Problem 33 - Chemistry in 30 days

Question

According to the adsorption theory of catalysis, the speed of the reaction increases because (a) adsorption lowers the activation energy of the reaction. (b) the concentration of reactant molecules at the active centers of the catalyst becomes high due to adsorption. (c) in the process of adsorption, the activation energy of the molecules becomes large. (d) adsorption produces heat which increases the speed of the reaction.

Step-by-Step Solution

Answer

(b) the concentration of reactant molecules at the active centers of the catalyst becomes high due to adsorption.

Step 1: Identify the Key Concept

Adsorption theory states that catalysis occurs when reactant molecules adhere to the surface of a catalyst, allowing the reaction to proceed more efficiently due to increased local concentration or altered activation energy.

Step 2: Analyze Each Option

(a) Adsorption could lower activation energy by facilitating interaction. (b) Adsorption indeed increases molecule concentration at active sites. (c) Activation energy doesn't become larger; it usually decreases. (d) Adsorption does release heat but this is a less direct reason for increased speed.

Step 3: Evaluate the Options

Option (b) aligns well with the concept of increased local concentration due to adsorption, which enhances reaction speed. Option (a) is also correct by considering lowered activation energy, but option (b) provides a more direct explanation.

Step 4: Select the Best Answer

The option most closely aligning with the central tenet of adsorption theory is (b) because it directly explains how increased reactant concentration at active centers improves reaction efficiency.

Key Concepts

Activation Energy

Activation energy is a fundamental concept in chemistry, especially when discussing reactions and catalysis. It's the minimum energy that reactant molecules need to successfully collide and lead to a reaction. Think of it as the energy hurdle that molecules must overcome to transform into products.
In relation to the adsorption theory of catalysis, lowering activation energy is a critical factor. When a catalyst is used, it provides an alternative pathway for the reaction. This pathway has a lower activation energy compared to the uncatalyzed reaction path. As a result, more molecules have enough energy to surpass this barrier, leading to an increased rate of reaction.
This doesn't mean the molecules need more energy, rather, the energy they already possess is sufficient because the barrier is lower. In simple terms, a catalyst makes it easier for molecules to react by lowering the energy hill they need to climb. Thus, reducing activation energy is one of the main ways a catalyst functions in speeding up reactions.

Reaction Speed

Reaction speed, often referred to as reaction rate, is a measure of how fast a reaction occurs. Two main factors that influence reaction speed are the concentration of reactants and the activation energy.
According to the adsorption theory, the reaction speed can be significantly affected by how reactants interact with a catalyst. Adsorption increases the concentration of reactant molecules on the catalyst's surface, essentially gathering more molecules in one area. This increases the likelihood of collisions between reactant molecules, thus enhancing the reaction speed.
In addition, by providing an alternative route with a lower activation energy, catalysts allow reactions to proceed faster. So, the dual effect of increasing reactant concentration and lowering activation energy explains why reactions speed up in the presence of a catalyst. Remember, a faster reaction doesn't mean the quality is compromised; it simply means that the transformation from reactants to products happens quicker.

Catalysis

Catalysis is a process that accelerates chemical reactions by introducing a catalyst. A catalyst is a substance that, although involved in the reaction, is not consumed by it. The magic of catalysis lies in its ability to transform the way reactants interact without undergoing permanent change.
The adsorption theory of catalysis explains one way this happens. It suggests that reactant molecules adhere to the surface of a catalyst, particularly on active centers where the reaction occurs. These active centers are tiny spaces or spots on the catalyst's surface that are especially reactive.
Once adsorbed, the concentration of reactants increases locally, making collisions more frequent and effective. Furthermore, typical energy requirements are altered, often lowered, allowing the reaction to proceed with greater ease.

• A catalyst doesn't change the equilibrium of the reaction; it merely helps reach that equilibrium faster.
• It's efficient, required in only small amounts relative to the reactants.
• While not consumed, it can be inactivated by poisons or degraded under certain conditions, impacting its effectiveness.

Understanding catalysis helps chemists design better processes to achieve desired reactions more efficiently and cost-effectively.