1. Midterm exam location
Odette Building, room 104
Feb. 3rd from 1 - 2:20 p.m.
March 10th from 1 - 2:20 p.m

2. Solutions contain more than one types of electrolytes
The ionic strength of the solution equals the sum of the ionic strength of each individual compound.
Example: Calculate the ionic strength of a solution that contains mol kg-1 K3[Fe(CN)6](aq), mol kg-1 NaCl(aq), and 0.03 mol kg-1 Ce(SO4)2 (aq).
Solution:
I (K3[Fe(CN)6]) = ½( 12*(0.05*3) + (-3)2*0.05)
= ½ ( ) = 0.3;
I (NaCl) = ½(12* (-1)2*0.04) = 0.04;
I (Ce(SO4)2) = ½(42* (-2)2*(2*0.03)) =0.36;
So, I = I(K3(Fe(CN)6]) + I(NaCl) + I(Ce(SO4)2)
= = 0.7

3. Calculating the mean activity coefficient
Example: Calculate the ionic strength and the mean activity coefficient of 2.0m mol kg-1 Ca(NO3)2 at 25 oC.
Solution: In order to calculate the mean activity coefficient with the eq. 10.3, one needs to know the ionic strength of the solution. Thus, the right approach is first to get I and then plug I into the equation 10.3.
I = ½(22* (-1)2*(2*0.002)) = ½ *6*0.002 = 0.006;
From Debye-Huckel limiting equation,
log(γ±) = - |2*1|*A*(0.006)1/2;
= - 2*0.509*0.0775;
= ;
γ± = 0.834;

4. Experimental test of the Debye-Hückel limiting law

5. Accuracy of the Debye-Hückel limiting law
Example: The mean activity coefficient in a mol kg-1 MnCl2(aq) solution is 0.47 at 25oC. What is the percentage error in the value predicted by the Debye-Huckel limiting law?
Solution: First, calculate the ionic strength
I = ½[22*0.1 + (-1)2*(2*0.1)] = 0.3
to calculate the mean activity coefficient.
log(γ) = -|2*(-1)| A* (0.3)1/2;
= - 2*0.509*0.5477 =
so γ = 0.277
Error = ( )/0.47 * 100%
= 41%

6. Extended Debye-Hückel law
B is an adjustable empirical parameter. It is different for each electrolyte.

7. Calculating parameter B
Example : The mean activity coefficient of NaCl in a diluted aqueous solution at 25oC is (at 10.0 mmol kg-1). Estimate the value of B in the extended Debye-Huckel law.
Solution: First calculate the ionic strength
I = ½[12* (-1)2*0.01] = 0.01
From equation
log(0.907) = - (0.509|1*(-1)|*0.011/2)/(1+ B*0.011/2)
B =

8. Half-reactions and electrodes
Two types of electrochemical cells:
1. Galvanic cell: is an electrochemical cell which produces electricity as a result of the spontaneous reactions occurring inside it.
2. Electrolytic cell: is an electrochemical cell in which a non-spontaneous reaction is driven by an external source of current.

9. Other important concepts include:
Oxidation: the removal of electrons from a species.
Reduction: the addition of electrons to a species.
Redox reaction: a reaction in which there is a transfer of electrons from one species to another.
Reducing agent: an electron donor in a redox reaction.
Oxidizing agent: an electron acceptor in a redox reaction.
Two type of electrodes:
Anode: the electrode at which oxidation occurs.
Cathode: the electrode at which reduction occurs

10. Typical Electrodes

11. Electrochemical cells
Liquid junction potential: due to the difference in the concentrations of electrolytes.
The right-hand side electrochemical cell is often expressed as follows:
Zn(s)|ZnSO4(aq)||CuSO4(aq)|Cu(s)
The cathode reaction (copper ions being reduced to copper metal) is shown on the right. The double bar (||) represents the salt bridge that separates the two beakers, and the anode reaction is shown on the left: zinc metal is oxidized into zinc ions

12. In the above cell, we can trace the movement of charge.
Electrons are produced at the anode as the zinc is oxidized
The electrons flow though a wire, which we can use for electrical energy
The electrons move to the cathode, where copper ions are reduced.
The right side beaker builds up negative charge. Cl- ions flow from the salt bridge into the zinc solution and K+ ions flow into the copper solution to keep charge balanced.
To write the half reaction for the above cell,
Right-hand electrode: Cu2+(aq) + 2e- → Cu(s)
Left-hand electrode: Zn2+(aq) e- → Zn(s)
The overall cell reaction can be obtained by subtracting left-hand reaction from the right-hand reaction:
Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq)

13. Expressing a reaction in terms of half-reactions
Example : Express the formation of H2O from H2 and O2 in acidic solution as the difference of two reduction half-reactions.
(in class discussion)
Redox couple: the reduced and oxidized species in a half-reaction such as Cu2+/Cu, Zn2+/Zn….
Ox + v e- → Red
The quotient is defined as: Q = aRed/aOx
Example: Write the half-reaction and the reaction quotient for a chlorine gas electrode.

14. Varieties of cells

15. Notation of an electrochemical cell
Phase boundaries are denoted by a vertical bar.
A double vertical line, ||, denotes the interface that the junction potential has been eliminated.
Start from the anode.
A general format:
Solid | gas phase | aqueous phase || aqueous phase | gas phase | solid

16. Cell Potential
Cell potential: the potential difference between two electrodes of a galvanic cell (measured in volts V).
Maximum electrical work : we,max = ΔG
Electromotive force, E,
Relationship between E and ΔrG:
ΔrG = -νFE
where ν is the number of electrons that are exchanged during the balanced redox reaction and F is the Faraday constant, F = eNA.
At standard conditions, this equation can be written as
ΔrGθ = -νFEθ