Electronegativity — Concepts, Formulas & Examples

Electronegativity is the tendency of an atom to attract shared electrons in a chemical bond. Linus Pauling’s scale is the most used. CBSE Class 11 and NEET test trends and effects on bond polarity and reactivity.

Core Concepts

Definition and scale

Linus Pauling scale — F is 4.0 (highest), Cs and Fr near 0.7 (lowest). Other scales — Mulliken (based on ionisation energy and electron affinity), Allred-Rochow.

Pauling’s method: He calculated electronegativity differences from bond energies. If the bond A-B is stronger than the average of A-A and B-B, the extra stability comes from the ionic character due to electronegativity difference.

\Delta\chi = \chi_A - \chi_B = 0.208\sqrt{\Delta}

Where $\Delta = E_{A-B} - \sqrt{E_{A-A} \times E_{B-B}}$ (all energies in kJ/mol). Here $\chi$ represents electronegativity on the Pauling scale.

Mulliken’s scale: Electronegativity = $\frac{\text{IE} + \text{EA}}{2}$ (average of ionisation energy and electron affinity). This gives values roughly 2.8 times the Pauling values.

Allred-Rochow scale: Based on the electrostatic force exerted by the nucleus on bonding electrons. $\chi = 0.359 \times \frac{Z_{eff}}{r^2} + 0.744$ where $r$ is in angstroms.

Periodic trends

Increases across a period (left to right) because effective nuclear charge increases and atomic size decreases. Decreases down a group because atomic size increases and effective pull weakens.

Across a period (say Period 2): Li (1.0) → Be (1.5) → B (2.0) → C (2.5) → N (3.0) → O (3.5) → F (4.0). The effective nuclear charge increases by roughly 1 unit per element, pulling shared electrons more strongly.

Down a group (say Group 17): F (4.0) → Cl (3.0) → Br (2.8) → I (2.5). Each new shell increases atomic radius, so the nucleus is farther from the bonding electrons and attracts them less.

Key values to remember:

Element	Pauling EN	Notes
F	4.0	Most electronegative element
O	3.5	Second most electronegative
N	3.0	Same as Cl
Cl	3.0	Most electronegative halogen after F
C	2.5	Key for organic chemistry
H	2.1	Reference point for many comparisons
Na	0.9	Typical electropositive metal

Noble gases generally do not have electronegativity values because they do not form bonds under normal conditions (though Xe compounds exist).

Effect on bond type

Small difference (less than 0.5) → non-polar covalent. Moderate (0.5 to 1.7) → polar covalent. Large (over 1.7) → ionic. Not a sharp boundary but a continuum.

\% \text{ionic character} = 16(\Delta\chi) + 3.5(\Delta\chi)^2

For NaCl: $\Delta\chi = 3.0 - 0.9 = 2.1$ . Percent ionic = $16(2.1) + 3.5(2.1)^2 = 33.6 + 15.4 = 49\%$ . Even NaCl is not 100% ionic — it has some covalent character.

Effect on acid strength

In oxyacids, higher electronegativity on central atom pulls electron density from the O-H bond, making it easier to release H $^+$ . HClO $_4$ is stronger than HClO because Cl has multiple oxygens.

Two rules for acid strength:

Same central atom, different number of oxygens (oxyacids): More oxygens = stronger acid. The extra electronegative O atoms pull electron density away from O-H, weakening it.

\text{HClO} < \text{HClO}_2 < \text{HClO}_3 < \text{HClO}_4

Same structure, different central atom: Higher electronegativity of central atom = stronger acid.

\text{HIO}_3 < \text{HBrO}_3 < \text{HClO}_3

For binary acids (hydrides) in the same period: acid strength increases with electronegativity. $\text{CH}_4 < \text{NH}_3 < \text{H}_2\text{O} < \text{HF}$ . Down a group, bond strength matters more than electronegativity, so HI > HBr > HCl > HF.

Effect on molecular properties

Polar bonds can add up to give polar molecules, which have higher boiling points and solubility in water. Non-polar molecules dissolve in non-polar solvents.

Electronegativity and inductive effect: In organic chemistry, electronegative atoms or groups attached to a carbon chain pull electron density toward themselves through sigma bonds. This is the -I effect (electron-withdrawing inductive effect). The strength of -I effect follows electronegativity: $\text{F} > \text{Cl} > \text{Br} > \text{I}$ .

The -I effect explains why chloroacetic acid ( $\text{ClCH}_2\text{COOH}$ ) is a stronger acid than acetic acid ( $\text{CH}_3\text{COOH}$ ). The electronegative Cl withdraws electron density from the O-H bond, making it easier to lose H $^+$ .

Electronegativity and oxidation state

The more electronegative atom in a bond is assigned the negative oxidation state. In H $_2$ O, oxygen is -2 (more electronegative) and hydrogen is +1. In OF $_2$ , oxygen is +2 because fluorine is more electronegative than oxygen — the only common compound where oxygen has a positive oxidation state.

Worked Examples

F has very high electronegativity. H-F is strongly polar, leading to strong hydrogen bonding between molecules. HF boils at 20°C; HCl boils at -85°C.

Each C=O is polar, but the molecule is linear and symmetric. The two bond dipoles cancel exactly, leaving no net molecular dipole.

For MgO: EN of Mg = 1.2, EN of O = 3.5. Difference = 2.3. Since $\Delta\chi > 1.7$ , the bond is predominantly ionic.

For HBr: EN of H = 2.1, EN of Br = 2.8. Difference = 0.7. Since $0.5 < \Delta\chi < 1.7$ , the bond is polar covalent.

For Cl $_2$ : EN of Cl = 3.0, EN of Cl = 3.0. Difference = 0. The bond is non-polar covalent.

HClO ( $pK_a \approx 7.5$ ), HClO $_2$ ( $pK_a \approx 2.0$ ), HClO $_3$ ( $pK_a \approx -1$ ), HClO $_4$ ( $pK_a \approx -10$ ).

Each additional oxygen atom stabilises the conjugate base by delocalising the negative charge over more electronegative O atoms. HClO $_4$ (perchloric acid) is one of the strongest acids known.

Acetic acid: $\text{CH}_3\text{COOH}$ , $pK_a = 4.76$ . Trichloroacetic acid: $\text{CCl}_3\text{COOH}$ , $pK_a = 0.65$ .

Three Cl atoms exert a strong -I effect, pulling electron density away from the COO $^-$ group. This stabilises the conjugate base (trichloroacetate ion), making the acid much stronger. Each Cl adds to the effect cumulatively.

Solved Problems (Exam Style)

Problem 1 (JEE Main pattern): Arrange the following in decreasing order of electronegativity: C, N, O, F, S.

F, O, N, C are all in Period 2 — electronegativity increases left to right: C (2.5) < N (3.0) < O (3.5) < F (4.0).

S is in Period 3, below O. It has lower electronegativity than O: S (2.5).

Decreasing order: F > O > N > C = S

Note that C and S have similar Pauling EN values (both 2.5), which is sometimes tested.

Problem 2 (NEET pattern): Which of the following has the most polar bond: HF, HCl, HBr, HI?

HF: $4.0 - 2.1 = 1.9$ . HCl: $3.0 - 2.1 = 0.9$ . HBr: $2.8 - 2.1 = 0.7$ . HI: $2.5 - 2.1 = 0.4$ .

HF has the most polar bond because fluorine has the highest electronegativity, giving the largest $\Delta\chi$ .

Problem 3 (CBSE Board): Calculate the percent ionic character of HCl bond given EN of H = 2.1 and Cl = 3.0.

$\Delta\chi = 3.0 - 2.1 = 0.9$

$\% \text{ionic} = 16(0.9) + 3.5(0.9)^2 = 14.4 + 2.835 = 17.2\%$

HCl is about 17% ionic and 83% covalent — a polar covalent bond, as we would expect.

Common Mistakes

Saying electronegativity increases down a group. It decreases.

Confusing electronegativity with electron affinity. Electronegativity is relative within a bond; electron affinity is an absolute energy release.

Writing that all polar bonds give polar molecules. Symmetry can cancel bond dipoles.

Assuming the most electronegative element has the highest electron affinity. Fluorine is the most electronegative, but chlorine has a higher electron affinity. This is because F is so small that electron-electron repulsion in the 2p subshell reduces its EA.

Forgetting that electronegativity applies only to bonded atoms. Isolated atoms do not have electronegativity. Noble gases (except in rare compounds) have no defined electronegativity on the Pauling scale.

Exam Weightage and Revision

This topic is a repeat performer in board papers and entrance exams. NEET typically asks one to two questions on the core mechanisms, CBSE boards give three to six marks, and state PMT papers often include a diagram-based long answer. The PYQs cluster around a small set of facts — lock those and you clear the topic.

JEE Main 2023 had a question on arranging elements by electronegativity. NEET 2022 tested the relationship between electronegativity and acid strength. CBSE boards ask about periodic trends in almost every paper — electronegativity is one of the top three properties tested.

When a question gives a scenario, identify the core mechanism first, then match it to the concepts above. Most wrong answers come from reading the scenario too quickly.

Three facts to lock — F is most electronegative, increases across period and decreases down group, affects bond type and acid strength.

Practice Questions

Q1. Why is oxygen more electronegative than nitrogen?

Oxygen has one more proton than nitrogen (8 vs 7), giving a higher effective nuclear charge. Both are in the same period, so they have similar atomic radii. The greater nuclear pull in oxygen attracts bonding electrons more strongly.

Q2. Arrange in order of increasing acid strength: CH $_3$ COOH, ClCH $_2$ COOH, Cl $_2$ CHCOOH, Cl $_3$ CCOOH.

Each additional Cl atom increases the -I effect, stabilising the conjugate base more. CH $_3$ COOH < ClCH $_2$ COOH < Cl $_2$ CHCOOH < Cl $_3$ CCOOH. The $pK_a$ values drop from 4.76 to 0.65 as you go from zero to three chlorine atoms.

Q3. Why does fluorine have lower electron affinity than chlorine despite being more electronegative?

Fluorine’s 2p orbitals are very small. Adding an extra electron creates significant electron-electron repulsion in this tiny space, reducing the energy released. Chlorine’s 3p orbitals are larger, so the incoming electron faces less repulsion. Electron affinity: Cl (349 kJ/mol) > F (328 kJ/mol).

Q4. What is the electronegativity of an element on Mulliken’s scale if its IE = 500 kJ/mol and EA = 300 kJ/mol?

Mulliken EN = $\frac{\text{IE} + \text{EA}}{2} = \frac{500 + 300}{2} = 400$ kJ/mol. To convert to the Pauling scale (roughly), divide by 230: $400/230 \approx 1.74$ .

Q5. Why is the O-H bond in water more polar than the N-H bond in ammonia?

Oxygen (EN = 3.5) is more electronegative than nitrogen (EN = 3.0). The electronegativity difference with H: O-H gives $3.5 - 2.1 = 1.4$ , N-H gives $3.0 - 2.1 = 0.9$ . The larger difference makes O-H more polar, which is why water has stronger hydrogen bonds than ammonia.

FAQs

Is electronegativity the same as electron affinity? No. Electronegativity is the tendency to attract shared electrons in a bond — it is a relative property that only applies when the atom is bonded. Electron affinity is the energy change when an isolated gaseous atom gains an electron — it is an absolute, measurable quantity.

Why do noble gases not have electronegativity values? Electronegativity is defined for atoms in bonds. Noble gases generally do not form bonds (their valence shell is full). Xenon does form some compounds (XeF $_2$ , XeF $_4$ ), so it can be assigned a Pauling EN of about 2.6.

Can electronegativity be negative? No. On the Pauling scale, all values are positive. The lowest values (around 0.7 for Cs and Fr) mean the atom has very little tendency to attract bonding electrons — it prefers to lose them and form cations.

How does electronegativity relate to oxidation number? In any bond, the more electronegative atom is assigned the electrons. This means it gets the negative oxidation state. In CO $_2$ , oxygen (-2) gets both shared pairs because it is more electronegative than carbon (+4).

Electronegativity is the single most useful concept for predicting chemical behaviour. Learn the trend and you can reason about any new molecule.