1. Dependence of cell potential on concentrations
Lab 10

2. Outline Purpose Electrochemistry Galvanic Cells Cell Potentials
Standard Reduction Potentials
The Nernst Equation
Graphical Nernstian Response
Procedure
Safety Concerns
Waste
Next Lab Reminder

3. Purpose
In this experiment students will construct half-cells of Cu2+ / Cu and Zn2+ / Zn in contact with KNO3 solution (salt bridge).
They will be able to show that there is a linear dependence of cell potential on concentration as per the Nernst equation.
The Nernst equation is used for calculations when non-standard conditions and/or concentrations are involved in a Voltaic setup.

4. Electrochemistry
A study of the interchange of electrical and chemical energy.
An electrical current can be established FROM a spontaneous chemical reaction.
Chemical change can be produced FROM an electrical current.

5. Galvanic (Voltaic) Cells
Galvanic cells use a redox reaction (chemical reaction) to generate an electrical current.
When both reagents are in the same solution, electrons are transferred directly when reagents collide, so no useful work is obtained (heat may be released).
When the reagents are separated, but connected through a salt bridge and metal electrodes, the electron transfer occurs through a wire and can, for example, run an electric motor (useful work obtained).

6. Galvanic (Voltaic) Cells
This is a traditional Galvanic cell setup. Ours will look slightly different.

7. Galvanic (Voltaic) Cells
Without a salt bridge:
Current flows from the anode to the cathode but builds up a negative charge (on the cathode).
Without a large external influx of energy, the current ceases its flow.
With a salt bridge:
Electrons are transferred from the reducing agent (anode) to the oxidizing agent (cathode).
The salt bridge ions neutralize the charge build-up (cations to the cathode, anions to the anode).
The circuit is complete, the net charge in each compartment becomes zero.
Current flows until the cell is discharged and equilibrium is reached. At that point, the components in the two cell compartments have the same free energy. (G = 0, Q = K, E = 0)

8. Cell Potential
Ecell (unit V) is the cell potential or electromotive force responsible for driving electrons from the reducing agent (anode) to the oxidizing agent (cathode)
We measure Ecell with a voltmeter which draws current through a known resistance.
When current flows through a wire, frictional heating results in lost energy.
A voltmeter therefore always reads a potential less than the maximum cell potential (E0cell). This occurs less so with digital voltmeters compared to analog voltmeters.

9. Standard Reduction Potentials
Half-reactions are written as REDUCTION reactions in reduction potential tables.
Each half-reaction has its own reduction potential, which can be positive, or negative, depending on how it compares to the standard hydrogen electrode:
2H+ + 2e-  H2 which has an E0 = 0.00 V
Our half-reactions are:
Zn2++ 2 e-  Zn E0 = V
Cu e-  Cu E0 = V

10. Standard Reduction Potentials
When the reduction potentials are added together, you get the standard reduction potential for the cell (E0).
A cell runs spontaneously in the direction that produces a positive cell potential. (E0 has to be positive for the reaction to work.)
Zn  Zn2++ 2 e V
Cu e-  Cu V
E0 = V
Both cell compartments must be in their standard states to obtain this “theoretical” value. (1 M, 1 atm, 25 C)
Experimentally we can find our E0cell value by plotting E, V vs. log Q and then solving for E when log Q = 0.

11. Standard Reduction Potentials
Because of nonstandard concentrations (and other conditions), experimentally:
Ecell < E0cell < E0
Ecell = the cell potential we will measure
E0cell = the experimental standard state potential difference from E,V vs. log Q. This is the largest potential we can possibly observe before the current flows.
E0 = the theoretical standard state potential difference (1.10 V)

12. The Nernst Equation
The Nernst equation demonstrates a linear relationship between galvanic cell potential and cell concentration.
Ecell = E0cell ln Q
where R = gas constant,
T = temperature in Kelvin
F = Faraday’s constant
n = number of mole electrons

13. The Nernst Equation
Adjusted for lab conditions (substituting in the values for R and F and 25C), with a few other conversions, we get:
Ecell = E0cell log Q
E = log Q + E0
y = m x b

14. Nernstian Response
A reversible electrode responds in a Nernstian fashion when E, V vs. log Q gives a straight line with a slope of
To calculate number of electrons transferred, we simply use:

15. Procedure Prepare your Cu2+ solutions.
Collect your Zn2+ and KNO3 solutions.
“Calibrate” your voltage probe.
Set up your experimental apparatus and perform your experiment as detailed in your lab manual.
Make up the required spreadsheet and graph based on your results.

16. Safety Concerns Reagents: Eye Contact: Skin Contact: Inhalation:
Cupric sulfate
Zinc sulfate
Potassium nitrate (1 M)
Copper / Zinc solids
Eye Contact:
Irritation, pain, redness, conjunctivitis, ulceration, mechanical harm, clouding of cornea
Skin Contact:
Irritation, redness, pain, itching
Inhalation:
Coughing, sore throat, shortness of breath, ulceration, methemoglobinemia, cyanosis, convulsions, tachycardia, dyspnea, dizziness, drowsiness, headache, perforation of the respiratory tract and death. Fumes from heating may cause symptoms similar to a cold.
Ingestion:
Burning of the mouth, esophagus, and stomach, hemorrhagic gastritis, nausea, vomiting, abdominal pain, metallic taste, tachycardia, hypotension, pulmonary edema, kidney damage, liver damage, hemorrhagic pancreatitis and diarrhea. Systemic copper poisoning with capillary damage, headache, cold sweat, weak pulse, CNS excitation, depression, jaundice, convulsions, blood effects, paralysis, coma and death.

17. Waste
Zinc solutions may go down the drain, flushed with a lot of water.
Copper solutions are toxic and MUST be disposed in the appropriate waste container in the fume hood.
KNO3 solutions may go down the drain.

18. lab 11 Reminder
Lab 11 next week.