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Electronegativity is the property that quantifies how strongly an atom in a covalent bond pulls the shared pair of electrons towards itself. The modern scale we use at A-Level was introduced by Linus Pauling in 1932, calibrated against bond-dissociation energies and assigning fluorine the maximum value 4.0. Pauling's insight transformed bonding from a binary ionic-versus-covalent classification into a continuum: at one extreme is the pure ionic limit (large electronegativity difference, Δχ, with electron transfer); at the other is the pure covalent limit (Δχ = 0, electrons shared equally). Almost every real bond sits between these extremes. This lesson establishes the periodic trends in electronegativity, the link between Δχ and bond polarity, the role of molecular geometry in deciding whether bond dipoles cancel, the quantitative measure of permanent dipole moment, and the way polarity rationalises solubility, volatility, infrared activity and the Period 3 chloride hydrolysis pattern.
Spec mapping (AQA 7405): This lesson maps to §3.1.3 (electronegativity and bond polarity). It cross-references lesson 0 (ionic bonding — pure ionic is the high-Δχ limit), lesson 1 (covalent bonding — pure covalent is the Δχ = 0 limit), lesson 2 (shapes of molecules — the symmetry that decides dipole cancellation), and lesson 4 (intermolecular forces — permanent dipole-dipole and hydrogen bonding originate here). It also previews §3.2.4 (Period 3 chlorides) and §3.3 (organic mechanisms — bond polarity governs electrophilic and nucleophilic sites). Refer to the official AQA specification document for the exact wording of each section.
Assessment objectives: AO1 — define electronegativity and recall typical Pauling values for H, C, N, O, F, Cl. AO2 — predict bond polarity from electronegativity differences; predict whether a molecule has a permanent dipole moment by combining bond polarity with molecular shape. AO3 — rationalise reactivity trends along Period 3 chlorides; predict hydrogen-bonding capability of unfamiliar molecules; explain why CO₂ is non-polar despite polar bonds while H₂O is strongly polar.
Key Definition: Electronegativity is the ability of an atom in a covalent bond to attract the bonding pair of electrons towards itself.
Two features of this definition deserve emphasis. First, electronegativity is a property of an atom in a bond, not of a free atom in isolation — the related but distinct properties for isolated atoms are ionisation energy (energy to lose an electron) and electron affinity (energy released on gaining one). Second, electronegativity is a relative quantity: the Pauling scale is dimensionless and calibrated against fluorine = 4.0.
The qualitative idea is straightforward. When two atoms share a pair of electrons, the more electronegative atom holds the shared pair closer to its own nucleus. The probability density of the bonding electrons is shifted, producing a partial negative charge (δ−) on the more electronegative atom and an equal partial positive charge (δ+) on its partner.
Pauling derived his scale from a thermochemical argument: the bond dissociation energy of A—B exceeds the geometric mean of A—A and B—B by an amount proportional to (χ_A − χ_B)². Setting fluorine to 4.0 fixes the scale.
| 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 |
| Iodine | I | 2.5 |
| Hydrogen | H | 2.1 |
| Phosphorus | P | 2.1 |
| Boron | B | 2.0 |
| Silicon | Si | 1.8 |
| Beryllium | Be | 1.5 |
| Magnesium | Mg | 1.2 |
| Lithium | Li | 1.0 |
| Sodium | Na | 0.9 |
| Potassium | K | 0.8 |
The six values to commit to memory for exam recall are H 2.1, C 2.5, N 3.0, O 3.5, F 4.0 and Cl 3.0. These cover almost every bond-polarity question on Papers 1, 2 and 3.
Key Point: Two coincidences in the Pauling table catch students out. (i) Nitrogen and chlorine share the value 3.0, so the N—Cl bond is essentially non-polar (Δχ ≈ 0). (ii) Carbon and sulfur both sit at 2.5, so C—S is also essentially non-polar — a fact that matters when interpreting the IR spectrum of thiols.
Pauling's is not the only scale. Mulliken defined electronegativity as the average of the first ionisation energy and the electron affinity, χ_M = (IE + EA)/2 — a property of the free atom rather than the bonded atom. Allred and Rochow derived electronegativity from the effective nuclear charge experienced at the covalent radius, χ_AR ∝ Z_eff/r². The three scales differ in absolute numerical value but agree closely on the ordering and on the periodic trends. For A-Level work, only the Pauling scale is required; the existence of alternatives is useful background and recurs in the Going Further section.
Electronegativity varies systematically across the periodic table:
Combining the two trends, the most electronegative element is fluorine (top-right of the main block, excluding the noble gases) and the least electronegative is caesium or francium (bottom-left). The trends correlate tightly with the trends in ionisation energy and electron affinity covered in the atomic-structure course.
| Element | Na | Mg | Al | Si | P | S | Cl |
|---|---|---|---|---|---|---|---|
| Electronegativity | 0.9 | 1.2 | 1.5 | 1.8 | 2.1 | 2.5 | 3.0 |
Z_eff rises from ~2.5 on Na to ~6.1 on Cl (Slater's rules); atomic radius contracts from 0.186 nm on Na to 0.099 nm on Cl. Both factors reinforce the upward trend.
The polarity of a single bond is governed by the electronegativity difference Δχ between the two bonded atoms. The standard rule-of-thumb table:
| Δχ | Bond character | Examples |
|---|---|---|
| 0 | Pure covalent (non-polar) | H—H, Cl—Cl, C—S |
| 0 < Δχ < 0.5 | Essentially non-polar | C—H (0.4), N—Cl (0.0) |
| 0.5 ≤ Δχ < 1.7 | Polar covalent | C—Cl (1.0), O—H (1.4), N—H (0.9), H—Cl (0.9) |
| Δχ ≥ 1.7 | Substantially ionic | Na—Cl (2.1), Mg—F (2.8) |
Important Caveat: The 1.7 boundary is a convenient bookkeeping device, not a sharp physical line. The transition from polar covalent to ionic is a continuum, not a binary switch. HF (Δχ = 1.9) sits just above the boundary yet behaves like a covalent molecule in many contexts; AgCl (Δχ ≈ 1.2) is conventionally treated as ionic but has substantial covalent character. The boundaries are useful, not absolute.
A polar bond is shown either by δ+/δ− labels on the bonded atoms, or by an arrow with a crossed tail pointing from δ+ to δ−. The H—Cl bond is the textbook example:
δ+ δ−
H ———→ Cl (Δχ = 0.9)
Other essential examples for the A-Level:
| Bond | Δχ | δ+ atom | δ− atom |
|---|---|---|---|
| O—H | 1.4 | H | O |
| N—H | 0.9 | H | N |
| C—Cl | 1.0 | C | Cl |
| C—Br | 0.7 | C | Br |
| C—F | 1.5 | C | F |
| C=O | 1.0 | C | O |
| H—F | 1.9 | H | F |
| H—Br | 0.7 | H | Br |
| H—I | 0.4 | H | I |
The δ+/δ− assignment is decided purely by who has the higher electronegativity — never by who is "bigger" or "smaller". This sounds obvious but a common error is to assume that the larger atom is δ−; for C—I, that gives the wrong answer (C is δ+, I is δ−, because χ(C) = 2.5 = χ(I), so C—I is in fact essentially non-polar — another trap).
A bond or molecular dipole moment μ is defined as the product of charge and the distance separating the partial charges:
μ = q × d
The SI unit is C m, but in practice chemists report dipole moments in debye (D), where 1 D = 3.336 × 10⁻³⁰ C m. For a molecule with several polar bonds, the net dipole moment is the vector sum of the individual bond dipoles — directions matter, and dipoles in opposite directions cancel.
| Molecule | μ / D | Comment |
|---|---|---|
| HF | 1.91 | Δχ = 1.9, very polar |
| H₂O | 1.85 | Bent, two O—H dipoles add to give a large net |
| NH₃ | 1.47 | Pyramidal, three N—H dipoles + lone pair |
| HCl | 1.08 | Δχ = 0.9 |
| HBr | 0.82 | Δχ = 0.7 |
| HI | 0.45 | Δχ = 0.4 |
| H₂S | 0.97 | Bent like H₂O but lower Δχ |
| CHCl₃ | 1.04 | Polar bonds + asymmetric tetrahedron |
| CH₂Cl₂ | 1.60 | Polar bonds, lower symmetry |
| CCl₄ | 0 | Polar bonds but tetrahedral symmetry cancels |
| CO₂ | 0 | Polar C=O bonds but linear → cancellation |
| BF₃ | 0 | Polar B—F bonds, trigonal planar → cancellation |
| CH₄ | 0 | Bonds essentially non-polar, symmetric anyway |
Notice the steady decline along HF → HCl → HBr → HI: as the halogen becomes less electronegative, Δχ shrinks, and the bond dipole falls. Bond length increases slightly down the group (which would, in isolation, raise μ), but the drop in q dominates the rise in d.
A polar molecule has a non-zero net dipole moment; a non-polar molecule has μ = 0. A molecule is polar if and only if both of the following hold:
Cancellation requires high symmetry. The geometries that produce a fully cancelled dipole when all surrounding atoms are identical are: linear (two bonds), trigonal planar (three bonds), tetrahedral (four bonds), trigonal bipyramidal (five bonds), and octahedral (six bonds). Any lone pair on the central atom — or any difference between the surrounding atoms — breaks that symmetry.
| Molecule | Shape | Surrounding atoms identical? | Polar? |
|---|---|---|---|
| HCl | Linear | n/a (heteronuclear diatomic) | Yes |
| H₂O | Bent (2 lp on O) | Yes | Yes (lp breaks symmetry) |
| NH₃ | Pyramidal (1 lp on N) | Yes | Yes (lp breaks symmetry) |
| CH₄ | Tetrahedral | Yes | No (Δχ ≈ 0 and symmetric) |
| CHCl₃ | Tetrahedral | No (3 Cl, 1 H) | Yes |
| CCl₄ | Tetrahedral | Yes | No (cancellation) |
| CO₂ | Linear (no lp on C) | Yes | No (cancellation) |
| SO₂ | Bent (1 lp on S) | Yes | Yes (lp breaks symmetry) |
| BF₃ | Trigonal planar | Yes | No (cancellation) |
| NF₃ | Pyramidal (1 lp on N) | Yes | Yes (lp breaks symmetry) |
CCl₄ is tetrahedral with four identical C—Cl polar bonds; the four bond dipoles point from C towards each Cl, and their vector sum is exactly zero. CCl₄ has μ = 0 — it is non-polar despite four polar bonds.
CHCl₃ is also tetrahedral, but one Cl is replaced by H. The three C—Cl dipoles still point outwards, while the single C—H dipole is much smaller (Δχ = 0.4 vs 1.0) and points the other way (from H δ+ towards C δ−, in this context). The three C—Cl dipoles no longer cancel each other in projection along the C—H axis, leaving a net dipole along that axis with μ = 1.04 D. CHCl₃ is polar.
The same logic applies to CH₂Cl₂ (polar, μ = 1.60 D) and CH₃Cl (polar, μ = 1.87 D): only the fully-substituted CCl₄ and the unsubstituted CH₄ are non-polar in this series.
Polar solutes dissolve well in polar solvents (water, ethanol) because permanent dipole-dipole and hydrogen-bonding interactions between solute and solvent compensate for the disruption of the solvent's own interactions. Non-polar solutes (hexane, iodine, CCl₄) dissolve in non-polar solvents because London forces dominate on both sides. The rule "like dissolves like" is shorthand for matching the polarity of solvent and solute — anchored on lesson 4 (intermolecular forces).
A polar molecule with permanent dipoles experiences both London dispersion forces and permanent dipole-dipole forces; the total intermolecular force is therefore stronger, and the boiling point is higher than for a non-polar molecule of comparable molecular mass. Compare propane (Mᵣ = 44, b.p. −42 °C, non-polar) with acetaldehyde (CH₃CHO, Mᵣ = 44, b.p. +20 °C, polar) — the polarity adds ~60 °C. The effect is even more dramatic when hydrogen bonding is possible (anchor for lesson 4).
A bond vibration produces an absorption in the infrared spectrum only if the vibration changes the molecular dipole moment. A symmetric stretch of a non-polar bond (e.g. the symmetric stretch of CO₂) is IR-inactive — neither bond polarity nor change of dipole are accessible. The asymmetric stretch of CO₂ is IR-active because the two C=O dipoles no longer cancel mid-vibration. This is why CO₂ is a greenhouse gas (it absorbs in the IR despite being non-polar overall) but N₂ and O₂ are not. The rule "only polar bond stretches are IR-active" is a useful shortcut, deepened in §3.3.15 organic spectroscopy.
The Period 3 chlorides illustrate the electronegativity-difference continuum in a single row of compounds:
| Chloride | Δχ | Bonding character | Hydrolysis on contact with water |
|---|---|---|---|
| NaCl | 2.1 | Ionic | Dissolves; neutral solution |
| MgCl₂ | 1.8 | Ionic (some covalent character) | Dissolves; slightly acidic |
| AlCl₃ | 1.5 | Polar covalent (significant) | Hydrolyses; acidic solution (HCl + Al(OH)₃) |
| SiCl₄ | 1.2 | Polar covalent | Hydrolyses violently to SiO₂·xH₂O + HCl |
| PCl₃ / PCl₅ | 0.9 | Polar covalent | Hydrolyses to H₃PO₃ / H₃PO₄ + HCl |
| SCl₂ | 0.5 | Polar covalent | Hydrolyses slowly |
| Cl₂ | 0 | Pure covalent | Disproportionates (HClO + HCl) |
The trend in Δχ predicts both the type of bonding and the chemical reactivity pattern. Ionic chlorides dissolve; polar covalent chlorides hydrolyse; this is one of the cleanest illustrations of the electronegativity-polarity continuum at A-Level (§3.2.4).
A charged rod (e.g. a comb rubbed on hair) held near a thin stream of liquid will deflect a polar liquid (water, ethanol) but not a non-polar liquid (hexane, CCl₄). The dipoles in the polar liquid align with the electric field; the bulk of the stream is pulled towards the rod. This is the canonical classroom demonstration of permanent molecular polarity.
Pauling's deepest insight was that ionic and covalent bonding are limits of a single phenomenon. Hannay and Smyth (1946) proposed an empirical formula relating the percentage ionic character of a bond to Δχ:
% ionic character ≈ 16·|Δχ| + 3.5·(Δχ)²
For HCl (Δχ = 0.9), this gives ~17% ionic character; for NaCl (Δχ = 2.1), ~71%. The formula is illustrative, not exact — modern calculations using molecular dipoles and quantum chemistry give somewhat different values — but the qualitative point holds: very few real bonds are 100% ionic or 100% covalent, and Δχ is the correct first variable to inspect.
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