl₂(s) | −641 |
| Alternative: Mg atomisation | Mg(s) → Mg(g) | +148 |
| Alternative: chlorine atomisation | Supply the equation in part a | Supply the value in part b |
| Alternative: cation formation | Supply the equation in part e | Use the relevant ionisation data |
| Alternative: anion formation | Supply the equation in part c | Supply the value in part d |
| Alternative: lattice formation | Supply the equation in part f | Unknown L |
- Part a: Write the equation, with state symbols, for the step that turns the chlorine in the elements into gaseous chlorine atoms for one mole of $\ce{MgCl2}$. (1 mark)
- Type the arrow as → or ->, and include every state symbol.
- Part b: Give the enthalpy change of the step in part a. (1 mark)
- Include the unit kJ/mol.
- Part c: Write the equation, with state symbols, for the step that forms the gaseous chloride ions from gaseous chlorine atoms. (1 mark)
- Type the arrow as → or ->, and write an electron as e-.
- Part d: Give the enthalpy change of the step in part c. (1 mark)
- Include the unit kJ/mol.
- Part e: Complete the directed cycle step from gaseous magnesium atoms to the required gaseous cations. Write the equation, including electrons and states. (1 mark)
- Write charges as 2+ and an electron as e-.
- Part f: Write the directed lattice-formation step, including every coefficient, charge and state. (1 mark)
- Part g: Construct the signed Hess equation before calculating. Use F for the formation enthalpy, A for Mg atomisation, I and J for its first and second ionisation energies, B for the Cl–Cl bond energy, E for the first electron affinity of Cl, and L for lattice formation enthalpy. (1 mark)
- Use these capital letters, =, +, -, and * for multiplication. E denotes the signed electron-affinity value supplied in the data.
- Part h: Calculate the lattice energy of $\ce{MgCl2}$. (1 mark)
- Include the unit kJ/mol.
- Part i: Give the lattice dissociation enthalpy of magnesium chloride. (1 mark)
- Include the unit kJ/mol.
- Part j: Another student atomises the elements, removes one electron from each Mg atom, then tries to form two chloride ions and combine the ions as MgCl₂. Which required step is missing before the chloride ions can be formed? Use your complete equation and result above to correct the energy account. (1 mark)
Reading a catalysed and an uncatalysed energy profile.
On an energy profile diagram for a reaction, the reactants are at $40\ \text{kJ mol}^{-1}$, the products at $95\ \text{kJ mol}^{-1}$ and the transition state of the uncatalysed reaction at $170\ \text{kJ mol}^{-1}$. A catalyst lowers the highest point of the pathway to $130\ \text{kJ mol}^{-1}$.
- Part a: Calculate $\Delta H$ for the forward reaction, with its sign, in kJ mol$^{-1}$. (1 mark)
- Part b: Calculate the activation energy of the uncatalysed forward reaction, in kJ mol$^{-1}$. (1 mark)
- Part c: Calculate the activation energy of the uncatalysed reverse reaction, in kJ mol$^{-1}$. (1 mark)
- Part d: Calculate the activation energy of the catalysed forward reaction, in kJ mol$^{-1}$. (1 mark)
- Part e: Calculate the activation energy of the catalysed reverse reaction, in kJ mol$^{-1}$. (1 mark)
- Part f: Which statement describes the effect of the catalyst on $\Delta H$? (1 mark)
Enthalpy change of neutralisation by calorimetry.
$50.0\ \text{cm}^3$ of $1.00\ \text{mol dm}^{-3}$ HCl is mixed with $50.0\ \text{cm}^3$ of $1.00\ \text{mol dm}^{-3}$ NaOH in an expanded polystyrene cup. The temperature of the mixture rises by $6.50\ \text{K}$. Assume the solution has density $1.00\ \text{g cm}^{-3}$ and specific heat capacity $4.18\ \text{J g}^{-1}\ \text{K}^{-1}$, and that the cup absorbs no heat.
- Part a: Calculate the heat released to the solution, in J. (1 mark)
- Part b: Calculate the amount of water formed, in mol. (1 mark)
- Part c: Calculate the enthalpy change of neutralisation, with its sign, in kJ mol$^{-1}$ to three significant figures. (1 mark)
- Part d: Some heat is lost to the surroundings. How does this affect the experimental value compared with the true value? (1 mark)
- Part e: Which change to the procedure best corrects for the heat lost during the experiment? (1 mark)
Enthalpy change of methanol synthesis from combustion data.
Find $\Delta H^\ominus$ for $\ce{CO(g) + 2H2(g) -> CH3OH(l)}$ using these standard enthalpy changes of combustion, in kJ mol$^{-1}$: $\ce{CO(g)}$ $-283$; $\ce{H2(g)}$ $-286$; $\ce{CH3OH(l)}$ $-726$.
- Part a: Calculate the total enthalpy change when 1 mol of CO and 2 mol of $\ce{H2}$ are burned completely, in kJ mol$^{-1}$. (1 mark)
- Part b: Calculate $\Delta H^\ominus$ for the formation of methanol from CO and $\ce{H2}$, in kJ mol$^{-1}$. (1 mark)
- Part c: Why is the enthalpy change of combustion of hydrogen multiplied by 2? (1 mark)
- Part d: Why is this enthalpy change found from a cycle rather than measured directly in a calorimeter? (1 mark)
Estimating an enthalpy change from bond energies.
$\ce{CH4(g) + Cl2(g) -> CH3Cl(g) + HCl(g)}$
Average bond energies, in kJ mol$^{-1}$: C–H 413, Cl–Cl 243, C–Cl 338, H–Cl 432.
- Part a: Which bonds must be counted in the calculation? (1 mark)
- Part b: Calculate the total bond energy of the bonds broken, in kJ mol$^{-1}$. (1 mark)
- Part c: Calculate the total bond energy of the bonds formed, in kJ mol$^{-1}$. (1 mark)
- Part d: Estimate $\Delta H$ for the reaction, with its sign, in kJ mol$^{-1}$. (1 mark)
- Part e: Why is this value only an estimate? (1 mark)
Born–Haber cycle for sodium chloride.
Data in kJ mol$^{-1}$: standard enthalpy change of formation of $\ce{NaCl(s)}$ $-411$; enthalpy change of atomisation of sodium $+108$; enthalpy change of atomisation of chlorine, $\ce{1/2Cl2(g) -> Cl(g)}$, $+121$; first ionisation energy of sodium $+496$; first electron affinity of chlorine $-349$.
Ionic radii, in nm: $\ce{Na+}$ 0.095; $\ce{Cl^-}$ 0.181; $\ce{Mg^{2+}}$ 0.065; $\ce{O^{2-}}$ 0.140.
- Part a: Calculate the sum of the enthalpy changes of all the steps in the cycle other than lattice formation, with its sign, in kJ mol$^{-1}$. (1 mark)
- Part b: Calculate the lattice energy of sodium chloride, with its sign, in kJ mol$^{-1}$. (1 mark)
- Part c: Which statement explains the sign of the lattice energy used here? (1 mark)
- Part d: Predict, with reasons, whether the lattice energy of magnesium oxide is more exothermic or less exothermic than that of sodium chloride. (2 marks)
Temperature and the spontaneity of a reaction.
A reaction has $\Delta H^\ominus=-75.0\ \text{kJ mol}^{-1}$ and $\Delta S^\ominus=-150\ \text{J mol}^{-1}\ \text{K}^{-1}$. Assume that $\Delta H^\ominus$ and $\Delta S^\ominus$ do not change with temperature.
- Part a: Calculate $\Delta G^\ominus$ at 298 K, with its sign, in kJ mol$^{-1}$. (1 mark)
- Part b: Calculate the temperature at which $\Delta G^\ominus=0$, in K. (1 mark)
- Part c: In which temperature range is the reaction spontaneous? (1 mark)
- Part d: $\Delta G^\ominus$ is negative at 298 K, but no reaction is observed when the reactants are mixed at 298 K. Which is the best explanation? (1 mark)