Ion

Last updated

An ion ( /ˈ.ɒn,-ən/ ) [1] is an atom or molecule with a net electrical charge.

Contents

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] while an anion is a negatively charged ion with more electrons than protons. [3] 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 atomic or monatomic ions, while two or more atoms form molecular ions or polyatomic ions. 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. [4] 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

The word ion was coined from Greek neuter present participle of ienai (Greek : ἰέναι), meaning "to go". A cation is something that moves down (Greek : κάτω pronounced kato, meaning "down") and an anion is something that moves up (Greek : ano ἄνω, 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) [5] 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. [6] [7] 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. [5]

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. [8] 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. [9] [10] [11]

Characteristics

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

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

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

Anions and cations

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. Ions.svg
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 less 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 (−) ( /ˈænˌ.ən/ ANN-eye-ən, from the Greek word ἄνω (ánō), meaning "up" [12] ) is an ion with more electrons than protons, giving it a net negative charge (since electrons are negatively charged and protons are positively charged). [13]

A cation (+) ( /ˈkætˌ.ən/ KAT-eye-ən, from the Greek word κάτω (káto), meaning "down" [14] ) is an ion with fewer electrons than protons, giving it a positive charge. [15]

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. [16]

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." [17]

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

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 extremely important[ citation needed ]; an example is energy that drives the breakdown of adenosine triphosphate (ATP)[ clarification needed ].

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

Schematic of an ion chamber, showing drift of ions. Electrons drift faster than positive ions due to their much smaller mass. Ion chamber operation.gif
Schematic of an ion chamber, showing drift of ions. Electrons drift faster than positive ions due to their much smaller mass.
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. Electron avalanche.gif
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. [4]

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

Chemistry

Denoting the charged state

Equivalent notations for an iron atom (Fe) that lost two electrons, referred to as ferrous. Ions notation.svg
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 notNa1+.

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. Ions notation2.svg
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+ example seen above is referred to as Fe(II) or FeII. The Roman numeral designates the formal oxidation state of an element, whereas the superscripted Indo-Arabic numerals denote the net charge. The two notations are, therefore, exchangeable for monatomic ions, but the Roman numerals cannot be applied to polyatomic ions. However, it is possible to mix the notations for the individual metal centre with a polyatomic complex, as shown by the uranyl ion example.

Sub-classes

If an ion contains unpaired 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

Formation of monatomic ions

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 and molecular ions

An electrostatic potential map of the nitrate ion (
2NO-3). The 3-dimensional shell represents a single arbitrary isopotential. Nitrate-ion-elpot.png
An electrostatic potential map of the nitrate ion (2NO3). The 3-dimensional shell represents a single arbitrary isopotential.

Polyatomic and molecular ions are often formed by the gaining or losing of elemental ions such as a proton, 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.

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. [18] 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

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

Common cations [19]
Common nameFormulaHistoric 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
Hydrogen 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+2mercurous
Common anions [19]
Formal nameFormulaAlt. name
Monatomic anions
Azide N3
Bromide Br
Carbide C
Chloride Cl
Fluoride F
Hydride H
Iodide I
Nitride N3−
Phosphide P3−
Oxide O2−
Sulfide S2−
Selenide Se2−
Oxoanions (Polyatomic ions) [19]
Carbonate CO2−3
Chlorate ClO3
Chromate CrO2−4
Dichromate Cr2O2−7
Dihydrogen phosphateH2PO4
Hydrogen carbonate HCO3bicarbonate
Hydrogen sulfate HSO4bisulfate
Hydrogen sulfite HSO3bisulfite
Hydroxide OH
Hypochlorite ClO
Monohydrogen phosphateHPO2−4
Nitrate NO3
Nitrite NO2
Perchlorate ClO4
Permanganate MnO4
Peroxide O2−2
Phosphate PO3−4
Sulfate SO2−4
Sulfite SO2−3
Superoxide O2
Thiosulfate S2O2−3
Silicate SiO4−4
Metasilicate SiO2−3
Aluminium silicate AlSiO4
Anions from organic acids
Acetate CH3COOethanoate
Formate HCOOmethanoate
Oxalate C2O2−4ethanedioate
Cyanide CN

See also

Related Research Articles

<span class="mw-page-title-main">Alkali metal</span> Group of highly reactive chemical elements

The alkali metals consist of the chemical elements lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr). Together with hydrogen they constitute group 1, which lies in the s-block of the periodic table. All alkali metals have their outermost electron in an s-orbital: this shared electron configuration results in their having very similar characteristic properties. Indeed, the alkali metals provide the best example of group trends in properties in the periodic table, with elements exhibiting well-characterised homologous behaviour. This family of elements is also known as the lithium family after its leading element.

<span class="mw-page-title-main">Chemistry</span> Scientific discipline

Chemistry is the scientific study of the properties and behavior of matter. It is a natural science that covers the elements that make up matter to the compounds made of atoms, molecules and ions: their composition, structure, properties, behavior and the changes they undergo during a reaction with other substances. Chemistry also addresses the nature of chemical bonds in chemical compounds.

<span class="mw-page-title-main">Chemical bond</span> Lasting attraction between atoms that enables the formation of chemical compounds

A chemical bond is a lasting attraction between atoms or ions that enables the formation of molecules and crystals. The bond may result from the electrostatic force between oppositely charged ions as in ionic bonds, or through the sharing of electrons as in covalent bonds. The strength of chemical bonds varies considerably; there are "strong bonds" or "primary bonds" such as covalent, ionic and metallic bonds, and "weak bonds" or "secondary bonds" such as dipole–dipole interactions, the London dispersion force and hydrogen bonding. Strong chemical bonding arises from the sharing or transfer of electrons between the participating atoms.

<span class="mw-page-title-main">Ionic bonding</span> Chemical bonding involving attraction between ions

Ionic bonding is a type of chemical bonding that involves the electrostatic attraction between oppositely charged ions, or between two atoms with sharply different electronegativities, and is the primary interaction occurring in ionic compounds. It is one of the main types of bonding, along with covalent bonding and metallic bonding. Ions are atoms with an electrostatic charge. Atoms that gain electrons make negatively charged ions. Atoms that lose electrons make positively charged ions. This transfer of electrons is known as electrovalence in contrast to covalence. In the simplest case, the cation is a metal atom and the anion is a nonmetal atom, but these ions can be of a more complex nature, e.g. molecular ions like NH+
4
or SO2−
4
. In simpler words, an ionic bond results from the transfer of electrons from a metal to a non-metal in order to obtain a full valence shell for both atoms.

In chemistry, a salt is a chemical compound consisting of an ionic assembly of positively charged cations and negatively charged anions, which results in a compound with no net electric charge. A common example is table salt, with positively charged sodium ions and negatively charged chloride ions.

An electrolyte is a medium containing ions that is electrically conducting through the movement of those ions, but not conducting electrons. This includes most soluble salts, acids, and bases dissolved in a polar solvent, such as water. Upon dissolving, the substance separates into cations and anions, which disperse uniformly throughout the solvent. Solid-state electrolytes also exist. In medicine and sometimes in chemistry, the term electrolyte refers to the substance that is dissolved.

<span class="mw-page-title-main">Ionic compound</span> Chemical compound involving ionic bonding

In chemistry, an ionic compound is a chemical compound composed of ions held together by electrostatic forces termed ionic bonding. The compound is neutral overall, but consists of positively charged ions called cations and negatively charged ions called anions. These can be simple ions such as the sodium (Na+) and chloride (Cl) in sodium chloride, or polyatomic species such as the ammonium (NH+
4
) and carbonate (CO2−
3
) ions in ammonium carbonate. Individual ions within an ionic compound usually have multiple nearest neighbours, so are not considered to be part of molecules, but instead part of a continuous three-dimensional network. Ionic compounds usually form crystalline structures when solid.

<span class="mw-page-title-main">Octet rule</span> Chemical rule of thumb

The octet rule is a chemical rule of thumb that reflects the theory that main-group elements tend to bond in such a way that each atom has eight electrons in its valence shell, giving it the same electronic configuration as a noble gas. The rule is especially applicable to carbon, nitrogen, oxygen, and the halogens; although more generally the rule is applicable for the s-block and p-block of the periodic table. Other rules exist for other elements, such as the duplet rule for hydrogen and helium, or the 18-electron rule for transition metals.

<span class="mw-page-title-main">Electrolytic cell</span> Cell that uses electrical energy to drive a non-spontaneous redox reaction

An electrolytic cell is an electrochemical cell that utilizes an external source of electrical energy to force a chemical reaction that would otherwise not occur. The external energy source is a voltage applied between the cell′s two electrodes; an anode and a cathode, which are immersed in an electrolyte solution. This is in contrast to a galvanic cell, which itself is a source of electrical energy and the foundation of a battery. The net reaction taking place in a galvanic cell is a spontaneous reaction, i.e, the Gibbs free energy remains -ve, while the net reaction taking place in an electrolytic cell is the reverse of this spontaneous reaction, i.e, the Gibbs free energy is +ve.

In physics, a charge carrier is a particle or quasiparticle that is free to move, carrying an electric charge, especially the particles that carry electric charges in electrical conductors. Examples are electrons, ions and holes. The term is used most commonly in solid state physics. In a conducting medium, an electric field can exert force on these free particles, causing a net motion of the particles through the medium; this is what constitutes an electric current. In conducting media, particles serve to carry charge:

A hydrogen ion is created when a hydrogen atom loses or gains an electron. A positively charged hydrogen ion (or proton) can readily combine with other particles and therefore is only seen isolated when it is in a gaseous state or a nearly particle-free space. Due to its extremely high charge density of approximately 2×1010 times that of a sodium ion, the bare hydrogen ion cannot exist freely in solution as it readily hydrates, i.e., bonds quickly. The hydrogen ion is recommended by IUPAC as a general term for all ions of hydrogen and its isotopes. Depending on the charge of the ion, two different classes can be distinguished: positively charged ions and negatively charged ions.

<span class="mw-page-title-main">F-center</span> Type of crystallographic defect

An F center or Farbe center is a type of crystallographic defect in which an anionic vacancy in a crystal lattice is occupied by one or more unpaired electrons. Electrons in such a vacancy in a crystal lattice tend to absorb light in the visible spectrum such that a material that is usually transparent becomes colored. The greater the number of F centers, the more intense the color of the compound. F centers are a type of color center.

<span class="mw-page-title-main">Valence electron</span> An outer shell electron which is associated with an atom

In chemistry and physics, a valence electron is an electron in the outer shell associated with an atom, and that can participate in the formation of a chemical bond if the outer shell is not closed. In a single covalent bond, a shared pair forms with both atoms in the bond each contributing one valence electron.

<span class="mw-page-title-main">Fajans' rules</span> Explains the covalent character in molecules

In inorganic chemistry, Fajans' rules, formulated by Kazimierz Fajans in 1923, are used to predict whether a chemical bond will be covalent or ionic, and depend on the charge on the cation and the relative sizes of the cation and anion. They can be summarized in the following table:

In chemical nomenclature, the IUPAC nomenclature of inorganic chemistry is a systematic method of naming inorganic chemical compounds, as recommended by the International Union of Pure and Applied Chemistry (IUPAC). It is published in Nomenclature of Inorganic Chemistry. Ideally, every inorganic compound should have a name from which an unambiguous formula can be determined. There is also an IUPAC nomenclature of organic chemistry.

An intramolecular force is any force that binds together the atoms making up a molecule or compound, not to be confused with intermolecular forces, which are the forces present between molecules. The subtle difference in the name comes from the Latin roots of English with inter meaning between or among and intra meaning inside. Chemical bonds are considered to be intramolecular forces which are often stronger than intermolecular forces present between non-bonding atoms or molecules.

An alkalide is a chemical compound in which alkali metal atoms are anions with a charge or oxidation state of −1. Until the first discovery of alkalides in the 1970s, alkali metals were known to appear in salts only as cations with a charge or oxidation state of +1. These types of compounds are of theoretical interest due to their unusual stoichiometry and low ionization potentials. Alkalide compounds are chemically related to the electrides, salts in which trapped electrons are effectively the anions.

This glossary of chemistry terms is a list of terms and definitions relevant to chemistry, including chemical laws, diagrams and formulae, laboratory tools, glassware, and equipment. Chemistry is a physical science concerned with the composition, structure, and properties of matter, as well as the changes it undergoes during chemical reactions; it features an extensive vocabulary and a significant amount of jargon.

<span class="mw-page-title-main">Chemical compound</span> Substance composed of multiple elements that are chemically bonded

A chemical compound is a chemical substance composed of many identical molecules containing atoms from more than one chemical element held together by chemical bonds. A molecule consisting of atoms of only one element is therefore not a compound. A compound can be transformed into a different substance by a chemical reaction, which may involve interactions with other substances. In this process, bonds between atoms may be broken and/or new bonds formed.

A metal ion in aqueous solution or aqua ion is a cation, dissolved in water, of chemical formula [M(H2O)n]z+. The solvation number, n, determined by a variety of experimental methods is 4 for Li+ and Be2+ and 6 for most elements in periods 3 and 4 of the periodic table. Lanthanide and actinide aqua ions have higher solvation numbers (often 8 to 9), with the highest known being 11 for Ac3+. The strength of the bonds between the metal ion and water molecules in the primary solvation shell increases with the electrical charge, z, on the metal ion and decreases as its ionic radius, r, increases. Aqua ions are subject to hydrolysis. The logarithm of the first hydrolysis constant is proportional to z2/r for most aqua ions.

References

  1. "Ion" Archived 2013-12-24 at the Wayback Machine entry in Collins English Dictionary .
  2. "Definition of CATION". www.merriam-webster.com. Archived from the original on 2021-10-06. Retrieved 2021-10-06.
  3. "Definition of ANION". www.merriam-webster.com. Archived from the original on 2021-10-06. Retrieved 2021-10-06.
  4. 1 2 3 Knoll, Glenn F (1999). Radiation detection and measurement (3rd ed.). New York: Wiley. ISBN   978-0-471-07338-3.
  5. 1 2 Frank A. J. L. James, ed. (1991). The Correspondence of Michael Faraday, Vol. 2: 1832-1840. p. 183. ISBN   9780863412493. Archived from the original on 2021-04-14. Retrieved 2020-10-16.
  6. Michael Faraday (1791-1867). UK: BBC.
  7. "Online etymology dictionary". Archived from the original on 2011-05-14. Retrieved 2011-01-07.
  8. "The Nobel Prize in Chemistry 1903". www.nobelprize.org. Archived from the original on 2018-07-08. Retrieved 2017-06-13.
  9. Harris, William; Levey, Judith, eds. (1976). The New Columbia Encyclopedia (4th ed.). New York City: Columbia University. p.  155. ISBN   978-0-231-03572-9.
  10. Goetz, Philip W. (1992). McHenry, Charles (ed.). The New Encyclopædia Britannica. Chicago: Encyclopaedia Britannica Inc. Vol. 1 (15 ed.). Chicago: Encyclopædia Britannica, Inc. p. 587. Bibcode:1991neb..book.....G. ISBN   978-0-85229-553-3.
  11. Cillispie, Charles, ed. (1970). Dictionary of Scientific Biography (1 ed.). New York City: Charles Scribner's Sons. pp. 296–302. ISBN   978-0-684-10112-5.
  12. Oxford University Press (2013). "Oxford Reference: OVERVIEW anion". oxfordreference.com. Archived from the original on 2017-01-18. Retrieved 2017-01-15.
  13. University of Colorado Boulder (November 21, 2013). "Atoms and Elements, Isotopes and Ions". colorado.edu. Archived from the original on February 2, 2015. Retrieved November 22, 2013.
  14. Oxford University Press (2013). "Oxford Reference: OVERVIEW cation". oxfordreference.com. Archived from the original on 2017-01-18. Retrieved 2017-01-15.
  15. Douglas W. Haywick, Ph.D.; University of South Alabama (2007–2008). "Elemental Chemistry" (PDF). usouthal.edu. Archived (PDF) from the original on 2011-12-04. Retrieved 2013-11-22.
  16. Purdue University (November 21, 2013). "Amino Acids". purdue.edu. Archived from the original on July 13, 2011. Retrieved November 22, 2013.
  17. Press, Frank; Siever, Raymond (1986). Earth (14th ed.). New York: W. H. Freeman and Company. p. 63. ISBN   0-7167-1743-3. OCLC   12556840.
  18. Chemical elements listed by ionization energy Archived 2009-03-30 at the Wayback Machine . Lenntech.com
  19. 1 2 3 "Common Ions and Their Charges" (PDF). Science Geek. Archived (PDF) from the original on 2018-02-18. Retrieved 2018-05-11.