1. Home
  2. >Chemistry
  3. >Energetics, Kinetics and Reaction Rate
Help

Activation energy

Get the most by viewing this topic in your current grade. Pick your course now.

?
Intros
Lessons
  1. What happens when molecules have activation energy?
  2. Potential energy changes in a chemical reaction
  3. The transition state / activated complex.
  4. Conditions of a successful collision.
?
Examples
Lessons
    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.
Topic Notes
?

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.

Understanding Potential Energy Curves

Potential energy curves are powerful tools in chemistry that illustrate the energy changes occurring during a chemical reaction. These curves provide a visual representation of how the potential energy of a system changes as reactants transform into products. Understanding potential energy curves is crucial for grasping the fundamental concepts of chemical kinetics and thermodynamics.

A typical potential energy curve plots the potential energy of a system against the reaction coordinate, which represents the progress of the reaction. The curve begins with the reactants at a higher energy level and ends with the products at a lower energy level, assuming an exothermic reaction. The shape of the curve reveals important information about the reaction's energetics and mechanism.

One of the key features of a potential energy curve is the activation energy barrier. This is represented by the peak of the curve and indicates the minimum energy required for the reaction to proceed. The height of this barrier determines the reaction rate a higher barrier means a slower reaction, while a lower barrier results in a faster reaction.

As the reaction progresses along the curve, we can observe the interplay between potential and kinetic energy. At the start of the reaction, the reactants possess mostly potential energy. As they begin to interact and overcome the activation energy barrier, some of this potential energy is converted into kinetic energy. This kinetic energy allows the molecules to collide with sufficient force to break existing bonds and form new ones.

At the top of the activation energy barrier, the system reaches a state known as the transition state. This is a high-energy, unstable configuration where old bonds are in the process of breaking, and new bonds are beginning to form. The transition state represents the point of maximum potential energy and minimum kinetic energy during the reaction.

As the system moves past the transition state and down the other side of the energy barrier, potential energy is further converted into kinetic energy. This release of energy can manifest as heat, light, or other forms of energy, depending on the specific reaction. The products formed at the end of the reaction typically have lower potential energy than the reactants, reflecting the overall energy change of the reaction.

The difference in potential energy between the reactants and products gives us the enthalpy change of the reaction. For exothermic reactions, this difference is negative, indicating a release of energy to the surroundings. Conversely, endothermic reactions show a positive enthalpy change, with the products having higher potential energy than the reactants.

Potential energy curves also help us understand the concept of reaction reversibility. In reversible reactions, the curve will show two activation energy barriers one for the forward reaction and one for the reverse reaction. The relative heights of these barriers determine the equilibrium position of the reaction.

It's important to note that potential energy curves can vary significantly depending on the specific reaction. Some reactions may have multiple peaks and valleys, representing intermediate steps or complex mechanisms. Others might show a gradual change in energy without a distinct activation barrier, typical of some catalyzed reactions.

In the context of chemical kinetics, potential energy curves help explain the temperature dependence of reaction rates. Higher temperatures provide molecules with more kinetic energy, making it easier for them to overcome the activation energy barrier. This is why many reactions proceed faster at elevated temperatures.

Understanding potential energy curves is also crucial in the field of catalysis in chemical reactions. Catalysts work by providing an alternative reaction pathway with a lower activation energy barrier. On a potential energy curve, this would be represented as a new curve with a lower peak, illustrating how catalysts can dramatically increase reaction rates without changing the overall energy change of the reaction.

In conclusion, potential energy curves are invaluable tools for visualizing and understanding the energetics of chemical reactions. They provide insights into reaction mechanisms, rates, and the interplay between kinetic and potential energy throughout the reaction process. By mastering the interpretation of these curves, chemists can gain a deeper understanding of chemical processes and design more efficient reactions and catalysts.

Activation Energy and Reaction Progress

Activation energy plays a crucial role in understanding reaction progress and the mechanisms behind chemical reactions. This concept is fundamental to chemical kinetics and thermodynamics, providing insights into how reactions occur and the energy barriers that must be overcome for a reaction to proceed. To comprehend the relationship between activation energy and reaction progress, it's essential to explore the concepts of activated complex, transition states, and intermediates.

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. As a reaction progresses, the system's energy changes, forming what is known as the reaction coordinate diagram or energy profile. This diagram illustrates the energy changes throughout the reaction, with the activation energy represented as the peak of the energy curve.

The activated complex, also known as the transition state, is a critical concept in understanding reaction progress. It represents the highest energy point along the reaction coordinate, where the reactants have reached a configuration that is neither fully reactants nor fully products. This unstable, short-lived species is at the top of the energy barrier and plays a pivotal role in determining the rate of the reaction. The formation of the activated complex is directly related to the activation energy; the higher the activation energy, the more difficult it is to form the activated complex and, consequently, the slower the reaction rate.

Transition states are crucial in understanding reaction mechanisms. They represent the point of maximum energy in the reaction coordinate, where bonds are in the process of breaking and forming. Unlike intermediates, transition states cannot be isolated or directly observed due to their extremely short lifetimes, typically lasting only femtoseconds. The structure and properties of transition states can be inferred through various experimental techniques and computational methods, providing valuable insights into reaction pathways and mechanisms.

Intermediates, on the other hand, are relatively stable species formed during a reaction that can often be isolated or detected. Unlike transition states, intermediates occupy local energy minima on the reaction coordinate diagram. They have a distinct chemical structure and can exist long enough to be observed or even isolated under certain conditions. Intermediates play a crucial role in multi-step reactions, providing insights into the reaction mechanism and helping to explain the overall reaction rate and product distribution.

To illustrate these concepts, let's consider the example of the SN2 reaction mechanism discussed in the video. In this reaction, a nucleophile attacks a carbon atom bonded to a leaving group. The reaction progresses through a transition state where the nucleophile and leaving group are partially bonded to the central carbon atom. This transition state represents the activated complex, with its formation requiring the system to overcome the activation energy barrier. The SN2 mechanism is a concerted process, meaning it occurs in a single step without forming a stable intermediate. This example highlights the difference between a transition state (the activated complex in the SN2 reaction) and an intermediate, which would be present in a multi-step reaction mechanism.

Another example from the video that illustrates these concepts is the addition of HBr to an alkene. This reaction proceeds through a carbocation intermediate, demonstrating a multi-step mechanism. The first step involves the formation of a carbocation, which is an intermediate species with a distinct structure and lifetime. The reaction then progresses through another transition state as the bromide ion attacks the carbocation to form the final product. This example showcases both transition states and an intermediate, highlighting their roles in reaction progress and mechanism.

Understanding the relationship between activation energy and reaction progress, along with the concepts of activated complex, transition states, and intermediates, is crucial for predicting and controlling chemical reactions. These principles find applications in various fields, including drug design, catalysis, and materials science. By manipulating activation energies through catalysts or changing reaction conditions, chemists can influence reaction rates and selectivity, leading to more efficient and sustainable chemical processes.

Successful Collisions and Reaction Mechanisms

Chemical reactions occur when molecules collide with each other, but not all collisions lead to successful reactions. Understanding the conditions necessary for successful molecular collisions is crucial in comprehending chemical reaction mechanisms and the overall process of chemical transformations. Two key factors play a vital role in determining whether a collision will result in a chemical reaction: sufficient energy and correct orientation of the colliding molecules.

Successful collisions require the molecules to possess enough energy to overcome the activation energy barrier. This energy is typically derived from the kinetic energy of the molecules as they move and collide. When molecules collide with sufficient energy, they can break existing chemical bonds and form new ones, leading to the formation of products. The concept of activation energy is fundamental in understanding why some reactions occur readily while others require additional energy input, such as heat or catalysts.

The correct orientation of molecules during collision is equally important for a successful reaction. Molecules must align in a specific way that allows for the breaking and forming of bonds. This orientation requirement is particularly crucial in complex molecules with multiple reactive sites. Even if molecules collide with sufficient energy, the reaction may not occur if they are not oriented correctly. This aspect of molecular collisions highlights the importance of spatial arrangements in chemical reactions and helps explain why some reactions proceed more slowly than others, despite having favorable energy conditions.

Chemical reaction mechanisms provide a detailed step-by-step description of how a chemical reaction occurs at the molecular level. These mechanisms take into account both the energy requirements and the molecular orientations necessary for successful collisions. By understanding reaction mechanisms, chemists can predict the outcomes of reactions, design more efficient processes, and develop new synthetic routes for creating desired products.

The relationship between successful collisions and activation energy is particularly significant. Activation energy represents the minimum energy required for a reaction to occur. When molecules collide with energy equal to or greater than the activation energy, and in the correct orientation, the reaction is likely to proceed. This concept explains why increasing temperature often accelerates reaction rates higher temperatures increase the average kinetic energy of molecules, leading to more collisions that exceed the activation energy threshold.

In the context of reaction mechanisms, successful collisions often involve the formation of transition states or intermediates. These are high-energy, unstable arrangements of atoms that exist briefly during the reaction process. The energy required to form these transition states is directly related to the activation energy of the reaction. By studying these transition states and intermediates, chemists can gain insights into the reaction pathway and potentially find ways to lower the activation energy, thereby increasing reaction efficiency.

The video content provides clear examples to illustrate these concepts. For instance, it might show animations of molecules colliding in various orientations, demonstrating how only certain alignments lead to successful reactions. It could also visualize the energy changes during a reaction, highlighting the activation energy barrier and how molecules with sufficient kinetic energy can overcome it. These visual representations help to clarify the abstract concepts of molecular collisions and reaction mechanisms.

Another important aspect covered in the video might be the role of catalysts in facilitating successful collisions. Catalysts work by providing an alternative reaction pathway with a lower activation energy, effectively increasing the number of successful collisions without changing the overall energy of the reactants or products. This concept ties together the ideas of collision theory, activation energy, and reaction mechanisms, showing how they are interconnected in the broader context of chemical kinetics.

Understanding successful collisions and reaction mechanisms has practical applications across various fields of chemistry. In organic synthesis, chemists use this knowledge to design reactions that maximize yield and minimize unwanted side products. In biochemistry, it helps explain enzyme function and the specificity of biological reactions. In industrial processes, this understanding guides the optimization of reaction conditions to improve efficiency and reduce costs.

In conclusion, successful molecular collisions leading to chemical reactions depend on both sufficient energy and correct orientation of the colliding molecules. These factors are intimately related to the concepts of activation energy and reaction mechanisms. By exploring these relationships, chemists gain a deeper understanding of chemical processes at the molecular level, enabling them to predict, control, and optimize reactions for a wide range of applications in science and industry.

Factors Affecting Activation Energy

Activation energy plays a crucial role in chemical reactions, determining the minimum energy required for a reaction to occur. Several factors can influence activation energy, ultimately affecting reaction rates. Understanding these factors is essential for comprehending the concepts presented in the video and their practical applications in chemistry.

Temperature is one of the most significant factors affecting activation energy. As temperature increases, molecules gain more kinetic energy, leading to more frequent and energetic collisions. This increased molecular motion reduces the activation energy barrier, making it easier for reactants to overcome and form products. Consequently, higher temperatures generally result in faster reaction rates. The Arrhenius equation quantifies this relationship, showing that reaction rates increase exponentially with temperature.

Catalysts are another critical factor influencing activation energy. These substances participate in reactions without being consumed, providing an alternative reaction pathway with lower activation energy. By lowering the energy barrier, catalysts increase the number of molecules with sufficient energy to react, thus accelerating the reaction rate. Enzymes in biological systems are excellent examples of catalysts that dramatically reduce activation energy for vital biochemical processes.

Concentration of reactants also impacts activation energy indirectly. While it doesn't change the activation energy itself, higher concentrations increase the frequency of molecular collisions. This increased collision frequency enhances the probability of successful reactions, effectively lowering the overall energy barrier for the reaction to proceed. In essence, concentration affects the pre-exponential factor in the Arrhenius equation, influencing reaction rates.

The surface area of solid reactants can influence activation energy in heterogeneous reactions. Increasing the surface area exposes more reactive sites, effectively reducing the energy barrier for the reaction. This principle is often applied in catalytic converters, where precious metals are finely divided to maximize their catalytic efficiency.

Pressure, particularly in gas-phase reactions, can affect activation energy. Higher pressures increase the concentration of gaseous reactants, leading to more frequent collisions and potentially lowering the effective activation energy barrier. This effect is particularly noticeable in reactions involving a change in the number of gas molecules.

The presence of inhibitors can increase activation energy. These substances interact with reactants or intermediates, making it more difficult for the reaction to proceed. By raising the energy barrier, inhibitors slow down reaction rates, a principle utilized in controlling unwanted reactions in various industrial processes.

Changes in activation energy significantly affect reaction rates. A decrease in activation energy, whether through catalysts, temperature increase, or other factors, leads to an exponential increase in reaction rate. This relationship is fundamental to understanding chemical kinetics and is elegantly described by the Arrhenius equation. Conversely, an increase in activation energy dramatically slows down reactions, sometimes to the point of practical cessation.

Relating these concepts to the video, it's crucial to understand that activation energy is not a fixed property but can be manipulated. The video likely emphasizes how various factors can be leveraged to control reaction rates in practical applications. For instance, the use of catalysts in industrial processes to enhance efficiency, or the careful control of temperature in chemical manufacturing to optimize yields.

In conclusion, activation energy is a dynamic concept influenced by multiple factors including temperature, catalysts, concentration, surface area, pressure, and inhibitors. Understanding these influences is key to controlling and predicting chemical reactions, with profound implications in fields ranging from industrial chemistry to biochemistry and environmental science. The interplay between these factors and activation energy forms the foundation of chemical kinetics, enabling scientists and engineers to design more efficient processes and develop innovative solutions to complex chemical challenges.

Applications of Activation Energy in Chemistry

Understanding activation energy concepts has far-reaching implications across various fields of chemistry and related sciences. These applications play a crucial role in designing chemical processes and developing new materials. As explained in the video, activation energy is the minimum energy required for a chemical reaction to occur. This fundamental concept serves as a cornerstone for numerous practical applications in both industry and research.

In the field of catalysis, knowledge of activation energy is essential for developing more efficient catalysts. Catalysts work by lowering the activation energy of a reaction, allowing it to proceed more quickly or at lower temperatures. This principle is applied in the production of many industrial chemicals, such as ammonia in the Haber process, where iron-based catalysts significantly reduce the activation energy required. Similarly, in the automotive industry, catalytic converters use precious metals to lower the activation energy for the conversion of harmful exhaust gases into less harmful substances.

The pharmaceutical industry heavily relies on activation energy concepts in drug design and development. Understanding the activation energy of drug-target interactions helps researchers create more effective medications. For instance, enzyme inhibitors are designed to increase the activation energy of certain biochemical reactions, thereby slowing down or preventing undesired processes in the body. This application of activation energy principles is crucial in developing treatments for various diseases, including cancer and viral infections.

In materials science, activation energy concepts guide the development of new materials with specific properties. For example, in the creation of advanced ceramics, understanding the activation energy of sintering processes allows scientists to optimize manufacturing conditions. This knowledge helps in producing ceramics with enhanced strength, durability, and other desired characteristics. Similarly, in polymer chemistry, activation energy principles are applied to control polymerization reactions, leading to the development of materials with tailored properties for specific applications.

The food industry also benefits from applications of activation energy. In food preservation, techniques such as flash freezing are based on manipulating the activation energy of chemical and biological processes that lead to food spoilage. By rapidly lowering the temperature, the activation energy for these degradation processes becomes insurmountable, effectively preserving the food. Additionally, in the design of food packaging materials, activation energy concepts are used to develop barriers that prevent oxidation and other chemical reactions that could compromise food quality.

These practical applications demonstrate the wide-ranging impact of activation energy concepts across various fields. By understanding and manipulating activation energy, scientists and engineers can design more efficient chemical processes, develop innovative materials, and create solutions to complex problems in areas ranging from environmental protection to healthcare. The fundamental principles explained in the video thus serve as a foundation for numerous advancements in chemistry and related sciences, highlighting the importance of this concept in both theoretical understanding and practical applications.

Conclusion

In conclusion, activation energy plays a crucial role in chemical reactions, determining the energy barrier that reactants must overcome to form products. The introductory video effectively visualizes this concept, illustrating how activation energy impacts reaction rates and the overall feasibility of chemical reactions. Key points discussed include the relationship between activation energy and temperature, catalysts' role in lowering activation energy, and the significance of the Arrhenius equation. Understanding activation energy is essential for predicting and controlling chemical reactions in various fields, from industrial processes to biological systems. Readers are encouraged to apply their newfound knowledge of activation energy to analyze different chemical reactions in their everyday lives and academic pursuits. By exploring this fundamental concept further, one can gain deeper insights into the intricacies of chemical kinetics and thermodynamics, opening doors to innovative applications in science and technology.

Activation Energy in Chemical Reactions

What happens when molecules have activation energy? Potential energy changes in a chemical reaction.

Step 1: Introduction to Activation Energy

Activation energy is a crucial concept in understanding chemical reactions. It refers to the minimum amount of energy required for reactants to transform into products. This energy is necessary to overcome the repulsion between molecules and initiate the reaction. In this section, we will explore the potential energy changes that occur during a chemical reaction, focusing on the transition from reactants to products.

Step 2: Understanding the Potential Energy Curve

The potential energy curve is a graphical representation that shows the energy changes during a chemical reaction. On the y-axis, we have potential energy, while the x-axis represents the reaction progress. The curve helps us visualize the energy states of reactants and products, as well as the transition states in between. Reactants are typically at a stable energy level, and products are also stable but at a different energy level. The transition between these states is where the activation energy plays a critical role.

Step 3: Initial Rise in Potential Energy

As reactant molecules approach each other, their electrons start to repel each other, causing an increase in potential energy. This repulsion reduces the kinetic energy of the molecules, which is then converted into potential energy. This conversion is why we see a rise in the potential energy curve as the reaction progresses from reactants towards the transition state. The faster the molecules move, the more kinetic energy they have, which can be converted into potential energy, increasing the likelihood of reaching the activation energy threshold.

Step 4: Formation of the Activated Complex

When the molecules reach a high enough potential energy, they form an activated complex, also known as the transition state. This is the peak of the potential energy curve. At this point, the chemical bonds in the reactants start to break, and new bonds begin to form, leading to the creation of products. The energy required to reach this state is the activation energy. If the molecules do not have enough energy to reach this state, the reaction will not proceed.

Step 5: Decrease in Potential Energy

After the activated complex is formed, the potential energy begins to decrease as the reaction progresses towards the formation of products. The repulsion between the newly formed product molecules causes them to move away from each other, converting potential energy back into kinetic energy. This decrease in potential energy is reflected in the downward slope of the potential energy curve.

Step 6: Reaching the Ground State

As the product molecules move further apart, the remaining potential energy is fully converted into kinetic energy, and the products reach their ground state. The ground state is the lowest energy state available for the products, where there is minimal repulsion between the molecules. This state is stable and represents the end of the reaction process. The potential energy of the products is typically lower than that of the reactants in an exothermic reaction, indicating that energy has been released during the reaction.

Step 7: Summary of Potential Energy Changes

In summary, the potential energy changes during a chemical reaction can be understood through the potential energy curve. The initial rise in potential energy is due to electron repulsion between reactant molecules, which is converted from kinetic energy. The peak of the curve represents the activated complex, where the activation energy is required to break and form chemical bonds. The subsequent decrease in potential energy occurs as the product molecules move apart, converting potential energy back into kinetic energy. Finally, the products reach their ground state, completing the reaction process.

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.

Prerequisite Topics for Understanding Activation Energy

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.

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.