Ion
An ion is an electrically charged particle produced when a neutral atom or molecule gains or loses one or more electrons, resulting in a net positive or negative charge.[1] These charged entities are fundamental to chemical bonding, electrical conductivity in solutions, and numerous natural processes.[2] Ions are classified by their charge and composition: cations carry a positive charge due to electron loss, while anions bear a negative charge from electron gain.[3] They can be monatomic, consisting of a single charged atom such as Na⁺ (sodium ion) or Cl⁻ (chloride ion), or polyatomic, involving groups of atoms like NH₄⁺ (ammonium) or SO₄²⁻ (sulfate).[4] Ionic compounds form when cations and anions combine in ratios that yield electrical neutrality, creating stable lattices essential for salts, minerals, and many industrial materials.[2] The concept of the ion emerged in the 19th century through electrochemical studies, with the term "ion" coined by Michael Faraday in 1834—derived from the Greek word ion meaning "to go"—to describe migrating charged particles during electrolysis.[5] Faraday also introduced terms like "cation" for positively charged ions (those moving toward the cathode) and "anion" for negatively charged ones.[6] This framework built on earlier observations of electrolysis by Humphry Davy and laid the groundwork for understanding atomic structure, later advanced by J.J. Thomson's 1897 discovery of the electron as the particle responsible for ionization.[7] Ions play critical roles across chemistry and biology: in aqueous solutions, they enable ionic bonding and conductivity, driving reactions in batteries, electroplating, and environmental processes.[2] Biologically, ions such as Na⁺, K⁺, Ca²⁺, and Cl⁻ are vital for nerve impulse transmission, muscle contraction, cellular metabolism, ATP production, and maintaining fluid balance in organisms.[2] Disruptions in ion transport, as seen in conditions like cystic fibrosis, underscore their physiological importance.[8] In concentrated electrolytes, ion interactions influence reaction rates and structures in both synthetic and natural systems.[9]Fundamentals
Definition and Charge
An ion is an atom or molecule that has a net electric charge resulting from the gain or loss of one or more electrons, distinguishing it from neutral atoms or molecules which have equal numbers of protons and electrons.[1] This charge imbalance arises when electrons are transferred during chemical interactions, leading to either a positive or negative charge on the particle.[10] The term "ion" derives from the Greek word iōn meaning "going" or "that which goes," coined by William Whewell in 1834 to describe the migrating charged entities observed by Michael Faraday in electrolytic solutions.[11] Ions are denoted using superscripts to indicate their charge magnitude and sign following the chemical symbol or formula; for monatomic ions, the convention omits the numeral "1" for single charges, as in Na⁺ for the sodium cation (which has lost one electron) or Cl⁻ for the chloride anion (which has gained one electron).[3] Polyatomic ions, consisting of covalently bonded atoms with an overall charge, use similar notation, such as OH⁻ for the hydroxide ion.[4] The presence of net charge on ions has fundamental implications for their behavior: in solutions, ions enable the conduction of electricity by migrating toward electrodes of opposite charge, classifying such solutions as electrolytes essential for processes like electrochemical cells.[12] Additionally, ions participate actively in chemical reactions, often through electrostatic attractions that form ionic compounds, unlike neutral species which require different interaction mechanisms.[13]Cations and Anions
Cations are positively charged ions that form when atoms, typically metals, lose one or more electrons from their valence shells.[2] Common examples include the hydrogen cation (H⁺), which exists as a proton or, in aqueous solutions, as the hydronium ion (H₃O⁺); the sodium cation (Na⁺); and the calcium cation (Ca²⁺).[14] Due to the loss of electrons, cations exhibit a smaller ionic radius compared to their parent neutral atoms, as the remaining electrons experience a higher effective nuclear charge that pulls them closer to the nucleus.[15] This reduction in size also decreases electron-electron repulsion, further contracting the ion.[16] Anions, in contrast, are negatively charged ions formed when atoms, usually nonmetals, gain one or more electrons to achieve a stable electron configuration.[2] Representative examples are the fluoride anion (F⁻), oxide anion (O²⁻), and chloride anion (Cl⁻).[17] Anions are larger than their corresponding neutral atoms because the added electrons increase electron-electron repulsion in the outer shell, while the nuclear charge remains unchanged, leading to greater spatial expansion.[15][18] The magnitude of charge on an ion influences its properties, with monovalent cations like potassium (K⁺) carrying a single positive charge and polyvalent cations like iron(III) (Fe³⁺) bearing multiple charges.[19] Higher charge magnitudes generally result in stronger interactions with surrounding molecules; for instance, in aqueous solutions, smaller monovalent cations such as lithium (Li⁺) form highly hydrated structures due to their high charge density, attracting a large number of water molecules into their solvation shell.[20] In ionic compounds, cations and anions combine in ratios that ensure overall charge neutrality, as seen in sodium chloride (NaCl), where one Na⁺ balances one Cl⁻ to form a neutral lattice.[3]Historical Development
Early Observations
Early observations of phenomena related to ions trace back to ancient times, with initial hints emerging from natural curiosities involving static electricity. Around 600 BCE, the Greek philosopher Thales of Miletus noted that amber, when rubbed, could attract lightweight objects such as feathers and straw, marking one of the earliest recorded instances of electrostatic attraction.[21] This observation, though not understood in terms of charged particles, represented a foundational recognition of electrical effects that would later connect to ionic behavior. In medieval alchemy, Arabic scholars advanced early experimental work on salts and their properties, including reactions that suggested decomposable components within compounds. Figures like Jabir ibn Hayyan (c. 721–815 CE), known as Geber in the Latin West, systematically studied salts, acids, and bases, preparing mixtures such as aqua regia capable of dissolving noble metals and observing effervescent or dissociative effects in saline solutions.[22] These alchemical investigations, focused on transformation and purification, provided precursors to recognizing ionic dissociation without the framework of electricity. The 18th century brought more structured inquiries into electrical phenomena. In the 1730s, French physicist Charles François du Fay distinguished two types of electricity—vitreous, produced by rubbing glass with silk, and resinous, from amber rubbed with fur—demonstrating that like charges repel and unlike attract, laying groundwork for understanding charge separation relevant to ions.[23] Meanwhile, in 1766, English chemist Henry Cavendish identified hydrogen as a distinct element by reacting metals with acids, observing the liberation of a flammable gas and inferring that acids contain hydrogen as a key constituent, an early insight into what would later be termed hydrogen ions.[24] The dawn of the 19th century enabled direct experimentation with ionic currents through electrical devices. In 1800, Italian physicist Alessandro Volta invented the voltaic pile, a stack of alternating metal discs and brine-soaked cloth that generated a steady electric current via chemical reactions in the electrolyte, producing ionic migration in solution.[25] Building on this, English chemist Humphry Davy in 1807 used electrolysis powered by large batteries to decompose molten salts, isolating reactive metals such as sodium from sodium hydroxide, revealing that salts could be broken into charged constituents at electrodes.[26] These experimental advances culminated in conceptual progress in the early 19th century, when Swedish chemist Jöns Jacob Berzelius developed his electrochemical dualism theory around 1811, positing that compounds consist of electropositive and electronegative parts, providing a theoretical basis for ionic conduction observed in earlier work. In the 1830s, Michael Faraday introduced the term "electrolyte" to describe substances that conduct electricity through the migration of their internal particles during decomposition, and his quantitative studies further refined these ideas.[27][28]Key Discoveries and Milestones
In the 1830s, Michael Faraday established foundational quantitative relationships in electrolysis through his two laws, which described the behavior of charged particles later termed ions. Faraday's first law states that the mass of a substance altered at an electrode during electrolysis is directly proportional to the quantity of electricity transferred. His second law posits that when the same quantity of electricity passes through different electrolytes, the masses of the substances deposited are proportional to their chemical equivalent weights. In his 1834 publication, Faraday coined the term "ion" to denote these migrating electropositive and electronegative particles responsible for conduction in electrolytes.[28] Building on Faraday's work, Svante Arrhenius advanced ion theory in 1884 by proposing that electrolytes dissociate into independent charged ions in aqueous solutions, even without an electric current, thereby explaining electrical conductivity and reaction rates. This electrolytic dissociation theory resolved discrepancies in osmotic pressure and conductivity observations, positing that ions exist as discrete entities capable of independent movement. Arrhenius's contributions were recognized with the 1903 Nobel Prize in Chemistry for his work on electrolyte solutions.[29][30] The early 20th century saw breakthroughs in understanding atomic structure through ion studies, particularly via mass spectrometry. In 1913, J.J. Thomson identified isotopes of stable elements, such as neon, by analyzing positive rays—streams of positively charged ions—in a modified cathode ray tube, revealing that elements could consist of atoms with different masses but identical chemical properties. Building on this, Francis Aston developed the first mass spectrograph in 1919, which precisely measured the mass-to-charge ratios of ions, confirming isotopes across multiple elements and enabling accurate atomic mass determinations. Aston's innovations earned him the 1922 Nobel Prize in Chemistry and laid the groundwork for modern isotopic analysis.[31][32] Post-1920s advancements integrated quantum mechanics into ion descriptions, applying the Schrödinger equation to model ionic wavefunctions and energies, particularly for simple systems like hydrogenic ions, which provided a theoretical framework for ionization potentials and spectral lines. In the 1950s, ion implantation emerged as a practical technique, accelerating ions into solid materials to dope semiconductors with precise control over impurity concentrations, revolutionizing device fabrication in electronics.[33][34] More recent milestones include progress in ion trapping for quantum computing, with NIST experiments from 1995 onward demonstrating entangled ion states and two-qubit gates using laser-cooled calcium and beryllium ions in Paul traps, advancing scalable quantum information processing. Additionally, since the 1980s, the GSI Helmholtz Centre has pioneered superheavy ion synthesis through heavy-ion collisions in accelerators like UNILAC and SIS, producing elements beyond uranium—such as bohrium (1981) and flerovium (co-discovered 1998)—and extending the periodic table up to oganesson (2006), with ongoing efforts through the 2020s at FAIR to explore island-of-stability predictions.[35][36]Physical Properties
Size and Mobility
Ions exhibit distinct physical sizes determined by their effective ionic radii, which vary based on charge, atomic number, and environment. Cations are generally smaller than their corresponding neutral atoms due to the loss of electrons, which increases the effective nuclear charge and pulls the remaining electrons closer to the nucleus; for example, the ionic radius of Li⁺ is 76 pm, compared to 152 pm for neutral Li.[37] In contrast, anions are larger than their parent atoms because additional electrons increase electron-electron repulsion, expanding the electron cloud; iodide ion (I⁻), for instance, has an ionic radius of 220 pm.[38] Across the periodic table, ionic radii decrease from left to right in a period for cations (due to increasing nuclear charge) and increase down a group (due to additional electron shells), while anions follow similar trends but with larger values overall.[39] In solutions, the effective size of ions often includes a hydration shell, where solvent molecules (e.g., water) surround the ion, significantly increasing its hydrodynamic radius compared to the bare ionic radius in crystals. For sodium ion (Na⁺), the crystal ionic radius is approximately 102 pm, but its hydrated radius in water is about 360 pm due to the tightly bound first solvation layer. This solvation affects ion transport, as the larger effective size influences interactions with the solvent. Ionic mobility refers to the ease with which ions move under an applied electric field, quantified by the drift velocity v_d = \mu E, where \mu is the mobility and E is the field strength. In electrolyte solutions, mobility governs ionic conductance, with smaller, less solvated ions exhibiting higher mobility; for instance, H⁺ and OH⁻ in water have anomalously high mobilities due to proton-hopping mechanisms via hydrogen bonds.[40] The Einstein relation connects mobility to diffusion, given by D = \frac{\mu k_B T}{q}, where D is the diffusion coefficient, k_B is Boltzmann's constant, T is temperature, and q is the ion charge, highlighting the thermal equilibrium between random diffusion and directed drift.[40] Key factors influencing mobility include solvent viscosity (higher viscosity reduces \mu by impeding motion) and solvation effects (stronger solvation shells increase effective size and drag).[41] In mass spectrometry, ion size and mass directly impact separation techniques that exploit mobility in fields. Ions are first accelerated by an electric field to gain kinetic energy \frac{1}{2} m v^2 = q V, where m is mass, v is velocity, q is charge, and V is accelerating voltage, resulting in lighter ions achieving higher speeds.[42] Subsequent deflection in magnetic or electric fields separates ions by their mass-to-charge ratio (m/z), as the radius of curvature in a magnetic field is r = \frac{m v}{q B}, with B the magnetic flux density; ions with lower m/z deflect more sharply.[43] Time-of-flight (TOF) analyzers measure mobility by timing ion travel over a fixed distance in a field-free region, where flight time t = \sqrt{\frac{2 d m}{q V}} (with d the drift length) allows determination of m/z, enabling analysis of ion masses from picometers-scale radii to complex molecules.[44] In plasmas, ion mobility is primarily limited by collisions with neutral particles or other charged species, reducing drift velocities compared to collisionless environments. The mean free path between collisions determines effective mobility, with higher densities leading to more frequent scattering and lower \mu.[45] A key parameter is the Debye length, which characterizes the spatial extent of electric field screening by mobile charges, given by \lambda_D = \sqrt{\frac{\epsilon_0 k_B T}{n e^2}}, where \epsilon_0 is the vacuum permittivity, T is the electron temperature, n is the electron density, and e is the elementary charge; this length scale (typically micrometers in laboratory plasmas) governs how quickly ions respond collectively to fields before screening occurs.[46]Occurrence in Nature
Ions are ubiquitous in Earth's atmosphere, primarily generated by cosmic rays that ionize nitrogen and oxygen molecules, producing primary ions such as N₂⁺ and O₂⁺, along with negative ions like O₂⁻ through electron attachment.[47][48] At sea level under fair-weather conditions, ion concentrations typically range from 200 to 800 negative ions per cubic centimeter and 250 to 1500 positive ions per cubic centimeter, averaging around 10³ ions/cm³ overall.[49] These ions play key roles in natural processes, including facilitating aerosol nucleation that contributes to cloud formation and aiding charge separation in thunderstorms that leads to lightning discharges.[50] In oceanic and geological environments, ions dominate dissolved compositions due to mineral weathering and evaporation cycles. Seawater contains high concentrations of Na⁺ (approximately 0.47 M) and Cl⁻ (approximately 0.54 M), accounting for over 85% of its total ionic content, with an overall salinity equivalent to about 0.5 M in major salts.[51] Geological processes, such as the dissolution of limestone (CaCO₃) by carbonic acid from rainwater, release Ca²⁺ ions into groundwater and rivers, contributing to hardness in natural waters and supporting ecosystems through mineral nutrient supply.[52] Biological systems rely on precise ion distributions across cell membranes for essential functions. In mammalian cells, intracellular K⁺ concentrations reach about 120 mM, while Na⁺ is low at around 10-15 mM; extracellularly, Na⁺ dominates at 145 mM and K⁺ is only 4 mM, establishing the electrochemical gradients vital for membrane potentials and nerve signaling.[53] In astrophysical contexts, ions abound in the interstellar medium (ISM), where ultraviolet radiation from stars photoionizes hydrogen and helium, yielding H⁺ and He⁺ as prevalent species amid low-density plasmas.[54] The solar wind, a stream of plasma from the Sun, consists predominantly of protons (95%) and alpha particles (He²⁺, 5%), carrying ionized material through the heliosphere and influencing planetary magnetospheres.[55] Recent James Webb Space Telescope (JWST) observations from 2022 to 2025 have enhanced understanding of ion abundances in exoplanet atmospheres, revealing significant ionization levels in ultra-hot Jupiters through detections of atomic and ionic lines, which suggest day-night variations in ion chemistry driven by stellar irradiation.[56]Chemical Aspects
Notation and Subtypes
Ions are denoted using standardized symbols that include the chemical formula with a superscript indicating the charge magnitude and sign, such as Cu²⁺ for the copper(II) ion or SO₄²⁻ for the sulfate ion.[57] According to IUPAC recommendations, the charge is represented as a right upper superscript following the formula, with the sign placed after the numeral (e.g., positive as ⁺, negative as ⁻), and for coordination entities, charge numbers may be specified in Arabic numerals within parentheses after the name, such as hexaamminecobalt(3+).[57] For elements with variable oxidation states, Roman numerals in parentheses denote the oxidation number, as in iron(II) for Fe²⁺ or iron(III) for Fe³⁺, ensuring unambiguous identification in chemical nomenclature.[58] Subtypes of ions are classified primarily as monatomic or polyatomic, with further distinctions based on composition and structure. Monatomic ions consist of a single atom with a net charge, named by modifying the element name—cations retain the elemental name with charge indication (e.g., Mg²⁺ as magnesium(2+)), while anions typically end in "-ide" (e.g., Cl⁻ as chloride(1−)).[57] Polyatomic ions, in contrast, involve multiple atoms and are named using systematic additive nomenclature or retained traditional names; for example, NH₄⁺ is ammonium or azanium, and PO₄³⁻ is phosphate or tetraoxophosphate(3−).[57][58] Oxoanions, a common polyatomic subtype containing oxygen, follow naming conventions that reflect central atom oxidation states, such as NO₃⁻ as nitrate (nitrogen in +5 state) or SO₄²⁻ as sulfate (sulfur in +6 state), with systematic alternatives like trioxonitrate(1−).[57] Acid-base related ions include the hydronium cation H₃O⁺, named oxidanium or oxonium, which represents protonated water, and the hydroxide anion OH⁻, named hydroxide or oxidanide.[57] Stable radical ions, such as the superoxide anion O₂⁻•, are denoted with a superscript dot to indicate the unpaired electron, and named using substitutive nomenclature like dioxide(1−).[57] Isotopic variants of ions specify the mass number as a left superscript on the elemental symbol, for instance ²³Na⁺ for the sodium-23 cation, following general nuclide notation rules to distinguish isotopes in chemical contexts.[57] Cluster ions, aggregates of multiple atoms or molecules with a net charge, employ additive or substitutive nomenclature, such as tetrachloridocuprate(2−) for [CuCl₄]²⁻ or tetraaluminide(2−) for Al₄²⁻, with 2013 IUPAC updates in the Blue Book emphasizing systematic names for preferred usage in complex structures like fullerenides.[57][59]Ionization Processes
Ionization processes refer to the mechanisms by which neutral atoms or molecules lose or gain electrons to form ions, typically requiring specific energy inputs under controlled conditions. For monatomic species, ionization primarily involves the removal of electrons from gaseous atoms, achieved through thermal, photo, or field methods. Thermal ionization occurs when atoms with low ionization potentials, such as alkali metals, are heated on a high-work-function metal surface like rhenium, promoting electron ejection due to thermal agitation exceeding the binding energy.[60] Photoionization happens when an atom absorbs a photon with energy equal to or greater than its ionization potential, ejecting an electron; the kinetic energy of the emitted electron is given by KE = h\nu - IP, where h\nu is the photon energy and IP is the ionization potential.[61] Field ionization employs strong electric fields, on the order of 10^9–10^10 V/m, to lower the potential barrier for electron tunneling from the atom, often facilitated by sharp tips in scanning tunneling microscopes (STMs); advances since the 1980s have integrated field ionization with STM for atomic-scale imaging and manipulation, enabling precise control in vacuum environments up to the 2020s.[62] The first ionization energy (IE₁), the minimum energy to remove the outermost electron from a neutral atom, exhibits periodic trends: it is highest for noble gases and decreases down a group due to increasing atomic radius and shielding, which reduces the effective nuclear attraction on valence electrons. For hydrogen, IE₁ is 13.59844 eV, while for sodium it is 5.139 eV, reflecting weaker binding in larger atoms.[63][64][65] Successive ionization energies increase sharply, as removing electrons from positively charged ions requires overcoming greater electrostatic repulsion; for sodium, the second ionization energy (IE₂) is 47.286 eV, over nine times IE₁, due to the stability of the noble gas core.[65] For polyatomic species, ionization often leads to fragmentation, known as dissociative ionization, where the molecule breaks into ionic and neutral fragments upon electron or photon absorption. A classic example is water, where electron impact or photoexcitation can produce H⁺ + OH• via cleavage of the O-H bond, with the process threshold around 18–20 eV depending on the state.[66] In gaseous environments, cluster ions form through sequential attachment of neutral molecules to a core ion, stabilized by evaporative cooling; these solvated clusters, like (H₂O)_n H⁺, mimic solution-phase behavior and are studied in supersonic expansions.[67] Anion formation typically proceeds via electron capture, an exothermic process for electronegative atoms where a free electron attaches to form a stable negative ion, releasing energy equal to the electron affinity. For fluorine, F + e⁻ → F⁻ is exothermic by 328 kJ/mol, driven by the high affinity of the compact 2p orbital.[68] Autoionization occurs in excited atomic or molecular states above the ionization continuum, where an electron from a lower orbital fills the excited-state vacancy, ejecting another electron without external energy input; this radiationless decay is prevalent in helium-like systems and influences spectral line broadening.[69] Multiphoton ionization, enabled by intense laser fields since the 1960s, allows atoms to absorb multiple photons whose combined energy exceeds the ionization potential, even if individual photons are below threshold; this nonlinear process is widely used for state-selective ionization in spectroscopy, with rates scaling as I^k where I is laser intensity and k is the number of photons.Bonding and Compounds
Ionic Bonding Mechanism
Ionic bonding primarily results from the electrostatic attraction between positively charged cations and negatively charged anions. This attraction is governed by Coulomb's law, which quantifies the force F between two point charges as F = k \frac{q_1 q_2}{r^2}, where k is the Coulomb constant ($8.99 \times 10^9 \, \mathrm{N \cdot m^2 / C^2}), q_1 and q_2 are the charges on the ions, and r is the distance between their centers.[70] In ionic compounds, these pairwise interactions extend throughout the crystal lattice, stabilizing the structure through a balance of attractive and repulsive forces.[71] In extended ionic crystals, the total electrostatic potential energy, known as the lattice energy U, accounts for interactions among all ions and is approximated by the Born-Landé equation: U = -\frac{N_A \alpha k q_1 q_2}{r_0} \left(1 - \frac{1}{n}\right), where N_A is Avogadro's number, \alpha is the Madelung constant (a structure-dependent factor accounting for the lattice geometry), r_0 is the equilibrium interionic distance, and n is the Born exponent (typically 7-12, reflecting short-range repulsion).[72] This equation highlights how lattice energy increases with ion charges and decreases with interionic distance, making compounds with small, highly charged ions particularly stable. While bonds with large electronegativity differences (e.g., \Delta EN = 2.1 for NaCl between Na (0.9) and Cl (3.0)) are considered purely ionic, partial covalent character can arise when a small, highly charged cation polarizes a large anion, as described by Fajans' rules. These rules predict greater covalency for cations with high charge density and anions with high polarizability.[73][74] Ionic crystals adopt specific lattice structures to maximize electrostatic attractions while minimizing repulsions, influenced by the radius ratio of cation to anion. For example, the rock salt (NaCl) structure features a face-centered cubic arrangement with coordination number 6, where each Na⁺ ion is surrounded by six Cl⁻ ions, and vice versa. In contrast, the cesium chloride (CsCl) structure has a body-centered cubic lattice with coordination number 8, suitable for larger cations like Cs⁺ relative to Cl⁻. These coordination numbers reflect optimal packing based on ion sizes, ensuring stability. The stability of ionic compounds in solution involves dissociation, where lattice energy must be overcome by solvation (hydration) energy. The Born-Haber cycle provides a thermodynamic pathway to calculate lattice energy indirectly by summing enthalpies of formation, sublimation, ionization, electron affinity, and dissociation steps, balancing the exothermic lattice formation against endothermic processes.[75] Hydration energy, arising from ion-dipole interactions with water, often compensates for lattice energy in soluble salts, leading to net exothermic dissolution. Since the 1990s, computational methods like density functional theory (DFT) have advanced predictions of ionic bonding by solving the Schrödinger equation approximately through electron density functionals, enabling accurate modeling of lattice energies and structures for complex ionic systems without empirical parameters.[76]Common Monatomic and Polyatomic Ions
Monatomic cations are positively charged ions derived from single atoms, commonly formed by metals losing electrons. Alkali metals from group 1 of the periodic table, such as lithium (Li⁺), sodium (Na⁺), and potassium (K⁺), typically form +1 cations by losing one valence electron.[77] These ions often originate from the dissolution of minerals, for instance, Na⁺ from halite (NaCl) deposits and K⁺ from sylvite (KCl).[78] Alkaline earth metals from group 2, including magnesium (Mg²⁺) and calcium (Ca²⁺), produce +2 cations by losing two electrons, sourced similarly from mineral weathering like dolomite for Mg²⁺ and Ca²⁺.[77] Transition metals exhibit variable charges; iron forms Fe²⁺ and Fe³⁺, while copper yields Cu⁺ and Cu²⁺, arising from ores such as hematite for iron ions.[10] Monatomic anions are negatively charged single-atom ions, primarily from nonmetals gaining electrons. Halides from group 17, including fluoride (F⁻), chloride (Cl⁻), bromide (Br⁻), and iodide (I⁻), acquire a -1 charge and are prevalent in ionic salts like sodium chloride (table salt). Oxide (O²⁻) and nitride (N³⁻) ions, with -2 and -3 charges respectively, occur in compounds such as calcium oxide and magnesium nitride, often from reactive metal-nonmetal combinations.[79] Polyatomic cations consist of multiple atoms with a net positive charge. The ammonium ion (NH₄⁺) forms through protonation of ammonia (NH₃) in aqueous solutions, acting as a stable cation in salts.[80] Hydronium (H₃O⁺) arises from protonation of water, representing the hydrated proton in acidic media.[80] Polyatomic anions feature covalently bound atoms with an overall negative charge. Common examples include sulfate (SO₄²⁻), nitrate (NO₃⁻), phosphate (PO₄³⁻), and carbonate (CO₃²⁻), which maintain internal covalent bonds while interacting ionically in compounds.[81] Organic polyatomic anions like acetate (CH₃COO⁻) derive from carboxylic acids and appear in salts such as sodium acetate.[82] Polyatomic ions exhibit stability due to internal covalent bonds that hold their structure intact, despite the overall ionic charge influencing reactivity in solutions and compounds.[83] For instance, ammonium nitrate (NH₄NO₃) serves as a key nitrogen fertilizer, providing NH₄⁺ and NO₃⁻ for plant uptake.[84] Phosphate ions (PO₄³⁻) were historically used in detergents to enhance cleaning but have been phased out due to environmental concerns like eutrophication.[85]| Ion Type | Examples | Charge | Common Sources/Roles |
|---|---|---|---|
| Monatomic Cations (Alkali) | Li⁺, Na⁺, K⁺ | +1 | Mineral dissolution (e.g., halite for Na⁺); electrolytes in biology |
| Monatomic Cations (Alkaline Earth) | Mg²⁺, Ca²⁺ | +2 | Weathering of dolomite; bone structure (Ca²⁺) |
| Monatomic Cations (Transition) | Fe²⁺/Fe³⁺, Cu⁺/Cu²⁺ | Variable | Ore minerals; redox processes |
| Monatomic Anions (Halides) | F⁻, Cl⁻, Br⁻, I⁻ | -1 | Sea salts; disinfectants (Cl⁻) |
| Monatomic Anions (Others) | O²⁻, N³⁻ | -2/-3 | Metal oxides/nitrides; ceramics |
| Polyatomic Cations | NH₄⁺, H₃O⁺ | +1 | Ammonia solutions; acids |
| Polyatomic Anions | SO₄²⁻, NO₃⁻, PO₄³⁻, CO₃²⁻, CH₃COO⁻ | Variable (-1 to -3) | Fertilizers, buffers; detergents (historical) |