When zinc sits in copper(II) sulfate solution, electrons jump straight from zinc atoms to copper ions touching the metal surface. The energy released just warms the solution slightly. Now imagine putting the zinc in one beaker and the copper ions in another, and connecting them with a wire. The electrons still want to move from zinc to copper(II), but now the only path is through the wire, and a flow of electrons through a wire is an electric current. That is the idea behind every battery: a galvanic cell (also called a voltaic cell) turns the energy of a spontaneous redox reaction into electrical energy by keeping the oxidation and reduction apart.
Parts of a Galvanic Cell
| Part | What it does in the Zn-Cu cell |
|---|---|
| Anode | The electrode where oxidation happens. . It is the negative electrode of a galvanic cell and loses mass as zinc dissolves. |
| Cathode | The electrode where reduction happens. . It is the positive electrode and gains mass as copper plates onto it. |
| External circuit | The wire (and any device, such as a voltmeter or bulb). Electrons always travel through it from anode to cathode, and never through the solutions. |
| Salt bridge or porous barrier | Contains a non-reacting electrolyte such as . Anions move toward the anode and cations move toward the cathode. Without it, the zinc beaker would build up extra and the copper beaker extra negative ions, and that charge imbalance would stop the current almost instantly. |
| Electrolytes | The solutions containing the ions of each half-cell, e.g. \text{Zn{(NO}_3)}_2 and \text{Cu{(NO}_3)}_2. Ions carry the charge through the solutions. |
Cell notation describes a whole cell in one line. The anode goes on the left and the cathode on the right. A single vertical line marks a phase boundary (such as solid metal touching solution), and a double line marks the salt bridge. If a half-reaction has no solid metal to use as an electrode, as with or , an unreactive electrode such as platinum or graphite is used and written at the outer end.
Standard Reduction Potentials and the Hydrogen Electrode
A voltmeter can only measure the difference in electrical potential between two half-cells, never one half-cell by itself. So chemists pick one half-cell as a reference and give it a value of exactly zero: the standard hydrogen electrode (SHE). It is a platinum electrode dipped in a solution with , with hydrogen gas bubbling over it at about 100 kPa, usually at 25 °C. Every other half-cell is connected to the SHE, and the measured voltage becomes that half-cell's standard reduction potential, . All values are written for the reduction direction.
- A more positive means the species is more easily reduced, so it is a stronger oxidizing agent (e.g. , about +2.87 V).
- A more negative means the reduction is less favourable, so the reverse reaction (oxidation of the metal) is favoured. The metal is a stronger reducing agent (e.g. Li, about −3.04 V).
- In a galvanic cell, the half-cell with the higher is the cathode; the one with the lower is the anode.
- Typical values used in this topic: +0.80 V, +0.34 V, −0.13 V, −0.26 V, −0.76 V, −1.66 V, −2.37 V. Values may differ slightly from the table you use.
Calculating Cell Potential and Predicting Spontaneity
For the zinc-copper cell, A positive cell potential means the reaction as written is spontaneous and the cell can do electrical work. A negative value means the reaction is not spontaneous as written; the reverse reaction is the spontaneous one. Standard potentials are for standard conditions (1 mol/L solutions, gases at about 100 kPa, 25 °C); a real cell's voltage drifts as concentrations change.