A polar covalent bond is a covalent bond in which electrons are shared unequally between two atoms with different electronegativities, giving the more electronegative atom a partial negative charge (δ−) and the other atom a partial positive charge (δ+).
A polar covalent bond forms when two atoms share a pair of electrons, but not fairly. One atom has a stronger pull on the shared electrons than the other. That pull is electronegativity, and per the CED (2.1.A.1), it increases left to right across a period and decreases down a group. When the two bonded atoms have noticeably different electronegativities, the electron density piles up near the more electronegative atom. That atom picks up a partial negative charge (written δ−), and its partner gets a partial positive charge (δ+).
Think of it as a tug-of-war where one side is clearly stronger but never actually wins. The rope (the electron pair) stays shared, just shifted toward the stronger atom. If the atoms have similar electronegativities, the bond is nonpolar covalent. The CED specifically calls out C–H bonds as effectively nonpolar even though carbon is slightly more electronegative than hydrogen (2.1.A.2). If the electronegativity difference is huge, like sodium and chlorine, the electron transfers completely and you get an ionic bond instead. Polar covalent sits in the middle of that bonding spectrum.
Polar covalent bonding lives in Topic 2.1 (Types of Chemical Bonds) and supports learning objective 2.1.A, which asks you to explain how the elements in a bond determine the type of bonding and the resulting properties. It also leans on Unit 1 ideas from Topic 1.8, since the periodic trends and Coulomb's law reasoning behind electronegativity are Unit 1 territory. This term is a gateway concept. Bond polarity feeds directly into dipole moments, molecular polarity, intermolecular forces (Unit 3), and solubility. If you can't identify which bonds are polar, you can't predict whether a molecule dissolves in water, has a high boiling point, or conducts electricity. A huge chunk of Units 2 and 3 quietly depends on this one idea.
Keep studying AP Chemistry Unit 2
Electronegativity (Units 1-2)
Electronegativity difference is the entire cause of bond polarity. Big difference means ionic, moderate difference means polar covalent, tiny difference means nonpolar. You justify electronegativity trends using the shell model and Coulomb's law from Unit 1, so a bonding question is often secretly a periodic trends question.
Dipole Moment (Unit 2)
Each polar bond creates a bond dipole, an arrow pointing toward the more electronegative atom. Add up all the bond dipoles in a molecule and you get the net dipole moment, which decides whether the whole molecule is polar.
Molecular Geometry (Unit 2)
Geometry decides whether bond dipoles cancel. CO₂ has two polar C=O bonds, but its linear shape makes them cancel, so the molecule is nonpolar. Polar bonds plus an asymmetric shape (like bent H₂O) gives you a polar molecule.
Coulomb's Law (Unit 1)
Why does fluorine pull electrons harder than carbon? Coulomb's law. Fluorine has a higher effective nuclear charge acting on valence electrons at a similar distance, so the attraction is stronger. This is the 'explain why' layer the AP exam loves to ask for.
Multiple-choice questions usually hand you pairs of elements and ask you to rank bond polarity, like identifying which pair forms the most nonpolar or least polar covalent bond. The move is always the same. Find each element on the periodic table, estimate electronegativity from position, and compare the differences. You don't get electronegativity values, so you must reason from periodic trends. Questions also test the C–H exception directly, like asking about electron distribution in propane (C₃H₈), where the right answer is that C–H bonds are effectively nonpolar. Property-based questions describe a mystery substance (low melting point, doesn't conduct, dissolves in nonpolar solvents) and ask you to connect those properties back to the bonding type. On FRQs, bond polarity shows up as a justification step. You'll use it to explain dipole moments, intermolecular forces, and solubility, and full credit requires citing the electronegativity difference, not just saying 'it's polar.'
A polar bond and a polar molecule are not the same thing, and the AP exam exploits this constantly. A polar bond is about one connection between two atoms with different electronegativities. A polar molecule requires polar bonds AND a geometry where the bond dipoles don't cancel. CO₂ is the classic trap. It has two polar C=O bonds, but the linear shape makes the dipoles cancel exactly, so CO₂ is a nonpolar molecule. Always check geometry before declaring a molecule polar.
A polar covalent bond shares electrons unequally because one atom has a higher electronegativity, creating partial charges (δ+ and δ−) on the bonded atoms.
Electronegativity increases left to right across a period and decreases down a group, so you can rank bond polarity from periodic table position alone.
Bonding is a spectrum from nonpolar covalent (similar electronegativities) to polar covalent (unequal sharing) to ionic (complete electron transfer).
The CED treats C–H bonds as effectively nonpolar even though carbon is slightly more electronegative than hydrogen, and this exception gets tested.
Polar bonds do not guarantee a polar molecule, because symmetric geometries like linear CO₂ make the bond dipoles cancel.
On FRQs, justify polarity claims by citing the electronegativity difference between the specific atoms, not just by labeling the bond polar.
It's a covalent bond where the shared electron pair sits closer to the more electronegative atom, giving that atom a partial negative charge (δ−) and the other atom a partial positive charge (δ+). It's covered in Topic 2.1 under learning objective 2.1.A.
Effectively nonpolar. The AP CED explicitly states that bonds between carbon and hydrogen are treated as nonpolar even though carbon is slightly more electronegative. This is why hydrocarbons like propane are nonpolar molecules.
In a polar covalent bond, electrons are still shared, just unequally (like H–Cl). In an ionic bond, the electronegativity difference is so large that the electron transfers completely, forming ions (like NaCl). Bonding is a continuum, and polar covalent sits between nonpolar covalent and ionic.
Yes. If the geometry is symmetric, the bond dipoles cancel and the molecule has no net dipole. CO₂ is the standard example, with two polar C=O bonds arranged linearly so the molecule is nonpolar overall.
Use periodic trends. Electronegativity increases toward fluorine (up and to the right), so the farther apart two elements are in that direction, the bigger the electronegativity difference and the more polar the bond. AP multiple-choice questions expect you to reason this way from the periodic table.