### Relative Atomic Mass (Ar)

- The size of atoms is so tiny that we can’t really compare their masses in conventional units such as kilograms or grams, so a unit called the relative atomic mass (
*A*_{r}) is used - The
**relative****atomic****mass**unit is equal to 1/12th the mass of a carbon-12 atom - All other elements are measured by comparison to the mass of a carbon-12 atom and since these are ratios, the relative atomic mass has no units
- For example, hydrogen has a relative atomic mass of 1, meaning that 12 atoms of hydrogen would have exactly the same mass as 1 atom of carbon

### Calculating Ar

- The relative atomic mass of each element is calculated from the
**mass number**and**relative abundances**of all the isotopes of a particular element - The equation below is used where the top line of the equation can be extended to include the number of different isotopes of a particular element present.
- So if there were 3 isotopes present then the top line of the equation would read:(% of isotope A x mass of isotope A) + (% of isotope B x mass of isotope B) + (% of isotope C x mass of isotope C)

#### Worked example

The table shows information about the isotopes in a sample of rubidium with 72% ^{85}Rb and 28% ^{87}Rb

Use information from the table to calculate the relative atomic mass of this sample of rubidium. Give your answer to one decimal place:

** **( 72 x 85 ) + ( 28 x 87 ) ÷ 100 = 85.6

Relative atomic mass = 85.6

#### Exam Tip

Isotopes are easy to recognize from their notation as they have the same symbol but different mass numbers.

For example, the two stable isotopes of copper are ^{63}Cu and ^{65}Cu