Relative masses
Relative masses
- Atoms are far too light to weigh in grams.
- So we compare every mass to one standard: carbon-12.
- This gives simple numbers for atoms, molecules and compounds.
The standard and $A_r$
- The unified atomic mass unit (u) is exactly $\tfrac{1}{12}$ of the mass of a carbon-12 atom.
- The relative atomic mass $A_r$ is the average mass of an element's atoms (over all isotopes, weighted by abundance), compared with $\tfrac{1}{12}$ of carbon-12.
- The relative isotopic mass is the mass of one isotope, same comparison.
Practice
The relative atomic mass (Ar) of an element is:
Ar is the average over all isotopes (weighted by abundance), measured against 1/12 of a carbon-12 atom.
$M_r$ and formula mass
- The relative molecular mass $M_r$ is the sum of the $A_r$ of all atoms in a molecule.
- The relative formula mass is the same idea for substances that aren't molecules (e.g. ionic compounds) — add the $A_r$ shown in the formula.
- Example: $M_r(\text{CO}_2) = 12 + (2 \times 16) = 44$.
Practice
The relative molecular mass (Mr) is found by:
Mr is the sum of the Ar values of every atom in the molecule.
Practice
What is the Mr of CO₂? (Ar: C = 12, O = 16)
Mr = 12 + (2 × 16) = 44.
Practice
For an ionic compound (not made of molecules), we use the:
Ionic compounds aren't molecules, so we add the Ar values in the formula to get the relative formula mass.
You've got it
Key idea
- masses are compared to $\tfrac{1}{12}$ of a carbon-12 atom (the unit u)
- $A_r$ = weighted average over isotopes; relative isotopic mass = one isotope
- $M_r$ = sum of $A_r$ of all atoms; formula mass = same for ionic compounds
- e.g. $M_r(\text{CO}_2) = 44$