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In chemistry, the valence (US spelling) or valency (British spelling) of an atom is a measure of its combining capacity with other atoms when it forms chemical compounds or . Valence is generally understood to be the number of that each atom of a given chemical element typically forms. are considered to be two bonds, to be three, to be four, to be five and to be six. In most compounds, the valence of hydrogen is 1, of oxygen is 2, of nitrogen is 3, and of carbon is 4. Valence is not to be confused with the related concepts of the coordination number, the oxidation state, or the number of for a given atom.
An alternative modern description is:
This definition differs from the IUPAC definition as an element can be said to have more than one valence.
In 1789, William Higgins published views on what he called combinations of "ultimate" particles, which foreshadowed the concept of valency bonds. If, for example, according to Higgins, the force between the ultimate particle of oxygen and the ultimate particle of nitrogen were 6, then the strength of the force would be divided accordingly, and likewise for the other combinations of ultimate particles (see illustration).
The exact inception, however, of the theory of chemical valencies can be traced to an 1852 paper by Edward Frankland, in which he combined the older radical theory with thoughts on chemical affinity to show that certain elements have the tendency to combine with other elements to form compounds containing 3, i.e., in the 3-atom groups (e.g., , , , etc.) or 5, i.e., in the 5-atom groups (e.g., , , , etc.), equivalents of the attached elements. According to him, this is the manner in which their affinities are best satisfied, and by following these examples and postulates, he declares how obvious it is that
This "combining power" was afterwards called quantivalence or valency (and valence by American chemists). In 1857 August Kekulé proposed fixed valences for many elements, such as 4 for carbon, and used them to propose structural formulas for many organic molecules, which are still accepted today.
Lothar Meyer in his 1864 book, Die modernen Theorien der Chemie, contained an early version of the periodic table containing 28 elements, for the first time classified elements into six families by their valence. Works on organizing the elements by atomic weight, until then had been stymied by the widespread use of equivalent weights for the elements, rather than atomic weights.
Most 19th-century chemists defined the valence of an element as the number of its bonds without distinguishing different types of valence or of bond. However, in 1893 Alfred Werner described transition metal coordination complexes such as , in which he distinguished principal and subsidiary valences (German: 'Hauptvalenz' and 'Nebenvalenz'), corresponding to the modern concepts of oxidation state and coordination number respectively.
For main-group elements, in 1904 Richard Abegg considered positive and negative valences (maximal and minimal oxidation states), and proposed Abegg's rule to the effect that their difference is often 8.
An alternative definition of valence, developed in the 1920's and having modern proponents, differs in cases where an atom's formal charge is not zero. It defines the valence of a given atom in a covalent molecule as the number of electrons that an atom has used in bonding:
or equivalently:
In this convention, the nitrogen in an ammonium ion bonds to four hydrogen atoms, but it is considered to be pentavalent because all five of nitrogen's valence electrons participate in the bonding.
In the 1930s, Linus Pauling proposed that there are also polar covalent bonds, which are intermediate between covalent and ionic, and that the degree of ionic character depends on the difference of electronegativity of the two bonded atoms.
Pauling also considered hypervalent molecules, in which main-group elements have apparent valences greater than the maximal of 4 allowed by the octet rule. For example, in the sulfur hexafluoride molecule (), Pauling considered that the sulfur forms 6 true two-electron bonds using sp3d2 hybrid atomic orbitals, which combine one s, three p and two d orbitals. However more recently, quantum-mechanical calculations on this and similar molecules have shown that the role of d orbitals in the bonding is minimal, and that the molecule should be described as having 6 polar covalent (partly ionic) bonds made from only four orbitals on sulfur (one s and three p) in accordance with the octet rule, together with six orbitals on the fluorines. Similar calculations on transition-metal molecules show that the role of p orbitals is minor, so that one s and five d orbitals on the metal are sufficient to describe the bonding. (2014). 9783527333158, Wiley – VCH. ISBN 9783527333158
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| chlorous acid | chloric acid | 1, 3, 5 and 7 |
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Many elements have a common valence related to their position in the periodic table, and nowadays this is rationalised by the octet rule. The Greek/Latin numeral prefix are used to describe ions in the charge states 1, 2, 3, and so on, respectively. Polyvalence or multivalence refers to Chemical species that are not restricted to a specific number of valence Chemical bond. Species with a single charge are univalent (monovalent). For example, the cation is a univalent or monovalent cation, whereas the cation is a divalent cation, and the cation is a trivalent cation. Unlike Cs and Ca, Fe can also exist in other charge states, notably 2+ and 4+, and is thus known as a Polyvalency (polyvalent) ion. Transition metals and metals to the right are typically multivalent but there is no simple pattern predicting their valency.
| + Valence adjectives using the suffix† |
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‡ As demonstrated by hit counts in Google web search and Google Books search corpora (accessed 2017).
§ A few other forms can be found in large English-language corpora (for example, *quintavalent, *quintivalent, *decivalent), but they are not the conventionally established forms in English and thus are not entered in major dictionaries.
The oxidation state of an atom in a molecule gives the number of it has gained or lost. In contrast to the valency number, the oxidation state can be positive (for an electropositive atom) or negative (for an electronegative atom).
Elements in a high oxidation state have an oxidation state higher than +4, and also, elements in a high valence state (hypervalent elements) have a valence higher than 4. For example, in , chlorine has 7 valence bonds (thus, it is heptavalent, in other words, it has valence 7), and it has oxidation state +7; in ruthenium tetroxide , ruthenium has 8 valence bonds (thus, it is octavalent, in other words, it has valence 8), and it has oxidation state +8.
In some molecules, there is a difference between valence and oxidation state for a given atom. For example, in disulfur decafluoride molecule , each sulfur atom has 6 valence bonds (5 single bonds with fluorine atoms and 1 single bond with the other sulfur atom). Thus, each sulfur atom is hexavalent or has valence 6, but has oxidation state +5. In the dioxygen molecule , each oxygen atom has 2 valence bonds and so is divalent (valence 2), but has oxidation state 0. In acetylene , each carbon atom has 4 valence bonds (1 single bond with hydrogen atom and a triple bond with the other carbon atom). Each carbon atom is tetravalent (valence 4), but has oxidation state −1.
| +Variation of valence vs oxidation state for bonds between two different elements |
* The perchlorate ion is monovalent, in other words, it has valence 1.
** Valences may also be different from absolute values of oxidation states due to different polarity of bonds. For example, in dichloromethane, , carbon has valence 4 but oxidation state 0.
*** Iron oxides appear in a crystal structure, so no typical molecule can be identified. In ferrous oxide, Fe has oxidation state +2; in ferric oxide, oxidation state +3.
| +Variation of valence vs oxidation state for bonds between two atoms of the same element |
The International Union of Pure and Applied Chemistry (IUPAC) has made several attempts to arrive at an unambiguous definition of valence. The current version, adopted in 1994:
Hydrogen and chlorine were originally used as examples of univalent atoms, because of their nature to form only one single bond. Hydrogen has only one valence electron and can form only one bond with an atom that has an incomplete Electron shell. Chlorine has seven valence electrons and can form only one bond with an atom that donates a valence electron to complete chlorine's outer shell. However, chlorine can also have oxidation states from +1 to +7 and can form more than one bond by donating valence electrons.
Hydrogen has only one valence electron, but it can form bonds with more than one atom. In the bifluoride ion (), for example, it forms a three-center four-electron bond with two fluoride atoms:
Another example is the three-center two-electron bond in diborane ().
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