Electronegativity and Bond Polarity — AQA A-level Chemistry

SPEC 3.1.3.6 • BONDING

Electronegativity & Bond Polarity

Understanding how atoms attract electrons and why some molecules are polar

What is Electronegativity?

Definition

Electronegativity is the power of an atom to attract the pair of electrons in a covalent bond towards itself.

When two atoms form a covalent bond, they share a pair of electrons. However, if one atom is more electronegative than the other, it will attract the shared electrons more strongly, pulling the electron density towards itself. This means the electron distribution in a covalent bond between elements with different electronegativities will be unsymmetrical.

Why Does This Happen?

This phenomenon arises from the positive nucleus's ability to attract the negatively charged electrons in the outer shells towards itself. Atoms with more protons (higher nuclear charge) and smaller atomic radii can attract bonding electrons more strongly.

The Pauling Scale

Electronegativity is measured on the Pauling scale, named after Linus Pauling. Values range from about 0.7 (caesium) to 4.0 (fluorine). Fluorine is the most electronegative element in the periodic table – it is best at attracting electron density towards itself when covalently bonded to another atom.

ElementFOClNCH

Electronegativity4.03.43.23.02.52.2

Factors Affecting Electronegativity

Three main factors determine how strongly an atom can attract bonding electrons:

1. Nuclear Charge

Attraction exists between the positively charged protons in the nucleus and negatively charged electrons in the energy levels. An increase in the number of protons leads to an increase in nuclear attraction for the outer electrons. Therefore, increased nuclear charge results in increased electronegativity.

2. Atomic Radius

The atomic radius is the distance between the nucleus and electrons in the outermost shell. Electrons closer to the nucleus are more strongly attracted towards its positive charge. Those electrons further away from the nucleus are less strongly attracted. Therefore, increased atomic radius results in decreased electronegativity.

3. Shielding

Filled inner energy levels can shield (mask) the effect of the nuclear charge, causing the outer electrons to be less attracted to the nucleus. The addition of extra shells and subshells in an atom will cause the outer electrons to experience less of the attractive force of the nucleus. Therefore, increased shielding results in decreased electronegativity.

Example: Sodium vs Caesium

Sodium (Period 3, Group 1) has a higher electronegativity than caesium (Period 6, Group 1). Although caesium has more protons, it also has many more electron shells, resulting in greater shielding and a much larger atomic radius. The increase in shielding and distance outweighs the increase in nuclear charge, so caesium attracts bonding electrons less strongly than sodium.

Link to Ionisation Energy

These are the same factors that affect ionisation energy! Atoms with high ionisation energies tend to have high electronegativities because both depend on how strongly the nucleus attracts electrons.

Trends in Electronegativity

Across a Period (L → R)

Electronegativity increases

Nuclear charge increases (more protons added)

Shielding remains relatively constant (no new shells added)

Atomic radius decreases (stronger attraction pulls electrons closer)

Greater attraction between nucleus and bonding electrons

Down a Group

Electronegativity decreases

Nuclear charge increases BUT...

Each element has an extra filled electron shell

This increases shielding significantly

Larger atomic radius = less attraction for bonding electrons

Overall Trend

The general trend is an increase in electronegativity towards the top right of the periodic table. Fluorine (F = 4.0) is the most electronegative element because it has a high nuclear charge, small atomic radius, and minimal shielding.

Worked Example01

Explain why chlorine is more electronegative than sodium.

1

Chlorine and sodium are in the same period (Period 3), so they have the same number of electron shells and similar shielding.

2

Chlorine has 17 protons; sodium has 11 protons. Chlorine has a higher nuclear charge.

3

Chlorine has a smaller atomic radius than sodium (fewer electrons, stronger pull inward).

4

The higher nuclear charge and smaller atomic radius mean chlorine attracts bonding electrons more strongly.

AnswerChlorine has a higher nuclear charge and smaller atomic radius than sodium, with similar shielding. This means the nucleus of chlorine attracts bonding electrons more strongly, giving it a higher electronegativity.

Polar and Non-Polar Bonds

When two atoms in a covalent bond have different electronegativities, the bonding electrons are not shared equally. This creates a polar bond.

Partial Charges (δ)

In a polar bond, the more electronegative atom gains a partial negative charge (δ−) and the less electronegative atom gains a partial positive charge (δ+). The symbol δ (delta) means "partial" – these are NOT full ionic charges.

Worked Example02

Draw the C–O bond showing partial charges and explain why it is polar.

1

Compare electronegativities: Oxygen (3.4) is more electronegative than carbon (2.5).

2

The bonding electrons are pulled towards oxygen, giving it a partial negative charge (δ−).

3

Carbon loses some electron density, giving it a partial positive charge (δ+).

AnswerThe C–O bond is polar because oxygen is more electronegative than carbon. Oxygen attracts the bonding electrons more strongly, becoming δ−, while carbon becomes δ+.

Polar vs Non-Polar Molecules

Having polar bonds does not automatically make a molecule polar. For a molecule to be polar, it must have:

Requirements for a Polar Molecule

Polar bonds – a difference in electronegativity between bonded atoms

Asymmetrical arrangement – the bond dipoles do NOT cancel out

If a molecule has polar bonds but they are arranged symmetrically, the individual dipoles cancel out, resulting in a non-polar molecule.

Common Exam Mistake

Students often assume that molecules with polar bonds are always polar. Always check the molecular shape! Symmetrical molecules like CO₂, BF₃, and CCl₄ have polar bonds but are non-polar overall because the dipoles cancel.

C–H Bonds

Carbon (2.5) and hydrogen (2.2) have very similar electronegativities. C–H bonds are treated as essentially non-polar at A-Level.

Examples: Polar & Non-Polar

MoleculeShapePolar Bonds?Symmetrical?Polar Molecule?

H₂OBentYes (O–H)NoYes

CO₂LinearYes (C=O)YesNo

CH₄TetrahedralNo (C–H ~non-polar)YesNo

CHCl₃TetrahedralYes (C–Cl)NoYes

Worked Example03

Explain why CO₂ is non-polar even though it contains polar C=O bonds.

1

Identify the bonds: CO₂ has two C=O double bonds. Oxygen is more electronegative than carbon, so each C=O bond is polar with δ+ on carbon and δ− on oxygen.

2

Determine the shape: CO₂ has 2 bonding pairs and 0 lone pairs on the central carbon. Using VSEPR, the shape is linear with a 180° bond angle.

3

Check for symmetry: The molecule is symmetrical – both oxygen atoms are on opposite sides of the carbon.

4

Consider the dipoles: The two bond dipoles point in exactly opposite directions and are equal in magnitude, so they cancel out.

AnswerCO₂ is non-polar because although it has polar C=O bonds, its linear shape means the two bond dipoles are equal and opposite. They cancel each other out, resulting in no net dipole moment.

Worked Example04

Deduce whether CHCl₃ (chloroform) is a polar or non-polar molecule. Explain your reasoning.

1

Identify the bonds: CHCl₃ has one C–H bond and three C–Cl bonds. Chlorine is more electronegative than carbon, so each C–Cl bond is polar (δ+ on C, δ− on Cl). The C–H bond is essentially non-polar.

2

Determine the shape: Carbon has 4 bonding pairs and 0 lone pairs. The shape is tetrahedral.

3

Check for symmetry: CHCl₃ is NOT symmetrical – there are three Cl atoms and one H atom around the central carbon. The three C–Cl dipoles do not cancel out with the C–H bond.

4

Conclusion: The dipoles do not cancel, so there is a net dipole pointing from H towards the three Cl atoms.

AnswerCHCl₃ is polar. It has a tetrahedral shape with three polar C–Cl bonds. Because the molecule is asymmetrical (three Cl atoms and one H atom), the bond dipoles do not cancel, giving a net dipole moment.

Determining Molecular Polarity

Follow this systematic approach for exam questions asking whether a molecule is polar:

3-Step Method

1Identify polar bonds – Look for atoms bonded to N, O, F, or Cl (the most electronegative elements)2Determine the shape – Use VSEPR theory (from Lesson 3.04) to find the molecular geometry3Check if dipoles cancel – Symmetrical arrangements cancel; asymmetrical arrangements give a net dipole

Symmetrical Shapes = Non-Polar

These symmetrical arrangements result in non-polar molecules (if all surrounding atoms are identical):

Linear (2 groups) – e.g., CO₂, BeCl₂

Trigonal planar (3 groups) – e.g., BF₃, AlCl₃

Tetrahedral (4 identical groups) – e.g., CH₄, CCl₄

Exam Tip

If a molecule has lone pairs on the central atom, it's usually asymmetrical and therefore polar. Examples: H₂O (2 lone pairs), NH₃ (1 lone pair).

Practice

Question 1: Arrange these elements in order of increasing electronegativity: K, Mg, S, Cl. Explain your reasoning.

1

Identify positions: K is in Period 4, Group 1. Mg is in Period 3, Group 2. S is in Period 3, Group 6. Cl is in Period 3, Group 7.

2

Apply trend across periods: For Period 3 elements (Mg, S, Cl), electronegativity increases left to right. So: Mg < S < Cl.

3

Apply trend down groups: K is below Na in Group 1. K has more electron shells than Mg, greater shielding, and larger atomic radius. K is less electronegative than Mg.

4

Combine: K has the lowest electronegativity (Period 4, Group 1), then Mg, then S, then Cl (highest).

AnswerK < Mg < S < Cl. Electronegativity increases across a period (more protons, similar shielding, smaller radius) and decreases down a group (more shells, greater shielding, larger radius). Chlorine is closest to fluorine, so has the highest electronegativity of these four.

Practice

Question 2: CCl₄ has polar C–Cl bonds but is a non-polar molecule. Explain why.

1

CCl₄ has four C–Cl bonds. Chlorine is more electronegative than carbon, so each bond is polar with δ+ on C and δ− on Cl.

2

The molecular shape is tetrahedral (4 bonding pairs, 0 lone pairs on carbon).

3

Because all four atoms bonded to carbon are identical (all Cl), the molecule is symmetrical.

4

The four bond dipoles point towards the corners of the tetrahedron and cancel each other out.

AnswerCCl₄ has a tetrahedral shape with four identical C–Cl bonds. The symmetrical arrangement means the four bond dipoles cancel out completely, resulting in no net dipole. Therefore CCl₄ is a non-polar molecule.

Practice

Question 3: Explain why the C–H bond is considered essentially non-polar, even though carbon and hydrogen have different electronegativities. What implication does this have for the polarity of methane (CH₄)?

1

Compare electronegativities: Carbon has an electronegativity of 2.5 and hydrogen has an electronegativity of 2.2 on the Pauling scale.

2

Calculate the difference: The difference is only 0.3, which is very small. For a bond to be significantly polar, a larger difference is needed.

3

Conclusion about C–H: Because the electronegativities are so similar, the bonding electrons are shared almost equally. The C–H bond is treated as essentially non-polar at A-Level.

4

Implication for CH₄: Since all four C–H bonds are essentially non-polar, there are no significant bond dipoles to consider. Combined with its symmetrical tetrahedral shape, methane is a non-polar molecule.

AnswerThe electronegativity difference between C (2.5) and H (2.2) is only 0.3, which is too small to create a significant dipole. The bonding electrons are shared almost equally, so C–H bonds are treated as non-polar. This means CH₄ has no polar bonds and is therefore a non-polar molecule.

Summary

Key Definitions

TermDefinition

ElectronegativityThe power of an atom to attract the pair of electrons in a covalent bond

Polar bondA covalent bond where electrons are shared unequally due to different electronegativities

Polar moleculeA molecule with polar bonds arranged asymmetrically, giving a net dipole

Key Exam Points

Electronegativity increases across a period and decreases down a group

Fluorine is the most electronegative element (4.0)

Use δ+ and δ− to show partial charges on polar bonds

A molecule is polar if it has polar bonds AND is asymmetrical

Symmetrical molecules (CO₂, BF₃, CCl₄) are non-polar despite having polar bonds

Connection to Next Lesson

In Lesson 3.06, we'll see how polarity affects intermolecular forces – polar molecules experience permanent dipole-dipole forces in addition to London forces.