AP Chemistry · Unit 1 · Topic 1.1
Moles and Molar Mass
A balance measures grams, but a chemical equation counts particles. The mole is the bridge between those two descriptions. Before reaching for a calculator, decide what you are counting: atoms, molecules, formula units, or ions. That decision determines which conversion is meaningful.
The target is an amount, not a new kind of mass
Explain how the same number of particles can correspond to different masses, interpret subscripts as particle ratios, and use units to defend a calculation. One mole contains exactly 6.02214076 × 1023 specified entities. A mole of helium atoms and a mole of carbon dioxide molecules have equal entity counts, but their masses differ because the entities differ.
Read the formula before calculating
Molar mass is the mass per mole, usually expressed in g mol−1. Build it from the entire formula, including parentheses. For Ca(NO3)2, one formula unit contains one calcium, two nitrogen, and six oxygen atoms. An ionic solid is an extended array, so calling those units separate calcium nitrate molecules would misrepresent its structure.
Three questions, three conversion factors
| Starting information | Wanted quantity | Conversion |
|---|---|---|
| Mass of substance | Amount of substance | Divide by its molar mass |
| Amount of substance | Number of entities | Multiply by Avogadro's constant |
| Amount of a compound | Amount of one constituent | Multiply by the formula subscript |
Keep substance names alongside units. “mol oxygen atoms” and “mol O2 molecules” are not interchangeable, even though both contain the word mole.
Worked example: a measured carbonate sample
A dry sample contains 5.30 g of pure Na2CO3. Using 105.99 g mol−1, first obtain formula-unit amount. Then use the sodium subscript.
The final value counts sodium ions in the solid. Multiplying by three instead would count oxygen atoms. The exact subscript does not reduce the three significant figures justified by the measured mass.
What a preparation error does
A student weighs a damp sample but treats every gram as anhydrous carbonate. The inferred carbonate amount is too high because some measured mass belongs to water. Drying to constant mass addresses that specific uncertainty; repeating the same damp weighing more precisely does not. Precision cannot repair a systematically wrong chemical assumption.
A useful reverse check
Reconstruct the mass from the calculated moles. Here, 0.0500 mol multiplied by about 106 g mol−1 gives about 5.30 g. Also ask whether the particle count should exceed one mole's count: this sample is only one twentieth of a mole of formula units, so its formula-unit count must be smaller than Avogadro's constant.
Avoid the common shortcut trap
Never divide by a formula subscript when you mean molar mass, and never use the atomic mass of a single element for an entire compound. A coefficient in a reaction changes how many moles react; it does not change the molar mass printed for the substance.
Practice, then check your reasoning
How many moles of oxygen atoms occur in 0.120 mol CO2?
Answer and reasoning: 0.240 mol oxygen atoms: every molecule contributes two oxygen atoms.
Two samples contain 0.20 mol He and 0.20 mol Ne. Which has more atoms?
Answer and reasoning: Neither. Their atom counts are equal; neon has the larger sample mass.
Scientific references: OpenStax: formula mass and the mole. OpenStax Chemistry 2e authors: Paul Flowers, Klaus Theopold, Richard Langley, and William R. Robinson. Consult source pages for their license terms. Accessed September 13, 2026. Explanations, comparison tables, and practice tasks were independently written for this site; numerical teaching datasets are illustrative unless explicitly identified otherwise.
Official curriculum reference: College Board AP Chemistry course page. The explanation and worked example are independently written for this study site.