2. Reversible Reactions 6-1
Reactions are spontaneous if DG is negative.
If DG is positive the reaction happens in the opposite direction.
2H2(g) + O2(g) ® 2H2O(g) + energy
2H2O(g) + energy ® 2H2(g) + O2(g)
2H2(g) + O2(g) H2O(g) + energy

3. Equilibrium
Forward reaction
H2O(g) + CO (g\ H2 (g) + CO2 (g)
Reverse reaction

4. Dynamic Equilibrium:
If kept in a closed system, then both reactions will occur. If this reaction is left then it will reach a point of equilibrium where the rate of the forward reaction equals the rate of the reverse reaction. So the concentrations of the reactants doesn’t change.
The equilibrium is dynamic, because even though the concentrations stay the same and nothing appears to be happening, both the forward and reverse reactions are continuing to occur.

5. [ ] of A decreases while [ ] of B increases till equilibrium is reached.
Equilibrium is reached when rate of forward reaction is the same as the reverse reaction.

6. Changes in reaction rates of the forward and reverse reactions for:
H2O + CO H2 +CO2
Rate of forward reaction decreases while reverse increases till the concentrations reach a level at which the rate of the forward and reverse reactions is the same. The system has reached equilibrium.

7. Equilibrium
When you first put reactants together the for ward reaction starts.
Since there are no products there is no reverse reaction.
As the forward reaction proceeds the reactants are used up so the forward reaction slows.
The products build up, and the reverse reaction speeds up.

8. Eventually you reach a point where the reverse reaction is going as fast as the forward reaction.
This is dynamic equilibrium.
The rate of the forward reaction is equal to the rate of the reverse reaction.
The concentration of products and reactants stays the same, but the reactions are still running.

9. Equilibrium
Equilibrium position- how much product and reactant there are at equilibrium.
Shown with the double arrow.
Reactants are favored
Products are favored
Catalysts speed up both the forward and reverse reactions so don’t affect equilibrium position.

10. A homogenous reaction is one in which all the substances are in the same state.
A heterogeneous reaction is one in which all the substances are not in the same state.

11. Measuring equilibrium
At equilibrium the concentrations of products and reactants are constant.
Keq is the equilibrium constant
Keq = [Products]coefficients [Reactants]coefficients
Square brackets [ ] means concentration in molarity (moles/liter)

12. Writing Equilibrium Expressions
General equation
aA + bB cC + dD
Keq = [C]c [D]d
[ A]a [B]b

13. N2 (g) + 3H2 (g) NH3 (g)

14. H2 (g) + F2 (g) 2HF(g)

15. Their concentrations don’t change
CaCO3 (s) CO2 (g) + CaO(s)
When writing equilibrium constant expressions for heterogeneous equilibria, you don’t include pure solids or pure liquids.
Their concentrations don’t change
Keq = [CO2 ][ CaO ] Keq = [CO2 ]
[ CaCO3 ]
Becomes 1 due to
solid phase

16. Keq = [H2]2[O2] Keq expression for this reaction:
2H2O (l) H2 (g) + O2(g)
Keq = [H2]2[O2]
H2O is a liquid and is not included in the equilibrium constant

17. Write the equilibrium expressions for the following reactions.
4NH3 (g) + 7O2 (g) 4NO2 (g) + 6H2O (g)
2H2O(g) H2(g) + O2(g)

18. Keq = [H2O]5 Keq expression for this reaction:
CuSO4 . 5H2O (s) CuSO4 (s) + 5H2O (g)
Keq = [H2O]5

19. Keq expression for this reaction:
2CaCO3(s) Ca(s) + 2CO2 (g) +O2(g)
2 H2 (g) + O2 (g) H2 O (l)

20. Try writing the equilibrium constant for:
N2 (g) H2 (g) 2NH3 (g)
C(s) H2 (g) CH4 (g)
CO(g) H2 (g) CO2 (g) + H2 (g)
C(s) H2O(g) CO (g) + H2 (g)
H2O(g) H2O(l)