Look at the bond between carbon and chlorine in chloromethane, CH₃Cl. Carbon and chlorine share one pair of electrons, but they do not share it evenly. Chlorine holds the pair more of the time. That one fact decides where chloromethane reacts.
Almost every reaction in this course starts from the same question: where do the electrons sit unevenly? This page gives you the tool that answers it.
What it is
Start with the simplest bond, H–H. The two atoms are identical, so they pull on their shared pair equally, and the pair sits in the middle. Now swap one hydrogen for fluorine. In H–F, fluorine pulls harder, and the shared pair spends more of its time near fluorine.
Electronegativity is how strongly an atom pulls on the shared electrons of a bond it is part of. Fluorine's electronegativity is high. Hydrogen's is lower.
The atom that holds the pair more of the time carries a small negative charge, written δ− ("delta minus"). Its partner carries a small positive charge, written δ+. These are partial charges: each is smaller than the charge on one whole electron. No electron has moved from one atom to the other. The pair just spends more of its time at one end.
Electronegativity describes an atom in a bond. Two other measures describe an atom on its own. Ionization energy is the energy needed to pull one electron off a single, unbonded atom. Electron affinity is the energy released when a single atom takes on one extra electron. Neither needs a partner. Electronegativity does.
The numbers on this page come from the Pauling scale, which the chemist Linus Pauling built from the energies of bonds. Because the scale comes from bond energies, it gives a value only for atoms that form bonds. Other scales are built in other ways, and they put the elements in much the same order.
The trend across the periodic table
Walk along the second row of the periodic table: carbon 2.55, nitrogen 3.04, oxygen 3.44, fluorine 3.98. Each step to the right raises the value. Now climb the halogen column, group 17, from the bottom: iodine 2.66, bromine 2.96, chlorine 3.16, fluorine 3.98. Each step up raises it too.
So electronegativity increases up and to the right on the periodic table. Fluorine, at the top right, is the most electronegative element of all. The noble gases in group 18 are left out, because they almost never form bonds.
Moving right along a row (a period) adds protons to the nucleus. The new electrons join the same shell, and electrons in one shell do little to block each other from the nucleus. So the pull on the valence electrons grows, and the atom even shrinks a little. A shared pair feels that stronger pull too.
Moving up a group removes a whole shell of electrons. It removes protons too, so the pull that gets through the inner electrons stays about the same. (Inner electrons blocking the nucleus is the shielding described in Orbitals.) What changes is distance. The valence electrons now sit much closer to the nucleus, so they feel its pull more strongly. Many books also credit less shielding, but the main cause is distance.
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 | Na | 0.93 |
EN is short for electronegativity. You do not need to memorize the numbers. You do need the order, as a reflex: F > O > Cl > N > Br > I ≈ S ≈ C > H. You will use it to decide where charge sits in every structure you draw. The periodic table tool gives the value for any element.
Two pairs in that order often surprise people. First, chlorine edges out nitrogen, 3.16 against 3.04, so the two are nearly tied. Second, carbon and hydrogen are close, 2.55 against 2.20. That is why this course treats a C–H bond as nonpolar, a simplification that holds up well in practice.
Iodine, sulfur and carbon are nearly tied as well (2.66, 2.58 and 2.55), so they share the ≈ signs in the order.
How big a difference makes a bond polar
Take the two values in a bond and subtract the smaller from the larger. For C–H that gives 2.55 − 2.20 = 0.35. For C–O it gives 3.44 − 2.55 = 0.89. For sodium chloride, Na–Cl, it gives 3.16 − 0.93 = 2.23. This difference is called ΔEN, read "delta E N" (Δ means "difference in").
The size of ΔEN sorts bonds into three rough bands.
| 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) |
- A nonpolar covalent bond shares its pair almost evenly.
- A polar covalent bond shares its pair unevenly, so it has a δ+ end and a δ− end.
- In a largely ionic bond, one atom has in effect taken the pair. Solid sodium chloride is made of Na⁺ and Cl⁻ ions, held together by the pull between opposite charges.
The cutoffs are conventions, not laws of nature. Polarity changes smoothly, and a bond at ΔEN = 1.6 is not a different kind of thing from one at 1.8. Some books draw the lines at slightly different numbers, such as 0.4 and 1.8. The bands move a little; the order of the bonds does not.
Because the bands are rough, a bond near a line can behave like its neighbor across it. C–Br has ΔEN = 0.41, just under 0.5, yet chemists treat it as polar: bromomethane, CH₃Br, is about as polar a molecule as chloromethane. The same goes for C–N (0.49) and even C–I (0.11): this course treats every bond from carbon to N, O or a halogen as polar. Bond polarity shows why the length of a bond matters too.
Inside an organic molecule, nearly every bond is covalent. The C–C and C–H bonds of the carbon framework are nonpolar. Bonds such as C–O, O–H and C–Cl are polar covalent. Ionic compounds do turn up, such as salts like sodium acetate (an ionic compound made from an acid), but the bonds within the carbon framework are covalent. The polar bonds are where reactions happen: polar enough to have a reactive end, covalent enough to stay part of one molecule.
Take chloromethane, CH₃Cl. It has three C–H bonds and one C–Cl bond. Work out ΔEN for each.
C–H: 2.55 − 2.20 = 0.35. That is under 0.5, so treat these three bonds as nonpolar.
C–Cl: 3.16 − 2.55 = 0.61. That bond is polar. Chlorine is the more electronegative atom, so chlorine is δ− and carbon is δ+.
So chloromethane has exactly one polar bond, and it leaves the carbon with less than its share of electron density. You have found the reactive site from a table of numbers, before learning a single reaction.
Two traps
Carbon tetrachloride, CCl₄, does the same thing in three dimensions. Its four C–Cl bonds point to the four corners of a tetrahedron (a pyramid with four triangular faces), spread evenly in space. Their four pulls balance, so the molecule is not polar.
A bond's polarity is a property of that one bond. A molecule's polarity comes from adding up all its bond pulls, each with its direction, and the molecule's shape decides whether they cancel. Molecular geometry explains these shapes, and Bond polarity adds up the pulls properly.
Fluorine's tight grip on the shared pair helps hold the C–F bond together, because the δ+ carbon and the δ− fluorine attract each other. So a polarity argument tells you where the charge sits and which end is electron-poor. It does not tell you how much energy it takes to break the bond. Both questions matter, and they have different answers.
The other way electronegativity shows up: induction
Start with acetic acid, the acid in vinegar, CH₃–COOH. Its –COOH group is a carbon double-bonded to one oxygen and single-bonded to an O–H. An acid is a molecule that gives away a hydrogen ion, H⁺, and acetic acid gives away the hydrogen of that O–H.
Now replace one hydrogen of the CH₃ with a chlorine, to get chloroacetic acid, ClCH₂–COOH. The chlorine is not bonded to the COOH group at all. Yet chloroacetic acid gives up its H⁺ far more readily than acetic acid does.
Chlorine pulls electron density out of its own C–Cl bond. The carbon it drains then pulls a little density from its next bond, toward the COOH group, and so on along the chain. The pull passes through the sigma bonds (the single bonds of the framework, from Bonding), and it weakens as it goes. This is the inductive effect, or induction.
Chemists measure acid strength as pKa, introduced in Hybridization: a lower pKa means a stronger acid, and each unit is a factor of ten. The first drawing compares chloroacetic acid with an acid whose chlorine sits further away. The ladder after it adds chlorines one at a time.
Induction adds up. Each extra chlorine lowers the pKa further, though by less each time: 1.9 units for the first, 1.6 for the second and 0.6 for the third. Three chlorines together lower it by about four units, which makes the acid about ten thousand times stronger.
Induction also dies off with distance, as the first drawing shows. In 4-chlorobutanoic acid the chlorine sits two bonds further from the COOH group than in chloroacetic acid. Its effect has almost vanished: pKa 4.52, against 4.76 for acetic acid.
Why does pulling electron density away help the acid give up H⁺? Once the H⁺ has gone, the oxygens are left carrying a negative charge. The chlorine draws some of that charge toward itself, which steadies it, and a charge that is steadied forms more readily. The Acidity Factors section of Acids & Bases works through this in full.
Fading with distance is also how you will tell induction apart from resonance, a second way molecules spread electrons, taught in Drawing Molecules & Moving Electrons. Resonance carries its effect along a connected run of double bonds and lone pairs, and it does not weaken bond by bond the way induction does.
What carries forward
Electronegativity is the first cause behind a great deal of this course. When a later explanation ends in "because oxygen is electronegative", it points back to this page.
- When formal charges must sit somewhere, it decides which of several possible structures is best, in the next two sections.
- It underlies nucleophiles and electrophiles, the names How Reactions Happen gives to the electron-rich and electron-poor partners.
- It drives acidity in Acids & Bases, where it is one of the four factors in the Acidity Factors section.
- It often matches the priority order for R/S assignment in Stereochemistry, though atomic number, not electronegativity, actually sets that order.
- It makes the carbon of a C=O group electron-poor, which drives Carbonyl Chemistry through Enolate Chemistry.