Atomic mass unit

Atomic mass units, abbreviated with the symbol u, are used to represent the mass of atomic particles. Scientists use these units to quantify atomic mass because atoms, the building blocks of everything that exists, are so small that using common units of measurement, such as grams, would not be efficient. Knowing the atomic mass of an element—a pure, naturally occurring physical substance—is helpful in determining exactly how much of that element is present in a chemical reaction.

The History of Atomic Mass Units

Atomic mass units are a fundamental component of chemistry. They originated in the studies of late eighteenth-century and early nineteenth-century European chemists. By 1800, these scientists understood that elements bonded chemically to become compounds, substances consisting of two or more elements. However, they could not determine exactly how much of one element was necessary to combine with another element. They could not possibly know this until they found the masses of these elements.

Soon, scientists began designing experiments to quantify the masses of the various elements they combined into compounds. From these tests, they discovered that all the new materials they produced possessed consistent compositions throughout. Because of this consistency, the scientists were able to determine the mass of the entire compound. By working backward, they could also find the approximate masses of each individual element.

A short time later, English chemist John Dalton improved this system by proposing that the atoms of elements actually combined to form these compounds. To help scientists understand and work with this novel concept, Dalton devised a scale of relative atomic masses, figures representing the masses of singular elemental atoms as they relate to one another. As an uncomplicated way of beginning this scale, he assigned the gas hydrogen an atomic mass of 1. Dalton intended this to become the standard upon which all other atomic masses were based.

How Atomic Mass Units Work

Although Dalton’s efforts in the early 1800s were vital to the development of contemporary atomic mass units, scientists later altered his proposed atomic scale to make it more accurate. Chemists rely on the fundamental idea of relative atomic masses as developed by Dalton. However, the standard was changed from hydrogen to carbon-12, and one atomic mass unit is defined as one-twelfth of the mass of a free, neutral carbon-12 atom.

The importance of isotopes in calculating an element’s relative atomic mass arises later in the process. The procedure begins with a study of the composition of an atom. All atoms contain particles called protons, neutrons, and electrons. Protons are positively charged and located inside an atom’s nucleus, or center. Neutrons share the nucleus with protons, but they have no charge. Electrons have negative charges and occupy regions around the nucleus called orbitals. An atom has the same number of positive protons and negative electrons. This means that all atoms are neutral.

The next step in discovering an element’s atomic mass is finding that element’s atomic number. An atomic number is the number of protons and electrons contained within one atom of an element. Atomic numbers can be found on a periodic table, a chart that shows every element’s chemical symbol and provides additional information about the elements. For example, the element sodium is represented by the symbol Na. A periodic table shows that sodium’s atomic number is 11, meaning that it has 11 protons and 11 electrons.

The mass number of an isotope is the total number of protons and neutrons in its nucleus. Although mass number and atomic mass are different concepts, the two figures are related. Because of this, mass numbers are helpful in understanding an element’s atomic mass. An isotope’s mass number is found by adding its protons and neutrons. An element’s atomic number always provides the exact amount of protons in one atom of the element. To find that element’s neutrons, its protons are subtracted from the element’s mass number, as given on a periodic table. For instance, in the case of one isotope of sodium, the atomic number of 11 is subtracted from the mass number 23 to yield 12 neutrons. The 11 protons are then added to the 12 neutrons to show that the atomic mass of this particular isotope of sodium is 23 u.

The fact that sodium possesses several isotopes is an important distinction to consider in finding the relative atomic mass of the element. Isotopes of the same element have the same amount of protons and electrons but different numbers of neutrons. This means that each isotope of an element will have a different mass number and atomic mass. To find the relative atomic mass of all the combined isotopes of one element, the atomic mass of each individual isotope must be found. Then, these numbers are combined as a weighted average to produce the relative atomic mass of the element as a whole. Because the figure is an average, it is usually a decimal rather than a whole number. The figure is then written using the atomic mass unit u, also called the dalton (Da) in honor of John Dalton and his work in this area.

Bibliography

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