Electronegativity and Polarity

This lesson covers electronegativity, polar bonds, dipole moments, and the distinction between polar and non-polar molecules. Understanding polarity is essential for explaining intermolecular forces and physical properties (AQA specification 3.1.3).

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Electronegativity

Key Definition: Electronegativity is the ability of an atom to attract the shared pair of electrons in a covalent bond towards itself.

Electronegativity is measured on the Pauling scale, which ranges from about 0.7 (francium) to 4.0 (fluorine).

Pauling Electronegativity Values

Element	Symbol	Electronegativity
Fluorine	F	4.0
Oxygen	O	3.5
Nitrogen	N	3.0
Chlorine	Cl	3.0
Bromine	Br	2.8
Carbon	C	2.5
Sulfur	S	2.5
Hydrogen	H	2.1
Phosphorus	P	2.1
Sodium	Na	0.9
Potassium	K	0.8

Trends in Electronegativity

Across a period (left to right): Electronegativity increases because:

• Nuclear charge increases
• Atomic radius decreases
• The bonding pair is more strongly attracted to the nucleus

Down a group: Electronegativity decreases because:

• Atomic radius increases
• More inner electron shells shield the nucleus
• The bonding pair is less strongly attracted to the nucleus

Exam Tip: Fluorine is the most electronegative element (4.0). The trend is often summarised as: electronegativity increases towards the top-right of the periodic table (excluding noble gases, which do not normally form bonds).

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Polar Bonds

When two atoms with different electronegativities form a covalent bond, the shared pair of electrons is pulled more towards the more electronegative atom. This creates a polar bond — a bond with a partial charge separation.

The more electronegative atom carries a partial negative charge (δ−) and the less electronegative atom carries a partial positive charge (δ+).

    δ+     δ−
    H ——→ Cl

The arrow (or δ+/δ− notation) indicates the direction of the dipole.

Non-Polar Bonds

If two identical atoms are bonded (e.g., H—H, Cl—Cl, O═O), the electronegativity difference is zero and the bond is non-polar. The electrons are shared equally.

If the electronegativity difference is very small (e.g., C—H, difference = 0.4), the bond is considered essentially non-polar.

Guidelines for Bond Type

Electronegativity Difference	Bond Type
0	Pure covalent (non-polar)
0.1 – 0.4	Weakly polar covalent
0.5 – 1.7	Polar covalent
> 1.7	Ionic (electron transfer)

Note: These boundaries are approximate. The transition from polar covalent to ionic is gradual, not abrupt. Some bonds with differences around 1.7 may be considered either highly polar covalent or ionic depending on context.

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Dipole Moments

A dipole moment is a measure of the polarity of a bond or molecule. It depends on:

1. The magnitude of the partial charges (δ+ and δ−)
2. The distance between the charges

Dipole moment (μ) = charge (q) × distance (d)

For a molecule, the overall dipole moment is the vector sum of all the individual bond dipoles.

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Polar and Non-Polar Molecules

A molecule may contain polar bonds but still be non-polar overall if the bond dipoles cancel due to molecular symmetry.

Polar Molecules (Net Dipole ≠ 0)

HCl:

    δ+     δ−
    H ———→ Cl

    Net dipole: →

One polar bond; the molecule has a net dipole moment. HCl is a polar molecule.

Water (H₂O):

        δ−
      O
     / \
   δ+H   Hδ+
    104.5°

Both O—H bonds are polar (pointing towards O). Because the molecule is bent, the two bond dipoles do not cancel. Water has a net dipole moment pointing from the H atoms towards the O atom. Water is a polar molecule.

Ammonia (NH₃):

       δ−
       N
      /|\
     H H H
     δ+ (all three)

Three polar N—H bonds in a pyramidal shape. The bond dipoles do not cancel. NH₃ is a polar molecule.

Non-Polar Molecules with Polar Bonds

Carbon dioxide (CO₂):

    δ−        δ+        δ−
    O ←——— C ———→ O

    Dipoles are equal and opposite → cancel → net dipole = 0