The standard enthalpy of formation of gaseous H 2 O at 298 K is − 241 .82 kJ mo l - 1 . Calculate ΔH f o at 373 K, given the following values of the molar heat capacities at constant pressure. Molar h

Question

The standard enthalpy of formation of gaseous $H_{2} O$ at $298 K$ is $- 241 .82 kJ$ $mo l^{- 1}$ . Calculate ${ΔH}_{f}^{o}$ at $373 K,$ given the following values of the molar heat capacities at constant pressure.
Molar heat capacity of $H_{2} (g) = 33.58 J K^{- 1} m o l^{- 1}$
Molar heat capacity of $H_{2} (g) = 28.84 J K^{- 1} m o l^{- 1}$
Molar heat capacity of $O_{2} (g) = 29.37 J K^{- 1} m o l^{- 1}$
Assume that the heat capacities are independent of temperature.

TestHub · Accepted Answer

The reaction is H 2 g + 1 2 O 2 g → H 2 O g Δ C p o = C p , m o H 2 O , g - C p , m o H 2 , g + 1 2 C p , m o O 2 = 33.58 - 28.84 + 1 2 29.37 = - 9.94 JK - 1 m o l - 1 Using Kirchhoff's equation, Δ H ° 373 K = Δ H ° 298 K + T 2 - T 1 Δ C p o = - 241.82 + 373 - 298 × - 9.94 = - 242.6 kJ mo l - 1

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