Question

Draw an energy diagram for a reaction with and without the presence of a catalyst that increases the rate of the reaction.

Short answer

In the energy diagrams, both reactions start and end at the same energy levels representing the energy of the reactants and products. The only difference is the peak (the activation energy) which is lower when a catalyst is present, thus indicating that the catalyst helps to increase the rate of the reaction.

Step-by-step solution

Drawing Energy Diagram without a Catalyst

Start by drawing an energy diagram for the reaction in absence of a catalyst. Label the vertical axis as 'Potential Energy' and the horizontal one as 'Reaction Progress'. Depict reactants at a certain energy level, and then show the products at a different energy level. Also, include a peak to represent the activation energy (E_a), which is the highest point in the diagram between reactants and products.

Drawing Energy Diagram with a Catalyst

Draw a similar energy diagram for the reaction with a catalyst. Label axes identically as before. The reactants and products will be at the exact same energy levels. However, the activation energy in this case will be lower, as the catalyst lowers the energy barrier for the reaction. Depict this by drawing a lower peak for the activation energy (E_a') on the energy diagram.

Comparison

Observe both diagrams side by side. The key difference to note is that the second diagram shows a decrease in the height of the energy barrier (activation energy), signaling an increased reaction rate in the presence of the catalyst.

Key concepts

Activation Energy

Activation energy is a crucial concept in chemistry, often depicted in energy diagrams. It represents the minimum amount of energy that reactant molecules must possess for a reaction to occur. Imagine trying to push a ball over a hill. The effort to get the ball to the top is similar to the energy required to kick-start a chemical reaction. In the energy diagram, activation energy is shown as the peak height that separates the reactants and the products. The higher the peak, the more energy is required for the reaction to proceed. In the absence of a catalyst, this energy barrier might be quite high, making the reaction slower or even preventing it from happening under certain conditions. Efficient reactions prefer a low activation energy, which allows for a greater number of molecules to have enough energy to react upon collision.

Potential Energy

Potential energy in the context of chemical reactions refers to the stored energy in reactants and products due to their positions, compositions, and arrangements of atoms. In energy diagrams, reactants are typically illustrated at one level of potential energy while products are at another. If the products lie at a lower energy level than the reactants, the reaction is exothermic, releasing energy. On the flip side, if the products are at a higher energy level, the reaction absorbs energy and is endothermic. The start and end points on the energy diagram reflect the potential energy of the reactants and products, respectively. This visual comparison helps students understand the energy changes during a reaction, influencing reactions' spontaneity and feasibility.

Reaction Progress

Reaction progress is another term that's often visualized on an energy diagram and illustrates the journey from reactants to products over time. On the horizontal axis of the diagram, the initial point represents the reactants before the reaction begins. As you move to the right, the progress of the reaction is shown. At the top of the activation energy peak, the reaction reaches its transition state—where old bonds are breaking and new bonds are forming. Following this peak, the diagram slopes down toward the energy level of the products, indicating the conclusion of the reaction. Essential for understanding the flow of the reaction, the reaction progress axis aids students in visualizing the step-by-step process of chemical transformations.

Catalyst Effects on Reactions

In chemistry, a catalyst is like a facilitator for a reaction, reducing the activation energy needed without being consumed in the process. Similar to providing a ramp for our aforementioned ball to easily roll over the hill, a catalyst offers an alternative pathway for the reaction with a lower energy barrier. This is depicted in an energy diagram as a peak that is noticeably lower than the one without a catalyst. The presence of a catalyst does not affect the initial or final energy states; it solely decreases the activation energy, meaning the same reaction can occur more readily. This catalytic influence can lead to a significant increase in reaction rates, which is why catalysts are invaluable in many industrial processes, biological systems, and research applications.