Calculate the average molar bond enthalpy of the carbon-hydrogen bond in a CH4 molecule.
The correct answer and explanation is:
Correct Answer:
The average molar bond enthalpy of the carbon-hydrogen (C–H) bond in a CH₄ molecule is approximately 412 kJ/mol.
Explanation:
Methane (CH₄) is a molecule consisting of one carbon atom bonded to four hydrogen atoms. Each of these bonds is a carbon-hydrogen (C–H) single covalent bond. To calculate the average bond enthalpy of the C–H bond in methane, the total energy required to break all four C–H bonds in one mole of methane molecules is used, then divided by four.
The reaction for the complete bond dissociation of methane is:
CH₄(g) → C(g) + 4H(g)
This process involves breaking all four C–H bonds in one mole of methane molecules, and the energy required to carry out this process is the total bond dissociation energy for methane. Experimental data shows that this total energy is approximately 1648 kJ/mol.
To find the average bond enthalpy per C–H bond, the total energy is divided by the number of bonds:
Average C–H bond enthalpy = 1648 kJ/mol ÷ 4 = 412 kJ/mol
This value represents the average energy needed to break one mole of C–H bonds in methane under standard conditions.
Bond enthalpy values are average quantities because actual bond strengths vary slightly depending on the molecular environment. For methane, all four C–H bonds are equivalent due to the molecule’s symmetrical tetrahedral structure, making this average very accurate and representative.
Understanding bond enthalpies is essential in thermochemistry because they allow prediction of the energy changes in chemical reactions. If the energy released in forming bonds is greater than that required to break them, the reaction is exothermic. If more energy is needed to break bonds than is released in forming them, the reaction is endothermic.