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[[File:Covalent.svg|thumb|180px|Four [[covalent bond]]s. Carbon has four valence electrons and here a [[Valence (chemistry)|valence]] of four. Each hydrogen atom has one valence electron and is univalent.]] | |||
In chemistry, a '''valence electron''' is an electron that is associated with an [[atom]], and that can participate in the formation of a [[chemical bond]]; in a single [[covalent bond]], both atoms in the bond contribute one valence electron in order to form a [[shared pair]]. The presence of valence electrons can determine the element's chemical properties and whether it may bond with other elements: For a [[main group element]], a valence electron can only be in the outermost [[electron shell]]. In a [[transition metal]], a valence electron can also be in an inner shell. | |||
An atom with a [[closed shell]] of valence electrons (corresponding to an [[electron configuration]] s<sup>2</sup>p<sup>6</sup>) tends to be [[inert gases|chemically inert]]. An atom with one or two valence electrons more than a closed shell is highly reactive, because the extra valence electrons are easily removed to form a positive [[ion]]. An atom with one or two valence electrons fewer than a closed shell is also highly reactive, because of a tendency either to gain the missing valence electrons (thereby forming a negative ion), or to share valence electrons (thereby forming a covalent bond). | |||
Like an electron in an inner shell, a valence electron has the ability to absorb or release energy in the form of a [[photon]]. An energy gain can trigger an electron to move (jump) to an outer shell; this is known as [[Excited_state#Atomic_excitation|atomic excitation]]. Or the electron can even break free from its associated atom's valence shell; this is [[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. | |||
==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. With the exception of groups 3–12 (the [[transition metal]]s), the number within the unit's place identifies how many valence electrons may be associated with a neutral atom of an element listed under that particular column. | |||
[[Image:Periodic table (polyatomic).svg|500px|thumb|right|The [[periodic table]] of the chemical elements]] | |||
{| class="wikitable" | |||
|- | |||
![[Periodic table group]]||Valence<br>electrons | |||
|- | |||
| Group 1 (I) ([[alkali metal]]s) || 1 | |||
|- | |||
| Group 2 (II) ([[alkaline earth metal]]s) || 2 | |||
|- | |||
| Groups 3-12 ([[transition metal]]s) || 2* 4s shell is complete and cannot hold any more electrons. | |||
|- | |||
| Group 13 (III) ([[boron group]]) || 3 | |||
|- | |||
| Group 14 (IV) ([[carbon group]]) || 4 | |||
|- | |||
| Group 15 (V) ([[nitrogen group|pnictogens]]) || 5 | |||
|- | |||
| Group 16 (VI) ([[chalcogen]]s) || 6 | |||
|- | |||
| Group 17 (VII) ([[halogen]]s) || 7 | |||
|- | |||
| Group 18 ([[noble gases]]) || 8** | |||
|} | |||
''* The general method for counting valence electrons is generally not useful for transition metals. Instead the modified [[d electron count]] method is used.'' | |||
''** Except for [[helium]], which has only two valence electrons. | |||
==Electron configuration== | |||
For a [[main group element]], the number of valence electrons that it may have depends on the [[electron configuration]] in a simple way. However, for a [[transition metal]], the relationship is more complex. | |||
For a main group element, a valence electron is defined as an electron that resides in the electronic shell of highest [[principal quantum number]] n.<ref>Petrucci R.H., Harwood W.S. and Herring F.G., General Chemistry (8th edn, Prentice-Hall 2002), p.339.</ref> For example the electronic configuration of 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 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 [[noble gas]] neon. | |||
However, this simple method does not work for transition metals, which have incomplete <math>n</math>d (i.e., 3d, 4d or 5d) subshells, whose energy is normally comparable with that of an <math>(n+1)</math>s electron. Instead, 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> For example, [[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 [[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 [[oxidation state]] as high as +7 (in the [[permanganate]] ion: {{chem|MnO|4|-}}). | |||
The further right of each transition metal series, the lower the energy of an electron in a d subshell and the less such an electron has the properties of a valence electron. Thus, although a [[nickel]] atom has, in principle, ten valence electrons (4s<sup>2</sup> 3d<sup>8</sup>), its oxidation state never exceeds four. | |||
Because the number of valence electrons which actually participate in chemical reactions is difficult to predict, the concept of the valence electron is less useful for a transition metal than for a main group element; as mentioned already, the [[d electron count]] provides a more useful tool for understanding the chemistry of a transition metal. | |||
==Chemical reactions== | |||
{{main|Valence (chemistry)}} | |||
The number of electrons in an atom's outermost valence shell governs its [[chemical bond|bonding]] behavior. Therefore, elements whose atoms can have the same number of valence electrons are grouped together in the [[periodic table]] of the elements. As a general rule, a [[main group element]] (except hydrogen or helium) tends to react to form a [[closed shell]], corresponding to the following [[electron configuration]]: s<sup>2</sup>p<sup>6</sup>. This tendency is called the [[octet rule]], because the bonded atoms shares eight valence electrons. | |||
The most [[reactivity (chemistry)|reactive]] kind of [[metallic element]] is an [[alkali metal]] of group 1 (e.g., [[sodium]] or [[potassium]]); this is because such an atom has only a single valence electron; during the formation of an [[ionic bond]] which provides the necessary [[ionization energy]], this one valence electron is easily lost to form a positive [[ion]] (cation) with a closed shell (e.g., Na<sup>+</sup> or K<sup>+</sup>). An [[alkaline 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., Mg<sup>2+</sup>). | |||
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 [[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). | |||
A [[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 [[covalent bond]]), or it can remove electrons from another atom (an [[ionic bond]]). The most reactive kind of nonmetal element is a [[halogen]] (e.g., [[fluorine]] (F) or [[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). | |||
Within each group of nonmetals, reactivity decreases with each lower rows 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 shell of a halogen is at a higher principal quantum number. | |||
In these simple cases where the octet rule is obeyed, the [[valence (chemistry)|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 which are [[Octet rule#Exceptions|exceptions]], and for which the valence is less clearly defined. | |||
==Electrical conductivity== | |||
Valence electrons are also responsible for the [[electrical conductivity]] of an element; as a result, an element may be classified as a [[metal]], a [[nonmetals]], or [[semiconductor]]s (or [[metalloid]]). | |||
A [[metal]] is an element with high electrical conductivity or malleability when in the [[solid]] state. In each row of the [[Periodic table (metals and non-metals)|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 [[ionization energy]], and in the solid state this valence electron is relatively free to leave one atom in order to associate with another nearby. Such a "free" electron can be moved under the influence of an [[electric field]], and its motion constitutes an [[electric current]]; it is responsible for the electrical conductivity of the metal. [[Copper]], [[aluminium]], [[silver]], and [[gold]] are examples of good conductors. | |||
A [[nonmetal]]lic element has low electrical conductivity; it acts as an [[insulator (electrical)|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 [[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 [[diamond]] (an [[allotrope]] of [[carbon]]) and [[sulfur]]. | |||
A solid compound containing metals can also be an insulator if the valence electrons of the metal atoms are used to form [[ionic bond]]s. For example, although elemental [[sodium]] is a metal, solid [[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 [[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 [[temperature]]. The typical elemental semiconductors are [[silicon]] and [[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. [http://dl.clackamas.cc.or.us/ch104-06/valence_electrons.htm Valence Electrons]. | |||
{{DEFAULTSORT:Valence Electron}} | |||
[[Category:Chemical bonding]] | |||
[[Category:Electron states]] | |||
Revision as of 00:32, 9 August 2013
In chemistry, a valence electron is an electron that is associated with an atom, and that can participate in the formation of a chemical bond; in a single covalent bond, both atoms in the bond contribute one valence electron in order to form a shared pair. The presence of valence electrons can determine the element's chemical properties and whether it may bond with other elements: For a main group element, a valence electron can only be in the outermost electron shell. In a transition metal, a valence electron can also be in an inner shell.
An atom with a closed shell of valence electrons (corresponding to an electron configuration s2p6) tends to be chemically inert. An atom with one or two valence electrons more than a closed shell is highly reactive, because the extra valence electrons are easily removed to form a positive ion. An atom with one or two valence electrons fewer than a closed shell is also highly reactive, because of a tendency either to gain the missing valence electrons (thereby forming a negative ion), or to share valence electrons (thereby forming a covalent bond).
Like an electron in an inner shell, a valence electron has the ability to absorb or release energy in the form of a photon. An energy gain can trigger an electron to move (jump) to an outer shell; this is known as atomic excitation. Or the electron can even break free from its associated atom's valence shell; this is 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.
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. With the exception of groups 3–12 (the transition metals), the number within the unit's place identifies how many valence electrons may be associated with a neutral atom of an element listed under that particular column.
.svg)
| Periodic table group | Valence electrons |
|---|---|
| Group 1 (I) (alkali metals) | 1 |
| Group 2 (II) (alkaline earth metals) | 2 |
| Groups 3-12 (transition metals) | 2* 4s shell is complete and cannot hold any more electrons. |
| Group 13 (III) (boron group) | 3 |
| Group 14 (IV) (carbon group) | 4 |
| Group 15 (V) (pnictogens) | 5 |
| Group 16 (VI) (chalcogens) | 6 |
| Group 17 (VII) (halogens) | 7 |
| Group 18 (noble gases) | 8** |
* The general method for counting valence electrons is generally not useful for transition metals. Instead the modified d electron count method is used.
** Except for helium, which has only two valence electrons.
Electron configuration
For a main group element, the number of valence electrons that it may have depends on the electron configuration in a simple way. However, for a transition metal, the relationship is more complex.
For a main group element, a valence electron is defined as an electron that resides in the electronic shell of highest principal quantum number n.[1] For example the electronic configuration of phosphorus (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 the molecule PF5; this configuration is normally abbreviated to [Ne] 3s2 3p3, where [Ne] signifies the core electrons whose configuration is identical to that of the noble gas neon.
However, this simple method does not work for transition metals, which have incomplete d (i.e., 3d, 4d or 5d) subshells, whose energy is normally comparable with that of an s electron. Instead, a valence electron for a transition metal is defined as an electron that resides outside a noble-gas core.[2] 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 gas 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 (4s2 3d5) outside the argon-like core; this is consistent with the chemical fact that manganese can have an oxidation state as high as +7 (in the permanganate ion: Template:Chem).
The further right of each transition metal series, the lower the energy of an electron in a d subshell and the less such an electron has the properties of a valence electron. Thus, although a nickel atom has, in principle, ten valence electrons (4s2 3d8), its oxidation state never exceeds four.
Because the number of valence electrons which actually participate in chemical reactions is difficult to predict, the concept of the valence electron is less useful for a transition metal than for a main group element; as mentioned already, the d electron count provides a more useful tool for understanding the chemistry of a transition metal.
Chemical reactions
Mining Engineer (Excluding Oil ) Truman from Alma, loves to spend time knotting, largest property developers in singapore developers in singapore and stamp collecting. Recently had a family visit to Urnes Stave Church. The number of electrons in an atom's outermost valence shell governs its bonding behavior. Therefore, elements whose atoms can have the same number of valence electrons are grouped together in the periodic table of the elements. As a general rule, a main group element (except hydrogen or helium) tends to react to form a closed shell, corresponding to the following electron configuration: s2p6. This tendency is called the octet rule, because the bonded atoms shares eight valence electrons.
The most reactive kind of metallic element is an alkali metal of group 1 (e.g., sodium or potassium); this is because such an atom has only a single valence electron; during the formation of an ionic bond which provides the necessary ionization energy, this one valence electron is easily lost to form a positive ion (cation) with a closed shell (e.g., Na+ or K+). An alkaline 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+).
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 principal 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).
A 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 covalent bond), or it can remove electrons from another atom (an ionic bond). The most reactive kind of nonmetal element is a halogen (e.g., fluorine (F) or chlorine (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).
Within each group of nonmetals, reactivity decreases with each lower rows 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 shell of a halogen is at a higher principal quantum number.
In these simple cases where the octet rule is obeyed, the 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 which are exceptions, and for which the valence is less clearly defined.
Electrical conductivity
Valence electrons are also responsible for the electrical conductivity of an element; as a result, an element may be classified as a metal, a nonmetals, or semiconductors (or metalloid).
A metal is an element with high electrical conductivity or malleability when in the 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 ionization energy, and in the solid state this valence electron is relatively free to leave one atom in order to associate with another nearby. Such a "free" electron can be moved under the influence of an electric field, and its motion constitutes an electric current; it is responsible for the electrical conductivity of the metal. Copper, aluminium, silver, and gold are examples of good conductors.
A nonmetallic 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 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 diamond (an allotrope of carbon) and sulfur.
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 sodium is a metal, solid 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 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 temperature. The typical elemental semiconductors are silicon and 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
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External links
- Francis, Eden. Valence Electrons.