Electronegativity — Foundations

If you had to reduce organic reactivity to a single question, it would be: where are the electrons unevenly distributed? Electronegativity is the answer to that question. Every polar bond, every partial charge, every arrow you will draw from an atom with electrons to spare to an atom short of them traces back to one atom pulling harder on a shared pair than its partner does.

What it is

Electronegativity measures how strongly an atom pulls on the electrons in a bond it is part of. It is a property of an atom in a bond, not of an isolated atom — which distinguishes it from related quantities like ionization energy or electron affinity that describe atoms on their own.

On the periodic table it increases up and to the right, excluding the noble gases. Fluorine, top right of the reactive elements, is the most electronegative element there is. Two trends combine to produce this. Moving right across a period adds protons to the nucleus without adding a new shell, so the nuclear pull on the valence electrons strengthens. Moving up a group puts the valence electrons in a smaller shell, closer to the nucleus and less shielded by inner electrons, so the pull is felt more strongly.

The trend, shaded by how hard each atom pulls. Two separate causes point the same way: going right adds protons without adding a shell, so the nucleus grips harder; going up puts the valence electrons in a smaller, less shielded shell, where they feel that grip more. Noble gases are left blank — they are not making bonds to pull on. You do not need the numbers, but you do need the order F > O > Cl > N > Br > I ≈ S ≈ C > H as a reflex.

In any bond between two different atoms, the shared electrons sit closer to whichever atom is more electronegative. That atom takes on a partial negative charge, written δ−; its partner takes a partial positive charge, δ+. These are partial charges, not full ones — no electron has actually been transferred, it is just spending more of its time at one end.

The Pauling scale, and the values worth knowing

Element	EN	Element	EN
F	3.98	C	2.55
O	3.44	S	2.58
Cl	3.16	I	2.66
N	3.04	H	2.20
Br	2.96	Li	0.98

You do not need to memorize the numbers, but you do need the ordering F > O > Cl > N > Br > I ≈ S ≈ C > H as a reflex, because you will use it to decide where charge sits in every structure you draw. You can look up any element in the periodic table tool.

Two entries in that list regularly surprise people. Nitrogen and chlorine are nearly tied (3.04 vs 3.16), so chlorine actually edges out nitrogen. And carbon and hydrogen are close (2.55 vs 2.20), which is why C–H bonds are treated as essentially nonpolar throughout the course — a convenient simplification that holds up well in practice.

How big a difference makes a bond polar

The useful quantity is the difference in electronegativity between the two bonded atoms, ΔEN.

ΔEN	Bond type	Example
< 0.5	Nonpolar covalent	C–H (0.35)
0.5 – 1.7	Polar covalent	C–Cl (0.61), C–O (0.89)
> 1.7	Largely ionic	Na–Cl (2.23)

The boundaries are conventions, not physical laws — polarity is a continuum, and a bond at ΔEN = 1.6 is not categorically different from one at 1.8. Nearly every bond in organic chemistry falls in the polar covalent band, which is the interesting middle ground: polar enough to have a reactive end, covalent enough to stay a molecule.

Worked example — reading a molecule for polarity

Take chloromethane, CH₃Cl. Compare each bond: C–H has ΔEN = 0.35, effectively nonpolar. C–Cl has ΔEN = 0.61, polar, with chlorine δ− and carbon δ+.

So there is exactly one meaningfully polar bond, and it leaves the carbon electron-poor. That carbon is the electron-poor site, and it is where anything electron-rich will go. You have located the reactive site of a molecule from a table of numbers, before knowing a single mechanism.

Chloromethane, read for polarity. Three C–H bonds are close enough in electronegativity to ignore; the one C–Cl bond is not, so chlorine holds the shared pair a little more of the time and carries a partial negative charge — δ−, a lean, not a transfer. The carbon left behind is δ+. That single conclusion, reached from a table of numbers before you have seen a mechanism, is where every substitution reaction in Module 6 begins.and that is the atom something electron-rich will attack

This is where the phrases "electron-poor" and "electron-rich" come from, and they are the vocabulary the rest of the course is written in. An electronegativity difference makes a bond polar, and a polar bond makes the carbon in C–Br electron-poor enough that something electron-rich will come and attack it — which is the entire premise of the substitution reactions in Module 6.

Two traps

Polar bonds do not guarantee a polar molecule. Carbon dioxide has two strongly polar C=O bonds, and a dipole moment of exactly zero — the molecule is linear, so the two bond dipoles point in opposite directions and cancel. Carbon tetrachloride does the same thing in three dimensions. Bond polarity is a property of a bond; molecular polarity is the vector sum, and geometry decides it. This is taken up properly in the Bond Polarity section.

Polarity is not the same thing as weakness. It is tempting to read a large ΔEN as "this bond is ready to come apart," and the halogens say otherwise. C–F is the most polar carbon–halogen bond and also the strongest, at about 116 kcal/mol against roughly 57 for C–I, the least polar of the four. The same tight grip on the shared electrons that pulls them toward fluorine is what holds the bond together. So a polarity argument tells you where the charge sits and which end is electron-poor; it does not tell you what it costs to break anything. Both questions matter, and they have different answers.

The other way electronegativity shows up: induction

An electronegative atom does not only polarize the bond it is in. It pulls electron density through the sigma framework from further away, weakening with each bond it has to travel through. This is the inductive effect, and it is cumulative.

The classic demonstration is acid strength. Acetic acid (CH₃COOH) has a pKa of 4.76. Replace one methyl hydrogen with chlorine and chloroacetic acid drops to 2.86 — over eighty times more acidic — because the chlorine pulls electron density away from the negatively charged oxygens left behind and helps stabilize that charge. Add two chlorines and it falls to 1.29; three, and trichloroacetic acid is at 0.65, a stronger acid than phosphoric acid. Move the chlorine further from the acid group and the effect fades quickly: 4-chlorobutanoic acid is barely more acidic than butanoic acid itself. Induction works through bonds and dies off over distance, which is the fingerprint that distinguishes it from resonance.

The inductive effect, measured. Each chlorine tugs electron density along the sigma bonds, away from the oxygens that will carry the negative charge once the acid gives up its proton — and a charge that is helped is a charge more willingly formed. Three chlorines move the pKa by four units, a factor of ten thousand. The last row is the tell: the same chlorine, three bonds further off, does almost nothing. Induction fades with distance — which is exactly how you tell it apart from resonance, which does not.Move the same chlorine three bonds away and the effect all but vanishes.

What carries forward

Electronegativity is the first cause behind an unreasonable amount of this course. It sets formal charge placement in the next two sections; it defines nucleophiles and electrophiles in Module 2; it drives acidity in Module 3 and is one of the four factors in the Acidity Factors section; it determines the priority order for R/S assignment in Module 5; and it is what makes carbonyl carbons electrophilic in Modules 9 through 11. When a later explanation bottoms out in "because oxygen is electronegative," this is the section it is pointing back at.