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Electron transfer from a neutral lithium (Li) atom on the left to a neutral fluorine (F) atom on the right would give a Li+ and F− ions.

An ion (/ˈaɪ.ɒn, -ən/)[1] is an atom or molecule with a net 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[2] (e.g. K+ (potassium ion)) while an anion is a negatively charged ion with more electrons than protons[3] (e.g. Cl− (chloride ion) and OH− (hydroxide ion)). Opposite electric charges are pulled towards one another by electrostatic force, so cations and anions attract each other and readily form ionic compounds. Ions consisting of only a single atom are termed monatomic ions, atomic ions or simple ions, while ions consisting of two or more atoms are termed polyatomic ions or molecular ions.

If only a + or − is present, it indicates a +1 or −1 charge, as seen in Na+ (sodium ion) and F−
 (fluoride ion). To indicate a more severe charge, the number of additional or missing electrons is supplied, as seen in O2−
2 (peroxide, negatively charged, polyatomic) and He2+ (alpha particle, positively charged, monatomic).[4]

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.[5] Ions are also created by chemical interactions, such as the dissolution of a 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

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The word ion was coined from neuter present participle of Greek ἰέναι (ienai), meaning "to go". A cation is something that moves down (Greek: κάτω, kato, meaning "down") and an anion is something that moves up (Greek: 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)[6] 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.[7][8] 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.[6]

Svante Arrhenius put forth, in his 1884 dissertation, the explanation of the fact that solid crystalline salts dissociate into paired charged particles when dissolved, for which he would win the 1903 Nobel Prize in Chemistry.[9] 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.[10][11][12]

Characteristics

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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 solvation shell around them. These solvated ions 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

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"Anion" redirects here. Not to be confused with the quasiparticle Anyon.
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 anion (−) (/ˈænˌaɪ.ən/ ANN-eye-ən, from the Greek word ἄνω (ánō), meaning "up"[13]) is an ion with more electrons than protons, giving it a net negative charge (since electrons are negatively charged and protons are positively charged).[14]

A cation (+) (/ˈkætˌaɪ.ən/ KAT-eye-ən, from the Greek word κάτω (kátō), meaning "down"[15]) is an ion with fewer electrons than protons, giving it a positive charge.[16]

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.[17]

Cations and anions are measured by their ionic radius and they differ in relative size: "Cations are small, most of them less than 10−10 m (10−8 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."[18]

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

Natural occurrences

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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 gemstones. In both inorganic and organic chemistry (including biochemistry), the interaction of water and ions is often relevant for understanding properties of systems; an example of their importance is in the breakdown of adenosine triphosphate (ATP), which provides the energy for many reactions in biological systems.

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Ions can be non-chemically prepared using various ion sources, usually involving high voltage or temperature. These are used in a multitude of devices such as mass spectrometers, optical emission spectrometers, particle accelerators, ion implanters, and ion engines.

As reactive charged particles, they are also used in air purification by disrupting microbes, and in household items such as smoke detectors.

As signalling and metabolism in organisms are controlled by a precise ionic gradient across 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

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Schematic of an ion chamber, showing drift of ions. Electrons drift faster than positive ions due to their much smaller mass.[5]
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, beta, 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.[5]

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

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Denoting the charged state

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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 Na+ and not Na1+.

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 Fe2+, Fe++, and Fe⊕⊕ shown in the figure, are thus equivalent.

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 spectroscopy; for example, the Fe2+ (positively doubly charged) example seen above is referred to as Fe(III), FeIII 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

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If an ion contains unpaired electrons, it is called a radical ion. Just like uncharged radicals, radical ions are very reactive. Polyatomic ions containing oxygen, such as carbonate and sulfate, are called oxyanions. 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

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Formation of monatomic ions

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Monatomic ions are formed by the gain or loss of electrons to the 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 configurations. 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

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:

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 bonds and combine to form sodium chloride, NaCl, more commonly known as table salt.

Formation of polyatomic ions

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An electrostatic potential map of the nitrate ion (2NO−3). 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, H+, in neutral molecules. For example, when ammonia, NH3, accepts a proton, H+—a process called protonation—it forms the ammonium ion, NH+4. Ammonia and ammonium have the same number of electrons in essentially the same 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 NH+3. However, this ion is unstable, because it has an incomplete valence shell around the nitrogen atom, making it a very reactive radical ion.

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

Formation of ions in nonpolar liquids

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Liquids with low dielectric constant (below 10) are not quite suitable for ions formation for several reasons. First of all, electrostatic attraction between cation and anion is much stronger than in water, which requires well developed solvating layer for preventing their immediate reaggregation. However, molecules of nonpolar liquids cannot create such layer due to lack of dipole moments. In addition, many electrolytes are not soluble in nonpolar liquids. Nevertheless, pioneering works by Onsager, Fuoss, Kraus in 20th century proved that ionization in nonpolar liquids is possible [19] .[20] Recent series of studies conducted by Dukhin and Parlia with wide variety of liquids and solutes (summarized in [21]) confirmed this conclusion and allowed formulation of the following concept for ionization in non-polar liquids, which is distinctively different from aqueous solutions. The solute substance must be amphiphile consisting of hydrophobic tail and polar head in order to create ions in nonpolar liquid. Existence of hydrophobic tail ensures solubility. Existence of polar head provides source for initial ions creation by dissociation. The most peculiar feature is formation of solvation layer around ions almost immediately after dissociation. Solvent molecules cannot build up such solvating layers. However, the neutral molecules of the solute do have some dipole moments at their polar heads. These dipoles would be attracted by primary ions right after dissociation. This attraction creates a layer of the neutral solute molecules around central ions, which can be considered as solvation layer. Such solvated ions look like charged inverse micelles.[22] Basically, solute amphiphilic molecules in nonpolar liquids are source of both, dissociation and self-solvation, which distinguishes this ionization from aqueous solutions dramatically. This concept led to creation of conductivity theory that fits experimental data for wide variety of nonpolar systems within up to 7 orders of magnitude.[23]

Ionization potential

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Main article: 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 nth ionization energy of an atom is the energy required to detach its nth electron after the first 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 orbitals 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 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 Cl−. Caesium has the lowest measured ionization energy of all the elements and helium has the greatest.[24] 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

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Main article: Ionic bond

Ionic bonding is a kind of chemical bonding 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 gases, 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.

Common ions

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Common cations[25]
Common name Formula Historic name
Monatomic cations
Aluminium Al3+
Barium Ba2+
Beryllium Be2+
Calcium Ca2+
Chromium(III) Cr3+
Copper(I) Cu+ cuprous
Copper(II) Cu2+ cupric
Gold(I) Au+ aurous
Gold(III) Au3+ auric
Hydron H+
Iron(II) Fe2+ ferrous
Iron(III) Fe3+ ferric
Lead(II) Pb2+ plumbous
Lead(IV) Pb4+ plumbic
Lithium Li+
Magnesium Mg2+
Manganese(II) Mn2+ manganous
Manganese(III) Mn3+ manganic
Manganese(IV) Mn4+
Mercury(II) Hg2+ mercuric
Potassium K+ kalic
Silver Ag+ argentous
Sodium Na+ natric
Strontium Sr2+
Tin(II) Sn2+ stannous
Tin(IV) Sn4+ stannic
Zinc Zn2+
Polyatomic cations
Ammonium NH+4
Hydronium H3O+
Mercury(I) Hg2+2 mercurous
Common anions[25]
Formal name Formula Alt. name
Monatomic anions
Bromide Br−
Carbide C−
Chloride Cl−
Fluoride F−
Hydride H−
Iodide I−
Nitride N3−
Phosphide P3−
Oxide O2−
Sulfide S2−
Selenide Se2−
Polyatomic anions
Azide N−3
Peroxide O2−2
Triiodide I−3
Oxoanions (Polyatomic ions)[25]
Carbonate CO2−3
Chlorate ClO−3
Chromate CrO2−4
Dichromate Cr2O2−7
Dihydrogen phosphate H2PO−4
Hydrogen carbonate HCO−3 bicarbonate
Hydrogen sulfate HSO−4 bisulfate
Hydrogen sulfite HSO−3 bisulfite
Hydroxide OH−
Hypochlorite ClO− chloroxide
Monohydrogen phosphate HPO2−4
Nitrate NO−3
Nitrite NO−2
Perchlorate ClO−4
Permanganate MnO−4
Peroxide O2−2
Phosphate PO3−4
Sulfate SO2−4
Sulfite SO2−3
Superoxide O−2
Thiosulfate S2O2−3
Silicate SiO4−4
Metasilicate SiO2−3
Aluminium silicate AlSiO−4
Anions from organic acids
Acetate CH3COO− ethanoate
Formate HCOO− methanoate
Oxalate C2O2−4 ethanedioate
Cyanide CN−

See also

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References

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Revisions and contributorsEdit on WikipediaRead on Wikipedia

Ion

View on Grokipedia
from Grokipedia
An ion is an atom or molecule that has a net electric charge due to the gain or loss of one or more electrons, resulting in a positively charged cation or a negatively charged anion. These charged particles are fundamental to chemical bonding, particularly in ionic compounds where oppositely charged ions are held together by electrostatic forces. The concept of the ion emerged in the 19th century during studies of electrolysis, with the term "ion" coined by in 1834 to describe particles that migrate under an , derived from the Greek word ion meaning "to go." This nomenclature, suggested with the help of , marked a shift from earlier atomic theories toward understanding charged species in solutions. By the late 1800s, the discovery of the by J.J. Thomson further clarified the mechanism by which atoms become ions through . Ions are classified as monatomic, consisting of a single charged atom such as Na⁺ (sodium cation) or (chloride anion), or polyatomic, involving a group of atoms with a net charge like (ammonium) or (sulfate). Cations form by losing electrons, typically from metals, while anions form by gaining electrons, often from nonmetals. In aqueous solutions, ions dissociate from ionic compounds, enabling conductivity and facilitating reactions in electrolytes. Ions play critical roles in chemistry, driving processes like , acid-base reactions, and chemistry, while in , they are essential for impulse transmission, , enzyme function, and maintaining cellular pH and osmotic balance. For instance, ions such as K⁺ and Na⁺ regulate membrane potentials in neurons, and Ca²⁺ triggers muscle contractions. Disruptions in ion transport, as seen in conditions like involving defective chloride channels, underscore their physiological importance.

Fundamentals

Definition and Charge

An is an atom or that bears a net , resulting from the gain or loss of one or more electrons. This charge imbalance distinguishes ions from neutral particles, as the unequal numbers of subatomic particles lead to an overall positive or negative electrical property. Positive ions, called cations, form when an atom or loses electrons, leaving it with more protons than electrons and thus a net positive charge. Conversely, negative ions, known as anions, arise when an atom or gains electrons, resulting in more electrons than protons and a net negative charge. These processes disrupt the electrical neutrality inherent in atoms, enabling ions to interact electrostatically with oppositely charged . In a neutral atom, the positive charge from an equal number of protons in the nucleus is balanced by the negative charge of an equal number of orbiting the nucleus, yielding no net charge. Ion formation occurs when this balance is altered, such as through , creating a charged entity where the proton-electron count is unequal. For instance, the sodium cation (Na⁺) exemplifies this by losing one from a neutral sodium atom, which has 11 protons and 11 , resulting in a +1 charge. Similarly, the chloride anion (Cl⁻) forms when a neutral atom, with 17 protons and 17 , gains one , producing a -1 charge.

Anions and Cations

Ions are classified into two primary types based on their charge: cations, which carry a positive charge due to the loss of one or more electrons, and anions, which carry a negative charge due to the gain of one or more electrons. This fundamental distinction leads to notable structural differences in their atomic or molecular architecture. Cations are typically smaller than their neutral parent atoms because the removal of electrons reduces electron-electron repulsion, allowing the nucleus to exert a stronger on the remaining electrons, pulling them closer to the nucleus. Conversely, anions are larger than their parent atoms as the addition of electrons increases electron-electron repulsion, decreasing the and causing the electron cloud to expand outward. In terms of behavior, cations and anions exhibit differences in mobility and reactivity within chemical environments, particularly in solutions. Ionic mobilities vary depending on ion , , hydration, and specific mechanisms. For example, due to the involving proton hopping, the hydronium ion (\ceH3O+\ce{H3O+}) has higher mobility than the hydroxide ion (\ceOH−\ce{OH-}) in at 25°C: approximately 3.63×10−33.63 \times 10^{-3} cm² s⁻¹ V⁻¹ for \ceH3O+\ce{H3O+} and 2.05×10−32.05 \times 10^{-3} cm² s⁻¹ V⁻¹ for \ceOH−\ce{OH-}. Anions, being larger and often more solvated, tend to be more stable in polar environments like , where their increased reduces reactivity toward certain nucleophilic or electrophilic interactions. Cations, with their electron-deficient nature, are generally more reactive, acting as Lewis acids that readily accept electron pairs from surrounding . Representative examples illustrate these characteristics across ion types. Common cations include metal ions such as calcium (\ceCa2+\ce{Ca^{2+}}), which is prevalent in biological systems and exhibits high reactivity in forming complexes, and sodium (\ceNa+\ce{Na+}), essential for signaling. Anions often derive from non-metals, like (\ceSO42−\ce{SO4^{2-}}), a involved in and more stable in acidic conditions due to its delocalized charge. These examples highlight how cations from metals tend toward higher charge densities and reactivity, while anions from non-metals provide structural stability in ionic compounds. A key behavioral distinction appears in , where an applied drives ion separation in molten or aqueous electrolytes. Cations migrate toward the (negative ), where they undergo reduction by gaining electrons, while anions move to the (positive ), undergoing oxidation by losing electrons; this directed migration, known as ionic conduction, depends on the ions' charge and mobility, with transference numbers quantifying their relative contributions to current flow. In natural settings, such as , this principle underlies the abundance of sodium cations and anions, which maintain electrolytic balance essential for marine ecosystems.

Historical Development

Early Observations

The earliest recorded observations of phenomena related to ions trace back to , where was noted around 600 BCE by the philosopher . He observed that rubbing —a fossilized —with or caused it to attract lightweight particles such as feathers, straw, or dust, demonstrating an attractive force without physical contact. This effect, now understood as triboelectric charging, represented the first empirical recognition of electrical properties in materials, though Thales attributed it to a soul-like animating force in the amber rather than discrete charged entities. In the , systematic investigations advanced these ancient insights into more structured concepts of . French Charles François de Cisternay du Fay, in 1733, conducted experiments showing that produced by rubbing glass (vitreous electricity) repelled similar charges but attracted those from rubbed resin or amber (resinous electricity), suggesting two distinct types of electrical fluid. Building on this, American , during the 1750s, proposed a unified theory in his correspondence and experiments, introducing the terms "positive" and "negative" charges to describe excesses and deficiencies of a single electrical fluid, respectively; he demonstrated these through kite experiments and studies that equated to electrical discharge. These developments shifted observations from qualitative attractions to a binary framework of opposing charges, laying groundwork for later ionic interpretations. The transition to chemical contexts occurred in the early 19th century through electrochemical experiments. In 1807–1808, British chemist used with a powerful to decompose molten compounds, isolating metallic sodium from and potassium from —highly reactive elements previously unknown in pure form. These decompositions implied the migration of charged species within the melt toward electrodes, as positive metals collected at the and oxygen at the , providing early evidence of electricity's role in separating atomic constituents despite the absence of a formal ion concept at the time.

Key Discoveries and Milestones

In the early 1830s, conducted systematic experiments on , culminating in the formulation of Faraday's laws, which established that the mass of a substance altered at an is directly proportional to the quantity of electricity passed through the and that the masses of different substances liberated by the same quantity of electricity are proportional to their chemical equivalent weights. These laws, published in 1834, provided the first quantitative foundation for understanding electrochemical processes. Faraday also introduced the term "ion" in the same year, deriving it from the Greek word iōn meaning "wanderer" or "going," to denote the charged particles that migrate toward electrodes during . Building on Faraday's empirical observations, advanced the theoretical understanding of ions in 1887 with his theory of electrolytic dissociation, proposing that electrolytes in spontaneously dissociate into free ions, thereby accounting for electrical conductivity and phenomena like and . This groundbreaking idea, detailed in his paper "Über die Dissociation der in Wasser gelösten Stoffe," resolved longstanding discrepancies between expected and observed properties of solutions and laid the groundwork for modern . A pivotal experimental breakthrough occurred in 1897 when J.J. Thomson investigated using deflection in electric and magnetic fields, discovering the as a constituent of all matter and determining its charge-to-mass ratio to be approximately 1.76 × 10^11 C/kg—about 1,800 times larger than that of a . Thomson's findings, reported in his paper "Cathode Rays" in , confirmed ions as composed of charged subatomic particles and shifted atomic theory toward a particulate model. The early 20th century saw further milestones in ion analysis with Francis Aston's invention of the mass spectrograph in 1919, an instrument that accelerated ions through electric and magnetic fields to separate them by , achieving resolutions sufficient to detect isotopic variations in elements like . Aston's device, described in "A Positive Ray Spectrograph," enabled precise mass measurements and confirmed Frederick Soddy's concept, revolutionizing . Concurrently, the brought quantum mechanical frameworks for ions, as Erwin Schrödinger's 1926 wave equation provided probabilistic descriptions of electron distributions in multi-electron atoms and ions, while Werner Heisenberg's offered complementary tools for calculating ionic energy levels and spectra.

Physical Properties

Size and Stability

The size of an ion, often expressed as its ionic radius, differs significantly from that of its parent neutral atom due to the gain or loss of electrons, which alters electron-electron repulsions and the effective nuclear charge experienced by the remaining electrons. Cations are generally smaller than their neutral counterparts because the removal of electrons reduces shielding, increasing the effective nuclear charge and drawing the electron cloud closer to the nucleus; for example, the atomic radius of neutral sodium (Na) is approximately 186 pm, while the ionic radius of Na⁺ (in six-fold coordination) is 102 pm. In contrast, anions are larger than their neutral atoms as the addition of electrons increases mutual repulsion among the outer electrons, expanding the electron cloud despite the same nuclear charge. These size variations are primarily governed by the (Z_eff), which quantifies the net positive charge felt by valence electrons after accounting for shielding by inner electrons. For cations, Z_eff rises upon electron loss, compressing the ion; for anions, although Z_eff remains similar, the extra electrons cause greater repulsion, leading to expansion. Ionic radii also depend on and the nature of surrounding ions, with values refined empirically from structural data. The stability of ions in solids is largely determined by , the exothermic released when gaseous ions assemble into a crystalline lattice, which increases with higher ion charges and smaller interionic distances due to stronger electrostatic attractions. Smaller ions contribute to higher , enhancing the stability of ionic compounds like NaCl. In solution, ion stability arises from solvation , where ions are stabilized by surrounding molecules (e.g., dipoles orienting toward the ion); this is greater for smaller, highly charged ions owing to closer approach of molecules. At higher concentrations or in low-dielectric , ions may form pairs or clusters to minimize free energy by reducing reorganization costs, with ion pairing more prevalent for large, low-charge ions like those in solutions. Quantum mechanical effects further influence ion stability through electron configurations, particularly when ions achieve a noble gas-like arrangement with filled valence shells, such as the [Ne] configuration of Na⁺ or [Ar] of Cl⁻, which minimizes energy due to complete octet stability and low reactivity. These configurations align with the octet rule, providing exceptional stability compared to ions with incomplete or expanded shells. Experimental determination of ionic radii relies on X-ray crystallography of ionic crystals, which measures interatomic distances in the lattice and assigns radii by assuming additivity (r_ion1 + r_ion2 = observed distance), as refined in systematic compilations from thousands of structures. This method provides precise values for various coordination environments, underpinning comparisons like those between Na⁺ and Cl⁻ in rock salt structures.

Natural Occurrences

Ions are ubiquitous in Earth's atmosphere, particularly in the , a region extending from about 50 to 1,000 kilometers above the surface where solar and radiation ionizes neutral atoms and molecules, producing a plasma consisting primarily of free and positive ions such as O⁺ and N₂⁺. This process varies diurnally, with higher electron densities on the dayside due to direct solar exposure. Additionally, strikes generate ions in the through electrical discharges that ionize air molecules, creating a conductive plasma channel of electrons and ions that facilitates charge equalization between cloud regions and the ground. In oceanic and geological environments, ions dominate the composition of natural waters through the dissolution of minerals in rocks and soils. , for example, has an average of about 3.5%, largely due to dissolved ions like sodium (Na⁺) and chloride (Cl⁻), which account for roughly 85% of all ionic content, with other major ions including magnesium (Mg²⁺), (SO₄²⁻), calcium (Ca²⁺), and (K⁺). These ions originate from the chemical weathering of continental minerals, such as feldspars and carbonates, which release cations and anions into rivers and that eventually mix with ocean basins. Biologically, ions are essential for cellular function and structural integrity. In living cells, sodium (Na⁺) and potassium (K⁺) ions maintain electrochemical gradients across membranes, enabling the propagation of impulses via action potentials, where rapid Na⁺ influx depolarizes the followed by K⁺ efflux for . Calcium ions (Ca²⁺) play a critical structural role in bones, where more than 99% of the body's calcium is stored as (Ca₁₀(PO₄)₆(OH)₂), providing rigidity and serving as a reservoir for systemic calcium . In cosmic settings, ions abound in the (), a dilute plasma of gas and dust between stars where cosmic rays and photons ionize and other elements, yielding species like H⁺, He⁺, and molecular ions such as H₃⁺ that influence cloud formation and star birth. The , a continuous outflow of coronal plasma from the Sun, carries primarily protons (H⁺ ions) and electrons at speeds of 300–800 km/s, with minor contributions from alpha particles (He²⁺) and heavier ions, forming a magnetized plasma that interacts with planetary magnetospheres and the . These environments highlight ions' prevalence in fully or partially ionized plasma states throughout the universe.

Chemical Aspects

Notation and Subclasses

The standard notation for ions employs superscripts to indicate charge, placed as a right upper index following the chemical symbol or formula, such as \ceH+\ce{H^+} for the hydrogen cation or \ceOH−\ce{OH^-} for the hydroxide anion, with the magnitude of the charge (unity omitted) preceding the sign.\] For coordination entities and complex ions, square brackets enclose the formula of the ion, with the charge indicated as a superscript outside the brackets, as in $ \ce{[Fe(CN)6]^4-} $ for the hexacyanidoferrate(4−) ion.\[ This convention ensures clarity in representing the structure and charge distribution, particularly for polyatomic species where ligands are listed alphabetically within the brackets before the central atom.$$] The evolution of ion notation traces back to Michael Faraday's introduction of the term "ion" in 1834 to describe charged particles in , though early representations lacked standardized symbols and relied on descriptive terms like "electropositive" or "electronegative" entities.[Inthe[1880s](/page/1880s),SvanteArrhenius′stheoryofelectrolyticdissociationhighlightedtheneedforchargenotation,leadingtoinitialproposalsinthe1890sby[WilhelmOstwald](/page/WilhelmOstwald),whousedsuperscriptdotsforcations(e.g.,Ba..)andprimesforanions(e.g.,PO4′′′),andby[WaltherNernst](/page/WaltherNernst),whoplaced+or−signsaboveortotherightofsymbols(e.g.,Ba++),apracticethatgainedtraction. In the [1880s](/page/1880s), Svante Arrhenius's theory of electrolytic dissociation highlighted the need for charge notation, leading to initial proposals in the 1890s by [Wilhelm Ostwald](/page/Wilhelm_Ostwald), who used superscript dots for cations (e.g., Ba..) and primes for anions (e.g., PO₄''' ), and by [Walther Nernst](/page/Walther_Nernst), who placed + or - signs above or to the right of symbols (e.g., Ba++), a practice that gained traction.] By the mid-20th century, algebraic notations like Ba+2 appeared sporadically, but IUPAC standardized the modern form with numbers preceding the charge sign (e.g., \ceBa2+\ce{Ba^2+}, \cePO43−\ce{PO4^3-}) in its 1950s guidelines, culminating in the comprehensive 2005 recommendations for inorganic nomenclature that formalized superscripts, brackets, and systematic naming.[$$ Common ions are often categorized by their charge (valency), with examples of metallic, non-metallic, and polyatomic ions listed below from the NCERT Class 9 Science textbook, Chapter 3, Table 3.6.
ValencyMetallic IonSymbolNon-metallic IonSymbolPolyatomic IonSymbol
1Sodium\ceNa+\ce{Na^+}Hydrogen\ceH+\ce{H^+}Ammonium\ceNH4+\ce{NH4^+}
1Potassium\ceK+\ce{K^+}Hydride\ceH−\ce{H^-}Hydroxide\ceOH−\ce{OH^-}
1Silver\ceAg+\ce{Ag^+}Chloride\ceCl−\ce{Cl^-}Nitrate\ceNO3−\ce{NO3^-}
1Copper (I)*\ceCu+\ce{Cu^+}Bromide\ceBr−\ce{Br^-}Hydrogen carbonate\ceHCO3−\ce{HCO3^-}
1Iodide\ceI−\ce{I^-}
2Magnesium\ceMg2+\ce{Mg^{2+}}Oxide\ceO2−\ce{O^{2-}}Carbonate\ceCO32−\ce{CO3^{2-}}
2Calcium\ceCa2+\ce{Ca^{2+}}Sulfide\ceS2−\ce{S^{2-}}Sulfite\ceSO32−\ce{SO3^{2-}}
2Zinc\ceZn2+\ce{Zn^{2+}}Sulfate\ceSO42−\ce{SO4^{2-}}
2Iron (II)*\ceFe2+\ce{Fe^{2+}}
2Copper (II)*\ceCu2+\ce{Cu^{2+}}
3Aluminium\ceAl3+\ce{Al^{3+}}Nitride\ceN3−\ce{N^{3-}}Phosphate\cePO43−\ce{PO4^{3-}}
3Iron (III)*\ceFe3+\ce{Fe^{3+}}
  • Some elements show more than one valency. A Roman numeral shows their valency in a bracket.
Ions are classified into subclasses based on their composition and electronic . Monatomic ions, also known as simple or atomic ions, consist of a single atom that has gained or lost electrons, exemplified by \ceLi+\ce{Li^+} ( cation) or \ceCl−\ce{Cl^-} ( anion), which follow predictable charges based on periodic table group trends.\] Polyatomic ions, in contrast, comprise two or more atoms covalently bonded together and acting as a single charged unit, such as $ \ce{NH4^+} $ ([ammonium](/page/Ammonium) cation) or $ \ce{SO4^2-} $ ([sulfate](/page/Sulfate) anion), where the overall charge results from the [net](/page/.net) [electron transfer](/page/Electron_transfer).\[ Molecular ions represent a of polyatomic ions that retain molecular stability in their charged form, like \ceH3O+\ce{H3O^+} ( cation), often encountered in aqueous solutions or gas-phase environments.$$] Radical ions form another subclass, characterized by an in addition to the net charge, making them highly reactive; these are denoted with a superscript dot to indicate the radical nature, as in \ceO2.−\ce{O2^{.-}} for the superoxide .[ Isotope-specific ions incorporate notation for isotopic [variants](/page/Mahindra_KUV100) by placing the [mass number](/page/Mass_number) as a left superscript before the element [symbol](/page/Symbol), such as $ \ce{^2H^+} $ for the deuteron, allowing distinction in spectroscopic or nuclear studies while adhering to general charge superscript rules.]

Ionic Bonding

Ionic bonding arises from the electrostatic attraction between oppositely charged ions, primarily cations and anions, which form when atoms gain or lose electrons to achieve stable electron configurations. This attraction is governed by Coulomb's law, which quantifies the force FF between two point charges as F=kq1q2r2F = k \frac{q_1 q_2}{r^2}, where kk is Coulomb's constant (8.99×109 N⋅m2/C28.99 \times 10^9 \, \mathrm{N \cdot m^2 / C^2}), q1q_1 and q2q_2 are the charges on the ions, and rr is the distance between their centers. In ionic compounds, these pairwise interactions extend throughout the crystal lattice, resulting in a strong, directional network that stabilizes the solid. Ionic compounds typically adopt ordered lattice structures to maximize attractive forces and minimize repulsion between like-charged ions. A classic example is (NaCl), which forms a rock salt structure where each Na⁺ ion is surrounded by six Cl⁻ ions, and vice versa, yielding a of 6 and an for efficient packing. This close-packed arrangement, often based on face-centered cubic lattices of anions with cations in interstitial sites, extends to other alkali halides like KCl and RbBr, promoting high stability through balanced electrostatic interactions. The properties of ionic compounds stem directly from their lattice energy and interionic forces. They exhibit high melting and boiling points—such as NaCl's melting point of 801°C—because significant energy is required to overcome the collective Coulombic attractions throughout the lattice. Ionic solids are hard yet brittle; applied stress can shift ion layers, aligning like charges and causing repulsion that shatters the crystal. Many are soluble in water due to the solvent's polarity, which allows hydration shells to stabilize separated ions, though solubility varies with lattice energy and ion size. While alkali halides represent prototypical ionic bonding, some compounds like aluminum chloride (AlCl₃) display partial covalent character owing to the high charge density of Al³⁺, leading to some electron sharing alongside electrostatic forces.

Formation Processes

Monatomic Ion Formation

Monatomic ions form when a single atom gains or loses one or more electrons, resulting in a net positive (cation) or negative (anion) charge, respectively. This process primarily involves the removal of electrons from neutral atoms to create cations or the addition of electrons to form anions, driven by energy inputs that overcome the atom's electron binding energies. The ease of formation depends on the atom's electron configuration and external conditions, such as temperature or electric fields. Electron removal mechanisms for cation formation include photoionization, where a photon of sufficient energy ejects an from an atom, and collisional ionization, which occurs in gases when high-energy particles like or ions transfer momentum to atomic during collisions. Photoionization is prominent in gaseous environments, such as in stellar atmospheres or laboratory plasmas, where or higher-energy radiation excites valence beyond the threshold. Collisional ionization, often thermal in nature, predominates in high-temperature plasmas or flames, where kinetic energy from surrounding particles facilitates ejection. These processes are governed by the ionization potential, serving as the minimum energy barrier for electron removal. For practical generation of metal cations, is a common electrochemical method, where an applied voltage drives oxidation at the , stripping s from metal atoms in molten salts or aqueous solutions to produce cations that migrate to the . In contrast, anion formation in typically involves , where free s attach to neutral halogen atoms, often in gas-phase or solvated environments, forming stable anions due to the ' high electron affinities. Periodic trends significantly influence monatomic cation formation: the ease decreases across a period from left to right due to rising effective nuclear charge, which increases ionization energy and tightens electron binding, while it increases down a group as atomic size grows and shielding effects reduce the pull on valence electrons, lowering ionization energies for larger atoms. For instance, alkali metals readily form cations in flames through thermal ionization, where high temperatures (around 2000 K) provide the energy to ionize sodium or potassium atoms, enabling atomic emission spectroscopy for detection. Similarly, halide anions form in aqueous solutions when halogens like chlorine react with water to produce hydrochloric acid, dissociating into solvated chloride ions stabilized by hydration shells.

Polyatomic Ion Formation

Polyatomic ions form when a group of covalently bonded atoms gains or loses electrons, resulting in a net charge on the molecular unit. This process often occurs through protonation, where a base accepts a proton (H⁺), or deprotonation, where an acid donates a proton, as described in Brønsted-Lowry acid-base theory. For instance, ammonia (NH3NH_3) undergoes protonation to form the ammonium ion: NH3+H+→NH4+NH_3 + H^+ \to NH_4^+. Similarly, acetic acid deprotonates in water to yield the acetate ion: CH3COOH⇌CH3COO−+H+CH_3COOH \rightleftharpoons CH_3COO^- + H^+. Negative polyatomic ions can also arise from electron addition to neutral molecules, creating an excess of electrons that imparts a negative charge while maintaining covalent bonds within the group. This mechanism contributes to the formation of stable anionic species, such as the ion (OH−OH^-), where an extra electron is associated with the oxygen atom. Many polyatomic ions achieve enhanced stability through , where electrons are delocalized across multiple atoms, distributing the charge evenly and lowering the overall energy. In the ion (NO3−NO_3^-), for example, the negative charge is spread over the three oxygen atoms via resonance structures, resulting in equivalent N-O bond lengths that are intermediate between single and double bonds. The ion (SO42−SO_4^{2-}) similarly exhibits resonance among its four oxygen atoms, stabilizing the 2- charge through delocalized electrons involving and oxygen. In analytical contexts like , polyatomic ions often form through fragmentation of larger ionized molecules, where unstable molecular ions break into charged fragments and neutral species. For example, the of organic compounds such as can produce polyatomic carbocations like C4H9+C_4H_9^+ (m/z 57), which are detected as peaks in the due to their relative stability. This process highlights how covalent bonds within polyatomic fragments persist despite the overall charge.

Ionization and Energy

Ionization Potential

The (IE), often referred to as the ionization potential, is defined as the minimum required to remove an from a neutral atom or in the gas phase to form a positive ion. This process is endothermic, with the IE representing the energy difference between the initial neutral species and the resulting cation plus the free electron. In the context of , where a ejects an , the at the threshold is expressed as [ IE = h\nu where $h$ is Planck's constant and $\nu$ is the frequency of the photon, assuming the ejected electron has zero kinetic energy.[](https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/Supplemental_Modules_%28Physical_and_Theoretical_Chemistry%29/Spectroscopy/Photoelectron_Spectroscopy/Photoelectron_Spectroscopy) More generally, for photoelectrons with measurable kinetic energy $KE_e$, the relation is $IE = h\nu - KE_e$.[](https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/Supplemental_Modules_%28Physical_and_Theoretical_Chemistry%29/Spectroscopy/Photoelectron_Spectroscopy/Photoelectron_Spectroscopy) Successive ionization energies describe the stepwise removal of multiple electrons from an atom, with each subsequent [IE](/page/.ie) increasing because the remaining electrons experience a stronger effective nuclear attraction after the previous removal. The first [IE](/page/.ie) corresponds to the transition from the neutral atom to the singly charged cation (e.g., Na → Na⁺ + e⁻), while higher-order IEs involve further [ionization](/page/Ionization) of the cation. For sodium, the first [IE](/page/.ie) is 496 kJ/mol (5.139 eV), whereas the second [IE](/page/.ie), which removes a [core electron](/page/Core_electron) from Na⁺, is significantly higher at 4562 kJ/mol (47.286 eV).[](https://www.physics.nist.gov/PhysRefData/Handbook/Tables/sodiumtable1.htm) These values highlight how successive IEs rise sharply, particularly when penetrating inner shells. Ionization energies are commonly measured using photoelectron [spectroscopy](/page/Spectroscopy) (PES), which involves irradiating the sample with photons of known energy and analyzing the kinetic energies of the emitted electrons to determine binding energies.[](https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/Supplemental_Modules_%28Physical_and_Theoretical_Chemistry%29/Spectroscopy/Photoelectron_Spectroscopy/Photoelectron_Spectroscopy) PES provides both adiabatic (minimum energy for ground-state transitions) and vertical (fixed geometry) IEs, offering insights into electronic structure. [Periodic trends](/page/Periodic_trends) in IE arise from atomic properties: IE generally decreases down a group due to larger atomic radii and increased shielding by inner electrons, making valence electrons easier to remove, while it increases across a period owing to higher [effective nuclear charge](/page/Effective_nuclear_charge) pulling electrons closer to the nucleus.[](https://chem.libretexts.org/Courses/Oregon_Institute_of_Technology/OIT%253A_CHE_202_-_General_Chemistry_II/Unit_3%253A_Periodic_Patterns/3.3%253A_Trends_in_Ionization_Energy) For anion formation, the analogous concept is [electron affinity](/page/Electron_affinity) (EA), which is the energy released when an [electron](/page/Electron) is added to a neutral atom (e.g., Cl + e⁻ → [Cl](/page/Chlorine)⁻). Unlike IE, EA is typically exothermic for nonmetals, reflecting the stability gained by achieving a filled octet. The EA of [chlorine](/page/Chlorine), for instance, is 349 kJ/mol (3.613 eV), indicating a strong tendency to form Cl⁻.[](https://webbook.nist.gov/cgi/cbook.cgi?ID=C22537151&Mask=20) This value underscores chlorine's high [electronegativity](/page/Electronegativity) and role in ionic compounds. ### Energy Requirements The energy required to form ions varies significantly depending on the environment, such as gas phase, solution, or solid state, due to interactions like [solvation](/page/Solvation) or lattice stabilization. In the gas phase, the [ionization energy](/page/Ionization_energy) represents the minimum energy to remove an [electron](/page/Electron) from an isolated atom, but in solution, solvation energy stabilizes the resulting ion, effectively lowering the ionization energy compared to the gas phase by several [electron](/page/Electron) volts for many [species](/page/Species).[](https://www.sciencedirect.com/science/article/abs/pii/S0020169314005829) For example, the gas-phase ionization energy of [water](/page/Water) is about 12.6 eV, but in [aqueous solution](/page/Aqueous_solution), the effective value decreases due to solvent reorganization around the [hydronium](/page/Hydronium) ion.[](https://webbook.nist.gov/cgi/cbook.cgi?ID=C7732185&Mask=20) In ionic solids, the overall energy for ion formation and lattice assembly is described by the Born-Haber cycle, which balances endothermic steps like sublimation, [ionization](/page/Ionization), and dissociation against the exothermic [lattice energy](/page/Lattice_energy) release.[](https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Crystal_Lattices/Thermodynamics_of_Lattices/Lattice_Energy:_The_Born-Haber_cycle) The cycle's [enthalpy](/page/Enthalpy) change for compound formation, ΔH_f, incorporates the [ionization energy](/page/Ionization_energy) (IE) of the metal, electron affinity (EA) of the [nonmetal](/page/Nonmetal), and [lattice energy](/page/Lattice_energy) (U), such that ΔH_f = ΔH_sub + (1/2)D + IE - EA - U for a typical MX salt like NaCl, where U often dominates and stabilizes the ionic structure.[](https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Crystal_Lattices/Thermodynamics_of_Lattices/Lattice_Energy:_The_Born-Haber_cycle) This framework reveals that high [lattice energies](/page/Lattice_energy) in solids with small, highly charged ions reduce the net energy barrier for [ionization](/page/Ionization) relative to isolated atoms.[](https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Crystal_Lattices/Thermodynamics_of_Lattices/Lattice_Energy:_The_Born-Haber_cycle) Calculations of ionization energies distinguish between adiabatic and vertical processes, influencing accuracy in different contexts. The adiabatic ionization energy is the difference between the ground-state energies of the neutral atom and ion, allowing nuclear relaxation, while the vertical ionization energy assumes fixed nuclear geometry and typically exceeds the adiabatic value by 0.1–1 eV due to Franck-Condon factors.[](https://cccbdb.nist.gov/adiabaticx.asp) Temperature affects these calculations by populating vibrational and rotational states, potentially reducing the effective energy threshold at higher temperatures (e.g., above 1000 K) through thermal excitation, while pressure in dense gases can induce pressure ionization, lowering the barrier via continuum lowering effects in plasmas.[](https://link.aps.org/doi/10.1103/PhysRevE.79.016409) Certain processes bypass standard energy requirements, serving as exceptions. Autoionization occurs in excited neutral atoms above the [ionization](/page/Ionization) threshold, such as [helium](/page/Helium) in the (2s^2) ^1S state, where internal rearrangement ejects an [electron](/page/Electron) without additional external energy input beyond the excitation to that state.[](https://www.sciencedirect.com/topics/physics-and-astronomy/autoionization) In plasmas, barrierless ionization can arise through mechanisms like field-enhanced tunneling or collisional processes in high-density environments, where collective effects eliminate discrete energy barriers.[](https://www.sciencedirect.com/science/article/abs/pii/S0370157398000179) Low-energy cosmic ionizations, such as those driven by [ultraviolet](/page/Ultraviolet) [radiation](/page/Radiation) in interstellar media, exemplify environmental influences on these exceptions.[](https://adsabs.harvard.edu/full/1921Obs....44..261M) Typical first ionization energies for the first 20 elements, measured in the gas phase under standard conditions, illustrate the scale of these requirements and periodic trends, with values increasing across periods due to effective nuclear charge.[](https://www.nist.gov/pml/ground-levels-and-ionization-energies-neutral-atoms) | Atomic Number | Element | Ionization Energy (eV) | |---------------|---------|------------------------| | 1 | H | 13.5984 | | 2 | He | 24.5874 | | 3 | Li | 5.3917 | | 4 | Be | 9.3227 | | 5 | B | 8.2980 | | 6 | C | 11.2603 | | 7 | N | 14.5341 | | 8 | O | 13.6181 | | 9 | F | 17.4228 | | 10 | Ne | 21.5645 | | 11 | Na | 5.1391 | | 12 | Mg | 7.6462 | | 13 | Al | 5.9858 | | 14 | Si | 8.1517 | | 15 | P | 10.4867 | | 16 | S | 10.3600 | | 17 | Cl | 12.97 | | 18 | Ar | 15.7596 | | 19 | K | 4.3407 | | 20 | Ca | 6.1132 | ## Applications and Detection ### Technological Uses Ions are integral to ion propulsion systems in [spacecraft](/page/Spacecraft), where accelerated charged particles generate efficient thrust for deep-space missions. NASA's Dawn mission, launched in 2007, employed gridded ion thrusters that ionized xenon gas to produce Xe⁺ ions, which were then electrostatically accelerated to speeds 7-10 times faster than those from chemical rockets.[](https://science.nasa.gov/mission/dawn/technology/ion-propulsion/) The spacecraft carried 425 kg of xenon propellant, a [chemically inert](/page/Chemically_inert) [noble gas](/page/Noble_gas) stored compactly due to its high density, enabling over 2,000 days of cumulative thrust—far exceeding prior missions like [Deep Space 1](/page/Deep_Space_1)—while consuming only about 3.25 mg of propellant per second at maximum power.[](https://science.nasa.gov/mission/dawn/technology/ion-propulsion/) This technology's high [specific impulse](/page/Specific_impulse) and adjustable thrust levels allowed Dawn to achieve multiple planetary orbits with minimal fuel, demonstrating ions' potential for sustained, low-thrust propulsion in interplanetary travel.[](https://science.nasa.gov/mission/dawn/technology/ion-propulsion/) In semiconductor fabrication, [ion implantation](/page/Ion_implantation) introduces [dopant](/page/Dopant) ions into crystalline materials to modify their electrical conductivity, forming the basis for devices like transistors. [Boron](/page/Boron) ions serve as p-type dopants, accepting electrons to create positively charged holes in [silicon](/page/Silicon) lattices, while [phosphorus](/page/Phosphorus) ions act as n-type dopants, donating excess electrons.[](https://www.iue.tuwien.ac.at/phd/wittmann/node7.html) By accelerating these ions at energies from 100 eV to 3 MeV and doses of 10¹¹ to 10¹⁶ ions/cm², precise depth profiles are achieved, enabling the creation of p-n junctions where p-type and n-type regions meet to establish built-in [electric fields](/page/Electric_Fields) essential for current rectification and amplification.[](https://www.iue.tuwien.ac.at/phd/wittmann/node7.html) Post-implantation annealing activates the dopants and repairs lattice damage, allowing carrier concentrations from 10¹³ to 10²¹ cm⁻³, which is critical for high-performance [CMOS](/page/CMOS) integrated circuits.[](https://www.iue.tuwien.ac.at/phd/wittmann/node7.html) Lithium-ion batteries harness the electrochemical migration of Li⁺ ions to store and deliver [electrical energy](/page/Electrical_energy), underpinning portable and grid-scale power systems. In a typical cell, the [anode](/page/Anode) (often [graphite](/page/Graphite)) releases Li⁺ ions during discharge, which diffuse through a liquid [electrolyte](/page/Electrolyte) to intercalate into the [cathode](/page/Cathode) (such as layered metal oxides), while [electron](/page/Electron)s flow externally via current collectors to power connected devices.[](https://www.energy.gov/energysaver/articles/how-lithium-ion-batteries-work) A porous separator prevents direct electron conduction within the battery, forcing the circuitous path that generates usable current, with the process reversing during charging to restore ion positions.[](https://www.energy.gov/energysaver/articles/how-lithium-ion-batteries-work) This reversible intercalation mechanism yields high [energy density](/page/Energy_density)—measured in watt-hours per kilogram—and rechargeability, making Li⁺-based batteries dominant in [consumer electronics](/page/Consumer_electronics), electric vehicles, and [renewable energy](/page/Renewable_energy) storage, though limited by factors like [electrolyte](/page/Electrolyte) stability and [dendrite](/page/Dendrite) formation.[](https://www.energy.gov/energysaver/articles/how-lithium-ion-batteries-work) Medical applications of ions include targeted cancer therapies and sterilization processes that leverage their energetic interactions with biological matter. Ion beam therapy accelerates heavy ions, such as carbon ions, to deposit energy precisely at tumor sites via the Bragg peak, where maximum ionization occurs at a selectable depth with rapid dose fall-off beyond, sparing adjacent healthy tissues more effectively than X-rays or protons.[](https://pmc.ncbi.nlm.nih.gov/articles/PMC5894097/) This approach enhances biological effectiveness—up to 3-4 times higher than conventional radiation—by inducing complex DNA damage resistant to cellular repair, proving particularly advantageous for radioresistant tumors like those in the pancreas or skull base, as demonstrated in treatments at Japan's National Institute of Radiological Sciences since 1994.[](https://pmc.ncbi.nlm.nih.gov/articles/PMC5894097/) ### Radiation Detection Methods Ionization chambers represent a foundational [principle](/page/Principle) in radiation detection, operating by measuring the [electric charge](/page/Electric_charge) produced when [ionizing radiation](/page/Ionizing_radiation) passes through a gas-filled volume, creating pairs of electrons and positive ions that are collected under an applied [electric field](/page/Electric_field). This direct collection of ion pairs allows for the quantification of [radiation](/page/Radiation) intensity, as the number of ion pairs is proportional to the energy deposited by the [radiation](/page/Radiation). Early developments in this method, including contributions from J.J. Thomson in understanding [ionization](/page/Ionization) processes around 1899, laid the groundwork for modern detectors.[](https://www.iaea.org/sites/default/files/publications/magazines/bulletin/bull23-4/23405043136.pdf) Geiger-Müller counters build on this principle but amplify the signal through Townsend [avalanche](/page/Avalanche)s, where initial ion pairs accelerate under a [high voltage](/page/High_voltage), ionizing additional gas molecules and creating a cascading avalanche of ions and electrons that results in a detectable pulse. These counters are particularly sensitive to beta and gamma [radiation](/page/Radiation), providing count rates rather than [energy](/page/Energy) [information](/page/Information) due to the saturating nature of the avalanche process. Scintillation detectors, in contrast, detect ions indirectly by converting the [energy](/page/Energy) from [radiation](/page/Radiation)-induced [ionization](/page/Ionization) and excitation in a [scintillator](/page/Scintillator) material—such as [sodium iodide](/page/Sodium_iodide)—into visible light photons, which are then captured by a [photomultiplier tube](/page/Photomultiplier_tube) to produce an electrical signal proportional to the radiation [energy](/page/Energy).[](https://indico.bnl.gov/event/9065/contributions/44421/attachments/32007/50796/chapt4.pdf)[](https://user-web.icecube.wisc.edu/~tmontaruli/801/lect13.pdf)[](https://pmc.ncbi.nlm.nih.gov/articles/PMC5843761/) Modern advances include [semiconductor](/page/Semiconductor) detectors, exemplified by silicon-based devices that detect charged particles through the generation of electron-hole pairs in a depleted [semiconductor](/page/Semiconductor) junction, offering superior energy resolution compared to gas-filled detectors due to the higher density of charge carriers. [Time-of-flight mass spectrometry](/page/Time-of-flight_mass_spectrometry) enhances ion identification by accelerating ions in a field-free drift region and measuring their arrival time at a detector, which depends on their [mass-to-charge ratio](/page/Mass-to-charge_ratio), enabling precise [speciation](/page/Speciation) of ions produced in [radiation](/page/Radiation) events. These techniques provide high-resolution [data](/page/Data) for complex ion mixtures.[](https://www.nrc.gov/docs/ml1122/ml11229a683.pdf)[](https://pmc.ncbi.nlm.nih.gov/articles/PMC11598097/) Such detection methods find critical applications in [nuclear physics](/page/Nuclear_physics), where semiconductor and time-of-flight systems identify and characterize ions from particle accelerators and reactions, facilitating studies of nuclear structure and reactions. In [environmental monitoring](/page/Environmental_monitoring), ionization chambers and scintillation detectors are employed to measure [radon](/page/Radon) ions and their decay products, assessing [indoor air quality](/page/Indoor_air_quality) and [radiation exposure](/page/Radiation_exposure) risks in real-time.[](https://www.nist.gov/programs-projects/applied-quantum-sensors-charged-particle-detection)[](https://pmc.ncbi.nlm.nih.gov/articles/PMC6749372/)

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