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Understanding Le Chatelier's Principle in Chemistry
Dive into Le Chatelier's Principle and master the art of predicting chemical equilibrium shifts. Learn how temperature, pressure, and concentration changes affect exothermic and endothermic reactions in real-world applications.

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Now Playing:Le chateliers principle – Example 0a
Intros
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  1. What happens at equilibrium?
  2. What happens at equilibrium?
    Recall dynamic equilibrium.
  3. What happens at equilibrium?
    What happens when changing conditions at equilibrium?
Examples
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  1. Apply Le Chatelier's principle to predict changes in equilibrium position.
    The reaction equation below shows an EXOTHERMIC reaction at equilibrium:

    2A(g) + B(g)  ⇌ \, \rightleftharpoons \, C(g) + D(g) \quad △H \triangle H = -89 kJ mol-1

    Predict and explain the change in equilibrium position with the following changes of conditions happening separately:
    1. An increase in pressure.
    2. A decrease in temperature.
    3. Addition of a catalyst.
    Introduction to dynamic equilibrium
    Jump to:NotesConceptFAQsPrerequisites
    Notes

    In this lesson, we will learn:

    • To recall Le Chatelier's principle when studying equilibria.
    • How to predict changes to equilibria given changes in reaction conditions.
    • How to explain changes in equilibrium position due to changes in reaction conditions.

    Notes:

    • Continuing from Dynamic Equilibrium, when more chemical reactions were studied, and more equilibria were found, chemists started changing the conditions (such as temperature and pressure) of the reaction to see what this did to the equilibrium they saw.
      This led to them finding Le Chatelier's principle:
      • When a reaction at equilibrium is disturbed by a change in conditions, the system will respond in a way that COUNTERACTS the disturbance and then re-establish a new equilibrium.

    • The system 'counteracts the disturbance' by favoring the forward or the backward reaction; it will do the opposite of whatever effect the disturbance had on the reaction! 'Favoring' a reaction means that reaction rate increases (causing a change in the ratio of reactants to products in the system), before a new equilibrium is reached as the two rates become equal again. Once this happens, the new ratio of reactant to product concentration won’t change unless equilibrium is disturbed again - in this case, the above happens again.
      • Remember, 'at equilibrium' only tells us that the forward and reverse reactions are happening at the same rate. Equilibrium does not tell us how much product and reactant there is in the system!

    • When a reaction at equilibrium is disturbed, it will work to re-establish equilibrium under these new conditions. If some facts about the reaction are known, the change to equilibrium position can be predicted from a change in conditions:
      • If temperature is increased:
        • In an exothermic reaction: The system will respond to favor the reverse reaction which leads to a greater concentration of reactants being observed in the reaction. An exothermic forward reaction, by definition, causes a net release of heat energy to the surroundings. The system therefore counteracts the disturbance of the temperature increase by temporarily favoring the reverse reaction, which will be the endothermic opposite of the exothermic forward reaction. This endothermic reverse reaction has a net absorbing of heat effect, absorbing the increase in temperature that initially disturbed the system. This is how Le Chatelier's principle works! When the reverse reaction is favored, we say the equilibrium has shifted to the left as the left-hand side of the reaction equation shows the reactants.
        • In an endothermic reaction: The system will respond by favoring the forward reaction. This is for the same reasons as above; the system will favor the endothermic forward reaction, which absorbs heat and counteracts the increase in temperature that it was disturbed by. Favoring the forward reaction means the equilibrium shifts to the right. In this case, more products will be produced in the system.
      • If temperature is decreased:
        • In an exothermic reaction: The system will respond to favor the forward reaction. This is because the exothermic forward reaction has a net releasing of heat effect, which will raise the temperature and counteract the decrease in temperature that initially disturbed it. In this case we say the equilibrium has shifted to the right and more products will be produced.
        • In an endothermic reaction: The system will respond to favor the reverse reaction. This is because the reverse reaction will be exothermic, causing a net release of heat energy to the system, raising the temperature that initially disturbed it. The equilibrium will shift to the left and more reactants will be present in the system.

      • If pressure is increased, count the number of moles of gas on both sides of the reaction equation:
        • When there are more moles of gas in the products than the reactants, the system will respond to favor the reverse reaction and the equilibrium will shift left. This has the effect of 'cutting back' the number of gas particles in the system as there are more of them in the products than the reactants. This leads to less gas particle collisions which decreases pressure, which counteracts the initial disturbance of a rise in pressure!
        • When there are less moles of gas in the products than the reactants, the system will respond to favor the forward reaction and the equilibrium will shift right. This counteracts the disturbance in the same way as explained above.
      • If pressure is decreased, count the number of moles of gas on both sides of the reaction equation:
        • When there are more moles of gas in the products than the reactants, the system will respond to favor the forward reaction and the equilibrium will shift right This has the effect of increasing the number of gas particles in the system as there are more gas moles in the products than the reactants. This leads to more gas particle collisions which increases pressure, counteracting the initial disturbance of a pressure decrease!
        • When there are less moles of gas in the products than the reactants, the system will favor the reverse reaction and the equilibrium will shift left. This leads to an increase in the number of gas particles in the system because there are more gas moles in the reactants than the products. This counteracts the initial disturbance of the pressure decrease.
      • If a catalyst is added to the reaction mixture, there is no change to the position of equilibrium. A catalyst does not change equilibrium; it simply allows equilibrium to be reached quicker.

    • The table below summarizes the changes to equilibrium position caused by a change in conditions:

      • Le Chatelier's principle dynamic equilibrium

    • Even though chemists know how to change a reaction at equilibrium to make the largest amount of product, there are practical issues with shifting the equilibrium.
      • Increasing the pressure on a closed system can be very expensive and there are safety considerations at very high pressure.
      • Many reactions, such as the Haber process, have an exothermic forward reaction. This strangely means that to shift the equilibrium to make more product, you need to cool the reaction down which will reduce reaction rate.
      Because of these and other factors, compromise conditions are often used in establishing the ideal conditions for a reaction at equilibrium.
    Concept

    Introduction to Le Chatelier's Principle

    Welcome to our exploration of Le Chatelier's Principle, a fundamental concept in chemistry that helps us understand how chemical systems respond to changes. This principle, named after French chemist Henry Louis Le Chatelier, is crucial for predicting the behavior of reactions at dynamic equilibrium. When we alter reaction conditions such as temperature, pressure, or concentration, Le Chatelier's Principle tells us how the system will adjust to counteract these changes. Our introductory video will guide you through this fascinating principle, making it easier to grasp its applications in real-world scenarios. You'll learn how chemists use this principle to optimize industrial processes and how it applies to natural phenomena. By understanding Le Chatelier's Principle, you'll gain valuable insights into the delicate balance of chemical reactions and how small changes can have significant effects. This knowledge is essential for anyone studying chemistry or related fields, providing a powerful tool for analyzing and controlling chemical systems.

    FAQs

    Here are some frequently asked questions about Le Chatelier's Principle:

    1. What is a forward reaction in chemical equilibrium?

      A forward reaction in chemical equilibrium is the process where reactants combine to form products. For example, in the reaction A + B C + D, the forward reaction is A + B C + D. At equilibrium, the forward reaction rate equals the reverse reaction rate.

    2. What happens to an exothermic reaction at equilibrium when heated?

      When an exothermic reaction at equilibrium is heated, Le Chatelier's Principle predicts that the equilibrium will shift towards the reactants (left) to counteract the temperature increase. This shift absorbs some of the added heat, as the reverse reaction is endothermic.

    3. How does equilibrium shift when temperature is increased for an endothermic reaction?

      For an endothermic reaction, increasing the temperature shifts the equilibrium towards the products (right). This shift favors the forward reaction, which absorbs heat, thus counteracting the temperature increase in accordance with Le Chatelier's Principle.

    4. What factors can favor a reverse reaction?

      Factors that can favor a reverse reaction include: decreasing the concentration of products, increasing the concentration of reactants, changing temperature (cooling for exothermic reactions, heating for endothermic reactions), and changing pressure (for gas reactions with more moles on the reactant side).

    5. How can you determine if a reaction is exothermic or endothermic at equilibrium?

      To determine if a reaction is exothermic or endothermic at equilibrium, observe how the equilibrium shifts with temperature changes. If increasing temperature favors the reverse reaction (shifts left), the forward reaction is exothermic. If increasing temperature favors the forward reaction (shifts right), the forward reaction is endothermic.

    Prerequisites

    To fully grasp Le Chatelier's principle, it's crucial to have a solid foundation in several key chemistry concepts. One of the most fundamental prerequisites is understanding dynamic equilibrium. This concept forms the basis of Le Chatelier's principle, as it describes the state in which forward and reverse reactions occur at equal rates, resulting in no net change in the concentrations of reactants and products.

    Another essential prerequisite is knowledge of the factors affecting reaction rate. Le Chatelier's principle deals with how chemical systems respond to changes in conditions, and understanding these factors is crucial for predicting and explaining these responses. This ties in closely with the concept of activation energy, which plays a significant role in determining how easily a reaction can proceed and how it might be affected by changes in temperature or the addition of a catalyst.

    A practical application of Le Chatelier's principle can be seen in the Haber process for ammonia production. This industrial process relies heavily on manipulating equilibrium conditions to maximize yield, making it an excellent real-world example of Le Chatelier's principle in action.

    Having a solid introduction to chemical reactions is also vital. This includes understanding how catalysts work in chemical reactions, as catalysts can significantly impact the rate at which equilibrium is established without affecting the equilibrium position itself.

    Lastly, knowledge of predicting chemical reactions using cell potential can provide valuable insights into the direction of spontaneous change in chemical systems, which is closely related to the principles outlined by Le Chatelier.

    By mastering these prerequisite topics, students will be well-equipped to understand and apply Le Chatelier's principle. This principle is not just an isolated concept but a powerful tool that integrates various aspects of chemical equilibrium and kinetics. It allows chemists to predict how chemical systems will respond to changes in concentration, pressure, temperature, and other factors.

    Understanding Le Chatelier's principle is crucial for many areas of chemistry, from industrial processes to environmental science. It helps explain phenomena such as why carbonated drinks fizz more when warm, how blood maintains its pH, and how to optimize industrial reactions for maximum efficiency. By building a strong foundation in these prerequisite topics, students will be better prepared to tackle more complex chemical systems and real-world applications of equilibrium principles.