Electrochemical cells convert chemical energy into electrical energy or use electrical energy to drive chemical change. They matter because batteries, corrosion, metal plating, and many sensors all depend on electron transfer between substances. In a galvanic cell, a spontaneous redox reaction produces a voltage that can power an external circuit.
Understanding the parts of the cell helps explain how charge moves through both wires and solutions.
A typical galvanic cell has two half-cells connected by a wire and a salt bridge. Oxidation occurs at the anode, where electrons are produced, and reduction occurs at the cathode, where electrons are consumed. Electrons flow through the external wire from anode to cathode, while ions move through the salt bridge to keep both solutions electrically neutral.
The cell voltage depends on the reduction potentials of the two half-reactions and tells how strongly the reaction tends to occur.
Understanding Electrochemical Cells
Standard reduction potentials are measured by comparing half-reactions with a reference electrode. Tables list every half-reaction as a reduction, even when a reaction will run in reverse inside a real cell. A more positive reduction potential means a stronger tendency to take electrons under standard conditions.
These conditions use solutions at one mole per litre, gases at a pressure of one bar, and a temperature often close to twenty five degrees Celsius. Potentials are not multiplied when the balanced equation is multiplied. Electrons must balance in the full redox equation, but voltage is an energy change per unit charge, so it stays the same.
Cell notation is a compact way to record what is happening in each half-cell. A single vertical line marks a boundary between different phases, such as a metal electrode touching its solution. A double vertical line represents the salt bridge or porous barrier.
The left side normally shows the oxidation half-cell, while the right side shows the reduction half-cell. Commas separate species in the same phase.
For example, an inert platinum electrode may be needed when all reacting species are dissolved or gaseous. Reading notation carefully prevents a common mistake, which is assuming that every electrode is a reacting metal.
Electrolytic cells use an external power supply to force a reaction that would not proceed by itself. This changes the signs of the electrodes compared with a galvanic cell. The anode is positive because the power supply pulls electrons away from it.
The cathode is negative because the supply delivers electrons to it. The names anode and cathode are still based on the reaction type, not on the sign. This rule helps with electroplating.
During copper plating, copper ions gain electrons at the object being coated. The object becomes the cathode. In industrial electrolysis, the amount of substance made depends on the charge passed through the cell and the number of electrons needed per particle.
Real cells rarely operate under standard conditions. As reactants are used up or products build up, the voltage changes. The Nernst equation connects this change to temperature, electron number, and the reaction quotient.
The reaction quotient compares product concentrations with reactant concentrations at that moment. A cell usually gives a lower voltage when products become more concentrated. Concentration cells show this effect clearly because they can produce a voltage using the same electrodes in solutions of different concentration.
When solving problems, first balance the overall reaction, identify the number of transferred electrons, and write the reaction quotient without solids or pure liquids. Keep track of units and signs, since one reversed half-reaction can reverse the predicted direction of the whole cell.
Key Facts
- Oxidation is loss of electrons, and reduction is gain of electrons.
- In a galvanic cell, the anode is negative and the cathode is positive.
- Electrons flow through the external wire from anode to cathode.
- Ecell = Ecathode - Eanode using standard reduction potentials.
- A positive Ecell means the redox reaction is spontaneous under the stated conditions.
- The salt bridge allows ion flow and prevents charge buildup in the half-cells.
Vocabulary
- Electrochemical cell
- A device that connects oxidation and reduction half-reactions to transfer electrons and energy.
- Anode
- The electrode where oxidation occurs and electrons are produced.
- Cathode
- The electrode where reduction occurs and electrons are consumed.
- Salt bridge
- A connection containing ions that completes the circuit by allowing charge balance between half-cells.
- Cell potential
- The voltage produced by an electrochemical cell due to the difference in reduction tendency between the two electrodes.
Common Mistakes to Avoid
- Calling the cathode the place where oxidation happens is wrong because reduction always occurs at the cathode in both galvanic and electrolytic cells.
- Reversing electron flow is wrong because electrons leave the anode and travel through the wire to the cathode in a galvanic cell.
- Adding the two reduction potentials without changing signs correctly is wrong because Ecell = Ecathode - Eanode when using standard reduction potentials.
- Forgetting the salt bridge is wrong because the cell quickly stops if charge buildup prevents further electron flow.
Practice Questions
- 1 A galvanic cell has Zn as the anode with Ered = -0.76 V and Cu as the cathode with Ered = +0.34 V. Calculate Ecell.
- 2 A cell reaction transfers 2 mol of electrons and has Ecell = 1.10 V. Using ΔG = -nFE with F = 96485 C/mol, calculate ΔG in kJ/mol reaction.
- 3 A student removes the salt bridge from a working galvanic cell. Explain what happens to the voltage and why the cell cannot keep operating.