In chemistry, valence, also known as valency or valency number, is a measure of the number of chemical bonds formed by the atoms of a given element. Over the last century, the concept of valence evolved into a range of approaches for describing the chemical bond, including Lewis structures (1916), valence bond theory (1927), molecular orbitals (1928), valence shell electron pair repulsion theory (1958) and all the advanced methods of quantum chemistry.
The etymology of the word "valence" is from 1425, meaning "extract, preparation," from Latin valentia "strength, capacity," and the chemical meaning referring to the "combining power of an element" is recorded from 1884, from German Valenz.
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 similarly for the other combinations of ultimate particles:
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 theories of free radicals and “type 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 three atom groups (e.g. NO3, NH3, NI3, etc.) or 5, i.e. in the five atom groups (e.g. NO5, NH4O, PO5, etc.), equivalents of the attached elements. It is in this manner, according to Franklin, that their affinities are best satisfied. Following these examples and postulates, Franklin declares how obvious it is that:
|“||A tendency or law prevails (here), and that, no matter what the characters of the uniting atoms may be, the combining power of the attracting element, if I may be allowed the term, is always satisfied by the same number of these atoms.||”|
This “combining power” was afterwards called quantivalence or valency (and valence by American chemists).
The concept was developed in the middle of the nineteenth century in an attempt to rationalize the formulae of different chemical compounds. In 1919, Irving Langmuir, borrowed the term to explain Gilbert N. Lewis's cubical atom model by stating that "the number of pairs of electrons which any given atom shares with the adjacent atoms is called the covalence of that atom." The prefix co-, e.g. co-author, means together, jointly, associated in action, partnered to a lesser degree, etc.,; thus a co-valent bond, essentially, means that the atoms share valence. Hence, if an atom, for example, had a +1 valence, meaning it was missing an electron, and another a -1 valence, meaning it had an extra electron, then a bond between these two atoms would result because they would be complementing or sharing their out of balance valence tendencies. Subsequently, it is now more common to speak of covalent bonds rather than "valence", which has fallen out of use in higher level work with the advances in the theory of chemical bonding, but is still widely used in elementary studies where it provides a heuristic introduction to the subject.
"Number of bonds" definition
The number of bonds formed by a given element was originally thought to be a fixed chemical property and in fact, in many cases, this is a good approximation. For example, in many of their compounds, carbon forms four bonds, oxygen two and hydrogen one. However it soon became apparent that, for many elements, the valence could vary between different compounds. One of the first examples to be identified was phosphorus, which sometimes behaves as if it has a valence of three and sometimes as if it has a valence of five. One method around this problem is to specify the valence for each individual compound: although it removes much of the generality of the concept, this approach has given rise to the idea of oxidation numbers (used in Stock nomenclature) and to lambda notation in the IUPAC nomenclature of inorganic chemistry.
- The maximum number of univalent atoms (originally hydrogen or chlorine atoms) that may combine with an atom of the element under consideration, or with a fragment, or for which an atom of this element can be substituted.
This definition reimposes a unique valence for each element at the expense of neglecting, in many cases, a large part of its chemistry.
The mention of hydrogen and chlorine is for historic reasons, although both in practice mostly form compounds in which their atoms form a single bond. Exceptions in the case of hydrogen include the ion [HF2]− and the various boron hydrides such as diborane: these are examples of three-center two-electron bonds. Chlorine forms a number of fluorides—ClF, ClF3 and ClF5—and its valence according to the IUPAC definition is hence five. Fluorine is the element for which the largest number of atoms combine with atoms of other elements: it is univalent in all compounds except the ion [H2F]+. In fact, the IUPAC definition can only be resolved by fixing the valences of hydrogen and fluorine as one, a convention which has been followed here.
Valences of the elements
Valences for the majority of elements are based on the highest known fluoride.
Other criticisms of the concept of valence
- The valence of an element is not always equal to its highest oxidation state: exceptions include ruthenium, osmium and xenon, which have valences of six (hexafluorides) but which form compounds with oxygen in the +8 oxidation state, and chlorine, which has a valence of five but a highest oxidation state of +7 (in perchlorates).
- The concept of "combination" cannot be equated with the number of bonds formed by an atom. In lithium fluoride (which has the NaCl structure), each lithium atom is surrounded by six fluorine atoms, whereas the valence of lithium is universally taken to be one, as the formula LiF would suggest.
- Valence - Online Etymology Dictionary.
- Partington, J.R. (1989). A Short History of Chemistry. Dover Publications, Inc. ISBN 0-486-65977-1.
- Franklin, E. (1852). Phil. Trans., vol. cxlii, 417.
- Pure Appl. Chem. 66: 1175 (1994).
- http://www.webelements.com/ (accessed 2006-02-20).
- In the gas phase, LiF does indeed exist as discrete diatomic molecules as the valences would suggest: Template:Cotton&Wilkinson6th
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