If the bonded atoms are non-metals and identical (as in N2 or O2), the electrons are shared equally between the two atoms, and the bond is called non-polar covalent. If the atoms are non-metals but differ (as in nitric oxide, NO), the electrons are shared unequally and the bond is called polar covalent -polar because the molecule has a positive and a negative electric pole much like the north and south poles of a magnet, and covalent because the atoms share electrons between them, even though unequally. These substances are not electrical conductors, nor do they have luster, ductility, or malleability.

   When a molecule of a substance contains atoms of both metals and non-metals, the electrons are more strongly attracted to the non-metals, which become negatively charged ions; the metals become positively charged ions. The ions then attract their opposites in charge, forming ionic bonds. Ionic substances conduct electricity when they are in the liquid state or in water solutions, but not in the crystalline state, because individual ions are too large to move freely through the crystal.

   Symmetrical sharing of electrons gives either metallic or non-polar covalent bonds; unsymmetrical sharing gives polar covalent bonds; electron transfer gives ionic bonds. The tendency for unequal distribution of electrons between pairs of atoms generally increases as they are farther apart in the periodic table.
For the formation of stable ions and of covalent bonds, the most common pattern is for each atom to achieve the same total number of electrons as the noble gas -Group 18 (or VIIIa)-element closest to it in the periodic table. The metals in Groups 1 (or Ia) and 11 (or Ib) of the periodic table tend to lose one electron to form singly positive ions; those in Groups 2 (or IIa) and 12 (or IIb) tend to lose two electrons to form doubly positive ions; and similarly for Groups 3 (or IIIb) and 13 (or IIIa). Likewise, the halogens, Group 17 (or VIIa), tend to gain one electron to form singly negative ions, and elements of Group 16 (or VIa) tend to gain two electrons to form doubly negative ions. As the net charge on an ion increases, however, the ion becomes less stable with respect to sharing electrons with other atoms, so most large apparent charges (as in MnO2, +4 and -2, respectively) would be minimized by covalent sharing of electrons.

   Covalent bonds form when both atoms lack the number of electrons in the nearest noble gas atom. Neutral chlorine atoms, for example, have one less electron per atom than do argon atoms (35 versus 36). When two chlorine atoms form a covalent bond sharing two electrons (one from each atom), both achieve the argon number of 36, Cl:Cl. It is common to represent a shared pair of electrons by a straight line between the atom symbols: Cl:Cl is written Cl-Cl.

   Similarly, atomic nitrogen is three electrons short of the neon number (ten), but each nitrogen can get the neon number if six electrons are shared between them: N…N or N؟N. This is called a triple bond. Sulphur, in the same way, can achieve the argon number by sharing four electrons in a double bond, S;S or SجS. In carbon dioxide, both the carbon (with six of its own electrons) and oxygen (with eight) achieve the neon number (ten) by sharing with double bonds: O=C=O. In all these bonding formulas, only the shared electrons are shown.

Valence 
   In most atoms, many of the electrons are so firmly attracted to their own nucleus that they can have no appreciable interaction with other nuclei. Only those electrons on the "outside" of an atom can interact with two or more nuclei. These are called valence electrons.

   The number of valence electrons in an atom is indicated by the atom's periodic table family (or group) number, using only the older Roman numeral designation. Thus we have one valence electron for elements in Groups 1 (or Ia) and 11 (or Ib). There are two valence electrons for elements in Groups 2 (or IIa) and 12 (or IIb), and four for elements in Groups 4 (or IVb) and 14 (or IVa). Each of the noble gases except helium (that is, neon, argon, krypton, xenon, and radon) has eight valence electrons. Elements in families (groups) near the noble gases tend to react to form noble gas sets of eight valence electrons. This is known as the Lewis Rule of Eight, which was enunciated by the American chemist Gilbert N. Lewis.

   The exception, helium (He), has a set of two valence electrons. Elements near helium tend to acquire a valence set of two: hydrogen by gaining one electron, lithium by losing one, and beryllium by losing two electrons. Hydrogen typically shares its single electron with one electron from another atom to form a single bond; such as in hydrogen chloride, H-Cl. The chlorine, originally with seven valence electrons, now has eight. These valence electrons can be shown as or . The structures of N2 and CO2 may now be expressed as or and or . These so-called Lewis structures show noble gas valence electron sets of eight for each atom. Probably 80 per cent of all covalent compounds can be reasonably represented by Lewis electron structures. The remainder, especially those containing elements in the central region of the periodic table, often cannot be described in terms of noble gas structures.

Resonance 
   An interesting extension of Lewis structures, called resonance, is found, for example, in nitrate ions, NO3-. Each N originally has five valence electrons, each O has six, plus one for the negative charge, or a total of 24 (5 + [3 × 6] + 1 = 24) valence electrons for four atoms. This is only an average of six valence electrons per atom, so covalent sharing must occur if the Lewis Rule of Eight is to apply. It is known that the nitrogen atom takes a central position surrounded by the three oxygen atoms, which can give an acceptable Lewis structure, except that there are three possible structures. Actually only one structure is observed. Each Lewis resonance structure suggests that two bonds should be single and one double. Experiments have shown, however, that all the bonds are actually identical in every respect, with properties intermediate between those observed for single and double bonds in other compounds. Modern theory suggests that a structure of localized, Lewis-type, shared electron bonds gives the general shape and symmetry of the molecule plus a set of delocalized electrons (shown by dotted lines) that are shared over the whole molecule.

[Contact us] [Our resources] [Feedback] [Terms of use]