At constant pressure,the combustion of 6.00 g of C₄H₁₀(g) releases 297 kJ of heat.What is ΔH for the reaction given below?
2 C₄H₁₀(g) + 13 O₂(g) → 8 CO₂(g) + 10 H₂O(l)

Question

At constant pressure,the combustion of 6.00 g of C₄H₁₀(g)releases 297 kJ of heat.What is ΔH for the reaction given below? 2 C₄H₁₀(g)+ 13 O₂(g)→ 8 CO₂(g)+ 10 H₂O(l)

Quizplus · Accepted Answer

First, we need to calculate the amount of moles of C₄H₁₀ that reacted: 6.00 g C₄H₁₀ x (1 mol C₄H₁₀/76.14 g C₄H₁₀) = 0.0788 mol C₄H₁₀ Now, we can use the stoichiometric coefficients to find the amount of heat released for the entire reaction: 2 mol C₄H₁₀ → 8 mol CO₂ + +7 mol H₂O ΔH = -297 kJ 0.0788 mol C₄H₁₀ → ΔH -297 kJ → (ΔH/0.0788 mol) = -3760.91 kJ/mol However, this is the ΔH for 2 moles of C₄H₁₀. To find the ΔH for 1 mole of C₄H₁₀, we divide by 2: ΔH = -3760.91 kJ/mol ÷ 2 = -1880.45 kJ/mol Finally, to find the ΔH for the entire reaction (per mole of C₄H₁₀), we multiply by the coefficient of C₄H₁₀ in the balanced equation: ΔH = -1880.45 kJ/mol x 2 = -3760.91 kJ/mol Therefore, the best choice is C).

At Constant Pressure,the Combustion of 6