Valence electron

In chemistry, valence electrons are the outermost electrons of an atom, which are important in determining how the atom reacts chemically with other atoms. Atoms with a complete shell of valence electrons (corresponding to an electron configuration s2p6) tend to be chemically inert. Atoms with one or two valence electrons more than a closed shell are highly reactive because the extra electrons are easily removed to form positive ions. Atoms with one or two valence electrons fewer than a closed shell are also highly reactive because of a tendency either to gain the missing electrons and form negative ions, or to share electrons and form covalent bonds.

Valence electrons have the ability, like electrons in inner shells, to absorb or release energy in the form of photons. This gain or loss of energy can trigger an electron to move (jump) to another shell or even break free from the atom and its valence shell. When an electron absorbs energy in the form of one or more photons, then it moves to a more outer shell depending on the amount of energy gained. (See also : electrons in an excited state). When an electron loses energy (photons), then it moves to a more inner shell.

Contents

The number of valence electrons

The number of valence electrons of an element is determined by its periodic table group (vertical column) in which the element is categorized. With the exception of groups 3–12 (transition metals), the number within the unit's place identifies how many valence electrons are contained within the elements listed under that particular column.

The periodic table of the chemical elements
Periodic table group Valence
electrons
Group 1 (I) (alkali metals) 1
Group 2 (II) (alkaline earth metals) 2
Groups 3-12 (transition metals) #*
Group 13 (III) (boron group) 3
Group 14 (IV) (carbon group) 4
Group 15 (V) (nitrogen group) 5
Group 16 (VI) (chalcogens) 6
Group 17 (VII) (halogens) 7
Group 18 (VIII or 0) (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.

Valence electrons and electron configuration

For main group elements, the number of valence electrons depends on the electron configuration in a simple way, but for transition metals the relationship is more complex.

For main group elements, valence electrons can be defined as those 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 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 the noble gas neon.

However this simple method does not work for transition metals, which have incomplete nd (i.e. 3d, 4d or 5d) subshells whose energy is normally comparable with that of the (n+1)s electrons. The valence electrons are instead defined as those 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 argon. In this atom, the 3d electrons have energies similar to those of the 4s electrons, and much higher than for the 3s and 3p electrons. In effect there are seven valence electrons (4s2 3d5) outside the argon-like core. This is consistent with the chemical fact that manganese can have oxidation states as high as +7 (in the permanganate ion MnO4-).

Towards the right of each transition metal series, the d electrons descend to lower energies and have less valence electron character. Thus although nickel has in principle ten valence electrons (4s2 3d8), the oxidation state never exceeds four. For zinc and succeeding elements, the 3d subshell is complete and the 3d electrons are considered core electrons.

Since the number of valence electrons which actually participate in chemical reactions is difficult to predict, the concept of valence electrons is less useful for transition metals than for main group elements. As mentioned above, the d electron count provides a more useful tool for the understanding of the chemistry of these elements.

Valence electrons in chemical reactions

Main article: Valence (chemistry)

The number of electrons in an atom's outermost valence shell governs its bonding behavior. Therefore, elements with the same number of valence electrons are grouped together in the periodic table of the elements. As a general rule, atoms of main group elements (except hydrogen and helium) tend to react to form a "closed" or complete shell, corresponding to an s2p6 electron configuration. This tendency is called the octet rule since the bonded atom has or shares eight valence electrons.

The most reactive metallic elements are the alkali metals of Group 1, for example sodium (Na) and potassium (K) whose atoms each have a single valence electron. This is easily lost to form a positive ion (cation) with a closed shell (Na+ or K+), during the formation of an ionic bond which provides the necessary ionization energy. The alkaline earth metals of Group 2, for example magnesium, are somewhat less reactive since each atom must lose two valence electrons to form a positive ion with a closed shell such as Mg2+.

Nonmetal atoms tend to attract additional valence electrons to attain a full valence shell. This can be achieved one of two ways: an atom can either share electrons with neighboring atoms, a covalent bond, or it can remove electrons from other atoms, an ionic bond. The most reactive non-metals are the halogens such as fluorine (F) and chlorine (Cl), which have electron configurations s2p5 and require only one additional valence electron for a closed shell. To form an ionic bond, a halogen atom can remove an electron from another atom to form an anion (F-, Cl-, etc.). To form a covalent bond, one electron from the halogen and one electron from another atom form a shared pair. For example in the molecule H-F, the line represents a shared pair of valence electrons, one from H and one from F.

In these simple cases where the octet rule is obeyed, the valence of an atom equals the number of electrons gained, lost or shared to form the stable octet. However there are also many molecules which are exceptions, and for which the valence is less clearly defined.

Valence electrons and electrical conductivity

The valence electrons are also responsible for the electrical conductivity of elements, which may be divided into metals, nonmetals, and semiconductors or metalloids.

Metals or metallic elements are elements with high electrical conductivity in the solid state. In each row of the periodic table the metals occur to the left of the nonmetals and thus have fewer valence electrons. The valence electrons which are present have small ionization energies, and in the solid state they are relatively free to leave one atom and move to its neighbour. These “free electrons” can move under the influence of an electric field and their motion constitutes an electric current. They are therefore responsible for the electrical conductivity of the metal. Copper, aluminium, silver and gold are examples of good conductors used widely in industry.

Nonmetallic elements have low electrical conductivity and act as insulators. They are found to the right of the periodic table with valence shells which are at least half full (except for boron). Their ionization energies are large so that electrons cannot leave an atom easily when an electric field is applied, and they conduct only very small electric currents. Examples of solid elemental insulators are diamond (an elemental form of carbon) and sulfur.

Solid compounds containing metals can also be insulators 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 cannot move easily in an electric field.

Semiconductors have an electrical conductivity intermediate between metals and nonmetals, and also differ from metals in that their conductivity increases with temperature. The typical elemental semiconductors are silicon and germanium with four valence electrons each. Their properties 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

  1. ↑ Petrucci R.H., Harwood W.S. and Herring F.G., General Chemistry (8th edn, Prentice-Hall 2002), p.339
  2. ↑ Miessler G.L. and Tarr, D.A., Inorganic Chemistry (2nd edn. Prentice-Hall 1999). p.48

External links

  1. Francis, Eden. Valence Electrons.