Key Concepts: Isotopes and Relative Atomic Mass

Isotopes and Relative Atomic Mass:

Isotopes:

Isotopes are atoms of the same element that have the same number of protons, but different numbers of neutrons. This means they have the same atomic number but different mass numbers.
For example, carbon-12 and carbon-14 are both isotopes of carbon. They both have 6 protons, but carbon-12 has 6 neutrons and carbon-14 has 8 neutrons.
Although isotopes of an element have different mass numbers, they have very similar chemical properties. This is because chemical reactions involve the electrons in an atom, not the protons or neutrons.
Some isotopes are unstable and emit radiation, making them radioactive. These isotopes can be used in medical imaging, in industry to check for cracks in metal, and in carbon dating to tell how old something is.

Relative Atomic Mass:

The relative atomic mass (Ar) of an element is an average of the mass numbers of the isotopes of the element, taking into account how much of each isotope there is.
It is nearly always a decimal number. For example, the Ar of chlorine is 35.5 because 75% of chlorine atoms are chlorine-35 and 25% are chlorine-37.
The relative atomic mass is measured on a scale where the atoms of the isotope carbon-12 are given a mass of exactly 12.
Whenever you might need to make calculations involving mass, it is essential to use relative atomic mass.
Understanding how to determine relative atomic mass from abundance of relative isotopes is crucial. You can find the relative atomic mass by using the formula: Ar = Σ (isotope abundance x isotope mass number) ÷ total abundance.

It is very important to remember these key concepts and to understand how isotopes and relative atomic mass are used in chemistry. Doing regular revision and practise questions can greatly aid in understanding and retention of these fundamentals.