General Chemistry · Electrochemistry and Thermodynamics
Redox Reactions and Thermodynamics
Reduction and oxidation move electrons, and that flow can light a bulb or drive a reaction. In this lab you assign oxidation numbers, identify what is oxidized and reduced in a galvanic cell, and find the cell potential, then turn to thermodynamics to calculate free energy and judge whether a reaction is spontaneous, finally linking the two through the relationship between cell potential and free energy. In every part you work out the answer yourself, then reveal the experimental value and compare.
Theory — Electron Transfer and Energy
Oxidation and reduction
A redox reaction is one in which electrons are transferred. Oxidation is the loss of electrons and reduction is the gain of electrons; a useful reminder is OIL RIG, oxidation is loss, reduction is gain. The species that loses electrons is the reducing agent, and the one that gains them is the oxidizing agent. We track the electrons with oxidation numbers, assigned by a short set of rules.
Rules for oxidation numbersA free element is 0. A monatomic ion equals its charge. Oxygen is usually −2, hydrogen usually +1. Group 1 metals are +1 and Group 2 are +2. The oxidation numbers in a neutral compound sum to 0; in an ion they sum to the ion charge.
An increase in oxidation number is oxidation; a decrease is reduction
Galvanic cells and cell potential
In a galvanic cell the two half-reactions are separated so the electrons travel through a wire, doing useful work. Oxidation happens at the anode and reduction at the cathode. Each half-reaction has a standard reduction potential; the half-cell with the higher value is reduced (the cathode), and the lower one is oxidized (the anode).
Standard cell potentialE°cell = E°cathode − E°anode
A positive E°cell means the cell reaction is spontaneous as written
Reduction half-reaction
E° (V)
Ag⁺ + e⁻ → Ag
+0.80
Cu²⁺ + 2e⁻ → Cu
+0.34
2H⁺ + 2e⁻ → H₂
0.00
Pb²⁺ + 2e⁻ → Pb
−0.13
Ni²⁺ + 2e⁻ → Ni
−0.25
Fe²⁺ + 2e⁻ → Fe
−0.44
Zn²⁺ + 2e⁻ → Zn
−0.76
Thermodynamics: free energy and spontaneity
Whether a reaction proceeds on its own is decided by the Gibbs free energy. It combines the heat exchanged, the enthalpy change ΔH, with the change in disorder, the entropy change ΔS, at temperature T in kelvin.
Because ΔS is usually given in joules and ΔH in kilojoules, divide the TΔS term by 1000 before subtracting. Redox and thermodynamics meet in one equation: the free energy of a cell reaction is set by its potential, where n is the moles of electrons transferred and F is the Faraday constant, 96485 coulombs per mole.
Free energy from cell potentialΔG° = −nFE°cell
A positive E°cell gives a negative ΔG°, confirming a spontaneous cell
Track the electrons
Oxidation numbers reveal which atom loses electrons (oxidized) and which gains them (reduced).
Higher potential reduces
The half-cell with the higher reduction potential is the cathode; the other is oxidized at the anode.
Free energy decides
A negative ΔG means the reaction is spontaneous, and it links directly to a positive cell potential.
Apparatus
The instruments you set up to build the cell, read its voltage, and measure heat changes.
Galvanic cell
Two half-cells joined by a salt bridge; electrons flow from the anode to the cathode through the wire.
Voltmeter
Reads the cell potential across the two electrodes in volts.
Digital multimeter
Measures voltage, current, or resistance; here it confirms the cell voltage.
Analytical balance
Weighs the electrodes before and after to track the mass change from electron transfer.
Coffee-cup calorimeter
Measures the heat released or absorbed by a reaction at constant pressure.
Thermometer
Records the temperature change used to find the enthalpy of a reaction.
Instructions — Running the Virtual Experiment
This is a predict, reveal, and compare lab. In every part you work out the answer yourself first, enter it, and only then does the simulation reveal the experimental value so you can check your work against it.
Part A — Oxidation Numbers (Oxidation Numbers tab)
1
Choose a compound and the element to assign. Using the rules, calculate the oxidation number of that element by hand, enter it (use a minus sign for negative values), and click Check. Work through at least five compounds.
Part B — Galvanic Cell and Cell Potential (Cell Potential tab)
1
Choose two electrode metals. Predict which one is oxidized (the anode), then calculate the cell potential from E°cathode − E°anode, enter both, and click Check. The voltmeter reading is then revealed for comparison.
Part C — Free Energy and Spontaneity (Free Energy tab)
1
Choose a reaction and a temperature. Calculate ΔG = ΔH − TΔS (remember to divide the TΔS term by 1000), predict whether the reaction is spontaneous, enter both, and click Check.
Part D — Free Energy from Cell Potential (ΔG from Eᵒ tab)
1
For a chosen cell, calculate ΔG° = −nFE°cell in kilojoules (with F = 96485 C/mol), enter it, and click Check to confirm that a positive cell potential gives a negative free energy.
For your reportInclude your oxidation-number assignments, the anode, cathode, and cell potential of the cells you built, your free-energy calculations with spontaneity, the link between ΔG° and E°, and screenshots.
Simulation — The Electrochemistry Bench
Redox and Thermodynamics Virtual LabCalculate first, then reveal and compare
Compound
Element
Your value
Actual
Choose a compound, calculate, and check.
Assign the oxidation number
Target element: Mn
Oxidation number— hidden
Electrodes
Metal 1 (E°)Zn (−0.76 V)
Metal 2 (E°)Cu (+0.34 V)
Build the cell
Anode (oxidized)— hidden
Cathode (reduced)— hidden
Voltmeter reading E°cell— hidden
Reaction
ΔH−92 kJ
ΔS−199 J/K
Temperature298 K
Free energy
ΔG— hidden
Spontaneity— hidden
Cell reaction
CellZn | Cu
Electrons transferred n2
E°cell1.10 V
Free energy from potential
ΔG°— hidden
Spontaneous?— hidden
Team Questions
Question 1. In a redox reaction, what is the term for the loss of electrons? (one word)
Question 2. What is the oxidation number of manganese in KMnO₄? (use + and a number)
Question 3. At which electrode does oxidation occur in a galvanic cell? (one word)
Question 4. For a Zn and Cu cell, what is the standard cell potential? (in volts, two decimals)
Question 5. Write the Gibbs free energy equation in terms of ΔH, T, and ΔS.
Question 6. If ΔG is negative, is the reaction spontaneous or nonspontaneous? (one word)
Question 7 — Challenge. A positive E°cell corresponds to a ΔG° that is positive or negative? (one word)
Example Lab Report
A worked example showing the expected format and the calculate, reveal, and compare workflow.
To assign oxidation numbers, identify oxidation and reduction in a galvanic cell and find its cell potential, to calculate free energy and judge spontaneity, and to relate cell potential to free energy.
Results (worked example)
Quantity
Value
Oxidation number of Mn in KMnO₄
+7
Zn–Cu cell: anode (oxidized)
Zn
Zn–Cu cell: cathode (reduced)
Cu
E°cell (0.34 − (−0.76))
1.10 V
ΔG for ΔH = −92 kJ, ΔS = −199 J/K at 298 K
−32.7 kJ (spontaneous)
ΔG° from −nFE° (n = 2)
−212.3 kJ
Zinc has the lower reduction potential, so it is oxidized at the anode, while copper is reduced at the cathode, giving a cell potential of 1.10 V. The free energy of the ammonia synthesis is negative at 298 K, so it is spontaneous, and the negative ΔG° from the cell potential confirms the Zn–Cu reaction is spontaneous.
Discussion and Conclusion
Each calculated value matched the experimental reading. Oxidation numbers identified the electron transfer, the higher reduction potential set the cathode, and a positive cell potential corresponded to a negative free energy. The free-energy calculations showed how spontaneity can depend on temperature, since a reaction with positive ΔH and positive ΔS becomes spontaneous only when T is large enough.
Practice Questions
Question 1
Find the oxidation number of chromium in K₂Cr₂O₇.
Hint: K is +1 (×2 = +2), O is −2 (×7 = −14); 2 + 2Cr − 14 = 0, so Cr = +6.
Question 2
For a cell made of silver and zinc, identify the anode and calculate E°cell.
Hint: Zn (−0.76) is oxidized at the anode, Ag (+0.80) is the cathode; E°cell = 0.80 − (−0.76) = 1.56 V.
Question 3
A reaction has ΔH = +178 kJ and ΔS = +161 J/K. Is it spontaneous at 298 K? At 1200 K?
