1. Chemical Equilibrium Honors Chapter 16

2. Equalized • All reactions want stability. This is obtained by chemicals moving back and forth until the reaction is stable. • Equilibria – plural for Equilibrium

3. What is Equilibrium Constant • All reactions contain a limited number of atoms in each direction. • Reactions can go Forward or in Reverse A + B AB *Remember: Vapor Pressure is noted by the number of molecules at equilibrium in the liquid or gas state.

4. Graph to Equilibrium [A] Concentration Equilibrium [B] Time

5. Rate of the Reactions Rate of the Forward Reaction Rate Rate of the Reverse Reaction Equilibrium Time

6. What is equilibrium?…… • At equilibrium the reaction is still proceeding • However, it is proceeding at the same rate in both directions (note the double arrow) • rate to the right = the rate to the left • Therefore, while individual molecules will react, the overall concentration of any reagent will remain constant…. • Some example reactions……... • THE REACTION DOES NOT STOP, keeps moving in both directions

7. The Law of Chemical Equilibrium • Remember: The reversible reaction achieves equilibrium when the rate of the forward reaction is equal to the rate of the reverse. • LAW OF MASS ACTION: the relative concentrations of reactants and products at equilibrium in terms of quantity called equilibrium constant.

8. Numerical (unitless) value used to evaluate the “status” or state of a reaction Determine which way it will progress Useful in all areas of chemistry and many other sciences Several different variations The Equilibrium Constant (K or Keq)

9. (homogenous reactions) Solution concentrations are in Molarity • (homogenous) Gas concentrations are in pressures (usually atm, but can be defined in other units) • (heterogenous) Solids have “unitless” concentrations of 1 • When the calculated K is at the “standard value”, the reaction is at equilibrium. • There is a unique K for each reaction and each temperature • Unique at pressures too, since they may vary concentrations!

10. Keq aA + bB cC + dD a, b, c, d represent coefficients ABCD represent compounds Keq = [C]c [D]d [A]a [B]b

11. What is the equilibrium expression for the reaction (homogenous reaction)2CO + O2 2CO2 Keq = [CO2]2 [CO]2[O2] Why no numbers? Because we do not know the Molarity of the compounds

12. What is the equilibrium expression for the homogenous reaction2CO + O2 2CO2 Keq = [CO2]2 [CO]2[O2] What if we have the molarity: Carbon Dioxide = 0.0100 M Carbon Monoxide = 0.0200 M Oxygen = 0.0200 M

13. What is the equilibrium expression for the Homogeneous reaction2CO + O2 2CO2 Keq = [0.0100M]2 [0.0200M]2[0.0200M] Keq = 12.5

14. Heterogenous reaction C (s) + H2O (g) CO (g) + H2 (g) Carbon is a solid and can not be measured in molarity. In this case, Carbon has a unit of 1. Keq = [CO][H2] [H2O]

15. Homogenous equilibrium: look at all compounds Heterogenous equilibrium: all solids have a unit of 1 since they can not be measured in M. Keq >> 1 the reaction is towards the products Keq << 1 the reaction is towards the reactants (barely moves) Manipulating Equilibrium Constants……...

16. If we disturb a reaction at equilibrium 1. change in Temperature 2. Change in Pressure 3. change in Concentration The reaction rates will shift to try to re-establish equilibrium concentrations of all reagents the rate in one direction will exceed the other Le Chatelier’s Principle

17. The Reaction Quotient • The reaction quotient (Q) is used to determine if the reaction is at equilibrium. • If Q < Keq then the equation falls toward the products. This means that reactant will react and produce more product, and less reactants are left over compared to equilibrium. • If Q > Keq then the equation falls toward the reactants. This means that more reactants are left over than products that could be made. • If Q = Keq then the equation is at equilibrium.

18. Changes in Concentration • All reactions will shift in reverse or forward to achieve equilibrium. Forward   Backward • If you add a concentration of a compound to the reaction, the reaction will try to balance out. Example: NO2  N2O4 If you add N2O4 to a stable reaction, the reaction will shift in reverse, changing N2O4 into NO2 until equalized.

19. Change in Pressure Take the equation 2NO2  N2O4 There are 2 moles to 1 mole ratio If the equation is at equilibrium, and pressure is applied decreasing the volume. Now the reaction must move in the direction to decrease the number of moles * The reaction will move in the forward dirction. If pressure is decrease, there is more volume. The reaction will move in the direction to increase the number of moles, it will now move in reverse.

20. Change in Temperature • Relates to Exothermic or endothermic reactions. • If a reaction is exothermic in one direction, then it is endothermic in the reverse. • If the reaction is exothermic (gives off heat) and heat is added, the reaction will move toward the reactants. • If the reaction is endothermic, and the reaction is cooled the reaction will move toward the products.