Electronegativity and Bond Polarity

Learning Objectives (OCR Spec 2.2.2)

By the end of this lesson you should be able to:

• Define electronegativity as the ability of an atom to attract the bonding pair of electrons in a covalent bond
• Describe trends in electronegativity across a period and down a group
• Use the Pauling scale to classify bonds as non-polar covalent, polar covalent or ionic
• Identify polar bonds and draw delta+/delta- charges or dipole arrows
• Decide whether a whole molecule is polar based on bond polarity and molecular shape

Defining Electronegativity

Electronegativity is the ability of an atom to attract the pair of electrons in a covalent bond toward itself. The concept was quantified by Linus Pauling in 1932; the most common scale bears his name.

Element	Pauling Electronegativity
F	3.98 (the most electronegative)
O	3.44
N	3.04
Cl	3.16
C	2.55
H	2.20
Si	1.90
Na	0.93
Cs	0.79

Trends in the Periodic Table

Electronegativity increases:

• Across a period (left to right): nuclear charge increases while atomic radius decreases, so the bonding pair is held closer and more strongly
• Up a group (bottom to top): atomic radius decreases and there is less shielding, so the bonding pair is held more strongly

Consequently the most electronegative element is fluorine (top right, excluding noble gases), and the least is caesium/francium (bottom left).

Noble gases are not typically assigned values because they rarely form bonds.

Non-polar vs Polar Bonds

When two atoms of similar electronegativity share electrons, the bonding pair sits symmetrically between them - a non-polar covalent bond. Examples include:

• H-H, Cl-Cl, N::N (same element - electronegativity difference = 0)
• C-H (difference = 0.35 - considered effectively non-polar)

When two atoms of different electronegativity bond, the more electronegative atom gains a greater share of the bonding electrons. It develops a partial negative charge (delta-) and the other atom a partial positive charge (delta+). This is a polar covalent bond with a dipole moment.

Bond	Difference	Classification
H-H	0	Non-polar
C-H	0.35	Essentially non-polar
C-Cl	0.61	Polar
H-Cl	0.96	Polar
O-H	1.24	Very polar
Na-Cl	2.23	Ionic (so polar that an ion forms)

As a rough guide:

• Difference < 0.4: non-polar covalent
• 0.4 - ~1.7: polar covalent

• 1.7: essentially ionic

Remember: this is a continuum, not three separate categories.

Representing Polar Bonds

There are two common conventions:

1. delta+/delta- charges placed above the atoms:

   delta+     delta-
      H --- Cl

2. Dipole arrow pointing from delta+ to delta- (the arrow has a cross at its tail):

      H --- Cl
      +---->

Polar Bonds vs Polar Molecules

A molecule is polar if it has an overall dipole moment. This depends on two things:

1. Presence of polar bonds
2. The symmetry (shape) of the molecule

Rule: If bond dipoles cancel by symmetry, the molecule is non-polar overall even though individual bonds are polar.

Examples of Non-Polar Molecules with Polar Bonds

CO2 (O=C=O, linear)

Each C=O bond is strongly polar (O delta-, C delta+). But the two dipoles point in exactly opposite directions and are equal in magnitude, so they cancel. CO2 is non-polar overall.

 delta-   delta+   delta-
   O === C === O        (linear, dipoles cancel)

CCl4 (tetrahedral)

Four polar C-Cl bonds arranged tetrahedrally cancel vectorially. CCl4 is non-polar.

BF3 (trigonal planar)

Three polar B-F bonds at 120 degrees cancel by symmetry. Non-polar overall.

Examples of Polar Molecules

H2O (bent, 104.5 degrees)

Two polar O-H bonds. Because the molecule is bent, the dipoles do not cancel - they add to give a net dipole pointing from the H side to the O. H2O is strongly polar.

NH3 (trigonal pyramidal, 107 degrees)

Three polar N-H bonds plus a lone pair on N. Dipoles add to give a net dipole along the lone pair direction. Polar.

CHCl3 (chloroform, tetrahedral)

Three C-Cl bonds (polar) plus one C-H bond. The four bonds are arranged tetrahedrally but are no longer identical - the C-H is much less polar than the C-Cl bonds - so dipoles do not cancel. Polar.

Deciding If a Molecule Is Polar

Use this algorithm: