Lesson overview
Relative masses: How heavy is a water molecule?
Blue crystals hold water you can’t see. Predict how much white solid 2.50 g of them will leave, heat them to constant mass on a digital balance, test the white solid with water, then repeat with 5.00 g and find the relative mass of one water molecule.
Suggested sequence
- Lesson 1 · How heavy is a water molecule? Relative atomic mass, relative molecular mass and relative formula mass. Read together and discuss (15–20 min), practise with self-checking (20–25 min).
- Lesson 2 · Where does the mass go? Reacting masses in simple proportions. Read together (10–15 min), show or explore the experiment (10 min), practise with self-checking (15–20 min).
- End of module · final check. 4 new questions, 20–30 min, without hints or notes; pupils check their answers once they have answered them all.
- One to two weeks later · return questions. 3 questions, 15 min, lesson closed.
Learning goals
Part 1: How heavy is a water molecule?
- Describe relative atomic mass, Ar, as the average mass of the isotopes of an element compared to 1/12th of the mass of an atom of 12C
- Define relative molecular mass, Mr, as the sum of the relative atomic masses. Relative formula mass, Mr, will be used for ionic compounds
- Calculate the relative atomic mass of an element from the relative masses and abundances of its isotopes
Part 2: Where does the mass go?
- Define relative molecular mass, Mr, as the sum of the relative atomic masses. Relative formula mass, Mr, will be used for ionic compounds
- Calculate reacting masses in simple proportions. Calculations will not involve the mole concept
Misconceptions to listen for
- “The relative atomic mass is the mass number of the atom.” — The mass number counts the protons and neutrons in one atom of one isotope. Ar is the average of all the isotopes, so it need not be a whole number.
- “Relative masses are in grams.” — Ar and Mr compare masses with 1/12 of the mass of a carbon-12 atom, so they have no unit.
- “Ar is the mass of the most common isotope.” — The most common isotope pulls the average towards it, but every isotope counts.
- “108 is the mass of a third isotope, silver-108.” — Silver has only two isotopes, 107 and 109. 108 is their average, not a third kind of atom.
- “The heavier isotope is always more common.” — If silver-109 were more common, the average would be above 108. Exactly halfway means about equal amounts.
- “The isotopes are equally common, because 69.7 is between 69 and 71.” — Equal amounts would give exactly 70. Any average lies between the isotopes; its position shows which one is more common.
- “Ar tells you nothing about how common each isotope is.” — It does: where the average lies between 69 and 71 shows which isotope is more common.
- “Row A.” — K2 means two potassium atoms: 2 × 39 + 16 = 94, not 39 + 16 = 55.
- “Row B.” — CH4 has four hydrogen atoms: 12 + 4 × 1 = 16. The small 4 counts the atoms of the element before it.
- “Row D.” — Potassium oxide is made of ions and has no molecules, so its Mr is a relative formula mass.
Scope
Relative masses (3.2 Core 1–3): relative atomic mass as the average mass of the isotopes of an element compared to 1/12 of the mass of an atom of carbon-12, relative molecular mass and relative formula mass as the sum of the relative atomic masses, and reacting masses in simple proportions, without the mole. The mole and molar mass, calculating relative atomic mass from isotope abundances, percentage yield and purity are left to other modules.