A student wants to determine the standard enthalpy of formation of butane, C4H10(g)\text{C}_4\text{H}_{10}(\text{g})C4H10(g), using a Hess's Law cycle as shown in the diagram below:
Given the following standard enthalpy of combustion data:
- ΔcHθ[C(s)]=−394 kJ mol−1\Delta_c H^\theta [\text{C(s)}] = -394\text{ kJ mol}^{-1}ΔcHθ[C(s)]=−394 kJ mol−1
- ΔcHθ[H2(g)]=−286 kJ mol−1\Delta_c H^\theta [\text{H}_2(\text{g})] = -286\text{ kJ mol}^{-1}ΔcHθ[H2(g)]=−286 kJ mol−1
- ΔcHθ[C4H10(g)]=−2877 kJ mol−1\Delta_c H^\theta [\text{C}_4\text{H}_{10}(\text{g})] = -2877\text{ kJ mol}^{-1}ΔcHθ[C4H10(g)]=−2877 kJ mol−1
What is the standard enthalpy of formation (ΔfHθ\Delta_f H^\thetaΔfHθ) of butane, and is this formation reaction endothermic or exothermic?
−129 kJ mol−1-129\text{ kJ mol}^{-1}−129 kJ mol−1, exothermic
+129 kJ mol−1+129\text{ kJ mol}^{-1}+129 kJ mol−1, endothermic
−5883 kJ mol−1-5883\text{ kJ mol}^{-1}−5883 kJ mol−1, exothermic
+2197 kJ mol−1+2197\text{ kJ mol}^{-1}+2197 kJ mol−1, endothermic