Define bond enthalpy.

State why bond-breaking is endothermic.

State Hess’s law.

State the condition that must be satisfied for Hess’s law to be applied to two alternative routes.

Write the equation, including state symbols, for the standard enthalpy change of formation of (\ce{CH3OH(l)}).

State the value of (\Delta H_f^\ominus) for (\ce{H2(g)}).

• H2

-> CH3CH2CH3(g)}]

The bonds changed are C=C and H–H broken; C–C and two C–H bonds formed. Use: C=C = 614, H–H = 436, C–C = 347, C–H = 414 kJ mol⁻¹.

Calculate the total energy required to break the bonds.

Calculate (\Delta H) for the reaction.

gives a value different from the experimental value.

State one assumption made when average bond enthalpies are used.

Explain why the calculated value may differ from the experimental value.

State the meaning of a weaker bond in terms of bond enthalpy.

Suggest how this affects the relative rates of nucleophilic substitution of bromoethane and chloroethane.

State the enthalpy change for (\ce{Q -> R}).

Determine the enthalpy change for (\ce{P -> R}).

• 1/2O2

-> CO2(g)}]

(\Delta H_f^\ominus): (\ce{CO(g)}) = −111 kJ mol⁻¹; (\ce{CO2(g)}) = −394 kJ mol⁻¹; (\ce{O2(g)}) = 0 kJ mol⁻¹.

State the expression used with formation data.

Calculate (\Delta H^\ominus).

Write the balanced equation for the standard enthalpy change of combustion of propan-1-ol, (\ce{C3H7OH(l)}), forming (\ce{CO2(g)}) and (\ce{H2O(l)}).

Explain why the coefficient of propan-1-ol must be one.

State the value of (\Delta H_f^\ominus) for graphite.

Explain why (\Delta H_f^\ominus) for diamond is not necessarily the same value.

Describe the general relationship shown by the graph.

Identify the carbon–halogen bond expected to require the least energy to break.

Explain why the use of average bond enthalpies may give an enthalpy change different from an experimental value.

Suggest how the trend helps explain the relative rates of nucleophilic substitution of chloroethane and iodoethane.

Identify the part of the diagram representing energy absorbed in bond-breaking.

Identify the part of the diagram representing energy released in bond-forming.

Determine whether the overall reaction is exothermic or endothermic.

Explain your answer to

in terms of bond energies.

Write an expression for (\Delta H) for (\ce{X -> Z}).

Calculate (\Delta H) for (\ce{X -> Z}).

• HBr

-> CH3CH2Br(g)}]

Average bond enthalpies: C=C = 614, H–Br = 366, C–C = 347, C–H = 414, C–Br = 276 kJ mol⁻¹.

Identify the bonds broken and formed that need to be considered.

Calculate (\Delta H) for the reaction.

• CH3COOH

-> CH3COOCH3

• H2O(l)}]

(\Delta H_c^\ominus): (\ce{CH3OH(l)}) = −726; (\ce{CH3COOH(l)}) = −875; (\ce{CH3COOCH3(l)}) = −1592 kJ mol⁻¹. Water is not combusted.

State the Hess expression using combustion data.

Calculate (\Delta H^\ominus).

(s)}):

[\ce{Na

-> Na(g)}]
[\ce{Na

-> Na+

• e-}]
  [\ce{1/2Cl2

-> Cl(g)}]
[\ce{Cl

• e- -> Cl-(g)}]

Name the first two processes.

Name the last two processes.

State two factors that influence lattice enthalpy.

Explain why (\ce{MgO}) has a larger lattice enthalpy than (\ce{NaCl}).

• Cl2

-> CH3Cl

• HCl(g)}]

Use the table to identify the bonds broken.

Use the table to identify the bonds formed.

Calculate (\Delta H) for the reaction.

State why this value is an estimate.

State the law that justifies using the cycle.

Use the cycle to write an expression for the direct enthalpy change (\Delta H_1).

Determine the unknown enthalpy change labelled (x).

Read from the table the temperature change for experiment 1.

Calculate the heat change, (q), for experiment 1 using (q=mc\Delta T).

Use the cycle to determine the decomposition enthalpy.

Suggest one reason why the experimental value may differ from a data-booklet value.

• 5O2

-> 4NO

• 6H2O(g)}]

State the value of (\Delta H_f^\ominus) for (\ce{O2(g)}).

Use the data to calculate (\Delta H^\ominus) for the reaction.

Explain why (\ce{H2O(g)}) and (\ce{H2O(l)}) would have different (\Delta H_f^\ominus) values.

State whether the reaction is exothermic or endothermic from your result.

Identify the standard state of carbon from the table.

Write the formation equation for (\ce{CO(g)}) using the standard state of carbon.

Use the table to determine (\Delta H^\ominus) for (\ce{C(s, diamond) -> C(s, graphite)}).

Explain why allotropes can have different standard enthalpies of formation.

(s)}), use these data:

(\Delta H_f^\ominus[\ce{LiF(s)}]) = −617 kJ mol⁻¹; atomization of Li = +161 kJ mol⁻¹; atomization of F = +79 kJ mol⁻¹; first ionization energy of Li = +520 kJ mol⁻¹; first electron affinity of F = −328 kJ mol⁻¹.

Lattice enthalpy is defined as (\ce{LiF

-> Li+

• F-(g)}).

Write the Hess expression for the lattice enthalpy.

Calculate the lattice enthalpy.

(s)}). Lattice enthalpy is defined as separation into gaseous ions.

Identify the step labelled (A).

Identify the step labelled (B).

Use the cycle to determine the lattice enthalpy of (\ce{KBr(s)}).

State one reason why (\ce{KBr}) has a smaller lattice enthalpy than (\ce{KF}).

• H2

-> C3H8(g)}]

State the Hess expression for calculating (\Delta H^\ominus) using combustion enthalpies.

Calculate (\Delta H^\ominus) for the reaction.

Explain why the order of reactants and products in the combustion expression is opposite to that used with formation enthalpies.

• Cl2

-> CH2ClCH2Cl(g)}]

Average bond enthalpies: C=C = 614, Cl–Cl = 242, C–C = 347, C–H = 414, C–Cl = 327 kJ mol⁻¹.

Identify the bonds broken and formed that should be considered in the calculation.

Calculate the enthalpy change for the reaction and explain why the value obtained from average bond enthalpies may differ from an experimental value.

State Hess’s law and explain why it applies to this system.

Determine (\Delta H) for (\ce{A -> D}) and discuss one experimental reason why a calorimetric Hess cycle may give an uncertain value.

Compare the bond lengths and bond enthalpies of C–Cl and C–I.

Explain how these differences influence the rates of nucleophilic substitution reactions of chloroalkanes and iodoalkanes, and relate your answer to bond-breaking and bond-forming in reactions.

(s)}). Lattice enthalpy is defined as:

[\ce{MgO

-> Mg^{2+}

• O^{2-}(g)}]

Identify why two ionization energy steps are included for magnesium.

Identify why two electron affinity steps are included for oxygen.

Use the cycle to determine the lattice enthalpy.

Explain why the second electron affinity of oxygen is positive.

Calculate the enthalpy change predicted by this simplified bond enthalpy model.

Evaluate the reliability of this prediction for a reaction carried out in aqueous solution.

• H2O

-> CO

• 3H2(g)}]

(\Delta H_f^\ominus): (\ce{CH4(g)}) = −75; (\ce{H2O(g)}) = −242; (\ce{CO(g)}) = −111; (\ce{H2(g)}) = 0 kJ mol⁻¹.

Define standard enthalpy change of formation and state why (\Delta H_f^\ominus) for (\ce{H2(g)}) is zero.

Calculate (\Delta H^\ominus) for the reaction and discuss the importance of state symbols in this calculation.

• 3H2

-> C6H12(l)}]

Standard enthalpies of combustion: (\ce{C6H6(l)}) = −3267 kJ mol⁻¹; (\ce{H2(g)}) = −286 kJ mol⁻¹; (\ce{C6H12(l)}) = −3920 kJ mol⁻¹.

Write the Hess expression used with combustion enthalpies and explain why (\ce{H2}) must be multiplied by 3.

Calculate (\Delta H^\ominus) for the hydrogenation reaction and evaluate one advantage of using combustion data rather than average bond enthalpies for this reaction.

Define standard enthalpy change of combustion and write the combustion equation for (\ce{C2H5OH(l)}).

Compare the use of formation data and combustion data to calculate reaction enthalpies, including the form of the Hess expressions and one reason why state symbols are important.

(s)}), defined as:

[\ce{CaO

-> Ca^{2+}

• O^{2-}(g)}]

Identify the types of enthalpy changes that must appear in the cycle to form gaseous (\ce{Ca^{2+}}) and (\ce{O^{2-}}) ions from the elements.

Explain how the lattice enthalpy is obtained from the cycle and why (\ce{CaO}) has a much larger lattice enthalpy than (\ce{KCl}).