Chemistry - A Molecular Science

H

F

I

F

I

H

Dc

= 1.8

Dc

= 1.3

Dc

= 0.5

I

Dc I

=0

a)

b)

c)

d)

dd+-HCl Figure 5.3 Charge distribution in some covalent bonds Red is used to show regions that carry negative charge, while blue is used for regions that carry positive charge. The dipole arrow is used to show direction only.
Figure 5.2 Bond dipole in HCl The bond dipole is shown as the blue arrow. Cl is more electronegative than H, so the H-Cl bond dipole points from H to Cl. A line perpendicular to the a

rrow is placed at the positive end

to produce a ‘+’ in the arrow.

as the more electronegative atom becomes electron rich and the less electronegative atom becomes electron poor. The result is that negative charge is produced on the more electronegative atom and positive charge on the less electronegative atom. The charge that is produced in a covalent bond is only a part of what it would be if the bond were ionic, so it is said to be a

partial charge,

which is represented with a

(delta). Thus, a covalent δ

bond involving atoms of different electronegativities has two electrical poles (

• and δ

+) δ

and is said to have a

bond

di

pole

. The bond dipole is represented by a vector pointing

from the center of positive charge toward the

center of negative charge with a line drawn

through the positive end to make a ‘+’ as shown in Figure 5.2. The strength of a bond dipole (the polarity of the bond) increases as the electronegativity difference (

Δχ

) between

the bound atoms increases. Consider the four molecules in Figure 5.3. Figure 5.3a

: The two iodine atoms in I

have identical electronegativities, so 2

Δχ

= 0 and

δ = 0.

The atoms attract the bonding electrons eq

ually and form a purely covalent bond. There

is no charge separation in purely covalent bonds (no red and blue regions in the figure).

Figure 5.3b

: Iodine (

χ = 2.7) is more electronegative than hydrogen (

χ = 2.2), so the bonding

electrons in an H-I bond are pulled away from the hydrogen and toward the iodine atom. As a result, the iodine atom is electron rich

and has a partial negative charge, while the

less electronegative hydrogen atom develops a partial positive charge. H-I has a bond dipole, but the bond is not very polar (as indicated by the very pale shading) because

Δχ

is only 0.5. The bond dipole in HI points from H to I.

Figure 5.3c

: Δχ

= 1.3 for an I-F bond, so it is more po

lar than an H-I bond (as indicated by the

darker shading). Fluorine is the more electronegative atom in this bond, so the bond dipole points toward fluorine.

Figure 5.3d

: Δχ

= 1.8 for an H-F bond, so the H-F bond is more polar than an I-F bond (darker

shading). The bond dipole points from H to F.

The bond between two atoms is purely covalent if their electronegativities are

identical, but it becomes increasingly polar as

the difference in their electronegativities

increases. At a sufficiently large

difference, the bond is so polar that it is considered to be

an ionic bond. The nature of the bond is sometimes given in terms of its

percent ionic

character

. The more polar a bond is, the higher is

its percent ionic character. An F-F

bond (

Δχ

= 0) is purely covalent and has no ionic character (0% ionic), while a RbF bond

(Δχ

= 3.2) is 90% ionic, which means that the charge (

) on the rubidium is +0.9. δ

As shown in Figure 5.4, the change from covalent to ionic is gradual with no distinct

line between the two bond types. In the fo

llowing discussion, we will assume that a bond

is covalent if it is less than 5% ionic, that a bond is polar covalent if it is between 5 and

Chapter 5 The Covalent Bond

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