Bond Polarity: Using Electronegativity Differences

When two atoms share a pair of electrons, they don’t always share it equally. If one atom pulls harder, the electrons spend more time near it, giving that end of the bond a small negative charge and the other end a small positive charge. This unequal sharing is called bond polarity, and it explains a huge range of chemistry: why water dissolves salt, why hydrogen chloride is an acid, and why some molecules stick together more strongly than others. The tool for predicting it is electronegativity.

Electronegativity: the pulling power

Electronegativity is the ability of an atom in a covalent bond to attract the shared pair of electrons towards itself. The most widely used scale was devised by Linus Pauling, running from about 0.7 to 4.0 (see the electronegativity trend and the electronegativity heatmap).

Pauling values for some common elements:

Element	Electronegativity	Element	Electronegativity
F	3.98	C	2.55
O	3.44	H	2.20
Cl	3.16	P	2.19
N	3.04	Mg	1.31
Br	2.96	Na	0.93
I	2.66	K	0.82
S	2.58	Cs	0.79

Two trends to remember:

• Electronegativity increases across a period: atoms have more protons but electrons are in the same shell, so the shared pair is pulled more strongly.
• Electronegativity decreases down a group: atoms get larger and the bonding electrons are further from the nucleus and more shielded.

Fluorine is the most electronegative element (see the most electronegative element); the least electronegative are caesium and francium at the bottom left.

Three kinds of bond

The difference in electronegativity (Δχ) between two bonded atoms tells you how the electrons are shared. A common rough guide:

Δχ	Bond type	Description
0 (or below about 0.4)	Non-polar covalent	Electrons shared (almost) equally
about 0.4 to 1.7	Polar covalent	Electrons shared unequally; partial charges
above about 1.7–2.0	Mainly ionic	Electron effectively transferred

These cut-offs aren’t sharp — bonding is a continuous spectrum from purely covalent to highly ionic, and different textbooks use slightly different boundaries (see ionic vs covalent bonds).

Partial charges: δ+ and δ−

In a polar bond, the more electronegative atom carries a partial negative charge, written δ− (“delta minus”). The less electronegative atom carries a partial positive charge, δ+. These are fractions of an electron’s charge, not whole charges like those on ions.

In hydrogen chloride:

δ+H–Clδ−

Another common notation is an arrow along the bond pointing towards the negative end, with a small cross at the positive end (+→).

Worked examples

1. Cl–Cl Δχ = 3.16 − 3.16 = 0 → non-polar covalent. Identical atoms always share equally.

2. C–H Δχ = 2.55 − 2.20 = 0.35 → essentially non-polar. This is why hydrocarbons such as methane and petrol are non-polar and don’t mix with water.

3. H–Cl Δχ = 3.16 − 2.20 = 0.96 → polar covalent. Cl is δ−, H is δ+.

4. O–H Δχ = 3.44 − 2.20 = 1.24 → polar covalent. O is δ−, H is δ+. The strong polarity of O–H bonds is behind water’s ability to form hydrogen bonds (see hydrogen bonding).

5. C–F Δχ = 3.98 − 2.55 = 1.43 → strongly polar covalent.

6. H–F Δχ = 3.98 − 2.20 = 1.78 → very polar; near the boundary but still covalent. HF is a molecular gas (boiling at 19.5 °C), not an ionic solid.

7. Na–Cl Δχ = 3.16 − 0.93 = 2.23 → ionic. Sodium chloride forms a lattice of Na⁺ and Cl⁻ (see ionic bonding explained).

8. Cs–F Δχ = 3.98 − 0.79 = 3.19 → strongly ionic; one of the most ionic compounds known.

9. C–Cl vs C–I C–Cl: 3.16 − 2.55 = 0.61 (polar). C–I: 2.66 − 2.55 = 0.11 (almost non-polar). Yet C–I bonds are the more reactive — a reminder that reactivity depends on bond strength as well as polarity (see bond length and bond strength).

How ionic is a bond? Percent ionic character

Pauling suggested estimating the percent ionic character of a bond from Δχ with the formula:

% ionic character ≈ [1 − e^(−(Δχ)²/4)] × 100

For H–Cl: (0.96)² / 4 = 0.23; e^(−0.23) = 0.79; % ≈ 21 %. For Na–Cl: (2.23)² / 4 = 1.24; e^(−1.24) = 0.29; % ≈ 71 %.

These numbers are estimates, but they make the key point: no bond is 100 % ionic, and even NaCl has some covalent character. A Δχ of about 1.7 corresponds to roughly 50 % ionic character, which is where the common cut-off comes from.

Polar bonds vs polar molecules

A molecule can contain polar bonds and still be non-polar overall, if the bond polarities cancel out because of the molecule’s symmetry:

• CO₂: each C=O bond is polar, but the molecule is linear (O=C=O), so the two bond dipoles point in opposite directions and cancel. CO₂ is non-polar.
• H₂O: each O–H bond is polar, and the molecule is bent, so the dipoles don’t cancel. Water is polar.
• CCl₄: four polar C–Cl bonds in a symmetrical tetrahedron cancel. Non-polar.
• CHCl₃ (chloroform): one C–H replaces a C–Cl, breaking the symmetry. Polar.

To decide whether a molecule is polar, you need both the bond polarities and the 3D shape (see VSEPR and molecular geometry, polar vs non-polar molecules and dipole moments).

Why bond polarity matters

• Solubility: “like dissolves like”. Polar and ionic substances dissolve in polar solvents such as water; non-polar substances dissolve in non-polar solvents such as hexane.
• Intermolecular forces: polar molecules attract each other through permanent dipole–dipole forces and, with N–H, O–H or F–H bonds, hydrogen bonds — raising boiling points (see intermolecular forces).
• Reactivity: the δ+ carbon in a C–Cl or C=O bond attracts electron-rich reagents (nucleophiles), which is the starting point of many organic reactions.
• Acidity: polar H–X bonds can release H⁺ in water; how easily depends on both polarity and bond strength.
• Bond strength: polarity adds an ionic contribution that strengthens bonds — H–F is one of the strongest single bonds.

Practice: rank these bonds by polarity

Use the electronegativity table to rank these bonds from least to most polar, and say which atom is δ− in each: C–O, N–H, C–N, H–Br, Mg–O.

Working:

Bond	Δχ	δ− atom
H–Br	2.96 − 2.20 = 0.76	Br
C–N	3.04 − 2.55 = 0.49	N
N–H	3.04 − 2.20 = 0.84	N
C–O	3.44 − 2.55 = 0.89	O
Mg–O	3.44 − 1.31 = 2.13	O (mainly ionic)

Order (least → most polar): C–N < H–Br < N–H < C–O < Mg–O.

Notice that C–N is only weakly polar, while the C–O bond is noticeably more polar — one reason the oxygen of alcohols and carbonyl compounds is such a common site of reaction. Magnesium oxide, with Δχ above 2, is best described as an ionic lattice of Mg²⁺ and O²⁻ rather than as molecules with a polar bond.

Common misconceptions

• “A polar bond has full charges.” Only partial charges, δ+ and δ−.
• “Polar bonds always make a polar molecule.” Symmetrical molecules like CO₂ and CCl₄ are non-polar.
• “There’s a sharp line between ionic and covalent.” It’s a continuous spectrum.
• “The more polar bond is always the more reactive.” Bond strength matters too.

Key takeaways

• Electronegativity measures an atom’s pull on shared electrons; it increases across a period and decreases down a group.
• Δχ ≈ 0: non-polar; about 0.4–1.7: polar covalent; above about 1.7–2.0: mainly ionic.
• The more electronegative atom gets δ−, the other δ+.
• A molecule’s overall polarity depends on bond polarities and shape.