The atomic weight of an element is the weighted average of the masses of the isotopes of that element. The weighted average is determined using the abundance and mass of each isotope. Most elements have more than one naturally occurring isotopes.

For example, there are two naturally occurring isotopes of copper, copper-63 and copper-65. The natural abundances of the isotopes is 69.2% and 30.8% respectively.

To determine the atomic weight:
Step 1: Multiply the mass number and the relative abundance (as a decimal). The mass of the electron is insignificant in this calculation and is not used.

isotope	mass number (amu)	x	abundance (as a decimal)	=	result
copper-63	63 amu	x	0.692	=	43.596
copper-65	65 amu	x	0.308	=	20.020
Step 2: Add up your results.			Atomic Weight		63.616

On your periodic table you can see that the atomic weight is 63.546. The reason for the difference between the actual atomic weight (as seen on the perioidc table) and the calculated atomic weight is due to the fact that the masses of the proton and neutron is not exactly 1 amu. The mass of a proton is approx 1.008 amu. The mass of a neutron is approximately 1.009 amu.