# 1.1.9 Relative Atomic Mass

### Relative Atomic Mass (Ar)

• The size of atoms is so tiny that we cannot really compare their masses in conventional units such as kilograms or grams, so a unit called the relative atomic mass (Ar) 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)

RAM Calculation Formula #### Worked Example

The table shows information about the Isotopes in a sample of rubidium with 72% 85Rb and 28% 87Rb Use information from the table to calculate the relative atomic mass of this sample of Rubidium. Give your answer to one decimal place: 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 63Cu and 65Cu ### Author: Francesca

Fran has taught A level Chemistry in the UK for over 10 years. As head of science, she used her passion for education to drive improvement for staff and students, supporting them to achieve their full potential. Fran has also co-written science textbooks and worked as an examiner for UK exam boards.
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