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Class 11  »  Chemistry  »  Thermodynamics

Bond Enthalpy - UNSOLVED PRACTICE SET

Class 11

Chapter: Thermodynamics | Topic: Bond Enthalpy

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Class 11

BOND ENTHALPY - UNSOLVED PRACTICE SET

Topic: Bond Enthalpy

Time: 40 mins | Marks: 30 | Difficulty: Medium

Multiple Choice Questions

Q1. Bond enthalpy is defined as:

  1. The energy required to form one mole of bonds
  2. The energy required to break one mole of bonds in gaseous molecules
  3. The energy released when one mole of bonds is formed
  4. The energy required to break all bonds in a solid

Q2. The bond enthalpy of H–H is 436 kJ/mol. This means:

  1. 436 kJ is released when 1 mole of H₂ is formed from H atoms
  2. 436 kJ is required to break 1 mole of H₂ into H atoms
  3. 436 kJ is required to break 1 mole of H₂ into H⁺ ions
  4. 436 kJ is released when 2 moles of H atoms combine

Q3. For a diatomic molecule, bond enthalpy is:

  1. Equal to the average bond enthalpy
  2. Different from the average bond enthalpy
  3. Always zero
  4. Equal to the lattice enthalpy

Q4. In the reaction H₂(g) + Cl₂(g) → 2HCl(g), the enthalpy change can be estimated using:

  1. Only bond enthalpies of reactants
  2. Only bond enthalpies of products
  3. Bond enthalpies of both reactants and products
  4. Only lattice enthalpies

Q5. The average bond enthalpy of C–H in methane is different from the bond dissociation enthalpy of the first C–H bond in methane because:

  1. Methane is a gas
  2. The four C–H bonds in methane are not identical in energy
  3. The average is taken over all four bonds
  4. Carbon has different hybridisation states

Q6. For an exothermic reaction, the sum of bond enthalpies of reactants is:

  1. Greater than the sum of bond enthalpies of products
  2. Less than the sum of bond enthalpies of products
  3. Equal to the sum of bond enthalpies of products
  4. Unrelated to the sum of bond enthalpies of products

Short Answer Questions

Q7. Define bond enthalpy. Why is bond enthalpy always a positive quantity?

Q8. Differentiate between bond dissociation enthalpy and average bond enthalpy with a suitable example.

Q9. Why is the average bond enthalpy of O–H in water different from the bond dissociation enthalpy of the first O–H bond in water? Explain.

Q10. Write the formula for calculating the enthalpy change of a reaction using bond enthalpies. Explain why this method gives only an approximate value.

Q11. The bond enthalpy of N≡N is very high (945 kJ/mol). What does this tell you about the reactivity of nitrogen gas?

Q12. Why can't bond enthalpies be used to calculate exact enthalpy changes for reactions involving ionic compounds? Explain.

Long Answer Questions

Q13. (a) Define bond enthalpy and explain why it is defined for gaseous molecules only.

(b) Explain the difference between bond dissociation enthalpy and average bond enthalpy using CH₄ as an example.

(c) The bond dissociation enthalpies for the four C–H bonds in methane are: 439, 458, 489, and 338 kJ/mol. Calculate the average C–H bond enthalpy.

Q14. (a) Describe how bond enthalpies can be used to estimate the enthalpy change of a reaction. Write the general formula and explain each term.

(b) Calculate the enthalpy change for the following reaction using bond enthalpies:

H2(g) + Cl2(g) → 2HCl(g)

Given: Bond enthalpies (kJ/mol): H–H = 436, Cl–Cl = 243, H–Cl = 431

(c) Compare the calculated value with the experimental value of –184.6 kJ and suggest reasons for any difference.

Q15. (a) Explain why the bond enthalpy of a double bond is not exactly twice the bond enthalpy of a single bond between the same atoms. Use C–C and C=C as examples.

(b) Arrange the following bonds in order of increasing bond enthalpy and justify your answer: C–C, C=C, C≡C

(c) A student argues that since the C≡C bond has the highest bond enthalpy, compounds containing triple bonds should be the most stable. Is this reasoning always correct? Explain with an example.

Numerical / Application-Based Problems

Q16. Calculate the enthalpy change for the combustion of methane using bond enthalpies:

CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)

Given bond enthalpies (kJ/mol):

C–H = 413, O=O = 498, C=O = 745, O–H = 463

Q17. The bond enthalpies of H–H, Cl–Cl, and H–Cl are 436, 243, and 431 kJ/mol respectively.

(a) Calculate ΔH for the reaction: H₂(g) + Cl₂(g) → 2HCl(g)

(b) Is this reaction exothermic or endothermic? Explain.

(c) The experimental value of ΔH for this reaction is –184.6 kJ. Why does the value calculated using bond enthalpies differ from the experimental value?

Q18. In the petrochemical industry, cracking of alkanes is an important process. Consider the cracking of propane:

C3H8(g) → C2H4(g) + CH4(g)

Given bond enthalpies (kJ/mol):

C–C = 347, C=C = 614, C–H = 413, H–H = 436

(a) Calculate the enthalpy change for this cracking reaction using bond enthalpies.

(b) Is this reaction exothermic or endothermic? What does this imply about the conditions needed for cracking?

(c) In India, many refineries use catalytic cracking to produce useful products like ethene and propene. Why is understanding the enthalpy change important for designing industrial reactors?

(d) Suggest one way to make this cracking process more energy-efficient in an industrial setting.


Total: 30 Marks | Time: 40 mins

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