Chemistry · Ch 5 — Chemical Bonding
Polarity of a Covalent Bond
Polarity of a Covalent Bond
A covalent bond's shared electron pair is not always shared perfectly equally between its two atoms — whether it is depends on whether the two bonded atoms have the same or different electronegativity. Between two atoms of the SAME element, or of very similar electronegativity (e.g. H–H, F–F, Cl–Cl), the shared pair is attracted equally by both nuclei and sits midway between them: this is a NONPOLAR covalent bond. Between two atoms of markedly DIFFERENT electronegativity (e.g. H–F, H–Cl), the shared pair is instead pulled noticeably closer to the more electronegative atom, separating the bond's overall charge into a partial negative charge (δ−) on the more electronegative atom and a partial positive charge (δ+) on the other — this charge separation is called a dipole, and the resulting bond is a POLAR covalent bond. A bond's polarity (the size of this charge separation) increases as the electronegativity difference between its two atoms increases; a difference of roughly 1.7 corresponds to about 50% ionic character within the bond. Dipole moment is a VECTOR quantity — conventionally drawn as a small arrow with its tail at the positive centre and its head pointing toward the negative centre, or as a crossed arrow above a Lewis structure — and for a POLYATOMIC molecule, the overall (net, resultant) dipole moment is the VECTOR SUM of every individual bond dipole in the molecule PLUS any contribution from lone pairs, so it depends jointly on the molecule's geometric shape (which fixes the DIRECTIONS the individual bond dipoles point in) and on the size of each bond dipole and lone pair. This is why some molecules built entirely from polar bonds (e.g. linear BeF2, trigonal-planar BF3) still end up with ZERO net dipole moment — their individual bond dipoles are symmetric enough to cancel exa …
What this figure shows. BeF2 is linear, so its two Be–F bond dipoles point in exactly opposite directions and are of equal magnitude; their vector sum is zero, so BeF2's net (molecular) dipole moment is zero despite each individual Be–F bond being polar. BF3 is trigonal planar and symmetric, with its three B–F bonds oriented 120° apart; any two of the three bond dipoles sum to a resultant exactly equal and opposite to the third, so the vector sum of all three is again zero — BF3's net dipole moment is likewise zero, des …
What this figure shows. Both NH3 and NF3 are pyramidal, with one lone pair on nitrogen, but their resultant dipole moments differ sharply: NH3 has C m (≈1.47 D) while NF3 has only C m (≈0.23 D). Because H is LESS electronegative than N, the three N–H bond dipoles point from H toward N; nitrogen's lone-pair orbital dipole points in that SAME general direction, so it ADDS to the bond dipoles' resultant, giving NH3 a comparatively large net dipole moment. Because F is MORE electronegative than N, the three N–F bond dipoles instead point from N toward F — the OPPOSITE direction to nitrogen's lone-pair orbital dipole — so the lone pair's contribution SUBTRACTS from (partially cancels) the …
Table 5.8, grouped by molecule type: Type AB — HF 1.91 D, linear; HCl 1.03 D, linear; HBr 0.79 D, linear; H2 0 D, linear. Type AB2 — H2O 1.85 D, bent; H2S 0.95 D, bent; CO2 0 D, linear. Type AB3 — NH3 1.47 D, trigonal pyramidal; NF3 0.23 D, trigonal pyramidal; BF3 0 D, trigonal planar. Type AB4 — CH4 0 D, tetrahedral; CHCl3 1.04 D, tetrahedral; CCl4 0 D, tetrahedral. In every zero-dipole case the molecule is symmetric (all B atoms/lone pairs arranged so bond dipoles cancel exactly); CHCl3's non-zero value (unlike CCl4's zero) is because replacing one Cl of C …