Electron in the outer shell of an atom's energy levels
Fourcovalent bonds. Carbon has four valence electrons and here avalence of four. Each hydrogen atom has one valence electron and is univalent.
Inchemistry andphysics,valence electrons areelectrons in the outermostshell of anatom, and that can participate in the formation of achemical bond if the outermost shell is notclosed. In a singlecovalent bond, a shared pair forms with both atoms in the bond each contributing one valence electron.
The presence of valence electrons can determine theelement'schemical properties, such as itsvalence—whether it may bond with other elements and, if so, how readily and with how many. In this way, a given element'sreactivity is highly dependent upon itselectronic configuration. For amain-group element, a valence electron can exist only in the outermostelectron shell; for atransition metal, a valence electron can also be in an inner shell.
An atom with a closed shell of valence electrons (corresponding to anoble gas configuration) tends to bechemically inert. Atoms with one or two valence electrons more than a closed shell are highly reactive due to the relativelylow energy to remove the extra valence electrons to form a positiveion. An atom with one or two electrons fewer than a closed shell is reactive due to its tendency either to gain the missing valence electrons and form a negative ion, or else to share valence electrons and form a covalent bond.
Similar to acore electron, a valence electron has the ability to absorb or release energy in the form of aphoton. An energy gain can trigger the electron to move (jump) to an outer shell; this is known asatomic excitation. Or the electron can even break free from its associated atom's shell; this isionization to form a positive ion. When an electron loses energy (thereby causing a photon to be emitted), then it can move to an inner shell which is not fully occupied.
The electrons that determinevalence – how an atom reacts chemically – are those with the highestenergy.
For amain-group element, the valence electrons are defined as those electrons residing in the electronic shell of highestprincipal quantum numbern.[1] Thus, the number of valence electrons that it may have depends on theelectron configuration in a simple way. For example, the electronic configuration ofphosphorus (P) is 1s2 2s2 2p6 3s2 3p3 so that there are 5 valence electrons (3s2 3p3), corresponding to a maximum valence for P of 5 as in themolecule PF5; this configuration is normally abbreviated to [Ne] 3s2 3p3, where [Ne] signifies the core electrons whose configuration is identical to that of thenoble gasneon.
However,transition elements have (n−1)d energy levels that are very close in energy to thens level.[2] So as opposed to main-group elements, a valence electron for a transition metal is defined as an electron that resides outside a noble-gas core.[3] Thus, generally, the d electrons in transition metals behave as valence electrons although they are not in the outermost shell. For example,manganese (Mn) has configuration 1s2 2s2 2p6 3s2 3p6 4s2 3d5; this is abbreviated to [Ar] 4s2 3d5, where [Ar] denotes a core configuration identical to that of the noble gasargon. In this atom, a 3d electron has energy similar to that of a 4s electron, and much higher than that of a 3s or 3p electron. In effect, there are possibly seven valence electrons (4s2 3d5) outside the argon-like core; this is consistent with the chemical fact that manganese can have anoxidation state as high as +7 (in thepermanganate ion:MnO− 4). (But note that merely having that number of valence electrons does not imply that the corresponding oxidation state will exist. For example,fluorine is not known in oxidation state +7; and although the maximum known number of valence electrons is 16 inytterbium andnobelium, no oxidation state higher than +9 is known for any element.)
The farther right in each transition metal series, the lower the energy of an electron in a d subshell and the less such an electron has valence properties. Thus, although anickel atom has, in principle, ten valence electrons (4s2 3d8), its oxidation state never exceeds four. Forzinc, the 3d subshell is complete in all known compounds, although it does contribute to the valence band in some compounds.[4] Similar patterns hold for the (n−2)f energy levels of inner transition metals.
Thed electron count is an alternative tool for understanding the chemistry of a transition metal.
The number of valence electrons of an element can be determined by theperiodic table group (vertical column) in which the element is categorized. In groups 1–12, the group number matches the number of valence electrons; in groups 13–18, the units digit of the group number matches the number of valence electrons. (Helium is the sole exception.)[5]
Helium is an exception: despite having a 1s2 configuration with two valence electrons, and thus having some similarities with the alkaline earth metals with theirns2 valence configurations, its shell is completely full and hence it is chemically very inert and is usually placed in group 18 with the other noble gases.
The valence shell is the set oforbitals which are energetically accessible for accepting electrons to formchemical bonds.
For main-group elements, the valence shell consists of thens andnp orbitals in the outermostelectron shell. Fortransition metals the orbitals of the incomplete (n−1)d subshell are included, and forlanthanides andactinides incomplete (n−2)f and (n−1)d subshells. The orbitals involved can be in an inner electron shell and do not all correspond to the same electron shell or principal quantum numbern in a given element, but they are all at similar energies.[5]
As a general rule, amain-group element (except hydrogen or helium) tends to react to form a s2p6electron configuration. This tendency is called theoctet rule, because each bonded atom has 8 valence electrons including shared electrons. Similarly, a transition metal tends to react to form a d10s2p6 electron configuration. This tendency is called the18-electron rule, because each bonded atom has 18 valence electrons including shared electrons.
The heavy group 2 elements calcium, strontium, and barium can use the (n−1)d subshell as well, giving them some similarities to transition metals.[7][8][9]
The number of valence electrons in an atom governs itsbonding behavior. Therefore, elements whose atoms have the same number of valence electrons are often grouped together in theperiodic table of the elements, especially if they also have the same types of valence orbitals.[10]
The mostreactive kind ofmetallic element is analkali metal of group 1 (e.g.,sodium orpotassium); this is because such an atom has only a single valence electron. During the formation of anionic bond, which provides the necessaryionization energy, this one valence electron is easily lost to form a positiveion (cation) with a closed shell (e.g., Na+ or K+). Analkaline earth metal of group 2 (e.g.,magnesium) is somewhat less reactive, because each atom must lose two valence electrons to form a positive ion with a closed shell (e.g., Mg2+).[citation needed]
Within each group (each periodic table column) of metals, reactivity increases with each lower row of the table (from a light element to a heavier element), because a heavier element has more electron shells than a lighter element; a heavier element's valence electrons exist at higherprincipal quantum numbers (they are farther away from the nucleus of the atom, and are thus at higher potential energies, which means they are less tightly bound).[citation needed]
Anonmetal atom tends to attract additional valence electrons to attain a full valence shell; this can be achieved in one of two ways: An atom can either share electrons with a neighboring atom (acovalent bond), or it can remove electrons from another atom (anionic bond). The most reactive kind of nonmetal element is ahalogen (e.g.,fluorine (F) orchlorine (Cl)). Such an atom has the following electron configuration: s2p5; this requires only one additional valence electron to form a closed shell. To form an ionic bond, a halogen atom can remove an electron from another atom in order to form an anion (e.g., F−, Cl−, etc.). To form a covalent bond, one electron from the halogen and one electron from another atom form a shared pair (e.g., in the molecule H–F, the line represents a shared pair of valence electrons, one from H and one from F).[citation needed]
Within each group of nonmetals, reactivity decreases with each lower row of the table (from a light element to a heavy element) in the periodic table, because the valence electrons are at progressively higher energies and thus progressively less tightly bound. In fact, oxygen (the lightest element in group 16) is the most reactive nonmetal after fluorine, even though it is not a halogen, because the valence shells of the heavier halogens are at higher principal quantum numbers.
In these simple cases where the octet rule is obeyed, thevalence of an atom equals the number of electrons gained, lost, or shared in order to form the stable octet. However, there are also many molecules that areexceptions, and for which the valence is less clearly defined.
Valence electrons are also responsible for the bonding in the pure chemical elements, and whether theirelectrical conductivity is characteristic of metals, semiconductors, or insulators.
MetallicNetwork covalentMolecularcovalentSingle atomsUnknownBackground color shows bonding of simple substances in theperiodic table. If there are several, the most stable allotrope is considered.
Metallic elements generally have highelectrical conductivity when in thesolid state. In each row of theperiodic table, the metals occur to the left of the nonmetals, and thus a metal has fewer possible valence electrons than a nonmetal. However, a valence electron of a metal atom has a smallionization energy, and in the solid-state this valence electron is relatively free to leave one atom in order to associate with another nearby. This situation characterisesmetallic bonding. Such a "free" electron can be moved under the influence of anelectric field, and its motion constitutes anelectric current; it is responsible for the electrical conductivity of the metal.Copper,aluminium,silver, andgold are examples of good conductors.
Anonmetallic element has low electrical conductivity; it acts as aninsulator. Such an element is found toward the right of the periodic table, and it has a valence shell that is at least half full (the exception isboron). Its ionization energy is large; an electron cannot leave an atom easily when an electric field is applied, and thus such an element can conduct only very small electric currents. Examples of solid elemental insulators arediamond (anallotrope ofcarbon) andsulfur. These form covalently bonded structures, either with covalent bonds extending across the whole structure (as in diamond) or with individual covalent molecules weakly attracted to each other byintermolecular forces (as in sulfur). (Thenoble gases remain as single atoms, but those also experience intermolecular forces of attraction, that become stronger as the group is descended: helium boils at −269 °C, while radon boils at −61.7 °C.)
A solid compound containing metals can also be an insulator if the valence electrons of the metal atoms are used to formionic bonds. For example, although elementalsodium is a metal, solidsodium chloride is an insulator, because the valence electron of sodium is transferred to chlorine to form an ionic bond, and thus that electron cannot be moved easily.
Asemiconductor has an electrical conductivity that is intermediate between that of a metal and that of a nonmetal; a semiconductor also differs from a metal in that a semiconductor's conductivity increases withtemperature. The typical elemental semiconductors aresilicon andgermanium, each atom of which has four valence electrons. The properties of semiconductors are best explained usingband theory, as a consequence of a small energy gap between avalence band (which contains the valence electrons at absolute zero) and aconduction band (to which valence electrons are excited by thermal energy).
^Tossell, J. A. (1 November 1977). "Theoretical studies of valence orbital binding energies in solid zinc sulfide, zinc oxide, and zinc fluoride".Inorganic Chemistry.16 (11):2944–2949.doi:10.1021/ic50177a056.
^abKeeler, James; Wothers, Peter (2014).Chemical Structure and Reactivity (2nd ed.). Oxford University Press. pp. 257–260.ISBN978-0-19-9604135.
