Polarity
As in the case of two people who decide to share some commodity for the purpose of achieving a larger goal and who must then decide the degree to which the common commodity will be shared equally or unequally between them (that commodity might be money, land, pizza, or an apartment), atoms that come together in covalent bonds also “negotiate” the degree to which their sharing of electron pair(s) will be equal or unequal. The nature and degree of sharing between the nuclei of two atoms in a covalent bond is determined by the relative difference in their respective electronegativities, with the atom of higher electronegativity getting the larger “share” of the electron pair(s). A polar bond is a dipole, with the positive end of the dipole at the less electronegative atom and the negative end at the more electronegative atom.
Nonpolar Covalent Bond
When atoms that have identical or nearly identical electronegativities share electron pair(s), they do so with equal distribution of the electrons. This is called a nonpolar covalent bond, and there is no separation of charge across the bond. Of course, only bonds between atoms of the same element will have exactly the same electronegativity and, therefore, share with perfectly equal distribution the pair(s) of electrons in the covalent bond. Examples of diatomic molecules include H2, Cl2, O2, and N2. At the same time, many bonds can be said to be approximately nonpolar. For example, the electronegativity difference between carbon and hydrogen is so sufficiently small that we can usually consider the C–H bond to be effectively nonpolar.
Polar Covalent Bond
Of course nobody really likes to share anything. We only say we do because we know that’s what expected of us. When nobody’s looking, a child will grab a few extra candies for herself, and even adults are known to fight over property lines, the last slice of pizza, or whether someone paid his fair share of the dinner tab. Atoms of elements that differ moderately in their electronegativities will share their electrons unevenly, resulting in polar covalent bonds. While the difference in their electronegativities (typically between 0.4 and 1.7 Pauling units) is not enough to result in the formation of an ionic bond, it is sufficient to cause a separation of charge across the bond, with the more electronegative element acquiring a greater portion of the electron pair(s) and taking on a partial negative charge,
Figure 3.1
A molecule that has such a separation of positive and negative charges is called a polar molecule. The dipole moment of the polar bond or polar molecule is a vector quantity, µ, defined as the product of the charge magnitude (q) and the distance between the two partial charges (r):
µ = qr
The dipole moment vector, represented by an arrow pointing from the positive to the negative charge, is measured in Debye units (coulomb-meter). Please note that the convention used by chemists for designating the direction of the dipole moment from positive to negative is the opposite of the convention used by physicists, who designate the direction of a dipole moment from negative to positive.
COORDINATE COVALENT BOND
In a coordinate covalent bond, the shared electron pair comes from the lone pair of one of the atoms in the molecule, while the other atom involved in the bond contributes nothing to the relationship. (This works great for atoms but might not be the best way to form a lasting marriage!) Once such a bond forms, however, it is indistinguishable from any other covalent bond. The distinction is only helpful for keeping track of the valence electrons and formal charges (see Figure 3.2). Coordinate covalent bonds are typically found in Lewis acid-base compounds (see Chapter 10, Acids and Bases). A Lewis acid is any compound that will accept a lone pair of electrons, while a Lewis base is any compound that will donate a pair of electrons to form a covalent bond; for example, as in the reaction between borontrifluoride (BF3) and ammonia (NH3) shown in Figure 3.2.
Figure 3.2
NH3 donates a pair of electrons to form a coordinate covalent bond; thus, it acts as a Lewis base. BF3 accepts this pair of electrons to form the coordinate covalent bond; thus, it acts as a Lewis acid. Lewis acids, incidentally, are commonly encountered in some often-tested organic chemistry reactions, such as anti addition across double bonds, and as catalysts in electrophilic aromatic substitutions (EAS).
COVALENT BOND NOTATION
The electrons that are involved in a covalent bond are in the valence shell and are called bonding electrons, while those electrons in the valence shell that are not involved in the covalent bond are called nonbonding electrons. The unshared electron pairs can also be called lone electron pairs, because they are associated only with one atomic nucleus. Because atoms can bond with other atoms in many different possible combinations, the Lewis structure system of notation has been developed to keep track of the bonded and nonbonded electron pairs. You may think of Lewis structures as a kind of bookkeeping. The number of valence electrons attributed to a particular atom in the Lewis structure of a molecule is not necessarily the same as the number of valence electrons the atom would have as an isolated atom, and the difference accounts for what is referred to as the formal charge of that atom in a particular Lewis structure. Often, more than one Lewis structure can be drawn for a molecule. If the possible Lewis structures differ in the way the atoms are connected (that is to say, they differ in their bond connectivity or arrangement), then the Lewis structures represent different possible compounds. If the Lewis structures show the same bond connectivity and differ only in the arrangement of the electron pairs, then these structures represent different resonance forms for a single compound. Lewis structures do not represent the actual or even theoretical geometry of a real compound. Their usefulness lies in showing you the different possible ways in which atoms may be combined to form different compounds or resonances of a single compound. When more than one arrangement can be made, you can assess the likelihood of each arrangement by checking the formal charges on the atoms in each arrangement. The arrangement that minimizes the number of formal charges, or minimizes the value of formal charges, is the arrangement that most likely represents the actual compound.