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Mastering Activation Energy in Chemical Reactions
Dive into the world of activation energy and discover its crucial role in chemical reactions. Learn how it influences reaction rates, the impact of catalysts, and its applications in various fields.

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Now Playing:Activation energy – Example 0a
Intros
  1. What happens when molecules have activation energy?
  2. What happens when molecules have activation energy?
    Potential energy changes in a chemical reaction
  3. What happens when molecules have activation energy?
    The transition state / activated complex.
Examples
    1. Explain using activation energy why incorrect orientation does not lead to successful collisions. Use the potential energy diagram to illustrate your answer.

    2. The reaction between chemicals A and B is exothermic A and is shown by the equation:
      2A + + B → A2_2B

      It is known that the activated complex forms an intermediate structure. Draw a potential energy curve for the reaction showing reactants, products and the intermediate structure.

Introduction to kinetics
Jump to:NotesConceptFAQsPrerequisites
Notes

In this lesson, we will learn:

  • To use the potential energy curve to describe reactant to product transition in chemical reactions.
  • The difference between a transition state and an intermediate in chemical reactions.
  • The definition of successful collision and the conditions for it to occur in a chemical reaction.
Notes:

  • In the 'Introduction to kinetics' (C12.1.1) lesson we introduced activation energy. It is the minimum energy that reactant molecules need to successfully collide and start changing from reactants to products. It is an obstacle that must be overcome for any reaction and rate of reaction to be measured.
    Here, we will talk about what happens to molecules when that obstacle is overcome.
  • The chemical reaction actually occurs at the peak of the curve marked Ea on an energy diagram – when the reactant molecules have enough potential energy to form an activated complex. This can either be a specific intermediate or just a transition state.
    • The transition state is a temporary high energy state seen during the breaking of reactant bonds and forming of product bonds – these are extremely short lived and can't be isolated as chemical substances on their own.
    • An intermediate is an unstable short-lived state the reactants take when reactant bonds have broken, but product bonds have not formed. Not all reactions have intermediates. Unlike transition states, intermediates can sometimes be isolated as they have a specific lifetime, though it may be extremely short, before the products are formed. An intermediate is shown by a small 'well' cut into the peak of the potential energy curve.
  • Potential energy diagrams during reactions can be explained in the following way. Events marked as points on the potential energy curve are highlighted in red:
    • Moving molecules have energy in the form of motion (AKA kinetic energy), so they will collide with each other. Remember that molecules at room temperature and pressure generally collide a few billion times per second!
    • Before they collide, when they get near each other repulsion between electrons in the separate molecules will start happening (in a similar way to when London forces start to form). This reduces the energy of motion in the molecules; this reduced 'kinetic energy' is stored as potential energy in the electrons. This is where an increase in potential energy begins in reactant molecules.
    • The faster the motion of the molecules originally was, the greater the increase in potential energy this will cause. If this rise in potential energy does not match the activation energy, nothing will happen. The molecules will eventually repel back, the excited electrons will repel each other again and send the unchanged reactant molecules back in motion in opposing directions.
    • If the molecules do reach a sufficiently high potential energy then the molecules, still close together, can form an activated complex or transition state. This is where the curve flattens out at the top of the potential energy 'hill'.
    • Once this occurs, the chemical bonds get rearranged. The reactant bonds are broken and the bonds of the product made, in a relatively short time period.
    • The newly made products begin repelling each other by their electrons, moving away from each other. This reduces potential energy in the electrons, which is now being used again as kinetic energy. This is where the curve begins dropping again, as PE decreases and kinetic energy increases.
    • As the molecules get further away, the remaining potential energy from the electron repulsion is 'converted' to kinetic energy until the products reach their 'ground state', their lowest energy state available, like in the reactants before the molecules collided.
  • Molecules that go through the reactant to product process described above have undergone a successful collision. There are two conditions for a successful collision. A molecule must have:
    • Sufficient energy – the activation energy required as explained above.
    • Correct orientation – the molecules have to be arranged in space in an ideal way that the reactant bonds will break and product bonds form when colliding. If they are not, the activation energy for the reaction will be significantly higher. Activation energy on PE curves is shown assuming there is perfect orientation – its lowest possible value.
  • In the lesson on the Boltzmann Distribution (C12.1.5) we saw the distribution of molecules and the kinetic energy they possess. The activation energy barrier is directly related to the rate of reaction through this – the lower the activation energy, the more likely any given molecule will possess enough kinetic energy that can be stored as potential energy when molecules approach and collide. Therefore, the lower the activation energy 'hill', the faster the reaction.
Concept

Introduction to Activation Energy

Activation energy is a crucial concept in understanding chemical reactions. It refers to the minimum energy required for a chemical reaction to occur. This energy barrier must be overcome for reactants to transform into products. The concept of activation energy helps explain why some reactions happen spontaneously while others require external energy input. Our introduction video provides a visual explanation of this fundamental principle, making it easier to grasp. The video illustrates how activation energy affects reaction rates and the role of catalysts in lowering this energy barrier. By understanding activation energy, chemists can predict and control reaction outcomes, which is essential in various fields, from industrial processes to biological systems. The visual representation in the video helps students and enthusiasts alike to better comprehend this abstract concept, making it an invaluable resource for learning about chemical kinetics and thermodynamics.

FAQs

  1. What is activation energy in chemistry?

    Activation energy is the minimum amount of energy required for a chemical reaction to occur. It represents the energy barrier that reactants must overcome to transform into products. This concept is fundamental in understanding reaction rates and the feasibility of chemical reactions.

  2. How does activation energy affect chemical reactions?

    Activation energy directly influences the rate of chemical reactions. Higher activation energy means fewer molecules have sufficient energy to react, resulting in slower reaction rates. Conversely, lower activation energy leads to faster reactions as more molecules can overcome the energy barrier.

  3. What is the role of catalysts in relation to activation energy?

    Catalysts lower the activation energy of a reaction without being consumed in the process. By providing an alternative reaction pathway with lower energy requirements, catalysts increase the rate of reaction without changing the overall energy change of the reaction. This makes catalysts crucial in many industrial and biological processes.

  4. How does temperature affect activation energy?

    While temperature doesn't change the activation energy itself, it affects the number of molecules with sufficient energy to overcome the activation energy barrier. Higher temperatures increase the average kinetic energy of molecules, leading to more successful collisions and faster reaction rates.

  5. What is the significance of the Arrhenius equation in understanding activation energy?

    The Arrhenius equation quantifies the relationship between activation energy, temperature, and reaction rate. It shows that reaction rates increase exponentially with temperature and decrease exponentially with activation energy. This equation is crucial for predicting how changes in conditions will affect reaction rates in various chemical processes.

Prerequisites

To fully grasp the concept of activation energy in chemistry, it's crucial to have a solid foundation in several prerequisite topics. Understanding these fundamental concepts will greatly enhance your ability to comprehend and apply the principles of activation energy in various chemical processes.

One of the key prerequisites is reaction mechanisms. These chemical reaction mechanisms provide insight into how molecules interact and transform during a reaction. By understanding the step-by-step processes involved in chemical reactions, you'll be better equipped to appreciate the role of activation energy in overcoming energy barriers.

Another important concept to grasp is electric potential and electric potential energy. This knowledge is particularly relevant when considering potential energy curves in the context of activation energy. These curves help visualize the energy changes that occur during chemical reactions and highlight the energy barrier that must be overcome for a reaction to proceed.

Understanding rotational kinetic energy and angular momentum is also beneficial when studying activation energy. This knowledge relates to kinetic energy and molecular collisions, which play a crucial role in providing the necessary energy for reactions to occur. The kinetic energy of colliding molecules must be sufficient to overcome the activation energy barrier for a reaction to take place.

Lastly, a solid introduction to chemical reactions is essential. This foundational knowledge helps in understanding how catalysts work in chemical reactions. Catalysts are substances that lower the activation energy of a reaction, allowing it to proceed more quickly or at lower temperatures. By grasping the basics of chemical reactions, you'll be better prepared to explore the role of activation energy and how it can be manipulated in various chemical processes.

By mastering these prerequisite topics, you'll develop a comprehensive understanding of the factors influencing activation energy and its significance in chemical reactions. This knowledge will enable you to analyze reaction rates, predict reaction outcomes, and appreciate the intricate balance of energy in chemical processes. As you delve deeper into the study of activation energy, you'll find that these foundational concepts continually resurface, reinforcing their importance in your chemistry education.