 As you look through the periodic table, you may have noticed that not all atomic masses are written as whole numbers. Elements can exist with slightly different numbers of neutrons. We call these isotopes of an element. It is particularly common for heavier elements, but a familiar, lighter example is chlorine. The two most common types of chlorine atom in the world around us are chlorine 35 and chlorine 37. If we took a sample of chlorine gas, we would find that it was made up of approximately 75% chlorine 35 and 25% of chlorine 37. There are other isotopes of chlorine, but these two are the most common. The relative abundance of the different isotopes, in other words how common they are, is always taken into consideration when calculating the relative atomic mass of an element. One way of visualizing this is to imagine that you have a sample of 100 naturally occurring chlorine atoms. Based on the abundance in the example already given, 75 of these would be chlorine 35 and 25 would be chlorine 37. Knowing this, we can work out the relative atomic mass. In other words, the weighted mean atomic mass. And we can do this using a formula as follows. 75 times 35 plus 25 times 37 all divided by 100 equals 35.5. And this explains why chlorine is listed as having a relative atomic mass of 35.5 in the periodic table. The relative atomic mass of an element is the weighted mean mass of an element and it is written in atomic mass units. Here's another example for you. A sample of bromine contains 50% bromine 79 and 50% bromine 81. What is its relative atomic mass? Pause the video. You may be able to work this out straight away. If not, write it down as in the chlorine example. The answer is 80. Detailed periodic tables of elements show that most relative atomic masses aren't actually whole numbers. This is all due to the naturally occurring isotopes of each element. So to recap, the relative atomic mass of an element is the weighted mean mass of the isotopes of an element. And by weighted, we mean that we are taking the relative abundance of each isotope into consideration. And this is why relative atomic masses aren't always whole numbers.