Electronegativity and Pauling values

Electronegativity is about attraction inside a covalent bond. When two atoms share a pair of electrons, the atom with the higher electronegativity attracts the bonding electrons more strongly.

The Pauling scale gives electronegativity as a relative number with no unit. A higher Pauling value means a stronger attraction for the bonding pair. You do not need to memorise a full table of values, but you must be able to interpret values when they are given.

Electronegativity increases towards fluorine in the periodic table. Fluorine has the highest electronegativity, so a bond to fluorine often has a strongly polarised bonding pair.

Worked example: interpreting Pauling values

A question gives these Pauling electronegativity values:

In an O-H bond, oxygen has the higher value, so oxygen attracts the bonding electrons more strongly than hydrogen. In a C-F bond, fluorine has the higher value, so fluorine attracts the bonding electrons more strongly than carbon.

The size of the difference also matters. The C-F difference is:

The C-H difference is:

The C-F bond is therefore much more polar than the C-H bond.

Using electronegativity to predict bond type

Electronegativity helps predict the type of bonding between two atoms. Treat it as a model, not as a sharp rule with a magic cut-off.

If two bonded atoms have the same or very similar electronegativities, the bonding electrons are shared almost equally. This is non-polar covalent bonding, as in Cl-Cl.

If the atoms have different electronegativities but are still sharing electrons, the bond is polar covalent. The bonding pair is drawn closer to the more electronegative atom, as in H-Cl or C-O.

If the difference in electronegativity is very large, electron transfer becomes more likely and the bonding is usually described as ionic. For example, sodium and chlorine have very different tendencies to attract electrons, so sodium chloride is treated as ionic.

Worked example: predicting the most polar bond

A question gives these values:

Compare H-Cl and H-Br:

H-Cl is predicted to be the more polar bond because it has the larger electronegativity difference.

Polar bonds and permanent dipoles

When covalently bonded atoms have different electronegativities, the bonding electrons are not shared equally. The more electronegative atom gains more electron density and becomes partially negative. The less electronegative atom becomes partially positive.

These are partial charges, written as delta negative and delta positive: δ- and δ+. They are not full ionic charges.

A polar bond has a bond dipole. The dipole points towards the more electronegative atom, because that end is partially negative.

For hydrogen chloride, chlorine is more electronegative than hydrogen:

The H-Cl molecule has a permanent dipole because its polar bond gives one end of the molecule a permanent partial positive charge and the other end a permanent partial negative charge.

Worked example: assigning partial charges

In a C-F bond, fluorine has a higher electronegativity than carbon. The bonding electrons are pulled closer to fluorine. Therefore:

The fluorine end is δ- and the carbon end is δ+.

Polar molecules and overall dipoles

A molecule can contain polar bonds without being a polar molecule. To decide whether the whole molecule is polar, you must consider both:

1. whether the molecule contains polar bonds
2. whether the directions of the bond dipoles cancel because of the molecular shape

Carbon dioxide and water show the difference clearly.

In CO2, each C=O bond is polar because oxygen is more electronegative than carbon. However, CO2 is linear. The two equal C=O dipoles point in opposite directions, so the dipoles cancel. CO2 has no overall dipole and is non-polar.

In H2O, each O-H bond is polar because oxygen is more electronegative than hydrogen. Water is bent, so the two O-H bond dipoles do not point in exactly opposite directions. The dipoles do not cancel, so H2O has an overall dipole and is polar.

The exam distinction is precise: dipoles cancel. Do not write that "polar bonds cancel" or "charges cancel". The bonds are still polar; it is their dipole effects that cancel in a symmetrical molecule.

Worked example: CF4

Fluorine is more electronegative than carbon, so each C-F bond in CF4 is polar:

CF4 is tetrahedral and symmetrical. The four C-F bond dipoles are arranged evenly around the carbon atom, so their effects cancel. CF4 has no overall dipole, so it is a non-polar molecule despite containing polar bonds.

Writing clear polarity explanations

A strong polarity explanation usually has two linked sentences:

1. State why the bond or bonds are polar, using electronegativity.
2. State whether the bond dipoles cancel, using the molecular shape.

For a polar molecule, the second sentence should say that the molecule is not symmetrical or that the dipoles do not cancel. For a non-polar molecule with polar bonds, the second sentence should say that the molecule is symmetrical or that the dipoles cancel.

Worked example: explaining SO2

Sulfur dioxide, SO2, is a bent molecule. Oxygen is more electronegative than sulfur.

Full explanation:

Oxygen is more electronegative than sulfur, so each S-O bond is polar. SO2 is bent/non-linear, so the S-O bond dipoles do not cancel. Therefore SO2 has an overall dipole and is a polar molecule.

Common error contrast

Incorrect: "SO2 is polar because the charges do not cancel."

Better: "SO2 is polar because its polar S-O bond dipoles do not cancel in the bent molecule."

This wording matters because partial charges are not separate full charges moving around the molecule. The model being assessed is permanent bond dipoles plus molecular shape.