View source for Physics:Valence electron

{{Short description|An electron in the outer shell of an atom's energy levels}}
[[File:Covalent.svg|thumb|180px|Four covalent bonds. Carbon has four valence electrons and here a valence of four. Each hydrogen atom has one valence electron and is univalent.]]

In [[Chemistry:Chemistry|chemistry]] and [[Physics:Physics|physics]], '''valence electrons''' are electrons in the outermost [[Physics:Electron shell|shell]] of an [[Physics:Atom|atom]], and that can participate in the formation of a [[Chemistry:Chemical bond|chemical bond]] if the outermost shell is not closed. In a single [[Chemistry:Covalent bond|covalent bond]], a shared pair forms with both atoms in the bond each contributing one valence electron.

The presence of valence electrons can determine the [[Chemistry:Chemical element|element]]'s [[Chemistry:Chemistry|chemical]] properties, such as its [[Chemistry:Valence|valence]]—whether it may bond with other elements and, if so, how readily and with how many. In this way, a given element's [[Chemistry:Reactivity|reactivity]] is highly dependent upon its [[Physics:Electron configuration|electronic configuration]]. For a [[Chemistry:Main-group element|main-group element]], a valence electron can exist only in the outermost [[Physics:Electron shell|electron shell]]; for a [[Chemistry:Transition metal|transition metal]], a valence electron can also be in an inner shell.

An atom with a closed shell of valence electrons (corresponding to a [[Physics:Noble gas configuration|noble gas configuration]]) tends to be chemically inert. Atoms with one or two valence electrons more than a closed shell are highly reactive due to the relatively [[Physics:Ionization energy|low energy]] to remove the extra valence electrons to form a positive [[Physics:Ion|ion]]. 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 a [[Physics:Core electron|core electron]], a valence electron has the ability to absorb or release energy in the form of a [[Physics:Photon|photon]]. An energy gain can trigger the electron to move (jump) to an outer shell; this is known as [[Physics:Excited state#Atomic excitation|atomic excitation]]. Or the electron can even break free from its associated atom's shell; this is [[Physics:Ionization|ionization]] 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.

==Overview==

===Electron configuration===

The electrons that determine [[Chemistry:Valence|valence]] – how an atom reacts chemically – are those with the highest [[Physics:Energy|energy]].

For a [[Chemistry:Main-group element|main-group element]], the valence electrons are defined as those electrons residing in the electronic shell of highest [[Physics:Principal quantum number|principal quantum number]] ''n''.<ref>{{cite book |last1 = Petrucci |first1 = Ralph H. |last2 = Harwood |first2 = William S. |last3 = Herring |first3 = F. Geoffrey |date=2002 |title = General chemistry: principles and modern applications |url = https://archive.org/details/generalchemistry00hill |url-access = registration |edition=8th |location=Upper Saddle River, N.J |publisher=Prentice Hall |isbn = 978-0-13-014329-7 |lccn=2001032331 |oclc=46872308 |page=[https://archive.org/details/generalchemistry00hill/page/339 339] }}</ref> Thus,  the number of valence electrons that it may have depends on the [[Physics:Electron configuration|electron configuration]] in a simple way. For example, the electronic configuration of [[Chemistry:Phosphorus|phosphorus]] (P) is 1s<sup>2</sup> 2s<sup>2</sup> 2p<sup>6</sup> 3s<sup>2</sup> 3p<sup>3</sup> so that there are 5 valence electrons (3s<sup>2</sup> 3p<sup>3</sup>), corresponding to a maximum valence for P of 5 as in the [[Physics:Molecule|molecule]] PF<sub>5</sub>; this configuration is normally abbreviated to [Ne] 3s<sup>2</sup> 3p<sup>3</sup>, where [Ne] signifies the core electrons whose configuration is identical to that of the [[Chemistry:Noble gas|noble gas]] [[Chemistry:Neon|neon]].

However, transition elements have (''n''−1)d energy levels that are very close in energy to the ''n''{{serif|s}} level.<ref>[http://www.chemguide.co.uk/atoms/properties/3d4sproblem.html THE ORDER OF FILLING 3d AND 4s ORBITALS]. chemguide.co.uk</ref> 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.<ref>Miessler G.L. and Tarr, D.A., Inorganic Chemistry (2nd edn. Prentice-Hall 1999). p.48.</ref> Thus, generally, the d electrons in transition metals behave as valence electrons although they are not in the outermost shell. For example, [[Chemistry:Manganese|manganese]] (Mn) has configuration 1s<sup>2</sup> 2s<sup>2</sup> 2p<sup>6</sup> 3s<sup>2</sup> 3p<sup>6</sup> 4s<sup>2</sup> 3d<sup>5</sup>; this is abbreviated to [Ar] 4s<sup>2</sup> 3d<sup>5</sup>, where [Ar] denotes a core configuration identical to that of the noble gas [[Chemistry:Argon|argon]]. 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 (4s<sup>2</sup> 3d<sup>5</sup>) outside the argon-like core; this is consistent with the chemical fact that manganese can have an [[Chemistry:Oxidation state|oxidation state]] as high as +7 (in the [[Chemistry:Permanganate|permanganate]] ion: {{chem|MnO|4|-}}). (But note that merely having that number of valence electrons does not imply that the corresponding oxidation state will exist. For example, [[Chemistry:Fluorine|fluorine]] is not known in oxidation state +7; and although the maximum known number of valence electrons is 16 in [[Chemistry:Ytterbium|ytterbium]] and [[Chemistry:Nobelium|nobelium]], 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 a [[Chemistry:Nickel|nickel]] atom has, in principle, ten valence electrons (4s<sup>2</sup> 3d<sup>8</sup>), its oxidation state never exceeds four. For [[Chemistry:Zinc|zinc]], the 3d subshell is complete in all known compounds, although it does contribute to the valence band in some compounds.<ref>{{cite journal |last1=Tossell |first1=J. A. |date=1 November 1977 |title=Theoretical studies of valence orbital binding energies in solid zinc sulfide, zinc oxide, and zinc fluoride |journal=Inorganic Chemistry |volume=16 |issue=11 |pages=2944–2949 |doi=10.1021/ic50177a056}}</ref> Similar patterns hold for the (''n''−2)f energy levels of inner transition metals.

The [[Chemistry:D electron count|d electron count]] is an alternative tool for understanding the chemistry of a transition metal.

===The number of valence electrons===

The number of valence electrons of an element can be determined by the periodic 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.)<ref name="KW">{{cite book |last1=Keeler |first1=James |last2=Wothers |first2=Peter |date=2014 |title=Chemical Structure and Reactivity |url= |edition=2nd |location= |publisher=Oxford University Press |pages=257–260 |isbn=978-0-19-9604135}}</ref>

{| class="wikitable"
!
! [[Chemistry:Alkali metal|1]]
! [[Chemistry:Alkaline earth metal|2]]
! colspan=14 |
! [[Chemistry:Group 3 element|3]]
! [[Chemistry:Group 4 element|4]]
! [[Chemistry:Group 5 element|5]]
! [[Chemistry:Group 6 element|6]]
! [[Chemistry:Group 7 element|7]]
! [[Chemistry:Group 8 element|8]]
! [[Chemistry:Group 9 element|9]]
! [[Chemistry:Group 10 element|10]]
! [[Chemistry:Group 11 element|11]]
! [[Chemistry:Group 12 element|12]]
! [[Chemistry:Boron group|13]]
! [[Chemistry:Carbon group|14]]
! [[Chemistry:Pnictogen|15]]
! [[Chemistry:Chalcogen|16]]
! [[Chemistry:Halogen|17]]
! [[Chemistry:Noble gas|18]]
|-
! [[Chemistry:Period 1 element|1]]
| bgcolor="{{element color|s-block}}" | H<br />1
|
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| bgcolor="{{element color|s-block}}" | He<br />2
|-
! [[Chemistry:Period 2 element|2]]
| bgcolor="{{element color|s-block}}" | Li<br />1
| bgcolor="{{element color|s-block}}" | Be<br />2
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| bgcolor="{{element color|p-block}}" | B<br />3
| bgcolor="{{element color|p-block}}" | C<br />4
| bgcolor="{{element color|p-block}}" | N<br />5
| bgcolor="{{element color|p-block}}" | O<br />6
| bgcolor="{{element color|p-block}}" | F<br />7
| bgcolor="{{element color|p-block}}" | Ne<br />8
|-
! [[Chemistry:Period 3 element|3]]
| bgcolor="{{element color|s-block}}" | Na<br />1
| bgcolor="{{element color|s-block}}" | Mg<br />2
|
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| bgcolor="{{element color|p-block}}" | Al<br />3
| bgcolor="{{element color|p-block}}" | Si<br />4
| bgcolor="{{element color|p-block}}" | P<br />5
| bgcolor="{{element color|p-block}}" | S<br />6
| bgcolor="{{element color|p-block}}" | Cl<br />7
| bgcolor="{{element color|p-block}}" | Ar<br />8
|-
! [[Chemistry:Period 4 element|4]]
| bgcolor="{{element color|s-block}}" | K<br />1
| bgcolor="{{element color|s-block}}" | Ca<br />2
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| bgcolor="{{element color|d-block}}" | Sc<br />3
| bgcolor="{{element color|d-block}}" | Ti<br />4
| bgcolor="{{element color|d-block}}" | V<br />5
| bgcolor="{{element color|d-block}}" | Cr<br />6
| bgcolor="{{element color|d-block}}" | Mn<br />7
| bgcolor="{{element color|d-block}}" | Fe<br />8
| bgcolor="{{element color|d-block}}" | Co<br />9
| bgcolor="{{element color|d-block}}" | Ni<br />10
| bgcolor="{{element color|d-block}}" | Cu<br />11
| bgcolor="{{element color|d-block}}" | Zn<br />12
| bgcolor="{{element color|p-block}}" | Ga<br />3
| bgcolor="{{element color|p-block}}" | Ge<br />4
| bgcolor="{{element color|p-block}}" | As<br />5
| bgcolor="{{element color|p-block}}" | Se<br />6
| bgcolor="{{element color|p-block}}" | Br<br />7
| bgcolor="{{element color|p-block}}" | Kr<br />8
|-
! [[Chemistry:Period 5 element|5]]
| bgcolor="{{element color|s-block}}" | Rb<br />1
| bgcolor="{{element color|s-block}}" | Sr<br />2
|
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| bgcolor="{{element color|d-block}}" | Y<br />3
| bgcolor="{{element color|d-block}}" | Zr<br />4
| bgcolor="{{element color|d-block}}" | Nb<br />5
| bgcolor="{{element color|d-block}}" | Mo<br />6
| bgcolor="{{element color|d-block}}" | Tc<br />7
| bgcolor="{{element color|d-block}}" | Ru<br />8
| bgcolor="{{element color|d-block}}" | Rh<br />9
| bgcolor="{{element color|d-block}}" | Pd<br />10
| bgcolor="{{element color|d-block}}" | Ag<br />11
| bgcolor="{{element color|d-block}}" | Cd<br />12
| bgcolor="{{element color|p-block}}" | In<br />3
| bgcolor="{{element color|p-block}}" | Sn<br />4
| bgcolor="{{element color|p-block}}" | Sb<br />5
| bgcolor="{{element color|p-block}}" | Te<br />6
| bgcolor="{{element color|p-block}}" | I<br />7
| bgcolor="{{element color|p-block}}" | Xe<br />8
|-
! [[Chemistry:Period 6 element|6]]
| bgcolor="{{element color|s-block}}" | Cs<br />1
| bgcolor="{{element color|s-block}}" | Ba<br />2
| bgcolor="{{element color|f-block}}" | La<br />3
| bgcolor="{{element color|f-block}}" | Ce<br />4
| bgcolor="{{element color|f-block}}" | Pr<br />5
| bgcolor="{{element color|f-block}}" | Nd<br />6
| bgcolor="{{element color|f-block}}" | Pm<br />7
| bgcolor="{{element color|f-block}}" | Sm<br />8
| bgcolor="{{element color|f-block}}" | Eu<br />9
| bgcolor="{{element color|f-block}}" | Gd<br />10
| bgcolor="{{element color|f-block}}" | Tb<br />11
| bgcolor="{{element color|f-block}}" | Dy<br />12
| bgcolor="{{element color|f-block}}" | Ho<br />13
| bgcolor="{{element color|f-block}}" | Er<br />14
| bgcolor="{{element color|f-block}}" | Tm<br />15
| bgcolor="{{element color|f-block}}" | Yb<br />16
| bgcolor="{{element color|d-block}}" | Lu<br />3
| bgcolor="{{element color|d-block}}" | Hf<br />4
| bgcolor="{{element color|d-block}}" | Ta<br />5
| bgcolor="{{element color|d-block}}" | W<br />6
| bgcolor="{{element color|d-block}}" | Re<br />7
| bgcolor="{{element color|d-block}}" | Os<br />8
| bgcolor="{{element color|d-block}}" | Ir<br />9
| bgcolor="{{element color|d-block}}" | Pt<br />10
| bgcolor="{{element color|d-block}}" | Au<br />11
| bgcolor="{{element color|d-block}}" | Hg<br />12
| bgcolor="{{element color|p-block}}" | Tl<br />3
| bgcolor="{{element color|p-block}}" | Pb<br />4
| bgcolor="{{element color|p-block}}" | Bi<br />5
| bgcolor="{{element color|p-block}}" | Po<br />6
| bgcolor="{{element color|p-block}}" | At<br />7
| bgcolor="{{element color|p-block}}" | Rn<br />8
|-
! [[Chemistry:Period 7 element|7]]
| bgcolor="{{element color|s-block}}" | Fr<br />1
| bgcolor="{{element color|s-block}}" | Ra<br />2
| bgcolor="{{element color|f-block}}" | Ac<br />3
| bgcolor="{{element color|f-block}}" | Th<br />4
| bgcolor="{{element color|f-block}}" | Pa<br />5
| bgcolor="{{element color|f-block}}" | U<br />6
| bgcolor="{{element color|f-block}}" | Np<br />7
| bgcolor="{{element color|f-block}}" | Pu<br />8
| bgcolor="{{element color|f-block}}" | Am<br />9
| bgcolor="{{element color|f-block}}" | Cm<br />10
| bgcolor="{{element color|f-block}}" | Bk<br />11
| bgcolor="{{element color|f-block}}" | Cf<br />12
| bgcolor="{{element color|f-block}}" | Es<br />13
| bgcolor="{{element color|f-block}}" | Fm<br />14
| bgcolor="{{element color|f-block}}" | Md<br />15
| bgcolor="{{element color|f-block}}" | No<br />16
| bgcolor="{{element color|d-block}}" | Lr<br />3
| bgcolor="{{element color|d-block}}" | Rf<br />4
| bgcolor="{{element color|d-block}}" | Db<br />5
| bgcolor="{{element color|d-block}}" | Sg<br />6
| bgcolor="{{element color|d-block}}" | Bh<br />7
| bgcolor="{{element color|d-block}}" | Hs<br />8
| bgcolor="{{element color|d-block}}" | Mt<br />9
| bgcolor="{{element color|d-block}}" | Ds<br />10
| bgcolor="{{element color|d-block}}" | Rg<br />11
| bgcolor="{{element color|d-block}}" | Cn<br />12
| bgcolor="{{element color|p-block}}" | Nh<br />3
| bgcolor="{{element color|p-block}}" | Fl<br />4
| bgcolor="{{element color|p-block}}" | Mc<br />5
| bgcolor="{{element color|p-block}}" | Lv<br />6
| bgcolor="{{element color|p-block}}" | Ts<br />7
| bgcolor="{{element color|p-block}}" | Og<br />8
|}

Helium is an exception: despite having a 1s<sup>2</sup> configuration with two valence electrons, and thus having some similarities with the alkaline earth metals with their ''n''s<sup>2</sup> 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.

==Valence shell==
The valence shell is the set of [[Physics:Atomic orbital|orbital]]s which are energetically accessible for accepting electrons to form [[Chemistry:Chemical bond|chemical bond]]s.

For main-group elements, the valence shell consists of the ''n''s and ''n''p orbitals in the outermost [[Physics:Electron shell|electron shell]]. For [[Chemistry:Transition metal|transition metal]]s the orbitals of the incomplete (''n''−1)d subshell are included, and for [[Chemistry:Lanthanide|lanthanide]]s and [[Chemistry:Actinide|actinide]]s 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 number ''n'' in a given element, but they are all at similar energies.<ref name=KW/>

{| class="wikitable"
|-
!Element type||[[Chemistry:Hydrogen|Hydrogen]] and [[Chemistry:Helium|helium]]||s- and p-blocks<br />([[Chemistry:Main-group element|main-group element]]s)||d-block<br />([[Chemistry:Transition metal|Transition metal]]s)||f-block<br />([[Chemistry:Lanthanide|Lanthanide]]s and [[Chemistry:Actinide|actinide]]s)
|-
!Valence orbitals<ref>{{cite journal
| title = Octacarbonyl Ion Complexes of Actinides [An(CO)8]+/− (An=Th, U) and the Role of f Orbitals in Metal–Ligand Bonding
| first1= Chaoxian |last1=Chi |first2=Sudip |last2=Pan | first3= Jiaye |last3=Jin |first4=Luyan |last4=Meng | first5= Mingbiao |last5=Luo |first6=Lili |last6=Zhao |first7=Mingfei |last7=Zhou |first8=Gernot |last8=Frenking
| journal = [[Chemistry:Chemistry: A European Journal|Chem. Eur. J.]]
| year = 2019
| volume = 25
| issue = 50
| pages = 11772–11784
| doi = 10.1002/chem.201902625
| pmid= 31276242 | pmc= 6772027 |doi-access=free }}</ref>
|
* 1s
|
* ''n''s
* ''n''p
|
* ''n''s
* (''n''−1)d
* ''n''p
|
* ''n''s
* (''n''−2)f
* (''n''−1)d
* ''n''p
|-
![[Chemistry:Electron counting|Electron counting]] rules
|Duet/Duplet rule
|[[Octet rule]]
|[[18-electron rule|18-electron rule]]
|32-electron rule
|}

As a general rule, a [[Chemistry:Main-group element|main-group element]] (except hydrogen or helium) tends to react to form a s<sup>2</sup>p<sup>6</sup> [[Physics:Electron configuration|electron configuration]]. This tendency is called the [[Octet rule|octet rule]], because each bonded atom has 8 valence electrons including shared electrons. Similarly, a transition metal tends to react to form a d<sup>10</sup>s<sup>2</sup>p<sup>6</sup> electron configuration. This tendency is called the [[18-electron rule|18-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.<ref>{{Greenwood&Earnshaw2nd|page=117}}</ref><ref>{{cite journal |last1=Zhou |first1=Mingfei |last2=Frenking |first2=Gernot |date=2021 |title=Transition-Metal Chemistry of the Heavier Alkaline Earth Atoms Ca, Sr, and Ba |url= |journal=Accounts of Chemical Research |volume=54 |issue=15 |pages=3071–3082 |doi=10.1021/acs.accounts.1c00277 |pmid=34264062 |s2cid=235908113 |access-date=}}</ref><ref>{{cite journal |last1=Fernández |first1=Israel |last2=Holzmann |first2=Nicole |last3=Frenking |first3=Gernot |date=2020 |title=The Valence Orbitals of the Alkaline-Earth Atoms |journal=Chemistry: A European Journal |volume=26 |issue=62 |pages=14194–14210 |doi=10.1002/chem.202002986 |pmid=32666598 |pmc=7702052 |doi-access=free }}</ref>

==Chemical reactions==
{{main|Chemistry:Valence}}
The number of valence electrons in an atom governs its [[Chemistry:Chemical bond|bonding]] behavior. Therefore, elements whose atoms have the same number of valence electrons are often grouped together in the [[Chemistry:Periodic table|periodic table]] of the elements, especially if they also have the same types of valence orbitals.<ref name=jensenlaw>{{cite web|url=http://www.che.uc.edu/jensen/W.%20B.%20Jensen/Reprints/081.%20Periodic%20Table.pdf|last1=Jensen|first1=William B.|title=The Periodic Law and Table|date=2000|archive-url=https://web.archive.org/web/20201110113324/http://www.che.uc.edu/jensen/W.%20B.%20Jensen/Reprints/081.%20Periodic%20Table.pdf |access-date=10 December 2022|archive-date=2020-11-10 }}</ref>

The most [[Chemistry:Reactivity|reactive]] kind of metallic element is an [[Chemistry:Alkali metal|alkali metal]] of group 1 (e.g., [[Chemistry:Sodium|sodium]] or [[Chemistry:Potassium|potassium]]); this is because such an atom has only a single valence electron. During the formation of an ionic bond, which provides the necessary [[Physics:Ionization energy|ionization energy]], this one valence electron is easily lost to form a positive [[Physics:Ion|ion]] (cation) with a closed shell (e.g., Na<sup>+</sup> or K<sup>+</sup>). An [[Chemistry:Alkaline earth metal|alkaline earth metal]] of group 2 (e.g., [[Chemistry:Magnesium|magnesium]]) is somewhat less reactive, because each atom must lose two valence electrons to form a positive ion with a closed shell (e.g., Mg<sup>2+</sup>).{{citation needed|date=April 2023}}

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 higher [[Physics:Principal quantum number|principal quantum number]]s (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|date=April 2023}}

A [[Chemistry:Nonmetal|nonmetal]] 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 (a [[Chemistry:Covalent bond|covalent bond]]), or it can remove electrons from another atom (an ionic bond). The most reactive kind of nonmetal element is a [[Chemistry:Halogen|halogen]] (e.g., [[Chemistry:Fluorine|fluorine]] (F) or [[Chemistry:Chlorine|chlorine]] (Cl)). Such an atom has the following electron configuration: s<sup>2</sup>p<sup>5</sup>; 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<sup>−</sup>, Cl<sup>−</sup>, 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|date=April 2023}}

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, the [[Chemistry:Valence|valence]] 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 are [[Octet rule#Exceptions|exceptions]], and for which the valence is less clearly defined.

==Electrical conductivity==
Valence electrons are also responsible for the bonding in the pure chemical elements, and whether their [[Physics:Electrical conductivity|electrical conductivity]] is characteristic of metals, semiconductors, or insulators.


[[Chemistry:Metal|Metal]]lic elements generally have high [[Physics:Electrical conductor|electrical conductivity]] when in the [[Physics:Solid|solid]] state. In each row of the periodic 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 small [[Physics:Ionization energy|ionization 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 characterises metallic bonding. Such a "free" electron can be moved under the influence of an [[Physics:Electric field|electric field]], and its motion constitutes an [[Physics:Electric current|electric current]]; it is responsible for the electrical conductivity of the metal. [[Chemistry:Copper|Copper]], [[Chemistry:Aluminium|aluminium]], [[Chemistry:Silver|silver]], and [[Chemistry:Gold|gold]] are examples of good conductors.

A [[Chemistry:Nonmetal|nonmetal]]lic element has low electrical conductivity; it acts as an insulator. 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 is [[Chemistry:Boron|boron]]). 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 are [[Chemistry:Diamond|diamond]] (an allotrope of [[Chemistry:Carbon|carbon]]) and [[Chemistry:Sulfur|sulfur]]. 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 by intermolecular forces (as in sulfur). (The [[Chemistry:Noble gas|noble gas]]es remain as single atoms, but those also experience intermolecular forces of attraction, that become stronger as the group is descended: helium boils at −269&nbsp;°C, while radon boils at −61.7&nbsp;°C.)

A solid compound containing metals can also be an insulator if the valence electrons of the metal atoms are used to form ionic bonds. For example, although elemental [[Chemistry:Sodium|sodium]] is a metal, solid [[Chemistry:Sodium chloride|sodium 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.

A [[Physics:Semiconductor|semiconductor]] 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 with [[Physics:Temperature|temperature]]. The typical elemental semiconductors are [[Chemistry:Silicon|silicon]] and [[Chemistry:Germanium|germanium]], each atom of which has four valence electrons. The properties of semiconductors are best explained using band theory, as a consequence of a small energy gap between a valence band (which contains the valence electrons at absolute zero) and a conduction band (to which valence electrons are excited by thermal energy).

==References==
{{reflist}}

==External links==
# Francis, Eden. [https://web.archive.org/web/20060115042442/http://dl.clackamas.cc.or.us/ch104-06/valence_electrons.htm Valence Electrons].

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