Enthalpy of Reactions (Formation, Combustion, Atomisation) - UNSOLVED PRACTICE SET
Chapter: Thermodynamics | Topic: Enthalpy of Reactions Formation Combustion Atomisation
ENTHALPY OF REACTIONS (FORMATION, COMBUSTION, ATOMISATION) - UNSOLVED PRACTICE SET
Topic: Enthalpy of Reactions Formation Combustion Atomisation
Multiple Choice Questions
Q1. The standard enthalpy of formation (ΔH°f) of an element in its standard state is:
- Always positive
- Always negative
- Zero
- Depends on the element
Q2. The enthalpy of combustion of a substance is always:
- Positive
- Negative
- Zero
- Either positive or negative
Q3. The enthalpy of atomisation of CH₄(g) refers to:
- CH₄(g) → C(g) + 4H(g)
- CH₄(g) → C(s) + 2H₂(g)
- C(s) + 2H₂(g) → CH₄(g)
- CH₄(g) → C(g) + 2H₂(g)
Q4. Which of the following reactions represents the standard enthalpy of formation of CO₂(g)?
- C(s) + O₂(g) → CO₂(g)
- CO(g) + ½O₂(g) → CO₂(g)
- C(g) + O₂(g) → CO₂(g)
- C(s) + 2O(g) → CO₂(g)
Q5. The enthalpy of atomisation of a diatomic molecule X₂ is:
- The energy required to form 1 mole of X₂ from its atoms
- The energy required to dissociate 1 mole of X₂ into its atoms
- The energy released when 1 mole of X₂ is formed from its elements
- The energy released when 1 mole of X₂ burns in oxygen
Q6. For the reaction: H₂(g) + ½O₂(g) → H₂O(l), ΔH° = –286 kJ/mol. This represents:
- Enthalpy of formation of H₂O(l)
- Enthalpy of combustion of H₂(g)
- Both (a) and (b)
- Neither (a) nor (b)
Short Answer Questions
Q7. Define standard enthalpy of formation. Why is it not possible to measure the enthalpy of formation of some compounds directly?
Q8. Differentiate between enthalpy of combustion and enthalpy of formation with one example for each.
Q9. What is meant by enthalpy of atomisation? Why is the enthalpy of atomisation of metals generally high?
Q10. Write thermochemical equations for:
(a) Standard enthalpy of formation of NH₃(g)
(b) Standard enthalpy of combustion of ethanol (C₂H₅OH)
Q11. The enthalpy of combustion of carbon (graphite) is –393.5 kJ/mol, while that of carbon (diamond) is –395.4 kJ/mol. Which allotrope is more stable? Explain.
Q12. Why is the enthalpy of atomisation of a diatomic molecule like Cl₂ different from the bond dissociation enthalpy of the same molecule? Explain.
Long Answer Questions
Q13. (a) Define the following terms with suitable examples:
(i) Standard enthalpy of formation (ΔH°f)
(ii) Standard enthalpy of combustion (ΔH°c)
(iii) Enthalpy of atomisation (ΔHat)
(b) Why is the enthalpy of formation of ozone (O₃) not zero, even though it contains only oxygen atoms?
(c) Explain why the enthalpy of combustion is always a negative quantity.
Q14. (a) Write balanced thermochemical equations for the standard enthalpy of formation of:
(i) H₂O(l)
(ii) CO₂(g)
(iii) NH₃(g)
(iv) CH₃OH(l)
(b) The standard enthalpy of formation of HCl(g) is –92.3 kJ/mol. What does this value tell you about the stability of HCl(g)?
(c) Calculate the enthalpy change for the reaction: 2H₂O₂(l) → 2H₂O(l) + O₂(g) using the following data:
ΔH°f [H₂O₂(l)] = –187.8 kJ/mol, ΔH°f [H₂O(l)] = –285.8 kJ/mol
Q15. (a) Define enthalpy of atomisation and explain its significance in calculating bond enthalpies.
(b) The enthalpy of atomisation of sodium metal is 108 kJ/mol. What does this value represent at the atomic level?
(c) Compare the enthalpy of atomisation of Na(s) and Cl₂(g). Which would be higher and why?
Numerical / Application-Based Problems
Q16. Calculate the standard enthalpy of formation of methane (CH₄) using the following data:
ΔH°c [C(graphite)] = –393.5 kJ/mol
ΔH°c [H₂(g)] = –285.8 kJ/mol
ΔH°c [CH₄(g)] = –890.3 kJ/mol
Q17. Given the following standard enthalpies of formation:
ΔH°f [CO₂(g)] = –393.5 kJ/mol
ΔH°f [H₂O(l)] = –285.8 kJ/mol
ΔH°f [C₂H₅OH(l)] = –277.7 kJ/mol
(a) Write the thermochemical equation for the combustion of ethanol.
(b) Calculate the standard enthalpy of combustion of ethanol.
(c) How much heat is released when 5.0 g of ethanol is burned completely?
(Molar mass of ethanol = 46 g/mol)
Q18. In rural India, many households still use wood (cellulose, approximately C₆H₁₀O₅) as a fuel for cooking. The enthalpy of combustion of cellulose is approximately –2800 kJ/mol.
(a) Write the balanced equation for the combustion of cellulose.
(b) Calculate the mass of wood needed to produce 1000 kJ of heat energy.
(c) A student suggests using LPG (liquefied petroleum gas, mainly propane C₃H₈) instead, which has ΔH°c = –2220 kJ/mol. Calculate the mass of propane needed to produce the same 1000 kJ of heat.
(d) Based on your calculations, which fuel is more efficient per gram? What environmental factors should be considered?
(Molar masses: Cellulose unit = 162 g/mol, Propane = 44 g/mol)