From the following bond energies : $H - H$ bond energy $: 431.37\, kJ\, mol ^{-1}$ $C = C$ bond energy $: 606.10\, kJ\, mol ^{-1}$ $C - C$ bond energy $: 336.49\, kJ\, mol ^{-1}$ $C - H$ bond energy $: 410.50\, kJ\, mol ^{-1}$ Enthalpy for the reaction, will be

Question

Which is an adiabatic process?

Answer 1

To solve this problem, we need to calculate the enthalpy change for a given chemical reaction using the bond energies provided. The bond enthalpy, or bond dissociation energy, is the energy required to break one mole of bonds in a gaseous substance.

Let us assume the given reaction involves the following bond-breaking and bond-forming processes:

- **Breaking Bonds:** - 1 mole of $H-H$ bonds - 1 mole of $C=C$ bonds - **Forming Bonds:** - 2 moles of $C-H$ bonds - 1 mole of $C-C$ bonds

First, we calculate the total energy required to break the bonds:

Energy to break 1 mole of $H-H$ bond: 431.37 \, \text{kJ/mol}
Energy to break 1 mole of $C=C$ bond: 606.10 \, \text{kJ/mol}

Total energy required for bond breaking:

431.37 + 606.10 = 1037.47 \, \text{kJ/mol}

Next, we calculate the energy released during the formation of bonds:

Energy released from forming 2 moles of $C-H$ bonds: 2 \times 410.50 = 821.00 \, \text{kJ/mol}
Energy released from forming 1 mole of $C-C$ bond: 336.49 \, \text{kJ/mol}

Total energy released for bond formation:

821.00 + 336.49 = 1157.49 \, \text{kJ/mol}

The enthalpy change for the reaction ($\Delta H$) is the difference between the energy required to break bonds and the energy released by bond formation:

\Delta H = \text{Energy of bonds broken} - \text{Energy of bonds formed} = 1037.47 \, \text{kJ/mol} - 1157.49 \, \text{kJ/mol}

Calculating this gives:

\Delta H = -120.02 \, \text{kJ/mol}

(Rounded, this becomes -120.0 kJ/mol)

Thus, the enthalpy change for the reaction is -120.0 \, \text{kJ/mol}. This implies the reaction is exothermic.

The correct answer is: -120.0 \, \text{kJ/mol}.

Answer 2