Electrochemical Cells Overview

A device that transforms chemical energy into electrical energy and vice versa. It consists of two electrodes immersed in an electrolyte.

The electrode where oxidation occurs. Electrons flow from the anode to the cathode.

The electrode where reduction occurs. It receives electrons from the anode.

A substance that allows ions to flow between electrodes in an electrochemical cell.

A chemical reaction where an atom, molecule, or ion loses electrons. It is characterized by an increase in oxidation state.

A chemical reaction where an atom, molecule, or ion gains electrons. It is characterized by a decrease in oxidation state.

An electrochemical cell that produces electrical energy from chemical reactions. The reactions are spontaneous.

An electrochemical cell that uses electrical energy to drive a non-spontaneous chemical reaction.

A simple galvanic cell where a zinc rod is dipped in a zinc sulfate solution and a copper rod is dipped in a copper sulfate solution. Electrons flow from zinc to copper.

In a galvanic cell, the anode is the negative terminal where oxidation takes place. The oxidation half-cell is the part where oxidation occurs.

In a galvanic cell, the cathode is the positive terminal where reduction takes place. The reduction half-cell is the part where reduction occurs.

In a galvanic cell, electrons flow from the anode (negative terminal) to the cathode (positive terminal).

A component in a galvanic cell that connects the two half cells, allowing ions to pass between them. It prevents intermixing of solutions and maintains electrical neutrality.

A shorthand notation representing the components of an electrochemical cell. The anode is on the left, the cathode on the right, separated by a double line representing the salt bridge.

The potential difference between an electrode and its electrolyte solution. It depends on the nature of the metal, ion concentration, and temperature.

The potential difference between the two electrodes in an electrochemical cell. It is expressed in volts.

A list of elements arranged in order of increasing reduction potential values. It is also known as the activity series.

Metals higher in the electrochemical series are stronger reducing agents, while those lower are stronger oxidizing agents. The series helps predict the spontaneity of reactions.

A formula used to calculate the electrode potential of a half-cell. It takes into account the standard electrode potential, concentration of ions, and temperature.

The electrode potential of an electrode under standard conditions: 1 M concentration, 298 K temperature, 1 atm pressure for gases.

At equilibrium, the cell reaction reaches a state where the cell potential is zero (E = 0), and the free energy change for the reaction is also zero (ΔG = 0).

A constant that expresses the ratio of products to reactants at equilibrium, indicating the extent to which a reaction proceeds to completion.

The Nernst equation can be used to calculate the equilibrium constant (K) of a reaction using the standard cell potential (E°) and the number of electrons transferred (n).

Using the Nernst equation, one can calculate the concentration of an ionic species in a cell if the concentration of the other species and the cell potential are known.

The Nernst equation can be used to determine the pH of a solution in an electrochemical cell by relating the cell potential, standard cell potential, and the concentration of hydrogen ions.

The electrochemical cell provides a means to measure the pH of a solution by relating the cell potential to the concentration of hydrogen ions.

The Nernst equation, used to calculate electrode potential, considers the influence of ion concentration on the cell potential by measuring the deviation from standard electrode potential.

A cell that uses the difference in concentration of a particular ion between two compartments to generate an electric potential.

A cell that uses the difference in partial pressure of a gas between two compartments to generate an electric potential.

An electrochemical cell where the electrode and electrolyte are in direct contact. It simplifies the setup and eliminates the need for a salt bridge.