Question

Which of the following statements is not correct? (a) The efficiency of a solid catalyst depends upon its surface area. (b) Catalyst operates by providing alternate path for the reaction that involves lower activation energy. (c) Catalyst lowers the energy of activation of the forwards direction without affecting the energy of activation of the backward direction. (d) Catalyst does not affect the overall enthalpy change of the reaction.

Short answer

Statement (c) is not correct, as a catalyst affects the activation energy of both the forward and backward reactions equally.

Step-by-step solution

Understanding Catalysts and Surface Area

The efficiency of a catalyst, especially a solid one, is indeed dependent on its surface area. This is because more surface area allows for more active sites where reactants can interact with the catalyst. Thus, statement (a) is a correct statement about the characteristics of a catalyst.

Catalysts and Activation Energy

Catalysts work by providing an alternative reaction pathway with a lower activation energy. This doesn't change the reactants or products or the overall enthalpy change of the reaction, but it does make it easier for the reaction to occur. Therefore, statement (b) accurately describes the function of a catalyst and is indeed correct.

Direction-Specific Effect on Activation Energy

Statement (c) suggests that a catalyst only lowers the energy of activation of the forward reaction without affecting the backward reaction. This is incorrect. A catalyst affects both the forward and backward reactions equally because it provides an alternative pathway for the reaction in BOTH directions. Thus, statement (c) is not a correct description of how catalysts function.

Catalysts and Enthalpy Change

Statement (d) says that a catalyst does not affect the overall enthalpy change of the reaction, which is true. A catalyst does not change the thermodynamics of a reaction; it only affects the kinetics by lowering the activation energy. Thus, statement (d) is correct.

Key concepts

Surface Area of Solid Catalysts

Catalysts play a crucial role in increasing the rate of chemical reactions without being consumed in the process themselves. When it comes to solid catalysts, their efficiency is significantly influenced by their surface area. A larger surface area means more space for reactants to adhere to and interact with the catalyst. This is quite analogous to having a larger workspace in which more tasks can be performed simultaneously.

Consider, for instance, a sponge. If you were to spread the sponge out, you'd expose more pores—the same goes for solid catalysts. By increasing the surface area, you provide more 'active sites' for the chemical reaction. These active sites are where the magic happens; reactants are temporarily bound, allowing the reaction to proceed at a faster rate than it would in the absence of the catalyst.

• High surface area = More active sites
• More active sites = Increased reaction rate

Activation Energy in Catalysis

Activation energy is the threshold energy that reactants must overcome to produce products. Think of it as the entry barrier for a party—if the barrier is too high, fewer guests arrive. In chemical reactions, if the activation energy is high, the reaction will be slower or might not proceed at all.

Catalysts are like VIP passes to this chemical 'party'. They provide an alternative pathway with a lower activation energy barrier. By doing so, they enable more reactant molecules to have enough energy to react. This doesn't mean that catalysts make reactions happen that wouldn't happen on their own. Instead, they simply make the reactions occur much faster.

Catalysts are not one-size-fits-all, though. Different reactions require different catalysts, and the choice of catalyst can drastically affect the effectiveness of the reaction. Importantly, the catalyst remains unchanged after the reaction, ready to assist in another reaction cycle.

Enthalpy Change and Catalysts

Enthalpy change represents the heat absorbed or released during a reaction, under constant pressure. It's a signifier of the overall energy exchange taking place in the reaction. One may think that anything influencing the rate of a reaction, like a catalyst, might also alter this energy landscape. However, that's not the case.

Catalysts are like shortcuts in a mountain trail—they make the journey quicker but don't change the height of the mountain itself. The total energy from the start to the end remains constant. So, whether a catalyst is present or not, the overall enthalpy change of the reaction—the difference in energy between reactants and products—remains the same. This is a key concept:

• Catalysts do not alter the energy content of reactants or products.
• Catalysts only affect how quickly equilibrium is reached.

This means that while catalysts are invaluable for increasing the rate of reactions, they do not change the 'energy budget' of the reactions themselves.