In chemistry, the valence or valency of an element is a measure of its combining power with other atoms when it forms chemical compounds or molecules. The concept of valence developed in the second half of the 19th century and helped successfully explain the molecular structure of inorganic and organic compounds. The quest for the underlying causes of valence led to the modern theories of chemical bonding, including the cubical atom (1902), Lewis structures (1916), valence bond theory (1927), molecular orbitals (1928), valence shell electron pair repulsion theory (1958), and all of the advanced methods of quantum chemistry.
The combining power, or affinity of an atom of a given element is determined by the number of hydrogen atoms that it combines with. In methane, carbon has a valence of 4; in ammonia, nitrogen has a valence of 3; in water, oxygen has a valence of 2; and in hydrogen chloride, chlorine has a valence of 1. Chlorine, as it has a valence of one, can be substituted for hydrogen, so phosphorus has a valence of 5 in phosphorus pentachloride, PCl5. Valence diagrams of a compound represent the connectivity of the elements, with lines drawn between two elements, sometimes called bonds, representing a saturated valency for each element. The two tables below show some examples of different compounds, their valence diagrams, and the valences for each element of the compound.
Valence only describes connectivity; it does not describe the geometry of molecular compounds, or what are now known to be ionic compounds or giant covalent structures. A line between atoms does not represent a pair of electrons as it does in Lewis diagrams.
An alternative modern description is:
This definition differs from the IUPAC definition as an element can be said to have more than one valence.
The etymology of the words valence (plural valences) and valency (plural valencies) traces back to 1425, meaning "extract, preparation", from Latin valentia "strength, capacity", from the earlier valor "worth, value", 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 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 theories of free radicals 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., NO3, NH3, NI3, etc.) or 5, i.e., in the 5-atom groups (e.g., NO5, NH4O, PO5, 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.
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 [Co(NH3)6]Cl3, 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.
The Rutherford model of the nuclear atom (1911) showed that the exterior of an atom is occupied by electrons, which suggests that electrons are responsible for the interaction of atoms and the formation of chemical bonds. In 1916, Gilbert N. Lewis explained valence and chemical bonding in terms of a tendency of (main-group) atoms to achieve a stable octet of 8 valence-shell electrons. According to Lewis, covalent bonding leads to octets by the sharing of electrons, and ionic bonding leads to octets by the transfer of electrons from one atom to the other. The term covalence is attributed to Irving Langmuir, who stated in 1919 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- means "together", so that a co-valent bond means that the atoms share a valence. Subsequent to that, it is now more common to speak of covalent bonds rather than valence, which has fallen out of use in higher-level work from the advances in the theory of chemical bonding, but it is still widely used in elementary studies, where it provides a heuristic introduction to the subject.
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 (SF6), 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 SF6 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.
|Group||Valence 1||Valence 2||Valence 3||Valence 4||Valence 5||Valence 6||Valence 7||Typical valencies|
|13 (III)||BCl3, AlCl3
|3 and 5|
|SO2||SO3||2 and 6|
|17 (VII)||HCl||HClO2||ClO2||HClO3||Cl2O7||1 and 7|
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 prefixes (mono-/uni-, di-/bi-, tri-/ter-, and so on) are used to describe ions in the charge states 1, 2, 3, and so on, respectively. Polyvalence or multivalence refers to species that are not restricted to a specific number of valence bonds. Species with a single charge are univalent (monovalent). For example, the Cs+ cation is a univalent or monovalent cation, whereas the Ca2+ cation is a divalent cation, and the Fe3+ 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 multivalent (polyvalent) ion. Transition metals and metals to the right are typically multivalent but there is no simple pattern predicting their valency.
|Valence||More common adjective‡||Less common synonymous adjective‡§|
|5-valent||pentavalent||quinquevalent / quinquivalent|
|multiple / many / variable||polyvalent||multivalent|
† The same adjectives are also used in medicine to refer to vaccine valence, with the slight difference that in the latter sense, quadri- is more common than tetra-.
‡ 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.
Because of the ambiguity of the term valence, nowadays other notations are used in practice. Beside the system of oxidation numbers as used in Stock nomenclature for coordination compounds, and the lambda notation, as used in the IUPAC nomenclature of inorganic chemistry, oxidation state is a more clear indication of the electronic state of atoms in a molecule.
The oxidation state of an atom in a molecule gives the number of valence electrons 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 can have a valence higher than four. For example, in perchlorates, chlorine has seven valence bonds and ruthenium, in the +8 oxidation state in ruthenium tetroxide, has eight valence bonds.
|Hydrogen chloride||HCl||H = 1 Cl = 1||H = +1 Cl = −1|
|Perchloric acid *||HClO4||H = 1 Cl = 7 O = 2||H = +1 Cl = +7 O = −2|
|Sodium hydride||NaH||Na = 1 H = 1||Na = +1 H = −1|
|Ferrous oxide **||FeO||Fe = 2 O = 2||Fe = +2 O = −2|
|Ferric oxide **||Fe2O3||Fe = 3 O = 2||Fe = + 3 O = −2|
* The univalent perchlorate ion (ClO4−) has valence 1.
** Iron oxide appears in a crystal structure, so no typical molecule can be identified.
In ferrous oxide, Fe has oxidation number II, in ferric oxide, oxidation number III.
|Chlorine||Cl2||Cl = 1||Cl = 0|
|Hydrogen peroxide||H2O2||H = 1 O = 2||H = +1 O = −1|
|Acetylene||C2H2||C = 4 H = 1||C = −1 H = +1|
|Mercury(I) chloride||Hg2Cl2||Hg = 2 Cl = 1||Hg = +1 Cl = −1|
Valences may also be different from absolute values of oxidation states due to different polarity of bonds. For example, in dichloromethane, CH2Cl2, carbon has valence 4 but oxidation state 0.
Frankland took the view that the valence (he used the term "atomicity") of an element was a single value that corresponded to the maximum value observed. The number of unused valencies on atoms of what are now called the p-block elements is generally even, and Frankland suggested that the unused valencies saturated one another. For example, nitrogen has a maximum valence of 5, in forming ammonia two valencies are left unattached; sulfur has a maximum valence of 6, in forming hydrogen sulphide four valencies are left unattached.
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 outer 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.
Maximum valences for the elements are based on the data from list of oxidation states of the elements.
Maximum valences of the elements
|Maximum valences are based on the List of oxidation states of the elements|
The Creutz–Taube ion is the metal complex with the formula [Ru(NH3)5]2(C4H4N2)5+. This cationic species has been heavily studied in an effort to understand the intimate details of inner sphere electron transfer, that is, how electrons move from one metal complex to another. The ion is named after Carol Creutz, who first prepared the complex, and her thesis advisor Henry Taube, who received a Nobel Prize in Chemistry for this and related discoveries on electron-transfer.Formal charge
In chemistry, a formal charge (FC) is the charge assigned to an atom in a molecule, assuming that electrons in all chemical bonds are shared equally between atoms, regardless of relative electronegativity. When determining the best Lewis structure (or predominant resonance structure) for a molecule, the structure is chosen such that the formal charge on each of the atoms is as close to zero as possible.
The formal charge of any atom in a molecule can be calculated by the following equation:
where V is the number of valence electrons of the neutral atom in isolation (in its ground state); N is the number of non-bonding valence electrons on this atom in the molecule; and B is the total number of electrons shared in bonds with other atoms in the molecule.Glossary of biology
Most of the terms listed in Wikipedia glossaries are already defined and explained within Wikipedia itself. However, glossaries like this one are useful for looking up, comparing and reviewing large numbers of terms together. You can help enhance this page by adding new terms or writing definitions for existing ones.
This glossary of biology terms is a list of definitions of fundamental terms and concepts of biology, its sub-disciplines, and related fields. For more specific definitions from other glossaries related to biology, see Glossary of ecology, Glossary of botany, Glossary of genetics, and Glossary of speciation.Inner sphere electron transfer
Inner sphere or bonded electron transfer is a redox chemical reaction that proceeds via a covalent linkage—a strong electronic interaction—between the oxidant and the reductant reactants. In Inner Sphere (IS) electron transfer (ET), a ligand bridges the two metal redox centers during the electron transfer event. Inner sphere reactions are inhibited by large ligands, which prevent the formation of the crucial bridged intermediate. Thus, IS ET is rare in biological systems, where redox sites are often shielded by bulky proteins. Inner sphere ET is usually used to describe reactions involving transition metal complexes and most of this article is written from this perspective. However, redox centers can consist of organic groups rather than metal centers.
The bridging ligand could be virtually any entity that can convey electrons. Typically, such a ligand has more than one lone electron pair, such that it can serve as an electron donor to both the reductant and the oxidant. Common bridging ligands include the halides and the pseudohalides such as hydroxide and thiocyanate. More complex bridging ligands are also well known including oxalate, malonate, and pyrazine. Prior to ET, the bridged complex must form, and such processes are often highly reversible. Electron transfer occurs through the bridge once it is established. In some cases, the stable bridged structure may exist in the ground state; in other cases, the bridged structure may be a transiently-formed intermediate, or else as a transition state during the reaction.
The alternative to inner sphere electron transfer is outer sphere electron transfer. In any transition metal redox process, the mechanism can be assumed to be outer sphere unless the conditions of the inner sphere are met. Inner sphere electron transfer is generally enthalpically more favorable than outer sphere electron transfer due to a larger degree of interaction between the metal centers involved, however, inner sphere electron transfer is usually entropically less favorable since the two sites involved must become more ordered (come together via a bridge) than in outer sphere electron transfer.List of English-language metaphors
A list of metaphors in the English language organised by type. A metaphor is a literary figure of speech that uses an image, story or tangible thing to represent a less tangible thing or some intangible quality or idea; e.g., "Her eyes were glistening jewels". Metaphor may also be used for any rhetorical figures of speech that achieve their effects via association, comparison or resemblance. In this broader sense, antithesis, hyperbole, metonymy and simile would all be considered types of metaphor. Aristotle used both this sense and the regular, current sense above.
With metaphor, unlike analogy, specific interpretations are not given explicitly.Noel Hush
Professor Noel Sydney Hush AO, DSc, FRS, FNAS, FAA, FRACI, FRSN is an Australian chemist at the University of Sydney.Peter Day (chemist)
Peter Day (born 20 August 1938 in Kent, England) is a British inorganic chemist and Emeritus Professor of Chemistry at University College London (UCL).Polyvalence
In chemistry, polyvalence or multivalence refers to atoms or molecules that exhibit more than one valence (chemistry).
Polyvalence or polyvalent may also refer to:
Polyvalence (music), the musical use of more than one key simultaneously
Polyvalent antibody, a group of antibodies that have affinity for various antigens
Polyvalent logic, a form of many-valued logic or probabilistic logic
Polyvalent vaccine, a vaccine that can vaccinate a person against more than one strain of a disease
Sala Polivalentă (disambiguation), various stadiums in Romania commonly translated as Polyvalent Hall
Snake antivenom that contains neutralizing antibodies against two or more species of snakesRichard Abegg
Richard Wilhelm Heinrich Abegg (January 9, 1869 – April 3, 1910) was a German chemist and pioneer of valence theory. He proposed that the difference of the maximum positive and negative valence of an element tends to be eight. This has come to be known as Abegg's rule. He was a gas balloon enthusiast, which caused his death at the age of 41 when he crashed in his balloon in Silesia.
Abegg received his PhD on July 19, 1891 as the student of August Wilhelm von Hofmann at the University of Berlin. Abegg learned organic chemistry from Hofmann, but one year after finishing his PhD degree he began researching physical chemistry while studying with Friedrich Wilhelm Ostwald in Leipzig, Germany. Abegg later served as private assistant to Walther Nernst at the University of Göttingen and to Svante Arrhenius at the University of Stockholm.
Abegg discovered the theory of freezing-point depression and anticipated Gilbert Newton Lewis's octet rule by revealing that the lowest and highest oxidation states of elements often differ by eight. He researched many topics in physical chemistry, including freezing points, the dielectric constant of ice, osmotic pressures, oxidation potentials, and complex ions.Univalent
Univalent may refer to:
Univalent function – a concept in mathematics;
Univalent foundations – a concept in mathematics;
Univalent relation R satisfies xRy ∧ xRz ⇒ y = z.
Valent may refer to:
Valency (linguistics)Vegan nutrition
Vegan nutrition refers to the nutritional and human health aspects of vegan diets. A well-planned, balanced vegan diet is suitable to meet all recommendations for nutrients in every stage of human life. Improperly planned vegan diets may be deficient in vitamin B12, vitamin D, calcium, iodine, iron, zinc and riboflavin (vitamin B2). Preliminary evidence from clinical research indicates that a vegan diet may lower the risk of cancer.Wolfgang Kaim
Wolfgang Kaim (born 13 May 1951 in Bad Vilbel, Germany) is a German chemist who is the chair of coordination chemistry at the University of Stuttgart. He is co-author of the internationally recognized book, Bioinorganic Chemistry which was awarded with the Literature Award of the German Chemical Industry.