Electronegativity and Bond Polarity

Spec Mapping — OCR H432 Module 2.2.2 — Bonding and structure, covering electronegativity as the ability of an atom to attract the bonding pair of electrons in a covalent bond, the trends in electronegativity across a period and down a group, the Pauling scale, the distinction between non-polar covalent, polar covalent, and ionic bonds, the representation of polar bonds with δ+ / δ− notation and dipole arrows, and the vector-sum approach for determining whether a molecule is polar (which depends on both bond polarity and molecular shape) (refer to the official OCR H432 specification document for exact wording).

Electronegativity ties the previous two lessons together. When two atoms covalently share electrons, the share is rarely perfectly equal — atoms with higher electronegativity pull the bonding pair toward themselves, generating a permanent bond dipole. Whether the molecule as a whole is polar then depends on a vector sum of all the bond dipoles, which in turn depends on the molecular shape (Lesson 8 VSEPR). This lesson formalises the electronegativity concept, presents the Pauling scale and its periodic-table trends, distinguishes polar bonds from polar molecules, and lays the foundation for intermolecular forces (Lesson 10) — without an electronegativity difference there would be no permanent dipoles, no hydrogen bonds, and no anomalous melting/boiling points of water (Lesson 11). The concept was introduced by Linus Pauling (1932) in The Nature of the Chemical Bond (Nobel Prize 1954); the scale is empirical, based on thermochemical bond-energy differences.

Key Definition: Electronegativity is the ability (or "power", in Pauling's phrasing) of an atom to attract the pair of electrons in a covalent bond toward itself. It is a relative property defined for atoms within molecules, not for free atoms.

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The Pauling electronegativity scale

Pauling defined electronegativity from thermochemical data: the bond enthalpy of a polar bond A-B exceeds the geometric mean of A-A and B-B by an amount that depends on the electronegativity difference. The scale runs from caesium (least electronegative, 0.79) to fluorine (most electronegative, 3.98) and is dimensionless.

Element	Pauling χ	Element	Pauling χ
F	3.98	C	2.55
O	3.44	H	2.20
Cl	3.16	P	2.19
N	3.04	Si	1.90
Br	2.96	Al	1.61
I	2.66	Mg	1.31
S	2.58	Na	0.93
Se	2.55	Cs	0.79

Noble gases are not typically assigned values because they rarely form chemical bonds (XeF₂, XeF₄, KrF₂ are the exceptions; for those compounds Xe and Kr behave as if they had moderate electronegativity, around 2.5–2.6).

Other scales exist — Mulliken (1934, average of ionisation energy and electron affinity), Allred-Rochow (1958, force-on-electron model), Allen (1989, weighted average of valence-shell ionisation energies) — but Pauling's is the standard A-Level reference.

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Periodic trends in electronegativity

flowchart TD
    A[Periodic trends] --> B[Across a period: left to right]
    A --> C[Down a group: top to bottom]
    B --> D[Nuclear charge increases]
    B --> E[Atomic radius decreases]
    D --> F[Bonding pair held more strongly]
    E --> F
    F --> G[Electronegativity INCREASES across a period]
    C --> H[Atomic radius increases]
    C --> I[Inner-shell shielding increases]
    H --> J[Bonding pair held less strongly]
    I --> J
    J --> K[Electronegativity DECREASES down a group]
    G --> L[F is the most electronegative element top-right corner]
    K --> M[Cs Fr are the least electronegative bottom-left]

Electronegativity increases:

• Across a period (left → right) — nuclear charge increases while atomic radius decreases. The bonding pair sits closer to a more highly charged nucleus and is pulled more strongly. Na (0.93) → Mg (1.31) → Al (1.61) → Si (1.90) → P (2.19) → S (2.58) → Cl (3.16).
• Up a group (bottom → top) — atomic radius decreases (fewer filled inner shells) and the bonding pair sits closer to the nucleus. F (3.98) > Cl (3.16) > Br (2.96) > I (2.66).

The combined trend places F at the top right and Cs/Fr at the bottom left as the extremes. Noble gases (He, Ne, Ar, Kr, Xe) are not assigned standard values.

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Non-polar vs polar covalent bonds

A covalent bond is non-polar when the two atoms have identical (or near-identical) electronegativity — the bonding pair sits symmetrically between them. Examples:

• Two atoms of the same element: H-H, Cl-Cl, O=O, N≡N (Δχ = 0).
• C-H (Δχ = 0.35) — conventionally treated as non-polar at A-Level because the difference is small.

A covalent bond is polar when the two atoms have different electronegativity. The more electronegative atom gains a partial (δ⁻) negative charge; the less electronegative atom gains a partial (δ⁺) positive charge. The δ symbols denote fractional charge (typically 0.1–0.5 e), not full ionic charges.

Bond	Δχ (Pauling)	Classification	Example use
H-H	0.00	Non-polar	Dihydrogen
C-H	0.35	Effectively non-polar	Hydrocarbons
C-N	0.49	Slightly polar	Amines
C-Cl	0.61	Polar	Haloalkanes
C-O	0.89	Polar	Alcohols, ethers
H-Cl	0.96	Polar	Hydrogen halides
C=O	0.89	Polar	Carbonyls
O-H	1.24	Very polar	Water, alcohols, carboxylic acids
N-H	0.84	Polar	Ammonia, amines
H-F	1.78	Very polar (still covalent)	Strongest H-X bond
Na-Cl	2.23	Ionic	NaCl lattice

As a rough A-Level guide:

• Δχ < 0.4 → non-polar covalent.
• 0.4 ≤ Δχ ≤ 1.7 → polar covalent.
• Δχ > 1.7 → essentially ionic.

These boundaries are guidelines, not laws — bonding is a continuum between purely covalent and purely ionic, and even "ionic" compounds (NaCl at 72 % ionic character per Pauling's formula) retain some covalent character.

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Representing polar bonds — δ⁺ / δ⁻ and dipole arrows

Two conventions are used in OCR mark schemes:

1. δ⁺ / δ⁻ labels on atoms δ⁺ δ⁻ H Cl δ⁻ on the more electronegative atom (Cl: 3.16 > H: 2.20)

2. Dipole arrow + cross-tail H Cl Arrow points TOWARDS the more electronegative atom Cross-tail (+) sits at the δ⁺ end

Both conventions are acceptable in OCR mark schemes; many candidates use both for clarity.

The δ⁻ always sits on the more electronegative atom; the arrow always points to the more electronegative atom (toward δ⁻).

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Polar bonds vs polar molecules — the symmetry rule

A molecule is polar if it has a non-zero net dipole moment. This requires two things:

1. At least one polar bond.
2. The bond dipoles do not cancel by symmetry.

Bond dipoles are vectors — they have magnitude and direction. The molecular dipole is the vector sum of all the bond dipoles. If the sum is zero (by molecular symmetry), the molecule is non-polar even though it contains polar bonds. If the sum is non-zero, the molecule is polar.

Non-polar molecules with polar bonds (dipoles cancel)

CO₂ (linear, 180°) O C O δ⁺ δ⁻ δ⁻ 2 equal opposite dipoles cancel → non-polar

CCl₄ (tetrahedral, 109.5°) C Cl Cl Cl Cl 4 equal tetrahedral dipoles cancel → non-polar

BF₃ (trigonal planar, 120°) B F F F 3 dipoles at 120° cancel → non-polar

H₂O (bent, 104.5°) — POLAR O H H Bent → dipoles add → net dipole points toward O

NH₃ (pyramidal, 107°) — POLAR N ·· H H H Pyramidal → dipoles add → net dipole along lone pair axis

CHCl₃ (tetrahedral) — POLAR C H Cl Cl Cl C-H weaker dipole than 3 × C-Cl → dipoles don't cancel

Polar molecules — dipoles do not cancel

• H₂O — bent (104.5°), 2 O-H dipoles add to give a net dipole bisecting the H-O-H angle, pointing from H side toward O. μ = 1.85 D.
• NH₃ — trigonal pyramidal (107°), 3 N-H dipoles plus the lone pair give a net dipole along the lone pair axis. μ = 1.47 D.
• CHCl₃ (chloroform) — tetrahedral but asymmetric (3 Cl + 1 H). The C-H bond is much less polar than the 3 C-Cl bonds, so dipoles do not cancel. μ = 1.04 D.
• HCl, HF, HBr, HI — diatomic with no cancellation possible. μ(HF) = 1.91 D, μ(HCl) = 1.08 D.

Non-polar molecules despite polar bonds

• CO₂ — linear (180°), 2 C=O dipoles point in opposite directions and cancel exactly. μ = 0.
• CCl₄ — tetrahedral with 4 identical C-Cl dipoles arranged symmetrically. μ = 0.
• BF₃ — trigonal planar (120°), 3 B-F dipoles cancel by 3-fold symmetry. μ = 0.
• SF₆ — octahedral, 6 S-F dipoles cancel by 6-fold symmetry. μ = 0.
• PCl₅ — trigonal bipyramidal, 5 P-Cl dipoles cancel. μ = 0.
• CH₄ — tetrahedral, but C-H dipoles are essentially zero in any case. μ ≈ 0.

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Deciding whether a molecule is polar — the algorithm

flowchart TD
    A[Identify all bonds in the molecule] --> B{Any electronegativity difference?}
    B -- "No" --> C[Non-polar molecule]
    B -- "Yes" --> D[At least one polar bond present]
    D --> E[Determine molecular shape using VSEPR]
    E --> F{All bond dipoles equal AND arranged symmetrically?}
    F -- "Yes" --> G[Vector sum = 0 → non-polar molecule]
    F -- "No" --> H[Vector sum ≠ 0 → POLAR molecule]
    G --> I[Examples: CO2 CCl4 BF3 SF6 PCl5]
    H --> J[Examples: H2O NH3 CHCl3 HCl SO2]

The crucial step is shape determination — without it, you cannot tell whether the bond dipoles cancel.

Worked example — SO₃ vs SO₂

SO₃ — S has 3 double bonds to O, no lone pairs. 3 regions → trigonal planar (120°). The three S=O dipoles point outward at 120° and cancel exactly by 3-fold symmetry. SO₃ is non-polar.

SO₂ — S has 2 double bonds to O + 1 lone pair. 3 regions → trigonal planar electron geometry, but 1 lone pair → bent molecular shape (~117°). The two S=O dipoles do not cancel; they add to give a net dipole pointing from S toward the bisector of the two O atoms. SO₂ is polar (μ = 1.63 D).