If the products of a certain reaction are favored (i. e., the ΔG value of the reaction is negative), one would expect a significant amount of product to form when the reactants are combined. Which of the choices could explain why such a reaction might produce very little product?

A. low activation energy, Ea

B. negative change in entropy, ΔS

C. high activation energy, Ea

D. positive change in enthalpy, ΔH

C. High activation Energy, Ea

Explanation:

The activation energy of a chemical reaction affects the rate at which the reaction proceeds regardless of its ΔG value. The higher the activation energy, the slower the reaction proceeds.

This occurs primarily due to the fact that the molecules of the reacting substances have to attain an energy value which is higher than the activation energy before a reaction starts. This is so even for reactions that are termed spontaneous. Molecules can only complete the reaction once they have reached the top of the activation energy barrier.

If the activation energy is high, only a few molecules will have enough energy to reach this barrier, thus slowing down the chemical reaction.

This makes option C the correct choice.

positive H and negative S

Explanation:

For a reaction to be spontaneous, the absolute best combination is a negative Delta H and a positive Delta S. When they are both positive, the reaction is only spontaneous at higher temperatures. When they are both negative, the reaction is only spontaneous at lower temperatures. and again if a catalyst is added to the reaction, the activation energy is lowered because a lower-energy transition state is formed. The catalyst does not affect the energy of the reactants or products (and thus does not affect ΔG).

So from these discussions

Ea does not affect G value at all (whether + Ea or - Ea).

And for product to be formed the reaction should be spontaneous, where H is negative and S positive else the reaction will yield low product.