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Electrochemistry Lab

Build galvanic and electrolytic cells from pairs of metals. Adjust ion concentrations and temperature to see how the Nernst equation modifies standard cell potential. Compute Gibbs free energy and determine spontaneity for any electrode combination.

Guided Experiment: Daniell Cell Investigation

If zinc is paired with copper using 1 M solutions at 25 C, what cell voltage do you predict? Which electrode will be oxidized?

Write your hypothesis in the Lab Report panel, then click Next.

Zn⁺⁺ (aq)ZnANODE (−)oxidationCu⁺⁺ (aq)CuCATHODE (+)reductionsalt bridgee⁻ flowE°cell = 1.100 VSpontaneousGalvanic Cell

Controls

Anode ion conc.1.000 M
Cathode ion conc.1.000 M
Temperature298.1 K

Results

E=1.1000−(8.314)(298.1)(2)(96485)ln⁡(1.000)=1.1000 VE = 1.1000 - \frac{(8.314)(298.1)}{(2)(96485)}\ln(1.000) = 1.1000 \text{ V}
E°cell
1.1000 V
E (Nernst)
1.1000 V
Q (reaction quotient)
1.0000
ΔG° (kJ/mol)
-212.27
n (electrons)
2
Spontaneous
Spontaneous
Half-Cell Reduction Potentials
Anode (oxidation half-cell):-0.7600 V
Cathode (reduction half-cell):0.3400 V
E°cell = E°cathode − E°anode:1.1000 V

Nernst Curve: E vs ln(Q)

The red dot marks the current operating point. The dashed line shows E°cell. Slope = −RT/nF.

Data Table

(0 rows)
#TrialAnodeCathode[Anode](M)[Cathode](M)E°cell(V)E (Nernst)(V)ΔG°(kJ/mol)Spontaneous
0 / 500
0 / 500
0 / 500

Reference Guide

Galvanic Cells

A galvanic cell converts chemical energy to electrical energy. It is spontaneous when E°cell is positive.

The anode undergoes oxidation (loses electrons). The cathode undergoes reduction (gains electrons). Electrons flow through the external circuit from anode to cathode. Ions flow through the salt bridge to maintain charge balance.

Ecell∘=Ecathode∘−Eanode∘E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}

If E°cell is positive, the reaction is spontaneous and the cell produces work. If negative, an external voltage must be applied (electrolytic cell).

Standard Reduction Potentials

Standard reduction potentials (E°) are measured at 25 °C, 1 M concentration, and 1 atm, relative to the standard hydrogen electrode (SHE, 0.00 V).

Half-Reaction E° (V)
Li⁺ + e⁻ → Li -3.04
Zn²⁺ + 2e⁻ → Zn -0.76
Fe²⁺ + 2e⁻ → Fe -0.44
Ni²⁺ + 2e⁻ → Ni -0.26
Sn²⁺ + 2e⁻ → Sn -0.14
Pb²⁺ + 2e⁻ → Pb -0.13
2H⁺ + 2e⁻ → H₂ 0.00
Cu²⁺ + 2e⁻ → Cu +0.34
Ag⁺ + e⁻ → Ag +0.80
Au³⁺ + 3e⁻ → Au +1.50

Nernst Equation

The Nernst equation adjusts cell potential for non-standard concentrations and temperatures.

E=E∘−RTnFln⁡QE = E^\circ - \frac{RT}{nF}\ln Q

At 25 °C this simplifies to:

E=E∘−0.0592nlog⁡10QE = E^\circ - \frac{0.0592}{n}\log_{10} Q

Where R = 8.314 J/(mol·K), T is temperature in Kelvin, n is electrons transferred, F = 96485 C/mol, and Q is the reaction quotient. When Q = 1, E equals E°. When Q > 1, E decreases; when Q < 1, E increases.

Gibbs Free Energy

The Gibbs free energy change links electrochemistry to thermodynamics. For a cell at standard conditions:

ΔG∘=−nFEcell∘\Delta G^\circ = -nFE^\circ_{\text{cell}}

When E°cell is positive, deltaG° is negative, confirming the reaction is spontaneous. Faraday's constant F = 96485 C/mol converts the electrical work (nEcell) to Gibbs energy in joules.

For the Daniell cell (Zn-Cu, E° = 1.10 V, n = 2):

ΔG∘=−(2)(96485)(1.10)=−212,267 J/mol\Delta G^\circ = -(2)(96485)(1.10) = -212{,}267 \text{ J/mol}