1. Changing Concentration
In chemical reactions…
Standard conditions = 1M for all
When reactants are >1M, it will increase product concentration and will increase the cell potential
When products are >1M, it will decrease the product concentration (oppose forward reaction), decreases cell potential
Mg(s) + 2H+(aq) -> Mg2+(aq) +H2(g)

2. Concentration Cell
A galvanic cell that has the same component on each side but at different concentrations
Causes a cell potential
Voltages are typically small
Ex: a cell has on its left side a 0.20 M Cu2+ solution and a M Cu2+ solution on the right side

3. Reaction Quotient Represented by Q Q = concentration of productsx
[reactants]y
*Solids and liquids are not included…
Multiply concentrations of products/reactants by one another
X and Y represent coefficients of each reactant and product respectively

4. Reaction Quotient Example (13.5)
H2(g) + I2(g) <-> 2HI(g)
[H2]o = 0.81 M, [I2]o = 0.44 M, [HI]o = 0.58 M
Q = __[HI]2_ = (0.58)2 = 0.94
[H2]1[I2]1 (0.81)(0.44)

5. Nernst Equation At 25°C…to find actual cell potential
Ecell = E°cell log(Q) n
n is number of moles of electrons transferred
Q is the reaction quotient (see previous slide)
E°cell is calculated using standard reduction potentials (learned earlier this chapter)

6. Nernst Equation Example
Calculate Ecell for a galvanic cell based on the following half-reactions at 25C.
(Eq 1) FeO H+ + 3e- -> Fe3+ + 4H2O
E° = V
(Eq 2) O2 + 4H+ + 4e- -> 2H2O
E° = V
[FeO42-] = 2.0 X 10-3 M, [Fe3+] = 1.0 X 10-3 M,
[O2] = 1.0 X 10-5 M, [H+] = 6.31 X 10-6
(see pg. 399 in study guide) E = 0.54 V