Electrochemistry Guide: Cells, Electrodes, and the Nernst Equation
Every time you start a car, use a smartphone, or recharge a laptop, you rely on electrochemistry. Batteries convert chemical energy into electrical energy through spontaneous redox reactions. Electroplating deposits a thin layer of silver onto a tarnished spoon. Electrolysis splits water into hydrogen and oxygen for clean fuel.
Electrochemistry bridges chemistry and electricity. It quantifies the relationship between chemical reactions and electrical energy, providing tools for energy storage, materials synthesis, and chemical analysis.
Electrochemical Cells
Electrochemical cells consist of two electrodes connected by an external circuit and immersed in electrolyte solutions. There are two types. Galvanic (voltaic) cells produce electrical energy from spontaneous chemical reactions. Electrolytic cells use electrical energy to drive nonspontaneous reactions.
Galvanic Cells
In a galvanic cell, oxidation occurs at the anode (negative electrode) and reduction occurs at the cathode (positive electrode). Electrons flow from the anode through the external circuit to the cathode. A salt bridge maintains charge neutrality by allowing ion movement between the two half-cells.
The classic Daniell cell has a zinc anode in ZnSO4 solution and a copper cathode in CuSO4 solution. Spontaneous reaction: Zn + Cu2+ → Zn2+ + Cu. Electrons flow from zinc to copper. The salt bridge (typically KNO3 in agar) allows NO3- to migrate toward the anode and K+ toward the cathode.
Cell Notation
Cell notation provides a shorthand for describing electrochemical cells. The Daniell cell is written: Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s). The anode is on the left, cathode on the right. Single vertical lines represent phase boundaries. Double vertical lines represent the salt bridge.
Standard Reduction Potentials
Every half-reaction has a characteristic tendency to occur as a reduction. The standard reduction potential (E°) measures this tendency relative to the standard hydrogen electrode (SHE), which is assigned E° = 0.00 V.
Tabulated E° values allow predicting reaction spontaneity. More positive E° means greater tendency for reduction. Cu2+ + 2 e- → Cu has E° = +0.34 V. Zn2+ + 2 e- → Zn has E° = -0.76 V. This means Cu2+ is more easily reduced than Zn2+.
Calculating Cell Potential
The standard cell potential is E°cell = E°cathode - E°anode. For the Daniell cell: E°cell = +0.34 - (-0.76) = +1.10 V. The positive value confirms the reaction is spontaneous.
For any spontaneous reaction, the cell potential is related to Gibbs free energy: ΔG° = -nFE°cell, where n is the number of electrons transferred and F is the Faraday constant (96,485 C/mol). A positive E°cell corresponds to negative ΔG°, confirming spontaneity. This connects to thermochemistry.
The Nernst Equation
Concentration affects cell potential. The Nernst equation relates cell potential to concentration: E = E° - (RT/nF) ln Q, where Q is the reaction quotient. At 25°C, using base-10 logarithms: E = E° - (0.0592/n) log Q.
For the Daniell cell with [Cu2+] = 0.010 M and [Zn2+] = 1.0 M: Q = [Zn2+]/[Cu2+] = 1.0/0.010 = 100. E = 1.10 - (0.0592/2) log(100) = 1.10 - 0.0296 × 2 = 1.04 V. The lower Cu2+ concentration reduces the cell potential.
The Nernst equation also explains how concentration cells work. Two half-cells with the same electrode but different ion concentrations produce a voltage. This principle is used in pH meters and ion-selective electrodes.
Batteries
Batteries are galvanic cells designed for practical use. Primary batteries are disposable (single use). Secondary batteries are rechargeable.
Lead-Acid Battery
Lead-acid batteries power most automobiles. The anode is lead, the cathode is lead dioxide, and the electrolyte is concentrated sulfuric acid. Each cell produces about 2.0 V, and six cells in series provide 12 V. The reaction is reversible, allowing recharging.
Anode: Pb + HSO4- → PbSO4 + H+ + 2 e-. Cathode: PbO2 + HSO4- + 3 H+ + 2 e- → PbSO4 + 2 H2O. Overall: Pb + PbO2 + 2 HSO4- + 2 H+ → 2 PbSO4 + 2 H2O.
Lithium-Ion Battery
Lithium-ion batteries have the highest energy density of common rechargeable batteries. Lithium ions move from anode to cathode during discharge and reverse during charging. The lightweight lithium provides high voltage (about 3.7 V per cell) and excellent energy-to-weight ratio.
Lithium-ion batteries power nearly all portable electronics — smartphones, laptops, tablets, and increasingly electric vehicles. Their development earned John Goodenough, Stanley Whittingham, and Akira Yoshino the 2019 Nobel Prize in Chemistry.
Fuel Cells
Fuel cells continuously consume fuel (typically hydrogen) and oxidant (oxygen from air) to produce electricity. Unlike batteries, they do not store energy — they convert fuel energy directly to electricity with high efficiency.
The hydrogen fuel cell uses a proton exchange membrane. Hydrogen is oxidized at the anode: 2 H2 → 4 H+ + 4 e-. Oxygen is reduced at the cathode: O2 + 4 H+ + 4 e- → 2 H2O. The only byproduct is water, making fuel cells a clean energy technology.
Electrolysis
Electrolytic cells use electrical energy to drive nonspontaneous reactions. Electrolysis of water: 2 H2O → 2 H2 + O2 requires applying a voltage greater than 1.23 V. Hydrogen gas forms at the cathode, oxygen at the anode.
Faraday’s laws of electrolysis relate the amount of substance produced to the quantity of charge passed. One mole of electrons carries 96,485 coulombs (one faraday). The mass of product = (I × t × M)/(n × F), where I is current in amperes, t is time in seconds, M is molar mass, and n is electrons transferred.
Industrial Electrolysis
The Hall-Héroult process produces aluminum metal by electrolysis of Al2O3 dissolved in molten cryolite. This process made aluminum affordable — before 1886, aluminum was more expensive than gold.
Chlor-alkali electrolysis produces chlorine gas, hydrogen gas, and sodium hydroxide from brine (NaCl solution). These chemicals are essential for water treatment, plastics manufacturing, and countless industrial processes. Electroplating deposits a thin metal coating using electrolysis, connecting to redox reactions.
Electrochemical Corrosion
Corrosion is the destructive electrochemical oxidation of metals. Rusting of iron involves iron oxidation at anodic sites and oxygen reduction at cathodic sites. The electrochemical nature of corrosion explains why salt accelerates rusting — it increases electrolyte conductivity.
Protection methods include painting (barrier coating), galvanizing (zinc coating), cathodic protection (sacrificial anodes), and alloying (stainless steel contains chromium that forms a protective oxide layer).
Electrochemical Sensors
Electrochemical sensors detect analytes by measuring electrical signals produced by redox reactions. The glucose meter used by diabetics measures the current produced by glucose oxidation catalyzed by glucose oxidase. The current is proportional to blood glucose concentration, enabling precise insulin dosing.
pH electrodes measure the potential difference between a reference electrode and a glass membrane sensitive to H+ concentration. The Nernst equation governs the response: each pH unit change produces about 59 mV at 25°C. Modern ion-selective electrodes can detect fluoride, calcium, potassium, and many other ions directly in solution.
Oxygen Sensors
Lambda sensors in automotive catalytic converters measure oxygen concentration in exhaust gases. A zirconia electrolyte produces a voltage that changes sharply at the stoichiometric air-fuel ratio. This signal allows the engine control unit to maintain optimal combustion conditions.
Clark-type oxygen electrodes measure dissolved oxygen in environmental and biological samples. Oxygen diffuses through a membrane and is reduced at a platinum cathode, producing a current proportional to oxygen concentration. These sensors are essential for monitoring aquatic ecosystem health.
Corrosion Prevention at Scale
Industrial corrosion prevention represents a significant engineering challenge. Pipelines are protected by a combination of coatings, cathodic protection, and corrosion inhibitors. Impressed current cathodic protection systems use an external power source to drive electrons into the structure, forcing it to act as a cathode.
Offshore oil platforms face aggressive corrosion from saltwater. Sacrificial aluminum anodes are attached to the platform structure. The aluminum corrodes instead of the steel, and the anodes are replaced periodically. Each platform may use hundreds of anodes weighing 50-100 kg each, replaced every 3-5 years.
Frequently Asked Questions
What is the difference between a galvanic and electrolytic cell? Galvanic cells produce electricity from spontaneous reactions. Electrolytic cells consume electricity to drive nonspontaneous reactions. Galvanic cells have positive cell potentials; electrolytic cells require external voltage.
How does a battery produce electricity? A battery converts chemical energy to electrical energy through spontaneous redox reactions. Electrons released at the anode flow through the external circuit to the cathode, providing useful electrical work.
What is the Faraday constant? The Faraday constant (96,485 C/mol) is the charge of one mole of electrons. It relates the amount of charge passed in electrolysis to the amount of chemical change.
Why do batteries have a specific voltage? The voltage is determined by the standard reduction potentials of the electrode materials and the concentrations of the electrolytes, as described by the Nernst equation.
Redox Reactions — Thermochemistry Guide — Acid-Base Chemistry