Answer:
The correct answer is:
Reaction B is more likely to occur at all than reaction A.
Explanation:
The activation energy in chemistry is the smallest amount of energy required to cause chemical or nuclear reaction in the reactants in chemical or nuclear systems. The activation energy is denoted by [tex]E_{a}[/tex], and it is measured in Joules (J), KiloJoules (KJ) or Kilocalories per mole (Kcal/mol)
The activation energy can be thought of simply as the minimum amount of energy required to overcome a barrier that prevents a reaction from occurring, hence, from our question, if Reaction A has a high activation energy, it means that the barrier to be overcome before a reaction will occur is large, meaning that the reaction system is more stable and the reaction is less likely to occur than Reaction B which has a low activation energy, meaning that just a relatively small amount of energy, when applied to the reaction system, will initiate a reaction, making it more likely to occur than reaction A.
You should also note that catalysts are substances that are capable of reducing the activation energy of a system, but remains unchanged at the end of the system.
The statement about reaction A and reaction B are true is: A. Reaction B is likely to occur at a faster rate than reaction A.
An activation energy can be defined as the minimum quantity of energy that must be provided for reacting chemical elements, so as to undergo a chemical reaction. Thus, the activation energy of a chemical reaction must first of all be reached by its combining chemical elements (reactants) before it can start or begin.
As a general rule, the lower the activation energy of a chemical reaction, the faster is the rate of a chemical reaction and vice-versa. This ultimately implies that, the activation energy of a chemical reaction should be lowered, in order for it to occur fast (quickly) enough over a short period of time.
In conclusion, reaction B is likely to occur at a faster rate than reaction A because of its low activation energy.
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