Electrochemistry: cell potential, Nernst equation, ΔG and equilibrium constant

Choose a metal pair, change the concentrations and the temperature and see how the cell potential follows the Nernst equation, with ΔG, the equilibrium constant K and the E against log Q graph. Free for every student.

Also searched as: nernst equation · nernst equation · cell potential · standard electrode potential · emf of a cell · gibbs energy and cell potential

About this experiment

Aim

To see how the potential of an electrochemical cell depends on the electrodes (standard potential E°), on the concentrations of the ions (Nernst equation) and on the temperature, and to relate E to ΔG and the equilibrium constant.

Theory

Each electrode has a standard electrode potential E° (at 1 mol/L and 298 K, measured against the standard hydrogen electrode). The standard cell potential is E°cell = E°(cathode) − E°(anode). For the Daniell cell it is 0.34 − (−0.76) = 1.10 V. Reduction takes place at the cathode and oxidation at the anode.

When the concentrations are not 1 mol/L, the potential changes according to the Nernst equation: E = E° − (RT/nF) ln Q, where Q is the reaction quotient. For M_a + C^n+ → A^n+ + C, Q = [A^n+]/[C^n+]. At 298 K this is E = E° − (0.0592/n) log Q. More product (a larger Q) gives a lower E; the cell stops at E = 0 (equilibrium).

The cell potential is connected with the Gibbs energy: ΔG = −nFE. A positive E means ΔG is negative and the reaction is spontaneous. At equilibrium E = 0, so ΔG° = −nFE° = −RT ln K and log K = nE°/0.0592 at 298 K. For the Daniell cell K is about 10³⁷.

Formulas used

• E°cell = E°cathode − E°anode
• E = E° − (RT/nF) ln Q = E° − (0.0592/n) log Q (298 K)
• ΔG = −nFE, ΔG° = −nFE°
• ΔG° = −RT ln K, log K = nE°/0.0592
• F = 96485 C/mol

Procedure: what to do in the virtual lab

1. Choose the Daniell cell with both concentrations 1 mol/L: read E° = 1.10 V.
2. Reduce the concentration of Cu²⁺ (cathode ion). Does E go up or down? Compare with the Nernst equation.
3. Increase the concentration of Zn²⁺ (anode ion) and watch E fall.
4. Change the temperature and see the small change in E and the change in ΔG.
5. Try the other pairs and note which metal is the anode.

Observations to make

• E changes by 0.0296 V when the concentration changes tenfold (n = 2).
• The graph of E against log Q is a straight line with a negative slope.
• The cell with the metals farthest apart in the electrochemical series (Mg–Cu) has the highest E°.

What you can change in this experiment

• Cell: Zn | Zn²⁺ ‖ Cu²⁺ | Cu (Daniell cell) / Zn | Zn²⁺ ‖ Ag⁺ | Ag / Mg | Mg²⁺ ‖ Cu²⁺ | Cu / Fe | Fe²⁺ ‖ Cu²⁺ | Cu
• log₁₀ of the concentration of the anode ion (mol/L)
• log₁₀ of the concentration of the cathode ion (mol/L)
• Temperature (K)

Viva questions on Electrochemistry

Questions an examiner may ask in the Class 12 chemistry practical. Run the experiment in the lab to find the answers yourself, then check them in the Learn tab.

1. State the Nernst equation.
2. How is the Gibbs energy of a cell reaction related to its EMF?
3. How is E°cell related to the equilibrium constant?
4. What happens to the EMF of a cell when it runs down?

Frequently asked questions

Syllabus: Class XII Chemistry, Chapter 2 — Electrochemistry, §2.2 (electrochemical cells), §2.3 (Nernst equation), §2.4 (Gibbs energy and EMF)

What you do in the lab

Change the controls and watch the chemistry behave in real time in 3D. Live readouts and graphs show the numbers behind it, and the Learn tab has the theory, procedure, observations and viva questions. Works on a phone or a laptop, in English and Hindi.

Topics: Nernst equation · cell potential · standard electrode potential · EMF of a cell · Gibbs energy and cell potential · equilibrium constant · electrochemistry · electrochemical series

More from this unit

• Conductivity of strong and weak electrolytes
• Electrolytic refining of copper
• Daniell cell and the Nernst equation
• Electrolysis of brine (chlor-alkali)
• Electrolysis of copper sulphate (inert electrodes)
• Electrolysis of potassium iodide
• Electrolysis of water
• Copper plating an iron key

Also try these chemistry experiments

• Faraday's first law: mass of copper deposited ∝ charge (Class 12)
• Electrochemical equivalent of silver (Class 12)
• Potassium ferrocyanide tests for Cu²⁺, Zn²⁺ and Fe³⁺ (Class 12)
• Group I cations: separating Pb²⁺ from Ag⁺ as chlorides (Class 12)