Basic Concepts of Chemical Bonding

Bond Polarity and Electronegativity


Bond Polarity and Electronegativity

The electron pairs shared between two atoms are not necessarily shared equally

Extreme examples:

1. In Cl2 the shared electron pairs is shared equally

2. In NaCl the 3s electron is stripped from the Na atom and is incorporated into the electronic structure of the Cl atom - and the compound is most accurately described as consisting of individual Na+ and Cl- ions

For most covalent substances, their bond character falls between these two extremes

Bond polarity is a useful concept for describing the sharing of electrons between atoms

Electronegativity

A quantity termed 'electronegativity' is used to determine whether a given bond will be nonpolar covalent, polar covalent, or ionic.

Electronegativity is defined as the ability of an atom in a particular molecule to attract electrons to itself

(the greater the value, the greater the attractiveness for electrons)

Electronegativity is a function of:

(Note that both of these are properties of the isolated atom)

For example, an element which has:

Will:

Such an atom will be highly electronegative

Fluorine is the most electronegative element (electronegativity = 4.0), the least electronegative is Cesium (notice that are at diagonal corners of the periodic chart)

General trends:

Electronegativity and bond polarity

We can use the difference in electronegativity between two atoms to gauge the polarity of the bonding between them

Compound

F2

HF

LiF

Electronegativity Difference

4.0 - 4.0 = 0

4.0 - 2.1 = 1.9

4.0 - 1.0 = 3.0

Type of Bond

Nonpolar covalent

Polar covalent

Ionic (non-covalent)

The sharing of electrons in HF is unequal: the fluorine atom attracts electron density away from the hydrogen (the bond is thus a polar covalent bond)

The H-F bond can thus be represented as:

A general rule of thumb for predicting the type of bond based upon electronegativity differences:


1996 Michael Blaber