What you'll be able to do: Use a table of standard reduction potentials to calculate a standard cell potential and predict whether a redox reaction is spontaneous.
Best after: Galvanic Cells and Cell Notation
A single number tells you whether a redox reaction will run on its own: the standard cell potential. Reading it off a table takes about thirty seconds once you know the two rules that govern the calculation.
These are recommended, not required. You can start this lesson at any time.
You cannot measure the potential of a single half-cell, only the difference between two. Chemists therefore fixed one half-cell as the zero point: the standard hydrogen electrode, 2 H⁺(1 M) + 2 e⁻ → H₂(1 atm), assigned exactly 0 V. Every tabulated value is that couple measured against hydrogen.
All values are written as reductions. A large positive value, such as +2.87 V for F₂, marks a powerful oxidizing agent that grabs electrons readily. A large negative value, such as -3.04 V for Li⁺, means the reduction is very unfavorable, which is another way of saying that lithium metal is an excellent reducing agent.
F₂ + 2 e⁻ → 2 F⁻, E° = +2.87 V
Given two half-cells, the one with the higher reduction potential wins the competition for electrons and becomes the cathode. The other one is forced to run backward as an oxidation and becomes the anode. This choice is exactly what makes E°cell positive.
E°cell = E°cathode - E°anode, where both numbers come straight from the table as reduction potentials. The subtraction already accounts for reversing the anode half-reaction, so flipping the sign yourself and then subtracting double-counts it. Adding a flipped anode value gives the same result, but mixing the two approaches is the most common source of sign errors.
E°cell = E°cathode - E°anode
If you multiply a half-reaction by 3 to balance electrons, the number of electrons changes but the potential does not. Potential is energy per unit charge, like temperature or density: doubling the amount of material does not change it. Only the total energy, which appears later as the free energy change, scales with the electron count.
Any oxidized species in the table can oxidize any reduced species that appears lower in the table. This is the reasoning behind the activity series: copper does not dissolve in HCl because Cu²⁺/Cu sits above H⁺/H₂, so hydrogen ions cannot oxidize copper, while zinc, which sits below, dissolves readily.
The voltage of a half-reaction written as a reduction, measured against the standard hydrogen electrode at standard conditions.
The half-reaction with the more positive E° is the reduction.
E°cell = E°cathode - E°anode, using tabulated reduction values.
E°cell > 0 means the reaction is spontaneous as written under standard conditions.
Scaling a half-reaction does not change its potential.
E°cell = E°cathode - E°anode
2 H⁺(1 M) + 2 e⁻ → H₂(1 atm), E° = 0 V
Given Cu²⁺ + 2 e⁻ → Cu, E° = +0.34 V and Zn²⁺ + 2 e⁻ → Zn, E° = -0.76 V, find E°cell for the zinc-copper cell and state whether it is spontaneous.
Try it first: Decide which half-cell is the cathode before touching the numbers.
0 of 3 steps revealed.
Will silver metal dissolve in 1 M HCl? Use Ag⁺ + e⁻ → Ag, E° = +0.80 V and 2 H⁺ + 2 e⁻ → H₂, E° = 0.00 V.
Try it first: Write the proposed reaction: is silver oxidized or reduced?
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For 2 Al(s) + 3 Cu²⁺ → 2 Al³⁺ + 3 Cu(s), with Al³⁺ + 3 e⁻ → Al at E° = -1.66 V and Cu²⁺ + 2 e⁻ → Cu at +0.34 V, find E°cell.
Try it first: Notice that the half-reactions had to be scaled by 2 and 3 to give 6 electrons.
0 of 3 steps revealed.
Why it's wrong: The subtraction already reverses that half-reaction, so the sign is applied twice.
Check instead: Plug in tabulated reduction values exactly as printed.
Why it's wrong: Potential is intensive; only free energy scales with the amount of reaction.
Check instead: Scale coefficients and electrons only, never the voltage.
Why it's wrong: What matters is the more positive value, not the larger absolute value.
Check instead: Compare positions on a number line, keeping signs.
Why it's wrong: The reverse reaction is spontaneous, and the forward direction can still be driven electrically.
Check instead: Read a negative value as pointing to the reverse direction.
Why it's wrong: Standard values assume 1 M and 1 atm; other conditions need the Nernst equation.
Check instead: Check the conditions before using E° directly.
No practice questions are available for this topic yet. You can still practice the whole unit.
Standard reduction potentials measure how strongly a species pulls electrons, all reported relative to the standard hydrogen electrode, which is defined as exactly 0 V. A more positive value means a stronger oxidizing agent, one more eager to be reduced. To build a spontaneous cell, the half-reaction with the higher (more positive) reduction potential runs as the reduction at the cathode, and the other is reversed to become the oxidation at the anode. The standard cell potential is E°cell = E°cathode - E°anode, using both values as tabulated for reduction and never flipping a sign for the reversed half-reaction. A positive E°cell means the reaction as written is spontaneous under standard conditions. Because potential is an intensity, not an amount, E° values are never multiplied when a half-reaction is scaled to balance electrons. The table also ranks reactivity: any species can oxidize anything that sits lower in reduction potential than itself.
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