Use the reaction and bond information to answer the question.

C2H6 → C2H4 + H2

Reactant bond energies: H–C = 413 kJ/mol, C–C single bond = 347 kJ/mol

Product bond energies: H–C = 413 kJ/mol, C=C double bond = 614 kJ/mol, H–H = 432 kJ/mol

Reactions often do not occur spontaneously (without the addition of energy) if it takes more energy to start the reaction than is released at the end. Based on the bond energies, would this reaction occur spontaneously? Why or why not?

No, because the reaction is exothermic.

Yes, because the reaction is exothermic.

Yes, because the reaction is endothermic.

No, because the reaction is endothermic

After calculating the enthalpy change for the reaction C₂H₆ → C₂H₄ + H₂, with the bond energies of C-C (347 kJ/mol), H-C (413 kJ/mol), C=C (614 kJ/mol), and H-H (432 kJ/mol), we found that the reaction does not occur spontaneously because it is endothermic (option D).

The reaction is:

C₂H₆ → C₂H₄ + H₂

We need to calculate the enthalpy for the products and reactants to know if this reaction will occur spontaneously.

Reactant

The bond enthalpy for reactants is given by:

[tex] \Delta H_{r} = \Delta H_{C_{2}H_{6}} [/tex]

Ethane (C₂H₆) has 1 C-C single bond and 6 H-C bonds, so:

[tex] \Delta H_{r} = \Delta H_{C-C} + 6*\Delta H_{H-C} [/tex]

The bond energies are:

ΔH (C-C) = 347 kJ/mol
ΔH (H-C) = 413 kJ/mol

Hence, the enthalpy for ethane is:

[tex] \Delta H_{r} = 347 kJ/mol + 6*413 kJ/mol = 2825 kJ/mol [/tex]

Products

The bond enthalpy for products is the following:

[tex] \Delta H_{p} = \Delta H_{C_{2}H_{4}} + \Delta H_{H_{2}} [/tex]

Hydrogen (H₂) has 1 H-H bond and ethylene (C₂H₄) has 1 C=C double bond and 4 H-C bonds, so:

[tex] \Delta H_{p} = \Delta H_{C=C} + 4*\Delta H_{H-C} + \Delta H_{H-H} [/tex]

The bond energies are:

ΔH (C=C) = 614 kJ/mol
ΔH (H-C) = 413 kJ/mol
ΔH (H-H) = 432 kJ/mol

Therefore, the bond enthalpy for products is:

[tex] \Delta H_{p} = 614 kJ/mol + 4*413 kJ/mol + 432 kJ/mol = 2698 kJ/mol [/tex]

Now, the enthalpy change for the reaction is:

[tex] \Delta H = \Delta H_{r} - \Delta H_{p} [/tex]

[tex] \Delta H = (2825 - 2698) kJ/mol = 127 kJ/mol [/tex]

We can see that the reaction is endothermic because the energy of the reactants is greater than the energy of the products (the enthalpy change is positive), which means that we need to apply energy for the reaction to take place.

Therefore, the answer is option D: No, because the reaction is endothermic.

Find more about bond energies here:

https://brainly.com/question/866298?referrer=searchResults

https://brainly.com/question/15803993?referrer=searchResults

I hope it helps you!