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Each positive nucleus is attracted toward the region of high electron density between them:
1s2 2s2 2p6 3s2 3p2 3p2 3p1 3p1 3p2 3p2 3s2 2p6 2s2 1s2
The bonded atoms come close enough together for their electron clouds to overlap.
Covalent bonds form between atoms when ionic or metallic bonding is unlikely because the gain or loss of electrons requires large amounts of energy.
The most obvious example is the bonding of nonmetal atoms with themselves (as in Cl2, H2, O2) and with each other (as in HCl, H2O, BrF3).
In bond formation, the semiconducting elements tend to behave like the nonmetals, forming covalent bonds in many compounds with nonmetals (eg. BCl3, SiF4).
The number of valence electrons and the octet rule govern the formation of many covalent compounds. For second-period elements (Li to F), which have only s and p orbitals available, 8 is the maximum number of valence electrons that can be accommodated.
Many stable compounds also exist in which the central atom is not surrounded by 8 electrons. Most of these fall into three categories;
1. Certain molecular compounds of beryllium, boron, or aluminium. Be has two valence electrons (2s2) while B (2s22p1) and Al (3s23p1) have three. Therefore, these elements can form compounds with two or three covalent bonds, respectively, with no remaining unshared electron pairs.
2. Certain molecular compounds of phosphorus, sulfur, chlorine, or other elements from the third period and beyond. In atoms with n = 3 or more, d orbitals are available. For elements of the 3rd to 6th periods, the outermost d subshells are empty. Therefore, these elements can accommodate more than 8 valence electrons if some of them utilise vacant outermost d subshells.
3. Compounds with unpaired electrons. Some compounds exist in which one or more electrons remain unpaired. In such cases the total of the valence electrons of the central atom and the atoms bonded to it is an odd number. Chlorine dioxide, for example, has a total of 19 valence electrons (6 from each of the two oxygen atoms and 7 from the chlorine atom).
A reasonable Lewis structure would be;
Co-ordinate Covalent Bonds
a single covalent bond in which both electrons in a shared pair come from the same atom. The donor atom provides both electrons to a co-ordinate covalent bond, and the acceptor atom accepts an electron pair for sharing a co-ordinate covalent bond.
For example: a hydrogen ion unites with an ammonia molecule by a co-ordinate covalent bond to form the ammonium ion;
All four hydrogen atoms and all four N-H bonds in the ammonium ion are found by experiment to be equivalent.
Frequently nonmetal atoms that are already part of molecules or ions form co-ordinate covalent bonds with metal atoms or ions, usually those of transition metals. For example, the nitrogen atom in ammonia can "co-ordinate" with the silver cation to form what is called a complex ion;
Neutral compounds with similar co-ordinate covalent bonding can also be formed, notably by carbon monoxide and metal atoms;
The atoms in polyatomic ions such as hydroxide ion, ammonium ion and sulfate ion, are held together by covalent bonds;
Polyatomic ions are charged because they have fewer or more electrons than are needed to balance the positive charges of the nuclei present in the ion.
Nonpolar and Polar Covalent Bonds
In molecules such as H2, Cl2, and N2, the electron density (the probability of finding the valence electrons in a given area) is equally divided between the two bonded atoms. In a covalent bond of this type - a nonpolar covalent bond - the electrons are equally shared.
Whenever atoms of the two different elements are covalently bonded, the sharing of the electrons becomes unequal, because no two atoms of different elements have exactly the same electron-attracting ability. The electron density around one atom becomes greater than that around the other.
How unequally the electrons are shared depends on the relative abilities of the two different atoms to attract electrons.
A covalent bond in which electrons are shared unequally is called a polar covalent bond; one atom acquires a partial negative charge (d -) and the other acquires a partial positive charge (d +).
These are not unit charges, but only represent a reorientation, a sort of pushing around, of the total electron density of the two bonded atoms. The entire molecule remains electrically neutral.
A polar molecule is a dipole - a pair of opposite charges of equal magnitude at a specific distance from each other.