Electrochemistry is the branch of science that deals with the relationship between chemical reactions and electricity. It works in two directions: chemical reactions can produce electrical current (as in a battery), and electrical current can drive chemical reactions that wouldn’t happen on their own (as in electroplating). This two-way exchange between chemical and electrical energy underpins everything from the battery in your phone to the way your nerve cells fire.
How It Works: Electrons on the Move
Every electrochemical process relies on a type of reaction called a redox reaction, short for reduction-oxidation. In these reactions, one substance loses electrons (oxidation) while another gains them (reduction). The movement of those electrons from one substance to another is what creates, or is created by, an electric current.
This happens at two surfaces called electrodes. The anode is where oxidation occurs, meaning atoms give up electrons. The cathode is where reduction occurs, meaning atoms or ions pick up electrons. Electrons travel through an external wire from the anode to the cathode, and that flow of electrons is the electrical current you can harness to do useful work. Meanwhile, inside the liquid or paste between the electrodes (the electrolyte), charged particles called ions migrate to complete the circuit. Without that ion movement through the electrolyte, the reaction would stall almost immediately because charge would build up with nowhere to go.
Two Types of Electrochemical Cells
Electrochemical setups fall into two main categories based on whether the reaction happens on its own or needs a push.
- Galvanic (voltaic) cells produce electricity from a spontaneous chemical reaction. Batteries are galvanic cells. The chemicals inside want to react, and the cell is designed so the electrons have to travel through your device on their way from one electrode to the other.
- Electrolytic cells use electricity to force a non-spontaneous reaction to occur. If you run current through water to split it into hydrogen and oxygen gas, that’s electrolysis. The reaction would never happen on its own; the external power source provides the energy.
The electrode labels stay the same in both types: oxidation at the anode, reduction at the cathode. What changes is the direction of electron flow relative to an outside power source. In a galvanic cell, the reaction itself pushes electrons out. In an electrolytic cell, an external battery or power supply pushes electrons in through the cathode.
Batteries and Energy Storage
Modern rechargeable batteries, like the lithium-ion cells in laptops and electric vehicles, are electrochemistry in action. When a lithium-ion battery discharges, lithium atoms nestled between layers of graphite at the negative electrode release an electron and become lithium ions. Those ions travel through the electrolyte to the positive electrode, typically a metal oxide material, where they slot into tiny channels or tunnels in the crystal structure. The electrons, meanwhile, flow through the external circuit powering your device.
Charging reverses this process. An applied voltage pushes lithium ions back out of the positive electrode and across to the graphite, where they re-intercalate (settle back between the carbon layers). This shuttling of lithium back and forth is why these batteries can be recharged hundreds of times. The voltage a cell produces depends on how strongly the two electrode materials want to exchange lithium, and it changes as the battery drains because the concentration of available lithium shifts at each electrode.
Fuel Cells: Turning Hydrogen Into Electricity
A hydrogen fuel cell is a galvanic cell that runs on a continuous fuel supply rather than stored chemicals. At the anode, hydrogen gas splits into hydrogen ions and electrons. The electrons travel through an external circuit (powering a motor, for instance), while the hydrogen ions pass through a membrane to the cathode. There, they meet oxygen from the air, combine with the arriving electrons, and form water. The only byproduct is water vapor, which is why fuel cells are attractive for clean energy.
Corrosion: Electrochemistry You Don’t Want
Rust is an electrochemical reaction that happens without any wires or batteries. When a droplet of water sits on an iron surface, the iron beneath the droplet dissolves, releasing electrons. Those electrons travel through the metal to the edge of the droplet, where dissolved oxygen from the air picks them up. The result is iron ions and hydroxide ions, which combine to form iron hydroxide. That compound quickly reacts with more oxygen to become the reddish-brown iron oxide we call rust.
This is essentially a tiny galvanic cell that forms spontaneously on the metal surface. The iron under the center of the droplet acts as the anode, and the area near the droplet’s edge acts as the cathode. Understanding corrosion as an electrochemical process is what allows engineers to prevent it, for example by coating metals with zinc (galvanizing), which preferentially corrodes instead of the underlying steel.
Electrochemistry in Your Body
Your nervous system runs on electrochemistry. A resting nerve cell maintains a voltage of about -70 millivolts across its membrane by keeping sodium ions concentrated outside and potassium ions concentrated inside. When a signal arrives, sodium channels open, and positively charged sodium ions rush into the cell, rapidly swinging the voltage toward +60 millivolts. This depolarization triggers neighboring channels to open in a chain reaction that races down the nerve fiber.
Almost immediately afterward, potassium channels open, and potassium ions flow out, dragging the voltage back down toward -85 millivolts. A protein pump then restores the original ion balance by pushing sodium back out and pulling potassium back in. The whole cycle takes just a few milliseconds, but it’s the fundamental mechanism behind every thought, sensation, and muscle movement.
Glucose Monitors and Biosensors
The blood glucose monitors used by millions of people with diabetes are electrochemical biosensors. A test strip contains an enzyme that reacts specifically with glucose in a drop of blood. That reaction produces a small molecule (in early designs, hydrogen peroxide) that gets oxidized at a tiny platinum electrode on the strip. The electrode detects the flow of electrons from that oxidation, and because the number of electrons is directly proportional to the amount of glucose present, the meter can display a concentration reading in seconds.
Newer generations of glucose sensors replaced the oxygen-dependent chemistry with synthetic molecules called redox mediators that shuttle electrons from the enzyme directly to the electrode. This made the sensors more reliable and less sensitive to variations in oxygen levels, which is one reason continuous glucose monitors have become practical for everyday use. The majority of commercial glucose biosensors today are electrochemical rather than optical, because they offer better sensitivity, lower cost, and simpler maintenance.
Faraday’s Laws: The Math Behind It
One of the most useful principles in electrochemistry is the direct relationship between electricity and the amount of chemical change it produces. Michael Faraday established that the mass of a substance deposited or dissolved at an electrode is directly proportional to the total electric charge passed through the system. If you double the current or run it for twice as long, you get exactly twice the material.
The key number is the Faraday constant: 96,485 coulombs per mole. That’s the total electric charge carried by one mole (about 602 billion trillion) of electrons. If you know the current, the time, and the number of electrons involved in the reaction, you can calculate exactly how much material will be produced. This relationship is the basis for industrial electroplating, metal refining, and the design of every battery, where engineers need to predict precisely how much energy a given amount of material can store.
Electrochemical Carbon Capture
One of the newer frontiers for electrochemistry is pulling carbon dioxide out of the atmosphere or the ocean. Several companies are developing systems that use electrochemical reactions to selectively capture CO₂ from air. Some approaches aim to do this at costs as low as $50 to $70 per ton of CO₂, with energy requirements around 1.0 gigajoule per ton. Other projects focus on extracting dissolved CO₂ from seawater at a target cost below $100 per ton, and one initiative in the Al Hajar mountains of Oman is working on electrochemically accelerated mineralization that could lock away carbon at a gigaton scale. These technologies are still scaling up, but they illustrate how broadly applicable electrochemistry has become, extending well beyond traditional batteries and industrial processes.