Electronegativity is an atom’s ability to attract shared electrons in a chemical bond. It helps explain why some bonds share electrons evenly while others pull electron density toward one atom. This idea matters because bond polarity affects molecular shape, solubility, boiling point, and many chemical reactions.
A small difference in electron attraction can change how a molecule behaves in water, in cells, or in a lab reaction.
When two atoms form a bond, the more electronegative atom pulls the shared electron cloud closer to itself. This gives that atom a partial negative charge, written δ-, while the other atom gets a partial positive charge, written δ+. The electronegativity difference, often written ΔEN, is used to estimate whether a bond is nonpolar covalent, polar covalent, or ionic.
Across the periodic table, electronegativity generally increases from left to right and decreases from top to bottom.
Understanding Chemistry: Electronegativity and Bond Polarity
The pull on bonding electrons comes from the positive nucleus inside each atom. A nucleus with more protons can attract electrons more strongly, but distance matters too. Electrons in inner shells block part of the nuclear pull from electrons involved in bonding.
This blocking effect is called shielding. As atoms get larger down a group, their outer bonding electrons sit farther from the nucleus and feel more shielding. Their attraction for shared electrons becomes weaker.
Across a period, proton number rises while the outer electrons remain in roughly the same shell. The stronger effective nuclear pull makes the atoms toward the right more able to draw bonding electrons closer.
Electronegativity values come from measurements and comparisons, not from a direct count of electrons. Different scales exist, so the numbers are best used as guides. Bonding is not divided by sharp walls into three completely separate types.
A bond with a large electronegativity difference may behave mostly like an ionic bond, yet some electron density can still remain between the atoms. The structure of a substance changes the result. In a solid crystal, oppositely charged ions are held in a repeating arrangement.
In a gas or liquid, individual molecules may show a range of unequal electron sharing. This is why the usual boundary values are useful estimates rather than absolute laws.
Bond polarity does not automatically mean that the whole molecule is polar. The directions of the bond pulls matter. Carbon dioxide has two polar carbon oxygen bonds, but its straight shape makes the pulls balance.
The molecule has no overall polar end. Water has polar oxygen hydrogen bonds in a bent shape, so the pulls do not cancel. Water molecules attract one another strongly through hydrogen bonding.
This helps explain water's unusually high boiling point and its ability to dissolve many ionic and polar substances. Oils contain many carbon hydrogen bonds with little polarity, so they do not mix well with water. Soap works because one part of each soap molecule interacts with water while another part interacts with oils.
Polarity helps predict where chemical reactions occur. A partially positive region can attract electron rich particles, while a partially negative region can attract electron poor particles. In biology, the polar shapes and charge patterns of molecules help enzymes recognize their reactants.
In batteries and salt solutions, charged particles move through polar liquids because the liquid molecules surround and stabilize them. When solving chemistry problems, first identify the bonded elements and compare their electronegativity values. Then mark the direction of electron pull before deciding whether the full molecular shape cancels it.
Keep partial charges separate from full ionic charges. A polar covalent bond has unequal sharing, not necessarily a complete transfer of an electron.
Key Facts
- Electronegativity is an atom’s ability to attract shared electrons in a chemical bond.
- ΔEN = |EN atom 1 - EN atom 2| measures the electronegativity difference between two bonded atoms.
- A nonpolar covalent bond usually has ΔEN from 0 to about 0.4, so electrons are shared nearly equally.
- A polar covalent bond usually has ΔEN from about 0.5 to 1.7, so electrons are shared unequally.
- An ionic bond usually has ΔEN greater than about 1.7, so electron transfer becomes a useful model.
- Electronegativity generally increases left to right across a period and decreases down a group.
Vocabulary
- Electronegativity
- Electronegativity is the ability of an atom to attract shared electrons in a chemical bond.
- Bond polarity
- Bond polarity is the uneven distribution of electron density between two bonded atoms.
- Partial charge
- A partial charge is a small charge imbalance in a polar bond, shown as δ+ or δ-.
- Dipole
- A dipole is a separation of positive and negative charge within a bond or molecule.
- Electron density
- Electron density describes where electrons are most likely to be found around atoms or within a bond.
Common Mistakes to Avoid
- Treating all covalent bonds as nonpolar is wrong because many covalent bonds have unequal electron sharing when the atoms have different electronegativities.
- Pointing the dipole arrow toward the less electronegative atom is wrong because electron density is pulled toward the more electronegative atom.
- Confusing partial charges with full ionic charges is wrong because δ+ and δ- show unequal sharing, not complete electron transfer.
- Using electronegativity trends backward is wrong because electronegativity generally increases across a period from left to right and decreases down a group.
Practice Questions
- 1 The electronegativity of H is 2.20 and the electronegativity of Cl is 3.16. Calculate ΔEN for the H-Cl bond and classify the bond as nonpolar covalent, polar covalent, or ionic.
- 2 The electronegativity of C is 2.55 and the electronegativity of O is 3.44. Calculate ΔEN for a C-O bond and identify which atom has δ-.
- 3 A molecule has two identical polar bonds arranged in opposite directions in a straight line. Explain why the molecule may be nonpolar even though each bond is polar.