Le Châtelier’s Principle

1. Le Châtelier’s Principle SCH4U0

2. Dynamic Equilibrium • We mentioned earlier that equilibria are dynamic, but what does this mean? • It means that the equilibrium can change over time • In practice this means that the “position” of the equilibrium can change if the conditions are changed • Meaning that the equilibrium concentrations (and sometimes K) will change if certain changes are made to the system • Le Châtelier came up with a simple principle to predict how an equilibrium would respond to specific changes

3. Le Châtelier’s Principle • Le Châtelier’s principle says that equilibria will respond to a stress by countering its effect • A stress is a change in concentration, temperature, or pressure • This means that if a reactant’s concentration is increased, the equilibrium will shift to lower that reactant’s concentration • Example: • If we add more into the system, the equilibrium concentrations will change so that the reactant concentration lowers and the product concentrations increase

4. Equilibrium Shift • We can see this while looking at the equilibrium law • If equilibrium has been established, and we add in more , we will no longer be at equilibrium • K is constant, so the concentrations alter to keep K the same Change by 0.99 M Increase to 4 M

5. Equilibrium Shift • When we altered the concentration of a reactant, the concentrations of everything else changed to establish a new equilibrium • With the same ratio of reactant/product concentrations • We say that the equilibrium has shifted to the right • Meaning that more products were made, and the reactants were used up a little more • If an equilibrium shifts to the left, the opposite occurs • The reactant concentrations increase, and the product concentrations lower

6. Using Le Châtelier’s Principle • Let’s use Le Châtelier’s principle to determine what is going to happen to the following equilibria when they are stressed • If we add more to the following systems, in which direction will they shift (or, what will happen to the )? Shift right or Decrease Shift left or Decrease

7. Pressure Changes • When the partial pressures of gaseous reactants/products is altered, the equilibrium will also shift. • Le Châtelier says that the EQ will shift to counter the pressure change • To increase P, the EQ will shift to the side with more moles of gas • Example • What happens when the total volume of the following systems is lowered (thus increasing all of the partial pressures via Boyle’s law- ) This will lower the total moles of gas – and lower the total pressure This will lower the total moles of gas – and lower the total pressure Shift left or Decrease Shift right or Decrease

8. Temperature Changes • Heat can be treated as a reactant/product in chemical reactions • Thus, if the temperature is altered, an equilibrium will shift to counter the change in heat • Examples: • What happens to the following equilibria when the temperature of the system is increased? This will take heat from the surroundings, thus cooling it This will take heat from the surroundings, thus cooling it Shift right or increase Shift left or Decrease

9. Changes That Don’t Alter the EQ • Some changes to the system do not have any effect on the equilibrium concentrations: • Catalysts • Adding catalysts does not alter the EQ concentrations • Inert Gases • Adding inert gases does not alter the partial pressures of the reactants/products, only the total pressure • These gases do not effect the EQ concentrations • Inert in this context means that it doesn’t react with anything in the system. The gases do not have to be inert with everything

10. Virtual Lab • At this point we will conduct a virtual lab to explore the use of Le Châtelier’s Principle Virtual Lab

11. Explaining Dynamic Equilibrium • Let’s discuss the underlying reasons behind the shifts in equilibrium predicted by Le Châtelier’s Principle • To analyze the underlying effects we will need to look at the rate laws of the forward/backward reactions and consider; • Activation energies • Collision probabilities • Average kinetic energies (Maxwell-Boltzmann distributions) • Concentrations • Rates

12. Concentration Changes • Concentration changes are fairly simple to explain, since the concentration of reactants/products is found in the rate laws • Example: • What happens when we increase the chlorine concentration? • To explain this, we need to analyze the rate laws Shift right or decrease

13. Concentration Changes • If we take this system at equilibrium and then increase the chlorine concentration; • The rate of the forward reaction will increase • The rate of the reverse reaction will be unchanged • We will no longer be at equilibrium! • The reactant concentrations will then decrease until the forward rate is again equal to the reverse rate

14. Pressure Changes • Pressure changes can be analyzed in a similar way to concentration, we just need to remember the ideal gas law • This means that for a single gas, its partial pressure is analogous to its concentration • Assuming T is constant • Because P and C are analogous, we can re-write our rate laws in terms of pressure

15. Pressure Changes

16. Pressure Changes • Now that we have the rate laws in terms of pressure we can analyze what happens when we alter the partial pressures • We already saw what happens when we add more/less of a substance with concentration changes • What happens when we lower the volume of the container? • All the gases will increase in concentration and partial pressure • Lets assume the volume is cut in half; all the pressures will double

17. Pressure Changes • So, when the volume was cut in half, all of the partial pressures doubled • This caused both rates to increase • But the forward rate increased by a larger amount than the reverse rate • Since the forward rate is faster; • The reactants will go down in concentration • The forward rate will slow, the reverse will speed up, until they are equal • Because the side with more moles of gas has larger exponents, it will always have a larger change in rate when the volume changes • Since all pressure will change by the same amount

18. Temperature Changes • Explaining the temperature change observations is a little trickier, but as long as we understand the distribution of particle kinetic energies in a system we can explain the EQ shifts. • Increasing the temperature will increase the rate of both reactions (forward and reverse) • But not by the same factor • The endothermic reaction will increase by a larger factor than the exothermic reaction

19. Temperature Changes • We should be able to tell that in a reaction, the exothermic direction will always have a smaller activation energy than the endothermic direction • Forward is exothermic • Reverse is endothermic

20. Temperature Changes Exothermic Endothermic ~50% ~75% ~5% ~50% T1 (green) < T2 (blue) There are 10 times more particles with enough energy to react There are 1.5 times more particles with enough energy to react

21. Temperature Changes • We saw from the Maxwell-Boltzmann distributions that when the temperature increases, the endothermic process increases the number of particles able to react by a larger fraction • This translates to a larger increase in rate (and rate constant) • Let’s assume the factors we saw on the graphs translate directly to rate increases • The reverse reaction is now faster than the forward reaction Endothermic Exothermic

22. Temperature Changes • The system was initially at equilibrium. After the temperature increase • The reverse reaction is faster than the forward reaction • This will cause the product concentrations to go down • They will be used faster than they are made • This will continue until; • The backward reaction slows down • The forward reaction speeds up • And both rates become equal again (equilibrium!)