Electrode potentials and the SHE

A galvanic cell is two half-cells (redox couples). In the Daniell cell, Cu²⁺ + 2e⁻ → Cu is the reduction half on copper and Zn → Zn²⁺ + 2e⁻ the oxidation half on zinc. A salt bridge joins the two solutions; if both electrodes share one electrolyte, no bridge is needed.

At each metal-solution interface, ions tend to deposit (charging the metal positive) while metal atoms tend to dissolve as ions (leaving electrons, charging it negative). The resulting potential difference is the electrode potential; with every species at unit concentration it is the standard electrode potential. By IUPAC convention, standard reduction potentials are called standard electrode potentials.

Anode = oxidation, negative with respect to its solution in a galvanic cell. Cathode = reduction, positive. Electrons leave the anode through the wire; conventional current runs the opposite way.

Cell notation writes the anode first, on the left, and the cathode last, on the right, with a single bar between metal and solution and a double bar for the salt bridge: Cu(s)|Cu²⁺(aq)||Ag⁺(aq)|Ag(s). Then E_cell = E_right − E_left = E_cathode − E_anode, and the emf is the value with no current drawn.

A single electrode's potential cannot be measured alone. The standard hydrogen electrode, Pt(s)|H₂(g, 1 bar)|H⁺(aq, 1 M), made of platinum coated with platinum black in acid with H₂ bubbling, is assigned 0 V at all temperatures.

Against the SHE (as anode), the Cu²⁺/Cu cell reads 0.34 V and the Zn²⁺/Zn cell −0.76 V, so E°(Daniell) = 0.34 − (−0.76) = 1.10 V. Positive E° for Cu means H⁺ cannot oxidise copper, so Cu does not dissolve in HCl (nitric acid attacks it through nitrate); negative E° for Zn means H⁺ does oxidise zinc.

Platinum or gold serve as inert electrodes: they carry electrons and offer a surface but do not react, as in Pt(s)|H₂(g)|H⁺(aq) and Pt(s)|Br₂(aq)|Br⁻(aq).

Down the table of standard potentials, from F₂/F⁻ (2.87 V) to Li⁺/Li (−3.05 V), oxidising power of the left-hand species falls and reducing power of the right-hand species rises. F₂ is the strongest oxidising agent and F⁻ the weakest reductant; Li is the strongest reducing agent in water and Li⁺ the weakest oxidant.

Cell measurements also give pH, solubility products, equilibrium constants and other thermodynamic data, and are used in potentiometric titrations.