Bond polarity — Foundations | Organic Chemistry

This section closes Module 1 by combining its two big ideas. Electronegativity told you which bonds are polar. Molecular geometry told you which directions those bonds point. Put them together and you can predict whether a whole molecule is polar — which determines what it dissolves in, what it boils at, and how it interacts with everything around it.

Bond dipoles

A bond is polar when its two atoms differ in electronegativity: the shared electrons sit closer to one side, leaving that atom δ− and its partner δ+. This separation of charge across a distance is a dipole, and it is a vector — it has both a size and a direction.

The convention is to draw a bond dipole as an arrow pointing toward the more electronegative atom, with a small cross bar at the positive end. In C–O the arrow points at oxygen; in C–Li, unusually, it points at carbon, because lithium is far less electronegative than carbon. That reversal is what makes organolithium reagents so useful: their carbon is the negative end of its bond, so it is the carbon that carries the spare electron density and the carbon that goes and attacks something else.

The drawing convention, and the one case that catches people out. The arrow always points at whichever atom pulls harder, with a cross bar at the end left short. In almost every bond you will draw, carbon is the δ+ end. In C–Li it is not — lithium is so much weaker a puller that the arrow reverses and the carbon carries the spare electron density. A carbon that is the negative end of its bond is a carbon that goes out and attacks things, which is the entire point of organolithium reagents.cross bar marks the δ+ end; the arrowhead marks δ−

A dipole's size scales with both the electronegativity difference and the bond length. This is why C–F, despite having the largest ΔEN of the carbon–halogen bonds, does not have the largest dipole moment — the C–F bond is short. C–Cl, with a smaller ΔEN but a longer bond, edges it out.

Molecular polarity depends on geometry too

A molecule's overall polarity is the vector sum of all its bond dipoles. Vectors of equal size pointing in opposite directions cancel. So a molecule built entirely from polar bonds can be completely nonpolar, provided its geometry is symmetric enough to cancel them.

The working procedure is: draw the Lewis structure, get the geometry from VSEPR, mark each bond dipole as an arrow, then add the arrows as vectors. Lone pairs contribute too — a lone pair is a concentration of electron density pointing away from the atom, so it acts like a dipole in its own right, which is part of why water is as polar as it is.

Worked examples — same bonds, opposite answers

CO₂. Linear, with two strongly polar C=O bonds pointing in exactly opposite directions. They cancel perfectly. Nonpolar, μ = 0 D exactly.

Water. Bent, with two O–H dipoles. Because the molecule is bent rather than linear, the two arrows do not oppose each other — they add to give a net dipole bisecting the H–O–H angle, reinforced by the two lone pairs pointing the other way. Polar, μ = 1.85 D.

CCl₄. Tetrahedral, with four identical polar C–Cl bonds pointing symmetrically to the corners of a tetrahedron. They sum to zero. Nonpolar, μ = 0 D.

CHCl₃. Also tetrahedral, but one chlorine has been replaced by a hydrogen. The symmetry is broken, the cancellation is incomplete, and the molecule is polar, μ = 1.04 D. One substitution, and chloroform behaves as a completely different kind of solvent from carbon tetrachloride.

Why the flat drawing of this one is hard to trust. A tetrahedron's four corners really do point symmetrically away from each other in space, so four identical bond dipoles really do sum to nothing — which is easy to assert on paper and easy to doubt. Turned solid, the symmetry is just visible. Replace one corner and it is equally visible that nothing cancels any more.

Four molecules, every one of them built from strongly polar bonds, and only two of them polar. Molecular polarity is the vector sum of the bond dipoles, so geometry does half the work: arrows of equal size pointing opposite ways cancel, and arrows that do not oppose add. Water is polar because it is bent, not because O–H is polar — and CCl₄ and CHCl₃, identical in every bond type, land on opposite sides of the answer because one is symmetric and the other is not.

Dipole moment: a measurable quantity

Molecular polarity is quantified as the dipole moment (μ), measured in debyes (D). Some values worth having a feel for: CO₂ 0 D, CCl₄ 0 D, toluene 0.36 D, chloroform 1.04 D, chloromethane 1.87 D, water 1.85 D, ethanol 1.69 D, acetone 2.88 D, DMSO 3.96 D, acetonitrile 3.92 D.

Two things in that list are worth a second look. Carbonyl compounds sit high, because a C=O bond is both strongly polar and, in most geometries, impossible to cancel — a molecule with one is polar almost regardless of what else it has.

And water, the substance most people would name first if asked for a polar liquid, is not at the top — DMSO and acetone both beat it. Dipole moment measures one specific thing: the charge separation within a single isolated molecule. What makes water behave the way it does as a solvent is something else, which the next few paragraphs get to.

"Like dissolves like." Polar solvents dissolve polar solutes because the favorable dipole–dipole interactions between them repay the cost of separating the solvent molecules from each other; nonpolar solutes cannot pay that cost, so they stay out. This is why oil and water separate, why a long hydrocarbon chain makes an alcohol progressively less water-soluble as it grows, and why extraction — the workhorse purification technique of every organic lab — works at all.

Why polarity shows up in physical properties

Polarity governs the strength of intermolecular forces, and intermolecular forces govern boiling point. Three tiers, weakest to strongest:

London dispersion forces act between all molecules, arising from momentary fluctuations in electron distribution. They are weak individually but scale with surface area, which is why pentane boils at 36 °C and decane at 174 °C despite both being nonpolar.

Dipole–dipole forces act between permanent dipoles and are stronger. Acetone (μ = 2.88 D, MW 58) boils at 56 °C, while butane (nonpolar, MW 58) boils at −1 °C. Nearly sixty degrees of difference at identical molecular weight, purely from polarity.

Hydrogen bonding is a special, unusually strong case of dipole–dipole attraction, available when hydrogen is bonded directly to N, O or F. Ethanol (MW 46) boils at 78 °C; dimethyl ether (MW 46, same formula, no O–H) boils at −24 °C. A hundred degrees from one hydrogen being on oxygen instead of on carbon.

Same formula, C₂H₆O. Same molecular weight. A hundred degrees apart. Ethanol’s hydrogen sits directly on oxygen, stripped of most of its electron density, so it can reach across to a lone pair on the next molecule — and pulling those molecules apart to boil them costs real energy. Dimethyl ether has the lone pairs to accept a hydrogen bond but no O–H to donate one, so its molecules barely hold on to each other. Hydrogen bonding needs both halves.

"Has polar bonds" and "is a polar molecule" are different claims. Carbon dioxide has two of the most polar bonds in common chemistry and a dipole moment of exactly zero. When a question asks about molecular polarity, the geometry is doing half the work — you cannot answer it from electronegativity alone.

What carries forward

Polarity is the practical face of Module 1. It decides solvent choice in every reaction you will run, and solvent choice is what tips SN1 against SN2 in Module 6. It explains the boiling-point and solubility trends you will be asked to rank for alcohols, ethers, aldehydes and carboxylic acids. And the δ+ carbon of a polar bond is, quite literally, the atom that gets attacked in the majority of the mechanisms ahead — the next module gives that idea its proper name: an electrophile.