1.5 · Bonding & Lewis

Electronegativity and bond polarity

From nonpolar covalent to ionic on one scale.

By the end you should be able to:

Classify bonds as nonpolar covalent, polar covalent or ionic, and locate partial charges

Key idea

Electronegativity and its trend

Electronegativity (EN, $\chi$ ) is an atom's ability to attract the shared electrons of a bond. On the Pauling scale it increases across a period (left to right) and up a group, so F (3.98) is the most electronegative element. Metals have low values; noble gases are usually not assigned one.

	Group 1	2	13	14	15	16	17
Period 2	Li 0.98	Be 1.57	B 2.04	C 2.55	N 3.04	O 3.44	F 3.98
Period 3	Na 0.93	Mg 1.31	Al 1.61	Si 1.90	P 2.19	S 2.58	Cl 3.16

H is 2.20. A useful order to memorise: F > O > Cl > N > Br > I > S ≈ C > H.

Formula

Classifying a bond by ΔEN

$\Delta\text{EN} = \left|\text{EN}_A - \text{EN}_B\right|$

ΔEN	Bond type	Examples
below 0.4	nonpolar covalent	Cl–Cl (0), C–H (0.35)
0.4 to below 1.8	polar covalent	H–Cl (0.96), O–H (1.24)
1.8 or more	ionic	Na–Cl (2.23), Li–F (3.00)

These cut-offs are approximate and vary between textbooks (some put the ionic boundary at 1.7 or 2.0). Bond character changes gradually, not in steps; this app uses 0.4 and 1.8.

Key idea

Partial charges and bond dipoles

In a polar covalent bond the more electronegative atom carries a partial negative charge ( $\delta^-$ ) and the other a partial positive charge ( $\delta^+$ ): $\mathrm{H}^{\delta+}\!-\mathrm{Cl}^{\delta-}$ .

The bond dipole is drawn as an arrow pointing toward the $\delta^-$ end, with a small cross at the $\delta^+$ tail (the usual chemistry convention).
A larger ΔEN gives larger partial charges and a larger bond dipole.
In an ionic bond the transfer is essentially complete: the higher-EN atom becomes the anion ( $\ce{Na+}$ $\ce{Cl-}$ ).

Method

Classifying a bond and placing δ+ and δ−

Look up both electronegativities.
Compute ΔEN as a positive difference.
Compare with the cut-offs: below 0.4 nonpolar covalent, 0.4 to below 1.8 polar covalent, 1.8 or more ionic.
For a polar bond put $\delta^-$ on the higher-EN atom and $\delta^+$ on the other; for an ionic bond the higher-EN atom is the anion.
To rank bonds by polarity, compare the ΔEN values directly: larger ΔEN, more polar.

Common mistake

Bond-polarity traps

Wrong: "a metal bonded to a nonmetal is always ionic." Right: check ΔEN. For Al–Cl it is $3.16 - 1.61 = 1.55$ , polar covalent, and gas-phase $\ce{AlCl3}$ is molecular.

Wrong: treating the cut-offs as sharp physical boundaries. Right: H–F (ΔEN = 1.78) counts as polar covalent and HF is molecular, yet it sits right next to the ionic cut-off; bonding changes continuously.

Wrong: "polar bonds mean a polar molecule." Right: $\ce{CO2}$ has polar C=O bonds but is nonpolar overall because the dipoles cancel (module 2).

Worked example

Worked example: classify and rank five bonds

Bond	ΔEN	Type	Negative end
C–H	2.55 − 2.20 = 0.35	nonpolar covalent	(C, barely)
N–H	3.04 − 2.20 = 0.84	polar covalent	N
H–Cl	3.16 − 2.20 = 0.96	polar covalent	Cl
O–H	3.44 − 2.20 = 1.24	polar covalent	O
Na–Cl	3.16 − 0.93 = 2.23	ionic	Cl (as $\ce{Cl-}$ )

Answer: polarity increases C–H < N–H < H–Cl < O–H < Na–Cl.

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