# Isotopes

## What is an isotope?

Element are made up of one type of atom, but the atoms of a element can come in slightly different forms.

Although atoms of the same element always have the same number of protons, they may have different numbers of neutrons. Atoms that differ in this way are called isotopes.

For example, two isotopes of carbon:

12                           mass number is different                      13

C                                                                  C

6                          atomic number is the same                      6

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## What is an isotope?

Element are made up of one type of atom, but the atoms of a element can come in slightly different forms.

Although atoms of the same element always have the same number of protons, they may have different numbers of neutrons. Atoms that differ in this way are called isotopes.

For example, two isotopes of carbon:

12                           mass number is different                      13

C                                                                  C

6                          atomic number is the same                      6

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## Properties of isotopes

The isotopes of an element are virtually identical in their chemical reactions.

This is because they all have the same number of protons and the same number of electrons.

The uncharged neutrons make little difference to chemical properties, but do affect physical properties such as melting point and density.

Natural samples of elements are often a mixture of isotopes.

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## What is relative atomic mass (r.a.m)?

The relative atomic mass of an element is the mass of its atoms compared with the mass of an atom of carbon-12. Each element has a different r.a.m. value.

Relative atomic mass is measured in atomic mass units (amu). The r.a.m. value for carbon-12 is 12 amu. The mass of a hydrogen atom is 1/12 the mass of carbon-12. So, the r.a.m. value for hydrogen is 1 amu.

A magnesium atom is twice as heavy as an atom of carbon-12. How many carbon atoms have the same mass as one magnesium atom? 2. What is the r.a.m. value of magnesium? 24 amu.

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## Where are r.a.m. values found?

The values of relative atomic mass are usually given in a data book or found in the periodic table, so you don't have to work them out or remember them at all.

12                      relative atomic mass

C         symbol

6                        atomic number

When looking up r.a.m. in the periodic table, remember that it is always the larger of the two numbers given

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## Why isn't r.a.m. always a whole number?

Relative atomic mass is not always a whole number. For example, the r.a.m. of chlorine is 35.5.

The standard r.a.m. value of each element is actually the average relative atimic mass, which takes all the isotopes of each element into account.

Chlorine has two isotopes:

chlorine-35 (75%) and chlorine-37 (25%).

average r.a.m. of chlorine = (35 x 75%) + (37 x 25%)

= (35 x 0.75) + (37 x 0.25)

= 26.25 + 9.25

= 35.5

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## Calculating average r.a.m. from isotopes

To calculate the average r.a.m. of a mixture of isotopes, multiply the percentage of each isotope by its relative atomic mass and then add these together.

Naturally-occurring bromine is composed of two isotopes: bromine-79 (50.5%) and bromine-81 (49.5%)

What is the average r.a.m. of naturally occuring bromine?

average r.a.m. = (79 x 50.5%) + (81 x 49.5%)

= (79 x 0.505) + (81 x 0.495)

= 39.895 + 40.095

= 79.99

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## Glossary of Keywords

atomic number - The number of protons in the nucleus of an atom, also known as the proton number.

electron - A negatively charged particle that orbits the nucleus of an atom.

isotopes - Different atoms have the same element. They have the same number of protons and electrons, but a different number of neutrons.

mass number - The number of protons and neutrons in the nucleus of an atom.

neutron - A neutral particle with a mass of 1. It is found in the nucleus of an atom.

proton - A positively-charged particle with a mass of 1. It is found in the nucleus of an atom.

relative atomic mass - The average mass of an element compared with 1/12 of the mass of carbon-12

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