Electrochemical Reactions
Electrochemistry studies the relationship between chemical reactions and electrical energy. Every battery you use, every rusting nail, and every electroplated surface involves electrochemical processes. At the heart of electrochemistry lies the transfer of electrons between chemical species — oxidation and reduction reactions harnessed to generate electricity or driven by it.
Oxidation and Reduction Revisited
Oxidation is the loss of electrons; reduction is the gain of electrons. These two half-reactions always occur together — you cannot have one without the other. The species that loses electrons is oxidized (and acts as the reducing agent), while the species that gains electrons is reduced (and acts as the oxidizing agent).
In the reaction between zinc metal and copper(II) sulfate solution:
Zn(s) → Zn²⁺(aq) + 2e⁻ (oxidation)
Cu²⁺(aq) + 2e⁻ → Cu(s) (reduction)
If zinc is simply dipped into copper sulfate solution, the electron transfer happens directly at the metal surface, releasing heat. But if we separate the two half-reactions into different compartments, we can force the electrons to travel through an external wire — generating usable electrical current.
Galvanic (Voltaic) Cells
A galvanic cell converts chemical energy into electrical energy through a spontaneous redox reaction. It consists of two half-cells, each containing an electrode immersed in an electrolyte solution.
Anode — The electrode where oxidation occurs. In a Zn/Cu cell, the zinc strip is the anode. Zinc atoms lose electrons and dissolve as Zn²⁺ ions. The anode carries a negative charge in a galvanic cell because it is the source of electrons.
Cathode — The electrode where reduction occurs. Copper ions in solution gain electrons at the copper strip and deposit as metallic copper. The cathode is positive because electrons flow toward it.
Salt bridge — A tube filled with an inert electrolyte (such as KNO₃ in agar gel) that connects the two half-cells. It allows ions to migrate between solutions, maintaining electrical neutrality. Without a salt bridge, charge buildup would quickly halt the reaction.
External wire — Electrons flow from the anode through the wire to the cathode, powering any device connected in the circuit.
Cell Notation
Electrochemists use a shorthand notation to describe cells:
Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)
The single vertical lines (|) represent phase boundaries, and the double vertical line (||) represents the salt bridge. The anode is written on the left, the cathode on the right.
Standard Reduction Potentials
Every half-reaction has a standard reduction potential (E°), measured in volts relative to the standard hydrogen electrode (SHE), which is assigned E° = 0.00 V by convention.
The standard cell potential is calculated as:
E°_cell = E°_cathode − E°_anode
For the Zn/Cu cell:
E°_cell = (+0.34 V) − (−0.76 V) = +1.10 V
A positive E°_cell indicates a spontaneous reaction under standard conditions. The more positive the cell potential, the greater the driving force for electron transfer.
The Nernst Equation
Under non-standard conditions (concentrations other than 1 M, pressures other than 1 atm), the cell potential deviates from E°. The Nernst equation quantifies this:
E = E° − (RT / nF) × ln Q
At 25 °C, this simplifies to:
E = E° − (0.0592 / n) × log Q
where n is the number of electrons transferred and Q is the reaction quotient. As a cell discharges and Q increases, E decreases until equilibrium is reached (E = 0, Q = K).
Electrolytic Cells
While galvanic cells harness spontaneous reactions, electrolytic cells use external electrical energy to drive non-spontaneous reactions. The key differences are:
- An external power source (battery or power supply) forces current through the cell
- The anode is positive and the cathode is negative (opposite to galvanic cells)
- The reaction would not occur without the input of electrical energy
Electrolysis of water splits water into hydrogen and oxygen gases when a voltage greater than 1.23 V is applied (in practice, about 1.8–2.0 V due to overpotential). This reaction is central to green hydrogen production.
Electroplating deposits a thin layer of metal onto a surface. The object to be plated serves as the cathode, and the plating metal (such as gold, chromium, or nickel) is either the anode or dissolved in the electrolyte. Faraday's laws of electrolysis relate the amount of metal deposited to the total charge passed through the cell.
Applications
Batteries are galvanic cells packaged for practical use. Primary batteries (alkaline, lithium) are single-use. Secondary batteries (lead-acid, lithium-ion) are rechargeable — during charging, the cell operates in electrolytic mode, reversing the discharge reaction.
Fuel cells continuously convert chemical fuel (typically hydrogen) and an oxidant (oxygen) into electricity, water, and heat. Unlike batteries, fuel cells do not store energy internally — they operate as long as fuel is supplied. Proton exchange membrane (PEM) fuel cells power vehicles and portable electronics.
Corrosion is an unwanted electrochemical process in which metals are oxidized by environmental agents (water, oxygen, salts). Iron rusting is a galvanic cell in miniature, with anodic and cathodic regions forming on the metal surface. Corrosion prevention strategies — galvanizing, cathodic protection, protective coatings — all exploit electrochemical principles.
Industrial electrolysis produces essential chemicals: chlorine and sodium hydroxide from brine (chlor-alkali process), aluminum from alumina (Hall-Heroult process), and copper refined to 99.99% purity for electrical wiring.