A student argues: 'Carbon's atomic mass is 12.011 amu, so 12.011 grams of carbon must contain exactly one carbon atom.' What is wrong with this claim?
ANothing is wrong — 12.011 grams of carbon does contain exactly one carbon atom.
BThe claim confuses amu (a unit for individual atoms) with grams; 12.011 grams of carbon actually contains 6.022 × 10²³ atoms — one mole.
CThe error is that carbon's atomic mass is not exactly 12.011 amu, only carbon-12 has exactly 12 amu.
DThe student should have said 12.011 grams contains 12 atoms, one per atomic mass unit.
The atomic mass unit (amu) is an extraordinarily tiny unit — one amu is about 1.66 × 10⁻²⁴ grams. A single carbon atom weighs roughly 2 × 10⁻²³ grams, not 12 grams. The key insight is that molar mass (in g/mol) is numerically equal to atomic mass (in amu), but the units are completely different. 12.011 g/mol means one mole — 6.022 × 10²³ atoms — weighs 12.011 grams, not that one atom does.
Question 2 Multiple Choice
What is the molar mass of glucose (C₆H₁₂O₆), given C = 12.011 g/mol, H = 1.008 g/mol, O = 16.00 g/mol?
C84.07 g/mol — summing the atomic masses without multiplying by subscripts
D96.06 g/mol — using only the heaviest atoms (6 × 16.00)
Molar mass for a compound is found by summing the molar masses of every atom in the chemical formula, accounting for subscripts. For C₆H₁₂O₆: 6 × 12.011 = 72.066, 12 × 1.008 = 12.096, 6 × 16.00 = 96.00, total = 180.16 g/mol. This value means one mole of glucose — 6.022 × 10²³ molecules — weighs 180.16 grams.
Question 3 True / False
The numerical value of an element's molar mass in g/mol is equal to its atomic mass in amu.
TTrue
FFalse
Answer: True
This numerical equality is the bridge between the atomic and laboratory worlds. Carbon's atomic mass is 12.011 amu, and its molar mass is 12.011 g/mol. The numbers match by definition — one mole (6.022 × 10²³ atoms) was defined precisely so that this equality holds. This is what allows chemists to convert seamlessly between counting atoms and weighing samples.
Question 4 True / False
The atomic mass listed on the periodic table for an element represents the mass of its most abundant naturally occurring isotope.
TTrue
FFalse
Answer: False
Atomic mass is a weighted average of the masses of all naturally occurring isotopes, weighted by their natural abundances. For carbon, the listed mass of 12.011 reflects ~98.9% carbon-12 (mass 12.000) and ~1.1% carbon-13 (mass 13.003). The value is not the mass of any single isotope — it is the average expected if you randomly sampled one atom from a natural mixture.
Question 5 Short Answer
Why does 12.011 grams of carbon contain 6.022 × 10²³ atoms rather than just one, even though carbon's atomic mass is 12.011 amu?
Think about your answer, then reveal below.
Model answer: Because amu and grams are vastly different units. One amu is about 1.66 × 10⁻²⁴ grams — a single carbon atom weighs only about 2 × 10⁻²³ grams. The mole was defined as exactly the number of atoms needed so that the mass in grams equals the atomic mass numerically. Molar mass (12.011 g/mol) means 12.011 grams PER MOLE — per 6.022 × 10²³ atoms — not per single atom.
This question targets the core confusion in the topic: the numerical equality between atomic mass (amu) and molar mass (g/mol) does not mean the units are interchangeable. The equality is a convenience of definition — Avogadro's number was chosen so the math works out neatly. Understanding this makes clear why molar mass serves as the essential conversion factor in stoichiometry: it connects the atomic-scale world (where we think in amu per atom) to the lab-scale world (where we measure in grams).