Catalyst: Videos & Practice Problems

A catalyst is a substance that accelerates a chemical reaction by lowering the activation energy without being consumed by the reaction itself. The energy diagram illustrates this by showing a lower energy barrier for the catalyzed reaction compared to the uncatalyzed one, indicating a reduced activation energy. The height of the activation energy is inversely related to the reaction speed; a lower activation energy means reactants can more readily overcome the energy barrier and transform into products. This concept is crucial in understanding how catalysts influence the dynamics of chemical reactions, making them faster and more efficient.

Catalyzed vs. Uncatalyzed Reactions

A catalyst is a substance that accelerates a chemical reaction by lowering the activation energy required for the reaction to occur, without being consumed in the process. In an energy diagram, the reactants are represented at the beginning, while the products are at the end. The difference in energy between these two points indicates the overall change in energy of the reaction.

Within the energy diagram, the highest point of the curve represents the transition state, which is crucial for understanding the activation energy. The activation energy (denoted as \(E_a\)) is the energy difference between the reactants and the transition state. In an uncatalyzed reaction, the activation energy is relatively high, as illustrated by the red curve, which we can label as \(E_{a \, \text{uncat}}\).

When a catalyst is introduced, the activation energy decreases, as shown by the blue curve, which represents the catalyzed reaction. This lower activation energy (denoted as \(E_{a \, \text{cat}}\)) allows the reactants to more easily reach the transition state and convert into products. Consequently, a lower activation energy results in a faster reaction rate, as the reactants can more readily overcome the energy barrier represented by the transition state.

In summary, the role of a catalyst is to facilitate the conversion of reactants to products by reducing the activation energy, thereby increasing the reaction rate. This fundamental concept is essential in understanding how catalysts function in various chemical processes.

Catalyzed vs. Uncatalyzed Reactions Example

In this scenario, we analyze a chemical reaction characterized by an enthalpy change (ΔH) of -20 kilojoules and an activation energy (Ea) of 40 kilojoules. The enthalpy change indicates that the products of the reaction are lower in energy than the reactants, specifically 20 kilojoules lower. This can be visualized by placing the reactants at 40 kilojoules and the products at 20 kilojoules on an energy diagram.

The activation energy represents the energy barrier that must be overcome for the reaction to proceed, which is the difference in energy between the reactants and the transition state. In this case, the transition state is positioned at 80 kilojoules, calculated as the sum of the reactants' energy (40 kilojoules) and the activation energy (40 kilojoules).

When a catalyst is introduced, it lowers the activation energy by 10 kilojoules, resulting in a new transition state energy of 70 kilojoules. The energy of the reactants remains unchanged at 40 kilojoules, and the products still sit at 20 kilojoules. To find the total difference in energy between the products and the transition state, we subtract the energy of the products from that of the transition state: 70 kilojoules (transition state) - 20 kilojoules (products) = 50 kilojoules.

Despite the introduction of the catalyst, the overall enthalpy change (ΔH) remains at -20 kilojoules, as it is a property of the reaction itself and not affected by the catalyst. The new activation energy can also be calculated as the difference between the new transition state and the reactants: 70 kilojoules - 40 kilojoules = 30 kilojoules.

In summary, the catalyst effectively reduces the activation energy and alters the energy landscape of the reaction, resulting in a total difference of 50 kilojoules between the products and the transition state, while maintaining the enthalpy change at -20 kilojoules.

Which of the following statements is true regarding the energy diagram provided?
i. The reaction is endothermic.
ii. Activation energy would be less than + 10 kJ after a catalyst is added.
iii. The reaction absorbs energy.
iv. Activation energy would be greater than + 10 kJ after a catalyst is added.

• A. I only
• B. II only
• C. I and III
• D. II, III and IV

Answer: (B)

Here’s what students ask on this topic:

What is a catalyst?

A catalyst is a substance that increases the rate of a chemical reaction without itself undergoing any permanent chemical change. It works by providing an alternative reaction pathway with a lower activation energy compared to the uncatalyzed mechanism. This means that more reactant particles have enough energy to react when a catalyst is present, leading to an increased rate of reaction.

Catalysts are crucial in various industrial and biological processes. In industry, they are used to enhance the efficiency of chemical production, such as in the synthesis of fertilizers, plastics, and fuels. In biological systems, enzymes act as natural catalysts that facilitate essential biochemical reactions necessary for life, such as digestion and metabolism.

Catalysts can be heterogeneous, where they are in a different phase than the reactants, or homogeneous, where they are in the same phase as the reactants. They are also characterized by their ability to be recovered and reused multiple times without losing their effectiveness.

How does a catalyst influence a chemical reaction?

A catalyst is a substance that increases the rate of a chemical reaction without being consumed in the process. It works by providing an alternative reaction pathway with a lower activation energy compared to the uncatalyzed mechanism. Activation energy is the minimum energy required for reactants to undergo a chemical reaction. By lowering this energy barrier, a catalyst allows more reactant particles to have enough energy to react when they collide, thus speeding up the reaction.

Catalysts do not alter the equilibrium position of a reaction; they only make it faster to reach. They are specific in nature, meaning that a particular catalyst will only work for a specific reaction or set of reactions. Catalysts are widely used in industry to increase the efficiency of chemical processes, leading to cost savings and lower energy consumption. In biological systems, enzymes act as natural catalysts to facilitate complex biochemical reactions under mild conditions, which are essential for life processes.

Which component is affected when a catalyst is added to a chemical reaction?

When a catalyst is added to a chemical reaction, it primarily affects the reaction rate. A catalyst speeds up the rate at which the reaction proceeds without being consumed in the process. It does this by providing an alternative reaction pathway with a lower activation energy than the uncatalyzed mechanism. This means that more reactant particles have the necessary energy to undergo the transformation into products at any given time, increasing the frequency of successful collisions. The catalyst remains chemically unchanged after the reaction, so it can continue to facilitate additional reactions. It's important to note that while a catalyst accelerates the reaction rate, it does not alter the equilibrium position of a reversible reaction; it simply helps the system reach equilibrium faster.