Dynamic equilibrium

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Intros
Lessons
  1. What is equilibrium?
  2. Can reactions reverse?
  3. Open and closed systems.
  4. Dynamic equilibrium.
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Examples
Lessons
  1. Understand how equilibrium and reversible reactions occur.
    The reaction between substances A and B to produce C and D is described below in an equation.

    A (g) + 2B (g)  \,   \, C (g) + D (g)

    The reaction takes place at high temperature and pressure with the container sealed. 
    1. Explain how sealing the reaction container could establish an equilibrium.
    2. Explain why leaving this reaction unsealed creates other practical problems.
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    Topic Notes
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    In this lesson, we will learn:

    • The definition of reversible reaction and dynamic equilibrium.
    • How the open and closed state of a system affect equilibrium.

    Notes:

    • We know chemical reactions as going from reactants to products, but many chemical reactions can go from products 'back' to reactants. Reactions which can go 'both ways' are called reversible reactions.
      • In the kinetics chapter, we learned about the activation energy barrier preventing reactants from forming products in chemical reactions. For a chemical change to occur, reactant particles need sufficient energy and correct orientation when colliding. These are the conditions of a successful collision.
        As long as the conditions for a successful collision are met, there is no reason why the transformation cannot go in the other way too! All that is needed is a certain activation energy.
    • In reversible reactions there are terms given to the 'direction' of the reaction which will be used in this chapter:
      • The forward reaction is the chemical change from reactants to products with respect to a given chemical reaction.
      • The reverse (aka backward) reaction is the reverse change from products to reactants.
    • In many cases, reversible reactions do not seem to be reversible because they are performed in open systems. Two more definitions for this chapter are below:
      • An open system is an environment where other substances or energy e.g. heat or light can enter and leave.
      • A closed system is an environment where substances and/or energy cannot enter and leave.
        • When a reaction takes place in an open system, the products escape or are removed from the reaction vessel to proceed with their intended use. The products are therefore removed from the conditions that could cause the reverse reaction to occur, and without the products available, the system will not be able to make the reverse reaction happen!
        • When a reaction takes place in a closed system, the products of the desired forward reaction cannot escape. This is often done when the desired products are gases so they are trapped in the reaction vessel and won't be lost. However, the products of the forward reaction are the reactants of the backward reaction – so this can start occurring!
    • Under some conditions, the rate of the reverse reaction equals the rate of the forward reaction, creating a balanced system of constant change. This is called dynamic equilibrium. This sometimes creates the appearance that the reaction has "stopped" but it has not – it is simply making products as quickly as it is re-making reactants, so the amounts of each do not change!
      • An analogy of this effect is filling a swimming pool which has a hole in it that is leaking water. If the pool is being filled by a hose at the same rate it's being drained by the hole, it is at equilibrium – constantly changing in both ways at the same rate!