Editing Electrochemical cell (section)

=== Galvanic cell ===
{{main|Galvanic cell}}
A galvanic cell (voltaic cell), named after [[Luigi Galvani]] ([[Alessandro Volta]]), is an electrochemical cell that generates electrical energy from spontaneous [[redox]] reactions.<ref name=":32">''Chemistry,'' Rice University, 2015. [Online]. Available: https://web.ung.edu/media/Chemistry2/Chemistry-LR.pdf</ref>
[[File:Galvanic cell with no cation flow.svg|thumb|Galvanic cell with no cation flow]]
A wire connects two different [[metal]]s (e.g. [[zinc]] and [[copper]]). Each metal is in a separate solution; often the [[Aqueous solution|aqueous]] [[Sulfate|sulphate]] or [[nitrate]] forms of the metal, however more generally metal salts and water which conduct [[Electric current|current]].<ref>{{Cite book |last=Ahmad |first=Dr. Zaki |url=http://worldcat.org/oclc/857524149 |title=Principles of corrosion engineering and corrosion control. |date=2013 |publisher=Butterworth-Heinemann |isbn=978-0-08-097134-6 |oclc=857524149}}</ref> A [[salt bridge]] or porous membrane connects the two solutions, keeping electric neutrality and the avoidance of charge accumulation. The metal's differences in oxidation/reduction potential drive the reaction until [[Chemical equilibrium|equilibrium]].<ref name=":23"/>

Key features:
* [[Spontaneous process|spontaneous reaction]]
* generates electric current
* current flows through a wire, and [[ion]]s flow through a salt bridge
* [[anode]] (negative), [[cathode]] (positive)

==== Half cells ====
Galvanic cells consists of two half-cells.  Each half-cell consists of an [[electrode]] and an [[electrolyte]] (both half-cells may use the same or different electrolytes).{{cn|date=December 2024}}

The chemical reactions in the cell involve the electrolyte, electrodes, and/or an external substance ([[fuel cell]]s may use [[Hydrogen|hydrogen gas]] as a [[Reagent|reactant]]). In a full electrochemical cell, species from one half-cell lose electrons ([[Redox|oxidation]]) to their electrode while species from the other half-cell gain electrons ([[Redox|reduction]]) from their electrode.{{cn|date=December 2024}}

A ''[[salt bridge]]'' (e.g., filter paper soaked in KNO<sub>3,</sub> NaCl, or some other electrolyte) is used to ionically connect two half-cells with different electrolytes, but it prevents the solutions from mixing and unwanted side reactions. An alternative to a salt bridge is to allow direct contact (and mixing) between the two half-cells, for example in simple [[electrolysis of water]].{{cn|date=December 2024}}

As electrons flow from one half-cell to the other through an external [[Electrical network|circuit]], a difference in charge is established. If no ionic contact were provided, this charge difference would quickly prevent the further flow of electrons. A salt bridge allows the flow of negative or positive ions to maintain a steady-state charge distribution between the oxidation and reduction vessels, while keeping the contents otherwise separate. Other devices for achieving separation of solutions are porous pots and gelled solutions. A porous pot is used in the [[Bunsen cell]].{{cn|date=December 2024}}

==== Equilibrium reaction ====
Each half-cell has a characteristic voltage (depending on the metal and its characteristic reduction potential). Each reaction is undergoing an [[Chemical equilibrium|equilibrium]] reaction between different [[oxidation states]] of the ions: when equilibrium is reached, the cell cannot provide further [[voltage]]. In the half-cell performing oxidation, the closer the equilibrium lies to the ion/atom with the more positive oxidation state the more potential this reaction will provide.<ref name=":23" /> Likewise, in the reduction reaction, the closer the equilibrium lies to the ion/atom with the more ''negative'' oxidation state the higher the potential.{{cn|date=December 2024}}

==== Cell potential ====
The cell potential can be predicted through the use of [[electrode potential]]s (the voltages of each half-cell). These half-cell potentials are defined relative to the assignment of 0 [[volt]]s to the [[standard hydrogen electrode]] (SHE). (See [[table of standard electrode potentials]]). The difference in voltage between electrode potentials gives a prediction for the potential measured. When calculating the difference in voltage, one must first rewrite the half-cell reaction equations to obtain a balanced oxidation-reduction equation.{{cn|date=December 2024}}

# Reverse the reduction reaction with the smallest potential (to create an oxidation reaction/overall positive cell potential)
# Half-reactions must be multiplied by integers to achieve electron balance.

Cell potentials have a possible range of roughly zero to 6 volts. Cells using water-based electrolytes are usually limited to cell potentials less than about 2.5 volts due to high reactivity of the powerful oxidizing and reducing agents with water which is needed to produce a higher voltage. Higher cell potentials are possible with cells using other [[solvent]]s instead of water. For instance, [[Lithium battery|lithium cells]] with a voltage of 3 volts are commonly available.{{cn|date=December 2024}}

The cell potential depends on the [[concentration]] of the reactants, as well as their type. As the cell is discharged, the concentration of the reactants decreases and the cell potential also decreases.{{cn|date=December 2024}}