Equilibrium Packet – General Chemistry

What is an equilibrium?

  • Up until now, we’ve treated all chemical reactions as if they’ve only proceeded in one direction, from A  B.
  • In the real world, reactions sometimes proceed in both directions (these are called reversible reactions). Instead of just going from A  B, they also go from B  A. When this happens, these are called equilibria.
  • Equilibria are expressed with a double-headed arrow. In the example of A  B and B  A, we write the equilibrium expression as A  B.

What are the properties of equilibria?

  • The rate of the forward reaction is the same as the rate of the reverse reaction.
  • Consider the equilibrium A  B. For this reaction to be at equilibrium, the rate of the reaction A  B has to be the same as the rate of the reaction B  A.
  • At equilibrium, the concentration of all the chemicals stays constant.
  • For a system to be at equilibrium, the concentrations of all of the chemicals must stay constant. (If they didn’t, that would suggest that the forward and reverse reaction rates weren’t the same).
  • Equilibria are dynamic processes.
  • Because the concentrations of the products and reactants don’t change in an equilibrium, it’s tempting to think that the reactions stop. They don’t.
  • The reason the concentrations of product and reactant don’t change is that the forward and reverse reactions occur at the same rate. For every molecule of A that converts to B, a molecule of B converts to A. As a result, the concentrations don’t change even though the reactions are still taking place.

You can change the position of an equilibrium (how much A and B are present) by using “Le Châtlier’s principle”:

  • As you know, chemists are interested not only in explaining why chemical reactions occur, but also how they can make the reactions proceed at higher rates.
  • We explain how chemical equilibria can be changed using Le Châtlier’s principle.
  • Definition: Whenever you change anything about an equilibrium, the equilibrium will shift to minimize the effect of whatever you did. This shift, while not returning the equilibrium to its original position, will establish a new equilibrium in which more of the product we want is formed.

Explaining Le Châtlier’s principle using examples: All of the examples below will be assuming that we use the reaction: A(s) + B(l) 2 C(g) ∆H = -100 kJ

  • Adding reactants pushes the equilibrium to the right.
  • If we add A to the reaction mixture, the equilibrium will want to get rid of it. Since the only way to do this is to use it up by making C, the equilibrium will shift toward the products.
  • Likewise, if we (for some reason) wanted to shift the equilibrium to the left, we’d add more C.
  • Removing products pushes the equilibrium to the right.
  • If we were to take away C as it’s formed, the equilibrium will want to replace the lost C by forming more of it. As a result, the equilibrium shifts toward products.
  • Likewise, if we (for some reason) wanted to shift the equilibrium to the left, we’d remove either A or B.
  • Changing the volume of the container in which the reaction is being performed will cause the equilibrium to shift in a way that decreases the pressure.
  • If we were to increase the pressure inside this container, the reaction will be pushed toward the reactants. This is because C is a gas and if C is converted to A and B (neither of which is a gas), the pressure inside the reaction vessel will be decreased.
  • Likewise, if we wanted to make more C, the way to do it would be to decrease the pressure inside the container so that more C will be formed to bring it back up.
  • Changing the temperature will shift an equilibrium such that the temperature decreases.
  • Instead of writing the heat of reaction as “∆H = [whatever]” after the equation, we can instead treat the energy as being either a product (if the reaction is exothermic) or as being a reactant (if the reaction is endothermic). When we do this we can see that changing the temperature will have the same effect as changing the concentrations of either products or reactants.
  • In our example, the reaction has an overall ∆H = -100 kJ. As a result, we can rewrite the equation as:

A(s) + B(l) 2 C(g) + 100 kJ

As a result, increasing the temperature (and hence, the energy)

will cause the equilibrium to be pushed to the left.

  • Likewise, we can push the equilibrium to the right by taking away energy (this will have the same effect as having removed a product).

Equilibrium Practice Problems – General

1)Explain why the reaction CH4 + 2 O2 CO2 + 2 H2O is normally not considered an equilibrium.

2)What is happening on a molecular level during an equilibrium process?

3)For the equilibrium A(g) 2 B(g) in which the process A  2 B is exothermic, sketch the energy diagram, including the products, reactants, ∆H, activation energy (for both the forward and reverse reaction).

4)For the reaction in problem 3, what would happen to the concentrations of A and B if we increased the pressure in the container that held the reaction? Explain why this happens.

5)For the reaction: 2 N2(g) + O2(g) 2 N2O(g)∆H = -

What would happen if I:

  • Increased the temperature?
  • Decreased the pressure by increasing the volume of the container?
  • Removed O2 from the mixture?
  • Added 0.50 atm of Xe gas?

6)For the reaction CH4(g) + 2 O2(g) CO2(g) + 2 H2O(g) ∆H = -

What would happen if I:

  • Decreased the temperature?
  • Increased the pressure by decreasing the volume of the container?
  • Added additional methane to the reaction?
  • Added a catalyst?

Honorh