Two kinds of electrochemical cell
If the oxidation and reduction of a redox reaction happen in separate places, the electrons must travel through a wire, and a current flows.
- A galvanic (voltaic) cell原电池 uses a thermodynamically favorable redox reaction (ΔG < 0) to produce electrical energy: chemical energy → electrical energy.
- An electrolytic cell电解池 uses an external power source to drive a thermodynamically unfavorable reaction (ΔG > 0): electrical energy → chemical energy.
In both, the electrodes are named by the reaction that happens there:
- Oxidation happens at the anode, and electrons leave it into the wire;
- Reduction happens at the cathode, which receives electrons from the wire.
The signs differ, though: in a galvanic cell the anode is negative (−) and the cathode positive (+); in an electrolytic cell the anode is connected to the positive (+) terminal of the source and the cathode to the negative (−) terminal.
Parts of a galvanic cell
Take iron and silver: an iron strip dips into Fe(NO₃)₂ solution and a silver strip into AgNO₃ solution. The two beakers are joined by a salt bridge盐桥 containing KNO₃ solution, and the two metals are joined by a wire.
- Anode (Fe): Fe → Fe²⁺ + 2e⁻, and the iron strip gets thinner;
- Cathode (Ag): Ag⁺ + e⁻ → Ag, and silver plates out, so the silver strip gains mass;
- Electrons flow through the wire from Fe to Ag;
- In the salt bridge, anions (NO₃⁻) move toward the anode side and cations (K⁺) toward the cathode side, balancing the extra positive charge building up at the anode and the positive charge being lost at the cathode, so both solutions stay electrically neutral. Without the salt bridge, the current would stop almost at once.
This cell can be written as Fe(s) | Fe²⁺(aq) ‖ Ag⁺(aq) | Ag(s): the anode goes on the left, “|” marks a boundary between phases, and “‖” stands for the salt bridge.
Choose two metals to build a galvanic cell, then switch to an electrolytic cell and see how each part changes (E° from OpenStax Appendix L):
Example 1: a Ni–Sn galvanic cell
A nickel strip dips into 1 M Ni(NO₃)₂ and a tin strip into 1 M Sn(NO₃)₂; they are connected by a KNO₃ salt bridge and a wire. E°: Ni²⁺/Ni −0.26 V, Sn²⁺/Sn −0.14 V. Identify the anode and cathode, write the electrode reactions, and describe the movement of electrons and ions and the changes in electrode mass.
-
Compare E° to decide what is reduced. The larger E°, the more readily the ion gains electrons: Sn²⁺/Sn (−0.14 V) is larger than Ni²⁺/Ni (−0.26 V), so Sn²⁺ is reduced and tin is the cathode:
-
The other electrode is oxidized. Nickel is the anode:
-
Overall reaction and E°cell. Each half-reaction transfers 2 electrons, so add them directly:
E°cell > 0, so the galvanic cell runs spontaneously. Consistent.
-
Movement of electrons and ions.
- Electrons: through the wire from Ni to Sn
- Salt bridge: NO₃⁻ moves toward the nickel side (balancing the Ni²⁺ formed), K⁺ toward the tin side (replacing the Sn²⁺ used up)
-
Mass and concentration changes. The nickel strip dissolves and loses mass; tin plates onto the tin strip, which gains mass. [Ni²⁺] increases and [Sn²⁺] decreases.
Example 2: turning it into an electrolytic cell
In the setup from Example 1, the wire is replaced by an external power source that drives the reverse reaction: Ni²⁺ + Sn → Ni + Sn²⁺. Which electrode should be connected to the positive terminal? What is the minimum voltage? Which electrode gains mass?
-
Decide what is oxidized and reduced now. In Ni²⁺ + Sn → Ni + Sn²⁺, Sn loses electrons and is oxidized, so the tin strip is the anode; Ni²⁺ gains electrons and is reduced, so the nickel strip is the cathode.
-
Connections. In an electrolytic cell the anode goes to the positive terminal and the cathode to the negative: positive terminal to the tin strip, negative to the nickel strip.
-
Minimum voltage. The standard cell potential in this direction (in V):
E°cell < 0 (ΔG° > 0), so it isn’t spontaneous, and the applied voltage must be more than 0.12 V.
-
Mass changes. Nickel plates out on the nickel strip, which gains mass; the tin strip dissolves and loses mass.
-
Check. Compared with the galvanic cell in Example 1, the electrode reactions, electron flow, and mass changes are all reversed. Consistent.
- The electrode where oxidation happens
Oxidation happens at the anode.
- The positive (+) electrode of a galvanic cell
In a galvanic cell the cathode is positive.
- The electrode connected to the positive terminal in an electrolytic cell
In an electrolytic cell the anode is connected to the positive terminal.
- The electrode in a galvanic cell where metal plates out and mass increases
Metal ions are reduced at the cathode.
- The electrode electrons leave to enter the wire
The anode loses electrons.
- The side that cations in the salt bridge move toward
Cations move toward the cathode side.
- The electrode where potassium forms during electrolysis of molten KBr
K⁺ + e⁻ → K is a reduction.
- The electrode where bromine forms during electrolysis of molten KBr
2Br⁻ → Br₂ + 2e⁻ is an oxidation.
Common mistakes
Mistake 1: “The anode is always positive.” In a galvanic cell the anode is negative; only in an electrolytic cell is it connected to the positive terminal. Identify the anode and cathode by whether oxidation or reduction happens there.
Mistake 2: “Electrons flow through the salt bridge.” Electrons flow only in the wire; ions move in the salt bridge, keeping both solutions electrically neutral.
Mistake 3: “A bigger electrode gives a higher voltage.” The voltage depends on which electrode reactions occur and on the ion concentrations, not on electrode size; a larger electrode just lets the cell run longer.
Quick check
After Fe(s) | Fe²⁺(aq) ‖ Ag⁺(aq) | Ag(s) has run for a while, 0.010 mol of Fe has dissolved from the iron electrode. How many moles of Ag have been deposited on the silver electrode in that time?
Each Fe loses 2 electrons and each Ag⁺ gains 1: 0.010 mol Fe gives up 0.020 mol of electrons, depositing 0.020 mol of Ag.
A galvanic cell is made from magnesium and tin (E°: Mg²⁺/Mg −2.37 V, Sn²⁺/Sn −0.14 V). Which way do electrons flow in the wire?
Mg has the smaller E°, so it is the anode and is oxidized; electrons flow from Mg through the wire to Sn.
A Ni–Cu galvanic cell uses a KNO₃ salt bridge, with nickel as the anode. As it operates, NO₃⁻ moves toward:
Ni²⁺ keeps forming on the anode side, adding positive charge, so anions move there to keep the solution neutral.
What is the fundamental difference between a galvanic cell and an electrolytic cell?
Both involve redox reactions; they differ in whether the reaction is spontaneous and in the direction of energy conversion.
Which observation best shows that a device is a galvanic cell?
A galvanic cell supplies electrical energy by itself.
For Zn(s) | Zn²⁺(aq) ‖ Pb²⁺(aq) | Pb(s), which statement is correct?
In cell notation the cathode is on the right; electrons flow from Zn to Pb, and cations move toward the Pb side.
AP-style free response
A student builds a galvanic cell from a lead strip in 1.0 M Pb(NO₃)₂, a silver strip in 1.0 M AgNO₃, and a KNO₃ salt bridge. E°: Pb²⁺/Pb −0.13 V, Ag⁺/Ag +0.80 V.
(a) Write the half-reaction at each electrode and the balanced overall reaction.
(b) Identify the anode and the cathode, and state which is the positive electrode.
(c) Calculate E°cell.
(d) Describe how K⁺ and NO₃⁻ move in the salt bridge, and state the purpose of the salt bridge.
(e) After the cell has run for a while, how have the ion concentrations in the two beakers changed? How have the electrode masses changed?
Model answer and scoring
(a) Anode: Pb → Pb²⁺ + 2e⁻; cathode: Ag⁺ + e⁻ → Ag; overall: Pb(s) + 2Ag⁺(aq) → Pb²⁺(aq) + 2Ag(s).
(b) Lead is the anode (negative), and silver is the cathode (positive).
(c) E°cell = 0.80 − (−0.13) = 0.93 V.
(d) K⁺ moves toward the silver side (cathode) and NO₃⁻ toward the lead side (anode). The salt bridge keeps both solutions electrically neutral so that the current can continue.
(e) [Pb²⁺] increases and [Ag⁺] decreases; the lead strip loses mass, and silver plates out on the silver strip, which gains mass.
Scoring: (a) both half-reactions and the overall reaction (coefficient 2 for silver); (b) electrodes and signs; (c) E°cell; (d) ion directions and purpose; (e) concentration and mass changes.
Summary
- Galvanic cell: ΔG < 0, chemical energy → electrical energy; electrolytic cell: ΔG > 0, electrical energy drives the reaction.
- Oxidation at the anode, reduction at the cathode; electrons flow through the wire from anode to cathode; in the salt bridge, cations move toward the cathode and anions toward the anode.
- In a galvanic cell the anode is negative; in an electrolytic cell the anode is connected to the positive terminal. Metal plates out on the cathode, which gains mass.
AP tip: To describe a cell, give in order: the anode and its half-reaction, the cathode and its half-reaction, the direction of electrons in the wire, the direction of ions in the salt bridge, and the changes in electrode mass. “Anode = oxidation” holds for both galvanic and electrolytic cells.