|Polar Bonds and Electronegativity|
|When two identical atoms form a covalent bond, as in H2
or Cl2, each has an equal share of the electron pair in the bond.
The electron density at both ends of the bond is the same, because the electrons
are equally attracted to both nuclei.
|When different kinds of atoms combine, as in HCl, the attractions
usually are not equal. Generally one of the nuclei attracts the electrons
more strongly than the other.
|The effect of unequal attractions for the bonding electrons is an
unbalanced distribution of electron density within the bond. It has been found
that a chlorine atom attracts the electrons more strongly than does a hydrogen.
In the HCl molecule, therefore, the electron cloud is pulled more tightly
around the Cl, and that end of the molecule experiences a slight buildup
of negative charge. The electron density that shifts toward the chlorine
atom is removed from the hydrogen, which causes the hydrogen end to acquire
a slight positive charge.
|In HCl the electron transfer is incomplete. The electrons are still
shared, but unequally. The charges on either end of the molecule are less
than full +1 and -1 charges; they are partial charges, normally indicated
by the lowercase Greek letter delta.
|A bond that carries partial positive and negative charges on opposite
ends is called a polar bond, or polar covalent bond. The term polar
comes from the notion of poles of opposite charge at either end of the bond.
Because there are two poles of charge involved, the bond is said to be a
|The polar bond in HCl causes the molecules as a whole to have opposite
charges on either end, so we say that HCl is a polar molecule. The HCl molecule
as a whole is also a dipole.
|The degree to which a covalent bond is polar depends on the relative
abilities of bonded atoms to attract electrons. The term that we use to describe
this relative attraction of an atom for the electrons in a bond is called
the electronegativity of the atom. In HCl the chlorine is more electronegative
than hydrogen. The electron pair of the covalent bond spends more of its
time around the more electronegative atom, which is why that end of the bond
acquires a partial negative charge.
|The concept of electronegativity has been put on a quantitative basis
- as numerical values have been assigned for each element. Refer to your
periodic table for these assigned values. This information is useful because
the difference in electronegativity values provides an estimate of
the degree of polarity of a bond. In addition, the relative magnitudes of
the electronegativities indicate which end of the bond carries the negative
charge. For instance, fluorine is more electronegative than chlorine. Therefore,
we would expect an HF molecule to be more polar than an HCl molecule. In addition,
hydrogen is less electronegative than either fluorine or chlorine, so in
both of these molecules the hydrogen bears the positive charge.
| .. ..
+ - + -
|The concept of electronegativity shows that there is no sharp dividing
line between ionic and covalent bonding.
|Ionic bonding and nonpolar covalent bonding simply represent the two
extremes. Ionic bonding occurs when the difference in electronegativity between
two atoms is very large; the more electronegative atom acquires essentially
complete control of the bonding electrons. In a nonpolar covalent bond, there
is no difference in electronegativity, so the pair of bonding electrons is
|The degree of polarity, which is the amount of ionic character of
the bond varies in a continuous way with changes in the electronegativity
difference. The bond becomes more than 50% ionic when the electro-negativity
difference exceeds 1.7
|Within the periodic table, electronegativity varies in a more of less
systematic way, and the trends follow those for the ionization energy (IE).
Atoms with large ionization energies also have large electronegativities.
An atom that has a small IE will give away an electron more easily than an
atom with a large IE, just as an atom with a small electronegativity will
lose its share of an electron pair more readily than an atom with a large