On March 20, 1995, the Japanese terrorist group Aum Shinrikyo (Sanskrit for “Supreme Truth”) released some sarin gas in the Tokyo subway system twelve people were killed, and thousands were injured (part (a) in the accompanying figure). We would not be far off if we limited our numbers to one or even two decimal places.Ĭhemistry Is Everywhere: Sulfur Hexafluoride Also, although each atom in a molecule is a particular isotope, we use the weighted average, or atomic mass, for each atom in the molecule.įor example, if we were to determine the molecular mass of dinitrogen trioxide, N 2O 3, we would need to add the atomic mass of nitrogen two times with the atomic mass of oxygen three times: This may seem like a trivial extension of the concept, but it is important to count the number of each type of atom in the molecular formula. The molecular mass is the sum of the masses of the atoms in a molecule. Now that we understand that atoms have mass, it is easy to extend the concept to the mass of molecules. (The record number is 10 isotopes for tin.) Table 3.5 Selected Atomic Masses of Some Elements Element Name This is, in part, the effect of an increasing number of isotopes as the atoms increase in size. Note that many of the atomic masses, especially the larger ones, are not very close to whole numbers. The atomic masses in Table 3.5 “Selected Atomic Masses of Some Elements” are listed to three decimal places where possible, but in most cases, only one or two decimal places are needed. Table 3.5 “Selected Atomic Masses of Some Elements” lists the atomic masses of some elements a more expansive table is in “Appendix A: Periodic Table of the Elements”. Virtually all elements exist as mixtures of isotopes, so atomic masses may vary significantly from whole numbers. Thus, we use 10.8 u for the atomic mass of boron. The atomic mass of boron is calculated similarly to what we did for our hypothetical example, but the percentages are different: Boron exists as about 20% boron-10 (five protons and five neutrons in the nuclei) and about 80% boron-11 (five protons and six neutrons in the nuclei). This example is similar to a real element. Note that no atom in our hypothetical element has a mass of 10.5 u rather, that is the average mass of the atoms, weighted by their percent occurrence.

The sum is the weighted average and serves as the formal atomic mass of the element. A weighted average is found by multiplying each mass by its fractional occurrence (in decimal form) and then adding all the products. What do we mean by a weighted average? Well, consider an element that consists of two isotopes, 50% with mass 10 u and 50% with mass 11 u. The atomic mass of an element is a weighted average of the masses of the isotopes that compose an element. Thus, although it is easy to speak of the mass of an atom, when talking about the mass of an element, we must take the isotopic mixture into account. What is the mass of an element? This is somewhat more complicated because most elements exist as a mixture of isotopes, each of which has its own mass. More exact masses are found in scientific references-for example, the exact mass of uranium-238 is 238.050788 u, so you can see that we are not far off by using the whole-number value as the mass of the atom. Thus, the mass of carbon-12 is about 12 u, the mass of oxygen-16 is about 16 u, and the mass of uranium-238 is about 238 u. There will not be much error if you estimate the mass of an atom by simply counting the total number of protons and neutrons in the nucleus (i.e., identify its mass number) and ignore the electrons. By this scale, the mass of a proton is 1.00728 u, the mass of a neutron is 1.00866 u, and the mass of an electron is 0.000549 u.

The atomic mass unit (u some texts use amu, but this older style is no longer accepted) is defined as one-twelfth of the mass of a carbon-12 atom, an isotope of carbon that has six protons and six neutrons in its nucleus. For macroscopic objects, we use units such as grams and kilograms to state their masses, but these units are much too big to comfortably describe the masses of individual atoms and molecules. Individual atoms and molecules, however, are very small, and the masses of individual atoms and molecules are also very small.

Express the masses of atoms and molecules.īecause matter is defined as anything that has mass and takes up space, it should not be surprising to learn that atoms and molecules have mass.

Atom vs molecule