Isotopes and relative atomic mass

Although all atoms of a particular element have the same number of protons, they can have different numbers of neutrons.
Atoms of the same element with the same number of protons but different numbers of neutrons are called isotopes.
Isotopes of an element have the same atomic number (number of protons) but different mass numbers (total number of protons and neutrons).
Because they have the same number of electrons and the same electron arrangement, isotopes of an element have the same chemical properties.
Their physical properties (like density and melting point) can be slightly different due to the difference in mass.
Example: Chlorine has two common isotopes:
- Chlorine-35: 17 protons, 18 neutrons (Mass number = 35)
- Chlorine-37: 17 protons, 20 neutrons (Mass number = 37)
- Both have 17 protons, so they are both chlorine.

The atomic masses you see on the periodic table are rarely whole numbers (even though protons and neutrons have a mass of approximately 1). This is because most elements exist naturally as a mixture of their isotopes.
The relative atomic mass (Ar) of an element is an average value that takes into account the masses of its different isotopes and how abundant each isotope is on Earth.
The relative atomic mass is a ratio compared to 1/12th the mass of a carbon-12 atom, so it has no units.

You need to be able to calculate the relative atomic mass of an element given the percentage abundance of its isotopes.
The formula used is:

Where Σ means "the sum of" (you do this calculation for each isotope and add the results together).

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