Two polar bonds. One nonpolar molecule, one wildly polar one.
CO2 has two polar bonds and is a nonpolar molecule. Water has two polar bonds and is one of the most polar molecules there is. Same ingredient, opposite result: the difference is shape. By the end, you can use bond polarity and shape to predict a molecule's polarity and the intermolecular forces it can have.
Core Claims
- Molecular Polarity: Depends on both bond polarity (ΔEN) and molecular shape.
- Symmetry & Cancellation: In symmetric shapes like linear CO2, opposing bond dipoles cancel out. This leaves the overall molecule nonpolar.
- Asymmetry & Reinforcement: Asymmetric shapes (like bent H2O) cause dipoles to reinforce, creating a net dipole and a polar molecule.
Dipole Cancellation vs. Reinforcement
Retrieval Checklist
- Draw vectors representing bond dipoles.
- Determine if molecular geometry cancels or reinforces dipoles.
- Predict whether a molecule is polar or nonpolar.
"The bonds are polar, so the molecule is polar." Sometimes.
Here is a trap worth naming: the bonds are polar, so the molecule is polar. Sometimes. Electronegativity difference (ΔEN) tells you whether a single bond dipole exists. That is one tug-of-war, judged bond by bond. Molecular polarity asks a different question. Do all of the molecule's tugs-of-war add up to a net pull? Pulls in opposite directions cancel. Pulls in the same general direction do not.
CO2 is the clean case. Each C=O bond really is polar: oxygen wins the tug-of-war, so oxygen ends up δ− and carbon δ+. But CO2 is a straight line, and the two oxygens sit on exactly opposite sides of carbon. Their pulls point in exactly opposite directions and cancel, not approximately, exactly. Nonpolar molecule, polar bonds and all.
Water refuses to line up the same way. Its two O–H bonds sit at about 104° to each other: bent, not straight. Both pulls point generally toward oxygen, so instead of canceling they gang up into a net dipole. The oxygen end of the molecule runs δ−, and the hydrogen end runs δ+. Water is polar because of its shape. It is not because O–H (ΔEN 1.24) is somehow “more polar” than C=O (ΔEN 0.89). Both bonds are genuinely polar. The difference is entirely geometric.
Molecular polarity takes two inputs: how polar the bonds are (ΔEN), and where they point (shape). Polar bonds in a symmetric arrangement usually cancel (giving a nonpolar molecule). Polar bonds in a lopsided arrangement don't (giving a polar molecule). Skip either input and you will get CO2 wrong, or water wrong, or both.
Saying “the bonds are polar, so the molecule is polar.” CO2 has two polar bonds and is a nonpolar molecule.
Watch the two-step check on methane. Then run it yourself on ammonia.
Same two questions every time: is the bond polar, and does the shape let those pulls cancel?
Five molecules. Same two-step check every time.
Pick a molecule. Judge its bond, then judge what its shape does to that bond's pull. The verdict appears only after you commit to both.
Do the bond pulls in this molecule cancel or reinforce?
Pick a molecule above to begin
- Lone pair, or the negative end (δ−)
- The positive end (δ+)
- Bond dipole: points from δ+ toward δ−
- Net dipole of the molecule
Practice classifying molecular polarity
Test your understanding of bond polarity and shape. Sort each molecule into the correct category below.
Fill the blanks from memory.
Type each blank from memory, then press Check. Press Show answers after your first check if you are stuck.
Electronegativity difference (ΔEN) tells you whether a single bond has a . Molecular shape decides whether those bond dipoles cancel or . A perfectly symmetric arrangement of polar bonds makes their pulls , giving a nonpolar molecule. A polar molecule picks up forces on top of the London dispersion forces every molecule has. Hydrogen bonding only switches on when hydrogen sits directly on nitrogen, oxygen, or .
Trace the whole chain before you write.
Shape → molecular polarity → IMF type → energy needed to separate molecules → physical state. A shortcut such as "water has O–H bonds, so it has hydrogen bonds" can give a correct answer for a flawed reason. It fails on harder molecules. Write out every link before you check.
Write your answer first. Then grade yourself.
Give yourself 1 mark for each idea you wrote down.
Methane (CH4) boils at −162°C. Ammonia (NH3) boils at −33°C (about 130 degrees higher) even though the two molecules have nearly identical mass (16.05 vs. 17.04 g/mol). Construct the full argument for the difference: bonds → shape → molecular polarity → IMF → boiling point.
- Bond polarities (1 mark): You say C–H is nonpolar (ΔEN = 0.35) and N–H is polar (ΔEN = 0.84).
- Shape to polarity (1 mark): You say CH4's symmetry cancels its pulls, but NH3's pyramid adds them into a net dipole.
- IMF difference (1 mark): You say CH4 has LDF only, while NH3 also has dipole-dipole forces and hydrogen bonding.
- Boiling point (1 mark): You say stronger attractions need more energy to break, so NH3 boils higher.
Microwave ovens operate on molecular polarity. Microwaves cook food by emitting radiation that matches the natural rotational frequency of polar water molecules. The waves make the water's net dipoles spin rapidly, generating heat through friction, while nonpolar containers (like plastics) are left cold and unaffected.
Adaptive Competency Practice: HS-PS1-3
Test your mastery of molecular polarity, dipole cancellation, and intermolecular forces one question at a time. The adaptive engine identifies individual skill gaps, provides tutoring hints on missed questions, and awards an official authenticated completion receipt.