The reaction is:
C₂H₄ + 3O₂ → 2CO₂ + 2H₂O; ΔH = 1410 kJ
When we reverse this reaction, the sign of the enthalpy change, ΔH, will be changed. The enthalpy change for the reversed reaction would be 1,410 kJ.
Next, we must also multiply the reaction by 2, so the final enthalpy change for the reverse reaction will be:
ΔH = 2,820 kJ
Answer:
Correct combinations are
For option A it is 2
For option B it is 4
For option D it is 3
For option C it is 1
Explanation:
For checking the spontaneity of a reaction, we have to check the sign of ΔG using the below formula
<h3>ΔG = ΔH - T×ΔS</h3>
where
ΔG is the change in Gibbs free energy
ΔH is the change in enthalpy
T is the temperature
ΔS is the change in entropy
<h3>For spontaneous reactions, ΔG must be less than zero and for non-spontaneous reactions ΔG must be greater than zero but for an equilibrium reaction ΔG must be equal to zero</h3>
So in case of 1 as ΔH and ΔS are positive if the temperature is above a certain value then ΔG will be less than zero
So in case of 2 as ΔH is negative and ΔS is positive then ΔG will always be less than zero at all temperatures
So in case of 3 as ΔH and ΔS are negative if the temperature is below a certain value then ΔG will be less than zero
So in case of 4 as ΔH is positive and ΔS is negative then ΔG will always be greater than zero but in reverse direction as ΔG is less than zero therefore in reverse direction the reaction will be spontaneous at all temperatures