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{{Short description|Particle, atom or molecule with a net electrical charge}}
{{Other uses|Ion (disambiguation)}}
{{pp-semi-indef}}
{{pp-move-indef|small=yes}}
{{EngvarB|date=September 2016}}
[[File:Ionic bonding.svg|thumb|right|300px|Electron transfer between [[lithium]] (Li) and [[fluorine]] (F). Forming an ionic bond, Li and F become Li{{sup|+}} and F{{sup|−}} ions.]]

An '''ion''' ({{IPAc-en|ˈ|aɪ|.|ɒ|n|,_|-|ən}})<ref>{{cite Collins Dictionary|ion}} {{Webarchive |url=https://web.archive.org/web/20131224095801/http://www.collinsdictionary.com/dictionary/english/ion |date=2013-12-24 |status=live}}.</ref> is an [[atom]] or [[molecule]] with a net [[electric charge|electrical charge]]. The charge of an [[electron]] is considered to be negative by convention and this charge is equal and opposite to the charge of a [[proton]], which is considered to be positive by convention. The net charge of an ion is not zero because its total number of electrons is unequal to its total number of protons.

A '''cation''' is a positively charged ion with fewer electrons than protons<ref>{{cite Merriam-Webster|cation |access-date=2021-10-06}} {{Webarchive |url=https://web.archive.org/web/20211006215047/https://www.merriam-webster.com/dictionary/cation  |date=2021-10-06 |status=live}}.</ref> while an '''anion''' is a negatively charged ion with more electrons than protons.<ref>{{cite Merriam-Webster|anion |access-date=2021-10-06}} {{Webarchive |url=https://web.archive.org/web/20211006215049/https://www.merriam-webster.com/dictionary/anion  |date=2021-10-06 |status=live}}.</ref> Opposite electric charges are pulled towards one another by [[electrostatic force]], so cations and anions attract each other and readily form [[ionic compound]]s.

Ions consisting of only a single atom are termed atomic or [[monatomic ion]]s, while two or more atoms form molecular ions or [[polyatomic ion]]s. In the case of physical ionization in a fluid (gas or liquid), "ion pairs" are created by spontaneous molecule collisions, where each generated pair consists of a free electron and a positive ion.<ref name="knoll">{{cite book |first=Glenn F.|last=Knoll |title=Radiation Detection and Measurement |edition=3rd |date=1999 |location=New York |publisher=[[Wiley (publisher)|Wiley]] |isbn=978-0-471-07338-3}}</ref> Ions are also created by chemical interactions, such as the [[dissolution (chemistry)|dissolution]] of a [[salt (chemistry)|salt]] in liquids, or by other means, such as passing a [[direct current]] through a conducting solution, dissolving an [[anode]] via [[ionization]].

==History of discovery{{anchor|History of discovery}}==
The word ''ion'' was coined from neuter present participle of 
Greek ἰέναι (''ienai''), meaning "to go". A cation is something that moves down ({{lang-el|κάτω}}, ''kato'', meaning "down") and an anion is something that moves up ({{lang-el|ἄνω}}, ''ano'', meaning "up"). They are so called because ions move toward the electrode of opposite charge. This term was introduced (after a suggestion by the English [[polymath]] [[William Whewell]])<ref name=whewell>{{cite book |url=https://books.google.com/books?id=9lknVoNGj30C&q=The%20Correspondence%20of%20Michael%20Faraday%20whewell&pg=PA183 |title=The Correspondence of Michael Faraday, Vol. 2: 1832–1840 |year=1991 |editor=Frank A. J. L. James |isbn=9780863412493 |page=183 |access-date=2020-10-16 |archive-date=2021-04-14 |archive-url=https://web.archive.org/web/20210414164907/https://books.google.com/books?id=9lknVoNGj30C&q=The%20Correspondence%20of%20Michael%20Faraday%20whewell&pg=PA183 |url-status=live }}</ref> by English physicist and chemist [[Michael Faraday]] in 1834 for the then-unknown species that ''goes'' from one [[electrode]] to the other through an aqueous medium.<ref>{{cite video | url=https://www.bbc.co.uk/history/historic_figures/faraday_michael.shtml | title=Michael Faraday (1791–1867) | publisher=[[BBC]] | location=UK}}</ref><ref>{{cite web | url=http://www.etymonline.com/index.php?term=ion | title=Online etymology dictionary | access-date=2011-01-07 | archive-date=2011-05-14 | archive-url=https://web.archive.org/web/20110514084635/http://www.etymonline.com/index.php?term=ion | url-status=live }}</ref> Faraday did not know the nature of these species, but he knew that since metals dissolved into and entered a solution at one electrode and new metal came forth from a solution at the other electrode; that some kind of substance has moved through the solution in a current. This conveys matter from one place to the other. In correspondence with Faraday, Whewell also coined the words ''[[anode]]'' and ''[[cathode]]'', as well as ''anion'' and ''cation'' as ions that are attracted to the respective electrodes.<ref name=whewell />

[[Svante Arrhenius]] put forth, in his 1884 dissertation, the explanation of the fact that solid crystalline salts [[Dissociation (chemistry)|dissociate]] into paired charged particles when dissolved, for which he would win the 1903 Nobel Prize in Chemistry.<ref>{{cite web|url=https://www.nobelprize.org/nobel_prizes/chemistry/laureates/1903/index.html|title=The Nobel Prize in Chemistry 1903|website=nobelprize.org|access-date=2017-06-13|archive-date=2018-07-08|archive-url=https://web.archive.org/web/20180708044958/https://www.nobelprize.org/nobel_prizes/chemistry/laureates/1903/index.html|url-status=live}}</ref> Arrhenius' explanation was that in forming a solution, the salt dissociates into Faraday's ions, he proposed that ions formed even in the absence of an electric current.<ref name="columbia">{{cite book|editor1-last=Harris|editor1-first=William|editor2-last=Levey|editor2-first=Judith|title=The New Columbia Encyclopedia|date=1976|publisher=[[Columbia University]]|location=New York City|isbn=978-0-231-03572-9|page=[https://archive.org/details/newcolumbiaencyc00harr/page/155 155]|edition=4th|url=https://archive.org/details/newcolumbiaencyc00harr/page/155}}</ref><ref name="EncBrit">{{cite book|editor1-last=McHenry|editor1-first=Charles|title=The New Encyclopædia Britannica|journal=Chicago: Encyclopaedia Britannica Inc|date=1992|publisher=[[Encyclopædia Britannica, Inc.]]|location=Chicago|isbn=978-0-85229-553-3|page=587|volume=1|edition=15|bibcode=1991neb..book.....G|last1=Goetz|first1=Philip W.}}</ref><ref name="SciBio">{{cite book|editor1-last=Cillispie|editor1-first=Charles|title=Dictionary of Scientific Biography|date=1970|publisher=[[Charles Scribner's Sons]]|location=New York City|isbn=978-0-684-10112-5|pages=296–302|edition=1}}</ref>

== Characteristics ==
Ions in their gas-like state are highly reactive and will rapidly interact with ions of opposite charge to give neutral molecules or ionic salts. Ions are also produced in the liquid or solid state when salts interact with solvents (for example, water) to produce ''solvated ions'', which are more stable, for reasons involving a combination of [[energy]] and [[entropy]] changes as the ions move away from each other to interact with the liquid. These stabilized species are more commonly found in the environment at low temperatures. A common example is the ions present in seawater, which are derived from dissolved salts.

As charged objects, ions are attracted to opposite electric charges (positive to negative, and vice versa) and repelled by like charges. When they move, their trajectories can be deflected by a [[magnetic field]].

Electrons, due to their smaller mass and thus larger space-filling properties as [[matter waves]], determine the size of atoms and molecules that possess any electrons at all. Thus, anions (negatively charged ions) are larger than the parent molecule or atom, as the excess electron(s) repel each other and add to the physical size of the ion, because its size is determined by its [[electron cloud]]. Cations are smaller than the corresponding parent atom or molecule due to the smaller size of the electron cloud. One particular cation (that of hydrogen) contains no electrons, and thus consists of a single proton – ''much smaller'' than the parent hydrogen atom.

===Anions and cations===
{{redirect-distinguish2|Anion|the quasiparticle [[Anyon]]}}
[[File:Ions.svg|thumb|upright=1.75|[[Hydrogen atom]] (center) contains a single [[proton]] and a single [[electron]]. Removal of the electron gives a cation (left), whereas the addition of an electron gives an anion (right). The [[hydrogen anion]], with its loosely held two-electron cloud, has a larger radius than the neutral atom, which in turn is much larger than the bare proton of the [[cation]]. Hydrogen forms the only charge-+1 cation that has no electrons, but even cations that (unlike hydrogen) retain one or more electrons are still smaller than the neutral atoms or molecules from which they are derived.]]

Anion (−) and cation (+) indicate the net electric charge on an ion. An ion that has more electrons than protons, giving it a net negative charge, is named an anion, and a minus indication "Anion (−)" indicates the negative charge. With a cation it is just the opposite: it has fewer electrons than protons, giving it a net positive charge, hence the indication "Cation (+)".

Since the electric charge on a proton is equal in magnitude to the charge on an electron, the net electric charge on an ion is equal to the number of protons in the ion minus the number of electrons.

An {{vanchor|anion}} (−) ({{IPAc-en|ˈ|æ|n|,|aɪ|.|ən}} {{Respell|ANN|eye|ən}}, from the Greek word ἄνω (''ánō''), meaning "up"<ref>{{cite web | url =http://oxfordindex.oup.com/view/10.1093/oi/authority.20110803095414154 | title =Oxford Reference: OVERVIEW anion | author =Oxford University Press | author-link =Oxford University Press | publisher =oxfordreference.com | date =2013 | access-date =2017-01-15 | archive-date =2017-01-18 | archive-url =https://web.archive.org/web/20170118065532/http://oxfordindex.oup.com/view/10.1093/oi/authority.20110803095414154 | url-status =live }}</ref>) is an ion with more electrons than protons, giving it a net negative charge (since electrons are negatively charged and protons are positively charged).<ref>{{cite web | url =http://ruby.colorado.edu/~smyth/G101-2.html | title =Atoms and Elements, Isotopes and Ions | author =University of Colorado Boulder | author-link =University of Colorado Boulder | publisher =colorado.edu | date =November 21, 2013 | access-date =November 22, 2013 | archive-date =February 2, 2015 | archive-url =https://web.archive.org/web/20150202061438/http://ruby.colorado.edu/~smyth/G101-2.html | url-status =live }}</ref>

A {{vanchor|cation}} (+) ({{IPAc-en|ˈ|k|æ|t|,|aɪ|.|ən}} {{Respell|KAT|eye|ən}}, from the Greek word κάτω (''káto''), meaning "down"<ref>{{cite web | url =http://oxfordindex.oup.com/view/10.1093/oi/authority.20110803095555447 | title =Oxford Reference: OVERVIEW cation | author =Oxford University Press | author-link =Oxford University Press | publisher =oxfordreference.com | date =2013 | access-date =2017-01-15 | archive-date =2017-01-18 | archive-url =https://web.archive.org/web/20170118065659/http://oxfordindex.oup.com/view/10.1093/oi/authority.20110803095555447 | url-status =live }}</ref>) is an ion with fewer electrons than protons, giving it a positive charge.<ref>{{cite web | url =http://www.usouthal.edu/geology/haywick/GY111/111-4.pdf | title =Elemental Chemistry | author =Douglas W. Haywick, Ph.D. | author2 =University of South Alabama | publisher =usouthal.edu | date =2007–2008 | author-link2 =University of South Alabama | access-date =2013-11-22 | archive-date =2011-12-04 | archive-url =https://web.archive.org/web/20111204134213/http://www.usouthal.edu/geology/haywick/GY111/111-4.pdf | url-status =live }}</ref>

There are additional names used for ions with multiple charges. For example, an ion with a −2 charge is known as a [[dianion]] and an ion with a +2 charge is known as a [[dication]]. A [[zwitterion]] is a neutral molecule with positive and negative charges at different locations within that molecule.<ref>{{cite web | url =http://chemed.chem.purdue.edu/genchem/topicreview/bp/1biochem/amino2.html#zwitter | title =Amino Acids | author =Purdue University | author-link =Purdue University | publisher =purdue.edu | date =November 21, 2013 | access-date =November 22, 2013 | archive-date =July 13, 2011 | archive-url =https://web.archive.org/web/20110713040227/http://chemed.chem.purdue.edu/genchem/topicreview/bp/1biochem/amino2.html#zwitter | url-status =live }}</ref>

Cations and anions are measured by their [[ionic radius]] and they differ in relative size: "Cations are small, most of them less than 10<sup>−10</sup> m (10<sup>−8</sup> cm) in radius. But most anions are large, as is the most common Earth anion, [[oxygen]]. From this fact it is apparent that most of the space of a [[crystal]] is occupied by the anion and that the cations fit into the spaces between them."<ref name="Press1986">{{cite book |last1=Press |first1=Frank |last2=Siever |first2=Raymond |title=Earth |edition=14th |publisher=[[W. H. Freeman and Company]] |publication-place=New York |year=1986 | isbn=0-7167-1743-3 | oclc=12556840 | page=63}}</ref>

The terms ''anion'' and ''cation'' (for ions that respectively travel to the anode and cathode during electrolysis) were [[#History of discovery|introduced by Michael Faraday in 1834]] following his consultation with [[William Whewell]].

===Natural occurrences===
Ions are ubiquitous in [[nature]] and are responsible for diverse phenomena from the luminescence of the Sun to the existence of the Earth's [[ionosphere]]. Atoms in their ionic state may have a different color from neutral atoms, and thus light absorption by metal ions gives the color of [[gemstone]]s. In both inorganic and organic chemistry (including biochemistry), the interaction of water and ions is extremely important{{citation needed|date=September 2020}}; an example is energy that drives the breakdown of adenosine triphosphate ([[Adenosine triphosphate|ATP]]).{{clarify|date=September 2020}}

== Related technology ==
Ions can be non-chemically prepared using various [[ion source]]s, usually involving high [[voltage]] or temperature. These are used in a multitude of devices such as [[Mass spectrometry|mass spectrometers]], [[optical emission spectrometer]]s, [[particle accelerators]], [[Ion implantation|ion implanters]], and [[Ion thruster|ion engines]].

As reactive charged particles, they are also used in [[Air ioniser|air purification]] by disrupting microbes, and in household items such as [[smoke detector]]s.

As signalling and metabolism in organisms are controlled by a precise ionic gradient across [[cell membrane|membranes]], the disruption of this gradient contributes to cell death. This is a common mechanism exploited by natural and artificial [[biocides]], including the [[ion channels]] [[gramicidin]] and [[amphotericin]] (a [[fungicide]]).

Inorganic dissolved ions are a component of [[total dissolved solids]], a widely known indicator of [[water quality]].

===Detection of ionizing radiation===
[[File:Ion chamber operation.gif|thumb|right|300px|Schematic of an ion chamber, showing drift of ions. Electrons drift faster than positive ions due to their much smaller mass.<ref name = "knoll"/>]]
[[File:Electron avalanche.gif|thumbnail|300px|Avalanche effect between two electrodes. The original ionization event liberates one electron, and each subsequent collision liberates a further electron, so two electrons emerge from each collision: the ionizing electron and the liberated electron.]]
The ionizing effect of radiation on a gas is extensively used for the detection of radiation such as [[alpha particle|alpha]], [[beta particle|beta]], [[gamma ray|gamma]], and [[X-rays]]. The original ionization event in these instruments results in the formation of an "ion pair"; a positive ion and a free electron, by ion impact by the radiation on the gas molecules. The [[ionization chamber]] is the simplest of these detectors, and collects all the charges created by ''direct ionization'' within the gas through the application of an electric field.<ref name = "knoll"/>

The [[Geiger–Müller tube]] and the [[proportional counter]] both use a phenomenon known as a [[Townsend avalanche]] to multiply the effect of the original ionizing event by means of a cascade effect whereby the free electrons are given sufficient energy by the electric field to release further electrons by ion impact.

==Chemistry==

=== Denoting the charged state ===
[[File:Ions notation.svg|thumb|right|Equivalent notations for an [[iron]] atom (Fe) that lost two electrons, referred to as [[ferrous]].]]

When writing the [[chemical formula]] for an ion, its net charge is written in superscript immediately after the chemical structure for the molecule/atom. The net charge is written with the magnitude ''before'' the sign; that is, a doubly charged cation is indicated as '''2+''' instead of '''+2'''. However, the magnitude of the charge is omitted for singly charged molecules/atoms; for example, the [[sodium]] cation is indicated as {{chem2|Na+}} and ''not'' {{chem2|Na(1+)}}.

An alternative (and acceptable) way of showing a molecule/atom with multiple charges is by drawing out the signs multiple times, this is often seen with transition metals. Chemists sometimes circle the sign; this is merely ornamental and does not alter the chemical meaning. All three representations of {{chem2|Fe(2+)}}, {{chem2|Fe^{++} }}, and {{chem2|Fe^{⊕⊕} }} shown in the figure, are thus equivalent.

[[File:Ions notation2.svg|thumb|Mixed Roman numerals and charge notations for the [[uranyl]] ion. The oxidation state of the metal is shown as superscripted Roman numerals, whereas the charge of the entire complex is shown by the angle symbol together with the magnitude and sign of the net charge.]]

Monatomic ions are sometimes also denoted with [[Roman numerals]], particularly in [[Spectroscopic notation#Ionization states|spectroscopy]]; for example, the {{chem2|Fe(2+)}} (positively doubly charged) example seen above is referred to as {{chem2|Fe(III)}}, {{chem2|Fe^{III} }} or Fe III (Fe I for a neutral Fe atom, Fe II for a singly ionized Fe ion). The Roman numeral designates the ''formal [[oxidation state]]'' of an element, whereas the superscripted Indo-Arabic numerals denote the net charge. The two notations are, therefore, exchangeable for monatomic ions, but the Roman numerals ''cannot'' be applied to polyatomic ions. However, it is possible to mix the notations for the individual metal centre with a polyatomic complex, as shown by the uranyl ion example.

=== Sub-classes ===
If an ion contains [[unpaired electron]]s, it is called a ''[[radical (chemistry)|radical]]'' ion. Just like uncharged radicals, radical ions are very reactive. Polyatomic ions containing oxygen, such as carbonate and sulfate, are called ''[[oxyanion]]s''. Molecular ions that contain at least one carbon to hydrogen bond are called ''organic ions''. If the charge in an organic ion is formally centred on a carbon, it is termed a ''[[carbocation]]'' (if positively charged) or ''[[carbanion]]'' (if negatively charged).

=== Formation ===
==== Formation of monatomic ions ====
Monatomic ions are formed by the gain or loss of electrons to the [[Valence electron|valence shell]] (the outer-most electron shell) in an atom. The inner shells of an atom are filled with electrons that are tightly bound to the positively charged [[atomic nucleus]], and so do not participate in this kind of chemical interaction. The process of gaining or losing electrons from a neutral atom or molecule is called ''ionization''.

Atoms can be ionized by bombardment with [[radiation]], but the more usual process of ionization encountered in [[chemistry]] is the transfer of electrons between atoms or molecules. This transfer is usually driven by the attaining of stable ("closed shell") [[electronic configuration]]s. Atoms will gain or lose electrons depending on which action takes the least energy.

For example, a [[sodium]] atom, Na, has a single electron in its valence shell, surrounding 2 stable, filled inner shells of 2 and 8 electrons. Since these filled shells are very stable, a sodium atom tends to lose its extra electron and attain this stable configuration, becoming a sodium cation in the process

:<chem>Na -> Na+ + e-</chem>

On the other hand, a [[chlorine]] atom, Cl, has 7 electrons in its valence shell, which is one short of the stable, filled shell with 8 electrons. Thus, a chlorine atom tends to ''gain'' an extra electron and attain a stable 8-[[electron configuration]], becoming a chloride anion in the process:

:<chem>Cl + e- -> Cl-</chem>

This driving force is what causes sodium and chlorine to undergo a chemical reaction, wherein the "extra" electron is transferred from sodium to chlorine, forming sodium cations and chloride anions. Being oppositely charged, these cations and anions form [[ionic bond]]s and combine to form [[sodium chloride]], NaCl, more commonly known as table salt.

:<chem>Na+ + Cl- -> NaCl</chem>

==== Formation of polyatomic and molecular ions ====
[[File:Nitrate-ion-elpot.png|thumb|right|An [[electric potential|electrostatic potential]] map of the [[nitrate ion]] ({{chem2|2NO3-}}). The 3-dimensional shell represents a single arbitrary [[isopotential]].]]

Polyatomic and molecular ions are often formed by the gaining or losing of elemental ions such as a proton, {{chem2|H+}}, in neutral molecules. For example, when [[ammonia]], {{chem2|NH3}}, accepts a proton, {{chem2|H+}}—a process called [[protonation]]—it forms the [[ammonium]] ion, {{chem2|NH4+}}. Ammonia and ammonium have the same number of electrons in essentially the same [[electron configuration|electronic configuration]], but ammonium has an extra proton that gives it a net positive charge.

Ammonia can also lose an electron to gain a positive charge, forming the ion {{chem2|NH3+}}. However, this ion is unstable, because it has an incomplete [[Valence electron|valence shell]] around the nitrogen atom, making it a very reactive [[radical (chemistry)|radical]] ion.

Due to the instability of radical ions, polyatomic and molecular ions are usually formed by gaining or losing elemental ions such as {{chem2|H+}}, rather than gaining or losing electrons. This allows the molecule to preserve its stable electronic configuration while acquiring an electrical charge.

==== Ionization potential ====
{{main|Ionization potential}}
The [[energy]] required to detach an electron in its lowest energy state from an atom or molecule of a gas with less net electric charge is called the ''ionization potential'', or ''ionization energy''. The ''n''th ionization energy of an atom is the energy required to detach its ''n''th electron after the first {{math|''n'' − 1}} electrons have already been detached.

Each successive ionization energy is markedly greater than the last. Particularly great increases occur after any given block of [[atomic orbital]]s is exhausted of electrons. For this reason, ions tend to form in ways that leave them with full orbital blocks. For example, sodium has one ''[[valence electron]]'' in its outermost shell, so in ionized form it is commonly found with one lost electron, as {{chem2|Na+}}. On the other side of the periodic table, chlorine has seven valence electrons, so in ionized form it is commonly found with one gained electron, as {{chem2|Cl-}}. Caesium has the lowest measured ionization energy of all the elements and helium has the greatest.<ref name="lenntech.com">[http://www.lenntech.com/Periodic-chart-elements/ionization-energy.htm Chemical elements listed by ionization energy] {{Webarchive|url=https://web.archive.org/web/20090330155714/http://www.lenntech.com/Periodic-chart-elements/ionization-energy.htm |date=2009-03-30 }}. Lenntech.com</ref> In general, the ionization energy of [[metals]] is much lower than the ionization energy of [[nonmetals]], which is why, in general, metals will lose electrons to form positively charged ions and nonmetals will gain electrons to form negatively charged ions.

===Ionic bonding===
{{main|Ionic bond}}

''Ionic bonding'' is a kind of [[chemical bond]]ing that arises from the mutual attraction of oppositely charged ions. Ions of like charge repel each other, and ions of opposite charge attract each other. Therefore, ions do not usually exist on their own, but will bind with ions of opposite charge to form a [[crystal lattice]]. The resulting compound is called an ''ionic compound'', and is said to be held together by ''ionic bonding''. In ionic compounds there arise characteristic distances between ion neighbours from which the spatial extension and the [[ionic radius]] of individual ions may be derived.

The most common type of ionic bonding is seen in compounds of metals and nonmetals (except [[noble gas]]es, which rarely form chemical compounds). Metals are characterized by having a small number of electrons in excess of a stable, closed-shell [[electronic configuration]]. As such, they have the tendency to lose these extra electrons in order to attain a stable configuration. This property is known as ''[[electropositivity]]''. Non-metals, on the other hand, are characterized by having an [[electron configuration]] just a few electrons short of a stable configuration. As such, they have the tendency to gain more electrons in order to achieve a stable configuration. This tendency is known as ''[[electronegativity]]''. When a highly electropositive metal is combined with a highly electronegative nonmetal, the extra electrons from the metal atoms are transferred to the electron-deficient nonmetal atoms. This reaction produces metal cations and nonmetal anions, which are attracted to each other to form a ''[[salt (chemistry)|salt]]''.

===Common ions===
{|
|-
|valign="top"|
{|class="wikitable"
|+Common cations<ref name=":0">{{Cite web|url=http://www.sciencegeek.net/Chemistry/chempdfs/CommonIons.pdf|title=Common Ions and Their Charges|website=Science Geek|access-date=2018-05-11|archive-date=2018-02-18|archive-url=https://web.archive.org/web/20180218214350/http://www.sciencegeek.net/Chemistry/chempdfs/CommonIons.pdf|url-status=live}}</ref>
|-
!style="text-align: left"|Common name
!style="text-align: left"|Formula
!style="text-align: left"|Historic name
|-
! colspan="3" style="background:aliceblue;"|''[[Monatomic ion|Monatomic]] cations''
|-
|[[Aluminium]]||{{chem2|Al(3+)}}||
|-
|[[Barium]]||{{chem2|Ba(2+)}}||
|-
|[[Beryllium]]||{{chem2|Be(2+)}}||
|-
|[[Calcium]]||{{chem2|Ca(2+)}}||
|-
|Chromium(III)||{{chem2|Cr(3+)}}||
|-
|Copper(I)||{{chem2|Cu+}}||cuprous
|-
|Copper(II)||{{chem2|Cu(2+)}}||cupric
|-
|[[Gold|Gold(I)]]||{{chem2|Au+}}||aurous
|-
|Gold(III)||{{chem2|Au(3+)}}||auric
|-
|[[Hydron (chemistry)|Hydron]]||{{chem2|H+}}||
|-
|Iron(II)||{{chem2|Fe(2+)}}||ferrous
|-
|Iron(III)||{{chem2|Fe(3+)}}||ferric
|-
|Lead(II)||{{chem2|Pb(2+)}}||plumbous
|-
|Lead(IV)||{{chem2|Pb(4+)}}||plumbic
|-
|[[Lithium]]||{{chem2|Li+}}||
|-
|[[Magnesium]]||{{chem2|Mg(2+)}}||
|-
|Manganese(II)||{{chem2|Mn(2+)}}||manganous
|-
|Manganese(III)||{{chem2|Mn(3+)}}||manganic
|-
|Manganese(IV)||{{chem2|Mn(4+)}}||
|-
|Mercury(II)||{{chem2|Hg(2+)}}||mercuric
|-
|[[Potassium]] ||{{chem2|K+}}||kalic <!-- The noun "kalium" is used more commonly than the adjective "kalic".-->
|-
|[[Silver]]||{{chem2|Ag+}}||argentous
|-
|[[Sodium]]||{{chem2|Na+}}||natric <!-- The noun "natrium" is used more commonly than the adjective "natric".-->
|-
|[[Strontium]]||{{chem2|Sr(2+)}}||
|-
|[[Tin]](II)||{{chem2|Sn(2+)}}||stannous
|-
|Tin(IV)||{{chem2|Sn(4+)}}||stannic
|-
|[[Zinc]]||{{chem2|Zn(2+)}}||
|-
! colspan="3" style="background:aliceblue;"|''[[Polyatomic ion|Polyatomic]] cations''
|-
|[[Ammonium]]||{{chem2|NH4+}}||
|-
|[[Hydronium]]||{{chem2|H3O+}}||
|-
|Mercury(I)||{{chem2|Hg2(2+)}}||mercurous
|}
|valign="top"|
{|class="wikitable"
|+Common anions<ref name=":0" />
|-
!style="text-align: left"|Formal name
!style="text-align: left"|Formula
!style="text-align: left"|Alt. name
|-
! colspan="3" style="background:aliceblue;"|''[[Monatomic ion|Monatomic]] anions''
|-
|[[Bromide]]||{{chem2|Br-}}||
|-
|[[Carbide]]||{{chem2|C-}}||
|-
|[[Chloride]]||{{chem2|Cl-}}||
|-
|[[Fluoride]]||{{chem2|F-}}||
|-
|[[Hydride]]||{{chem2|H-}}||
|-
|[[Iodide]]||{{chem2|I-}}||
|-
|[[Nitride]]||{{chem2|N(3-)}}||
|-
|[[Phosphide]]||{{chem2|P(3-)}}||
|-
|[[Oxide]]||{{chem2|O(2-)}}||
|-
|[[Sulfide]]||{{chem2|S(2-)}}||
|-
|[[Selenide]]||{{chem2|Se(2-)}}||
|-
! colspan="3" style="background:aliceblue;"|''[[Polyatomic]] anions''
|-
|[[Azide]]||{{chem2|N3-}}||
|-
|[[Peroxide]]||{{chem2|O2(2-)}}||
|-
|[[Triodide]]||{{chem2|I3(-)}}||
|-
! colspan="3" style="background:aliceblue;"|''Oxoanions ([[Polyatomic ion]]s)<ref name=":0" />''
|-
|[[Carbonate]]||{{chem2|CO3(2-)}}||
|-
|[[Chlorate]]||{{chem2|ClO3-}}||
|-
|[[Chromate ion|Chromate]]||{{chem2|CrO4(2−)}}||
|-
|[[Dichromate]]||{{chem2|Cr2O7(2-)}}||
|-
|[[Dihydrogen phosphate]]||{{chem2|H2PO4-}}||
|-
|[[Hydrogen carbonate]]||{{chem2|HCO3-}}||bicarbonate
|-
|[[Hydrogen sulfate]]||{{chem2|HSO4-}}||bisulfate
|-
|[[Hydrogen sulfite]]||{{chem2|HSO3-}}||bisulfite
|-
|[[Hydroxide]]||{{chem2|OH-}}||
|-
|[[Hypochlorite]]||{{chem2|ClO-}}||
|-
|[[Monohydrogen phosphate]]||{{chem2|HPO4(2-)}}||
|-
|[[Nitrate]]||{{chem2|NO3-}}||
|-
|[[Nitrite]]||{{chem2|NO2-}}||
|-
|[[Perchlorate]]||{{chem2|ClO4-}}||
|-
|[[Permanganate]]||{{chem2|MnO4-}}||
|-
|[[Peroxide]]||{{chem2|O2(2-)}}||
|-
|[[Phosphate]]||{{chem2|PO4(3-)}}||
|-
|[[Sulfate]]||{{chem2|SO4(2-)}}||
|-
|[[Sulfite]]||{{chem2|SO3(2-)}}||
|-
|[[Superoxide]]||{{chem2|O2-}}||
|-
|[[Thiosulfate]]||{{chem2|S2O3(2-)}}||
|-
|[[Silicate]]||{{chem2|SiO4(4-)}}||
|-
|[[Metasilicate]]||{{chem2|SiO3(2-)}}||
|-
|[[Aluminium silicate]]||{{chem2|AlSiO4-}}||
|-
! colspan="3" style="background:aliceblue;"|''Anions from organic acids''
|-
|[[Acetate]]||{{chem2|CH3COO-}}||ethanoate
|-
|[[Formate]]||{{chem2|HCOO-}}||methanoate
|-
|[[Oxalate]]||{{chem2|C2O4(2-)}}||ethanedioate
|-
|[[Cyanide]]||{{chem2|CN-}}||
|}
|}

==See also==
{{Div col}}
* [[Air ioniser]]
* [[Aurora]]
* [[Electrolyte]]
* [[Gaseous ionization detector]]
* [[Ioliomics]]
* [[Ion beam]]
* [[Ion exchange]]
* [[Ionizing radiation]]
* [[Stopping power (particle radiation)|Stopping power of radiation particles]]
{{Div col end}}

==References==
{{Reflist}}

{{Branches of chemistry}}
{{Galvanic cells|state=collapsed}}
{{Authority control}}

[[Category:Ions| ]]
[[Category:Physical chemistry]]
[[Category:Charge carriers]]