The atomic mass is the number of grams of the element in one mole of atoms of the element. This is a very useful property when it comes to practical calculations, as it allows easy conversion between the mass and moles of a given quantity of atoms or molecules of the same type.
Note that the relative atomic masses listed on the periodic table are average values for the associated element. Chemical elements have different isotopes - chemical forms that differ in mass because of the addition or subtraction of one or more neutrons to the atom’s nucleus. [4] X Research source Thus, the relative atomic mass listed on the periodic table is suitable as an average value for atoms of a certain element, but not as the mass of a single atom of that element. Relative atomic masses, as listed on the periodic table, are used to calculate molar masses for atoms and molecules. Atomic masses, when expressed in amu, as on the periodic table, are technically unitless. However, by simply multiplying an atomic mass by 1 g/mol, a workable quantity is obtained for an element’s molar mass - the mass (in grams) of one mole of an element’s atoms. For example, the atomic mass of iron is 55. 847 amu, which means one mole of iron atoms would weigh 55. 847 grams.
Because it’s an average of several different types of isotopes, the value on the periodic table isn’t the exact value for any single atom’s atomic mass. The atomic masses for individual atoms must be calculated by taking into account the exact number of protons and neutrons in a single atom.
Let’s say that we’re working with the carbon atom. Carbon always has 6 protons, so we know its atomic number is 6. We can also see on the periodic table that the square for carbon (C) has a “6” at the top, signifying that carbon’s atomic number is 6. Note that an element’s atomic number doesn’t have any direct bearing on its relative atomic mass as listed on the periodic table. Though, especially among elements at the top of the periodic table, it may seem that an atoms’ atomic mass is about twice its atomic number, atomic mass isn’t ever calculated by doubling an element’s atomic number.
The number of neutrons can be determined by the isotope designation of the element. For example, carbon-14 is a naturally occurring radioactive isotope of carbon-12. You will often see an isotope designated with the number as a superscript before the element symbol: 14C. The number of neutrons is calculated by subtracting the number of protons from the isotope number: 14 – 6 = 8 neutrons. Let’s say the carbon atom we’re working with has six neutrons (12C). This is by far the most common isotope of carbon, accounting for nearly 99% of all carbon atoms. [6] X Research source However, about 1% of carbon atoms have 7 neutrons (13C). Other types of carbon atoms with more or less than 6 or 7 neutrons exist in very small amounts.
Our carbon atom has 6 protons + 6 neutrons = 12. The atomic mass of this specific carbon atom is 12. If it was a carbon-13 isotope, on the other hand, we would know that it has 6 protons + 7 neutrons = an atomic weight of 13. The actual atomic weight of carbon-13 is 13. 003355[8] X Research source , and is more precise because it was determined experimentally. Atomic mass is very close to the isotope number of an element. For basic calculation purposes, isotope number is equal to atomic mass. When determined experimentally, the atomic mass is slightly higher than the isotope number due to the very small mass contribution from electrons.
For our purposes, let’s say we’re working with the isotopes carbon-12 and carbon-13.
Let’s say that the abundance of carbon-12 is 99% and the abundance of carbon-13 is 1%. Other carbon isotopes do exist, but they exist in quantities so small that, for this example problem, they can be ignored.
Our sample contains carbon-12 and carbon-13. If carbon-12 makes up 99% of the sample and carbon-13 makes up 1% of the sample, multiply 12 (the atomic mass of carbon-12) by 0. 99 and 13 (the atomic mass of carbon-13) by 0. 01. A reference book will give percent proportions based on all the known amounts of an element’s isotopes. Most chemistry textbooks include this information in a table at the end of the book. A mass spectrometer can also yield the proportions for the sample being tested.
In our example, 12 x 0. 99 = 11. 88 for carbon-12, while 13 x 0. 01 = 0. 13 for carbon-13. The relative atomic mass of our example is 11. 88 + 0. 13 = 12. 01.
