Tenth Grade · Science · week 20

Molar mass, percent composition, and empirical formulas

Calculate molar mass from a chemical formula, use it to convert between grams and moles, calculate percent composition, and determine an empirical or molecular formula.

Molar mass is the mass in grams of one mole of a substance, and you calculate it by adding up the atomic masses (from the periodic table) of every atom in the formula. Water, H2O, has a molar mass of 2(1.01) + 16.00 = 18.02 g/mol. Molar mass is the bridge between the mole (a count) and the gram (a mass you can actually measure on a balance): grams = moles x molar mass, and moles = grams / molar mass.

Percent composition tells you what fraction of a compound's total mass comes from each element: percent = (mass of that element in the formula / molar mass of the whole compound) x 100. In water, hydrogen is (2.02 / 18.02) x 100 = 11.21%, and oxygen is (16.00 / 18.02) x 100 = 88.79% -- and the percentages for every compound must always add up to 100%.

An empirical formula is the simplest whole-number ratio of atoms in a compound, and you can find it from percent composition: treat the percentages as grams (assuming a 100 g sample), convert each element's grams to moles, then divide every mole value by the smallest one to get a whole-number ratio. If you also know the compound's actual molar mass, you can find the true molecular formula: divide the actual molar mass by the empirical formula's mass to see how many times bigger the real molecule is, then multiply every subscript in the empirical formula by that factor.

📋 For the teacher
Teach it (10–15 min):
  • Treat molar mass calculation as a simple, repeatable checklist: list every element, multiply each by its subscript, add the atomic masses from the periodic table, and add it all up -- slowing down to write each step prevents most arithmetic slips.
  • Point out explicitly that "grams = moles x molar mass" and "moles = grams / molar mass" are the same relationship rearranged, exactly like the density triangle from earlier this year -- one relationship, two directions.
  • For empirical formula problems, the "divide everything by the smallest" step is where most mistakes happen; have them double-check that the smallest value in their own list divides to exactly 1.00 before trusting the rest of the ratios.
  • The glucose example (CH2O empirical, C6H12O6 molecular) is a genuinely satisfying one to sit with -- point out that glucose, corn syrup ingredients, and other "simple sugars" share this same empirical formula family.
Talk about it:
  • Why is molar mass described as the bridge between moles and grams?
  • What does percent composition tell you about a compound?
  • Why might two different compounds share the same empirical formula but have different molecular formulas?

Mastery looks like: They correctly calculate molar mass, convert between grams and moles, calculate percent composition, and determine an empirical formula (and molecular formula, given the actual molar mass) from percent composition.

Common stumbles: Forgetting to multiply an atomic mass by its subscript when finding molar mass; dividing by the largest mole value instead of the smallest when finding an empirical formula ratio.

Homework for this lessonType it, write it with a stylus, or print it for pencil & paper.open →
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