What is electronegativity and how is it measured in Chemistry?

Electronegativity is a measure of an atom's ability to attract and hold onto electrons within a chemical bond. It is a dimensionless quantity, typically measured on the Pauling scale, where fluorine is assigned the highest value of 4.0. Understanding electronegativity is crucial for Cambridge International A Levels as it helps predict the nature of chemical bonds.

How does electronegativity affect the type of chemical bonding?

Electronegativity differences between atoms determine the type of bond formed. If the difference is large (greater than 1.7 on the Pauling scale), the bond is likely ionic, as electrons are transferred. If the difference is small, the bond is covalent, with electrons shared more equally. This concept is fundamental in AS-Level Physical Chemistry under the Cambridge International A Levels curriculum.

Why is electronegativity important in predicting molecular polarity?

Electronegativity is key to predicting molecular polarity because it influences how electrons are distributed in a molecule. When atoms with different electronegativities form a bond, the electrons are drawn more towards the more electronegative atom, creating a dipole. Understanding this helps students in Cambridge International A Levels to determine the polarity of molecules, which affects their physical properties and reactivity.

What is the trend of electronegativity across the periodic table?

Electronegativity generally increases across a period from left to right and decreases down a group in the periodic table. This trend is due to increasing nuclear charge and decreasing atomic radius across a period, and increasing atomic radius and electron shielding down a group. Recognizing these trends is essential for students studying Chemical Bonding in AS-Level Physical Chemistry.

How does electronegativity relate to bond strength and bond length?

Electronegativity influences both bond strength and bond length. Generally, a greater difference in electronegativity between two atoms results in a stronger bond due to increased ionic character. Conversely, bonds between atoms with similar electronegativities tend to be longer and weaker. This relationship is a critical aspect of understanding chemical bonding in the Cambridge International A Levels Chemistry syllabus.

Why is "Defining electronegativity as 'attracting electrons' generally." a common mistake in Electronegativity and bonding?

Why it happens: Dropping the 'bonding/shared pair' detail. How to avoid it: Say it attracts the bonding (shared) pair of electrons in a covalent bond. Source: 9701 Examiner Reports

Why is "Confusing electronegativity with electron affinity." a common mistake in Electronegativity and bonding?

Why it happens: Both involve attracting electrons. How to avoid it: Electronegativity = pull on a bonding pair within a bond; electron affinity = energy change when a gaseous atom gains an electron.

Why is "Explaining the group trend without mentioning shielding." a common mistake in Electronegativity and bonding?

Why it happens: Only citing atomic radius. How to avoid it: Combine increased radius AND increased shielding outweighing nuclear charge.

Why is "Treating 1.8 as an exact ionic/covalent boundary." a common mistake in Electronegativity and bonding?

Why it happens: Memorising a number instead of reasoning. How to avoid it: Bonding is a continuum — argue from the size of the supplied ΔEN.

Why is "Assuming polar bonds always give a polar molecule." a common mistake in Electronegativity and bonding?

Why it happens: Ignoring molecular symmetry. How to avoid it: If the bond dipoles cancel by symmetry (CO₂, CCl₄), the molecule is non-polar. Source: 9701 Examiner Reports 2022-2024

Chemical BondingCambridge International A Level Chemistry 9701

Electronegativity and bonding

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Electronegativity and Bonding — Cambridge International AS & A Level Chemistry 9701 (2025-2027 syllabus)

Define electronegativity, explain what controls it, account for its periodic trends, and use Pauling electronegativity differences to predict whether a bond is ionic or covalent.

What you’ll learn

Mapped to the Cambridge International A Level 9701 syllabus (2025-2027).

3.1.1 — Define electronegativity as the power of an atom to attract electrons to itself.
3.1.2 — Explain the factors influencing electronegativity in terms of nuclear charge, atomic radius and shielding by inner shells and sub-shells.
3.1.3 — State and explain the trends in electronegativity across a period and down a group.
3.1.4 — Use differences in Pauling electronegativity values to predict the formation of ionic and covalent bonds.

Defining electronegativity

It's how strongly an atom pulls the shared pair of electrons in a covalent bond towards itself.

Definition (learn this exactly). Electronegativity is the power of an atom to attract the bonding pair of electrons (shared electrons) in a covalent bond towards itself.

It is a property of an atom within a bond, not of an isolated atom.
The most common scale is the Pauling scale, a dimensionless number — fluorine is the highest at $4.0$ .
A bigger electronegativity means the atom pulls shared electrons more strongly, making the bond polar.

Electronegativity = power to attract the shared (bonding) electrons.
Property of a bonded atom.
Pauling scale; F is the most electronegative (4.0).

Factors and periodic trends

Same three factors as ionisation energy: nuclear charge, radius and shielding.

Electronegativity is set by how strongly the nucleus can pull on the bonding electrons — the same three factors that control ionisation energy:

Nuclear charge — more protons → stronger pull → higher electronegativity.
Atomic radius — smaller atom → bonding pair closer to nucleus → higher electronegativity.
Shielding — more inner shells → weaker net pull → lower electronegativity.

Across a period (left → right): electronegativity increases.

Nuclear charge increases and atomic radius decreases (shielding roughly constant), so the nucleus attracts the bonding pair more strongly.

Down a group: electronegativity decreases.

Atomic radius increases and shielding increases, outweighing the higher nuclear charge, so the bonding pair is attracted less strongly.

Overall pattern. Electronegativity is highest at the top right (excluding noble gases): the order to remember is F > O > N ≈ Cl.

Factors: nuclear charge ↑, radius ↓, shielding.
Across period: increases. Down group: decreases.
Top-right elements most electronegative (F highest).

Predicting bond type from electronegativity difference

Big difference → ionic; small → polar covalent; zero → non-polar covalent.

The difference in electronegativity ( $\Delta$ EN) between two bonded atoms predicts the type of bond:

$\Delta$ EN	Bond type	Example
Zero (same element)	Non-polar covalent	$\text{Cl}_2$ , $\text{O}_2$
Small (roughly < 1.8)	Polar covalent	$\text{HCl}$ , $\text{H}_2\text{O}$
Large (roughly > 1.8)	Ionic	$\text{NaCl}$ , $\text{MgO}$

Same electronegativity → electrons shared equally → pure (non-polar) covalent.
Moderate difference → electrons pulled towards the more electronegative atom → polar covalent (a $\delta+$ / $\delta-$ dipole).
Very large difference → the more electronegative atom effectively takes the electrons → ionic.

Worked. Using Pauling values Na 0.9 and Cl 3.0: $\Delta\text{EN} = 2.1$ (large) → ionic (NaCl). For H 2.1 and Cl 3.0: $\Delta\text{EN} = 0.9$ (small) → polar covalent (HCl).

(Bonding is a continuum: there is no exact cut-off, so use the difference as a guide and quote the Pauling values given in the question. The presence of covalent character in ionic compounds is not assessed at 9701.)

ΔEN ≈ 0 → non-polar covalent.
ΔEN small → polar covalent.
ΔEN large (> ~1.8) → ionic.
Use the Pauling values supplied.

How it’s examined

Electronegativity is tested through the definition (1-2 marks), trend explanations that reuse the IE factors, and 'predict the bond type' questions using supplied Pauling values. It also underpins polarity (3.6) and many organic mechanism questions. Examiner reports flag the vague definition (omitting 'bonding/shared pair') and trend answers that forget shielding.

Sources: Cambridge International AS & A Level Chemistry 9701 syllabus (2025-2027), section 3.1; 9701 Examiner Reports 2022-2024; 9701/22 May/Jun 2024 question paper and mark scheme. Last reviewed 2026-05-20.

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Step-by-step worked examples — Electronegativity and bonding

Step-by-step solutions to past-paper-style questions on electronegativity and bonding, written exactly the way a tutor would explain them at the board.

1Identifying the most electronegative element (2 marks)
Getting started• trend
Question
From Na, P, Cl and F, state the most electronegative element and explain your choice.
Step-by-step solution
1. Step 1
  Electronegativity rises across a period and up a group, so the top-right element wins.
2. Step 2
  F is the most electronegative element (smallest, least shielded; Pauling 4.0).
Answer
Fluorine — the most electronegative element on the Pauling scale.
Examiner tip
Fluorine is the benchmark: nothing is more electronegative. Cl is second among these.
2Ordering atoms by electronegativity (2 marks)
Getting started• trend
Question
Place Na, Cl, Si and S in order of increasing electronegativity, and explain the trend.
Step-by-step solution
1. Step 1
  All in Period 3 — electronegativity increases left to right.
2. Step 2
  Order: Na < Si < S < Cl (increasing nuclear charge, decreasing radius).
Answer
Na < Si < S < Cl.
3Predicting bond type from Pauling values (3 marks)
Building confidence• bond type
Question
Using Pauling electronegativities (Mg 1.3, O 3.4, H 2.2, S 2.6), predict and justify the bond type in (a) MgO and (b) H₂S.
Step-by-step solution
1. Step 1
  MgO: $\Delta\text{EN} = 3.4 - 1.3 = 2.1$ (large).
  $\Delta\text{EN}(\text{MgO}) = 2.1 \Rightarrow \text{ionic}$
2. Step 2
  H₂S: $\Delta\text{EN} = 2.6 - 2.2 = 0.4$ (small).
  $\Delta\text{EN}(\text{H–S}) = 0.4 \Rightarrow \text{polar covalent}$
Answer
MgO: ionic (large ΔEN). H₂S: polar covalent (small ΔEN).
Examiner tip
Quote the actual difference and link its size to the bond type — don't just name the bond.
4Assigning bond polarity (3 marks)
Building confidence• polarity, dipole
Question
For the bonds H–Cl, C–Cl and C–F, mark the δ+ and δ− atoms and rank the bonds by polarity. (EN: H 2.2, C 2.6, Cl 3.2, F 4.0)
Step-by-step solution
1. Step 1
  The more electronegative atom is δ−. H–Cl: H $^{δ+}$ Cl $^{δ-}$ ; C–Cl: C $^{δ+}$ Cl $^{δ-}$ ; C–F: C $^{δ+}$ F $^{δ-}$ .
2. Step 2
  Bigger ΔEN = more polar.
  $\Delta\text{EN}: \text{C–Cl}\,0.6 < \text{H–Cl}\,1.0 < \text{C–F}\,1.4$
Answer
All have the halogen as δ−; polarity increases C–Cl < H–Cl < C–F.
Examiner tip
δ− always sits on the more electronegative atom; the size of ΔEN sets how polar the bond is.
5Explaining the down-group trend (3 marks)
Stretch• trend
Question
Explain why fluorine is more electronegative than iodine. (3)
Step-by-step solution
1. Step 1
  Smaller atom: F has a smaller atomic radius, so the bonding pair sits closer to the nucleus.
2. Step 2
  Less shielding: F has fewer inner shells screening the nucleus.
3. Step 3
  Stronger attraction: the nucleus attracts the bonding pair more strongly → higher electronegativity.
Answer
F is smaller with less shielding, so its nucleus attracts the bonding pair more strongly than iodine's.
6Polar bonds but a non-polar molecule (4 marks)
Stretch• polarity, shape
Question
Both CO₂ and H₂O contain polar bonds. Explain why CO₂ has no overall dipole but H₂O does. (4)
Step-by-step solution
1. Step 1
  Both have polar bonds: O is more electronegative than C and H, so each bond has a dipole.
2. Step 2
  CO₂ is linear (O=C=O): the two bond dipoles are equal and point in opposite directions, so they cancel.
3. Step 3
  H₂O is bent (104.5°): the two bond dipoles do not point opposite, so they add to a net dipole.
4. Step 4
  Conclusion: CO₂ is non-polar (dipoles cancel by symmetry); H₂O is polar.
Answer
Symmetry decides: CO₂'s linear shape cancels the dipoles; water's bent shape does not.
Examiner tip
A molecule is only polar if the bond dipoles do NOT cancel — always check the shape, not just the bonds.

Key Formulae — Electronegativity and bonding

The formulae you need to memorise for electronegativity and bonding on the Cambridge International A Level 9701 paper, with every variable defined in plain English and a note on when to use it.

Electronegativity difference
$\Delta\text{EN} = \text{EN}(\text{more EN atom}) - \text{EN}(\text{less EN atom})$
$\Delta\text{EN}$
difference in Pauling electronegativity
When to use
Predicting bond character: ΔEN ≈ 0 → non-polar covalent; small → polar covalent; large (≳1.8) → ionic. Treat it as a continuum, not a hard cut-off.

Key Definitions and Keywords — Electronegativity and bonding

Definitions to memorise and the exact keywords mark schemes credit for electronegativity and bonding answers — sharpened from recent examiner reports for the 2026 Cambridge International A Level 9701 sitting.

Electronegativity
Examiner keyword
The power of an atom to attract the bonding pair of electrons in a covalent bond towards itself.
Pauling scale
Examiner keyword
A dimensionless scale of electronegativity; fluorine has the highest value (4.0).
Polar covalent bond
Examiner keyword
A covalent bond between atoms of different electronegativity, giving an uneven charge distribution (δ+ / δ−).
Non-polar covalent bond
Examiner keyword
A bond between atoms of equal (or near-equal) electronegativity, so the bonding pair is shared evenly (e.g. Cl–Cl).
Dipole (bond dipole)
Examiner keyword
A separation of charge along a bond, with a δ+ and a δ− end, caused by the electronegativity difference.

Common Mistakes and Misconceptions — Electronegativity and bonding

The traps other students keep falling into on electronegativity and bonding questions — taken from recent Cambridge International A Level 9701 examiner reports and mark schemes — and how to avoid them.

✕Defining electronegativity as 'attracting electrons' generally.
9701 Examiner Reports
Why it happens
Dropping the 'bonding/shared pair' detail.
How to avoid it
Say it attracts the bonding (shared) pair of electrons in a covalent bond.
✕Confusing electronegativity with electron affinity.
Why it happens
Both involve attracting electrons.
How to avoid it
Electronegativity = pull on a bonding pair within a bond; electron affinity = energy change when a gaseous atom gains an electron.
✕Explaining the group trend without mentioning shielding.
Why it happens
Only citing atomic radius.
How to avoid it
Combine increased radius AND increased shielding outweighing nuclear charge.
✕Treating 1.8 as an exact ionic/covalent boundary.
Why it happens
Memorising a number instead of reasoning.
How to avoid it
Bonding is a continuum — argue from the size of the supplied ΔEN.
✕Assuming polar bonds always give a polar molecule.
9701 Examiner Reports 2022-2024
Why it happens
Ignoring molecular symmetry.
How to avoid it
If the bond dipoles cancel by symmetry (CO₂, CCl₄), the molecule is non-polar.

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Electronegativity and bonding — frequently asked questions

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Electronegativity is set by how strongly the nucleus can pull on the bonding electrons — the same three factors that control ionisation energy:

Nuclear charge — more protons → stronger pull → higher electronegativity.
Atomic radius — smaller atom → bonding pair closer to nucleus → higher electronegativity.
Shielding — more inner shells → weaker net pull → lower electronegativity.

Across a period (left → right): electronegativity increases.

Nuclear charge increases and atomic radius decreases (shielding roughly constant), so the nucleus attracts the bonding pair more strongly.

Down a group: electronegativity decreases.

Atomic radius increases and shielding increases, outweighing the higher nuclear charge, so the bonding pair is attracted less strongly.

Overall pattern. Electronegativity is highest at the top right (excluding noble gases): the order to remember is F > O > N ≈ Cl.

\Delta

Bond type

Example

Zero (same element)

Non-polar covalent

\text{Cl}_2

\text{O}_2

Small (roughly < 1.8)

Polar covalent

\text{HCl}

\text{H}_2\text{O}

Large (roughly > 1.8)

Ionic

\text{NaCl}

\text{MgO}

Electronegativity and bonding

Short Study Notes in the form of Slides

Short Study Notes — Electronegativity and bonding

Detailed Notes

What you’ll learn

How it’s examined

1Identifying the most electronegative element (2 marks)

2Ordering atoms by electronegativity (2 marks)

3Predicting bond type from Pauling values (3 marks)

4Assigning bond polarity (3 marks)

5Explaining the down-group trend (3 marks)

6Polar bonds but a non-polar molecule (4 marks)

Electronegativity difference

Electronegativity

Pauling scale

Polar covalent bond

Non-polar covalent bond

Dipole (bond dipole)

✕Defining electronegativity as 'attracting electrons' generally.

✕Confusing electronegativity with electron affinity.

✕Explaining the group trend without mentioning shielding.

✕Treating 1.8 as an exact ionic/covalent boundary.

✕Assuming polar bonds always give a polar molecule.

Practice questions

Past paper style quiz

Assess your understanding

Video lesson

Electronegativity and bonding — frequently asked questions

Electronegativity and bonding

Short Study Notes in the form of Slides

Short Study Notes — Electronegativity and bonding

Detailed Notes

What you’ll learn

How it’s examined

1Identifying the most electronegative element (2 marks)

2Ordering atoms by electronegativity (2 marks)

3Predicting bond type from Pauling values (3 marks)

4Assigning bond polarity (3 marks)

5Explaining the down-group trend (3 marks)

6Polar bonds but a non-polar molecule (4 marks)

Electronegativity difference

Electronegativity

Pauling scale

Polar covalent bond

Non-polar covalent bond

Dipole (bond dipole)

✕Defining electronegativity as 'attracting electrons' generally.

✕Confusing electronegativity with electron affinity.

✕Explaining the group trend without mentioning shielding.

✕Treating 1.8 as an exact ionic/covalent boundary.

✕Assuming polar bonds always give a polar molecule.

Practice questions

Past paper style quiz

Assess your understanding

Video lesson

Electronegativity and bonding — frequently asked questions