Electronegativity Difference and Bond Character
Covalent bond - a shared pair of electrons between two atoms
Electronegativity - the ability of an atom to attract electrons in a covalent bond.
Electronegativity Difference, E, E = E1-E2,
For example, each H—–O bond in H2O has:
E = EO - EH
E = 3.44 - 2.10
E = 1.44
As the electronegativity difference grows from 0.00 to >3.00, the sharing of electrons becomes more skewed. First
one atom becomes partially negative, -, while the other becomes partially positive, +, and the bond is polarized, a
polar covalent bond. Then as the sharing becomes more and more unequal, the electrons reside only on one atom.
The atoms are now ionized (+ or -), and the bond is called ionic. This is shown in Table 1.
Table 1
↔ polar covalent
nonpolar covalent
Electronegativity
Difference, E
Percent Ionic Character
Percent Covalent
Character
↔
ionic
0.00
0.65
0.94
1.19
1.43
1.67
1.91
2.19
2.54
3.03
0%
10 %
20 %
30 %
40 %
50 %
60 %
70 %
80 %
90 %
100 %
90 %
80 %
70 %
60 %
50 %
40 %
30 %
20 %
10 %
For a shared pair of electrons, if one atom is able to attract the electrons to itself (more electronegative) that
atom will begin to become negatively charged, -, (a negative pole) while the other atom (least electronegative)
begins to become positively charge, +, (a positive pole). The two atoms become a dipole (meaning 2 poles), and
the bond will become a polar covalent bond. [see Table 2]
If the difference in attracting the electrons, E, is so great, then one atom may just take the electrons for
itself. This stops any sharing of electrons, and the bond is an ionic bond. The atom that took the electrons is the
anion (negative), and the atom that lost the electrons is the cation (positive). [see Table 2]
If the difference in attracting the electrons, E, is very small, then the sharing remains relatively equal and no
charges develop. No developed charges means there are no pole, which makes the bond a nonpolar covalent bond. In
the special case that the electronegativity difference, E, is zero, then no atom attracts the electrons to itself and the
sharing is perfectly equal. Such a bond is called a pure covalent bond and is nonpolar also. [see Table 2]
Table 2
Electronegativity Difference, E

E = 0.00
0.00 < E < 0.65
0.65 < E < 1.67
1.67 < E
Polarity
nonpolar
nonpolar
polar
Bond Type
pure covalent
nonpolar covalent
polar covalent
ionic
More on Dipoles
 An electric dipole consists of two opposite charges that are the same magnitude and separated in space.
 Unequal sharing of electrons (E > 0) results in one end of the bond being negative and one end being
positive. This is a dipole (2 poles, one negative, one positive)
 The atom with the greater electronegativity becomes partially negative, -,
 The atom with the smaller electronegativity becomes partially positive, +.
 Drawing an arrow from + to - shows the electric field between the charges.
Practice Problem: Hydrogen fluoride, HF
o Draw the Lewis Structure
o Determine the polarity and bond type for each bond.
o Draw the dipole for each bond.
HF, would be written as H—– F
Since E = 1.10, the bond is polar covalent.
E = 3.98 - 2.10 = 1.10}
F has the greater electronegative so it is partially negative, -, and H with the smaller electronegativity is
partially positive, +.
Final Answer:
H —–: F
+ → - polar covalent
Problem Set on Determining the Type of Bond
Name:
Period:
Directions
For each formula,
Date:
Chemistry
Dr. Mandes, Ph.D.
• Draw the Lewis Structure and geometry
(2 pts. each)
• Calculate the electronegativity difference, E, for each bond
(2 pts. each)
• Use the E to determine the bond type for each bond
(1 pt. each)
• For each bond, place the electron pair closest to the most electronegative atom (1 pt. each)
• For each bond, draw a “-“ or a “–“ at the end with the electrons and a “+“ or a “+“ at the end
without the electrons. If equal sharing, draw nothing.
(1 pt. each)
1. CBr4
2. NH3
3. H2O
4. KCl (draw the “K” like “H”)