Atomic Mass and Weight

Atomic mass and weight

The atomic mass of an atom (i.e., a specific isotope of an element) is measured in comparison with the mass of one atom of carbon-12 (¹²C) that is assigned a mass of 12 atomic mass units (amu). Atomic mass is sometimes erroneously confused with atomic weight—the obsolete term for relative atomic mass. Atomic weights, however, are still listed on many Periodic tables.

A mole of any element or compound (i.e., 6.022×10²³—Avogadro's number—atoms or molecules) weighs its total unit atomic mass (formerly termed atomic weight) in grams. For example, water (H₂O) has a molar mass (the mass of 6.022×10²³ water molecules) of approximately 18 grams (the sum of 2 hydrogen atoms, each with an atomic mass of 1.0079 amu, bonded with one oxygen atom with an atomic mass of 15.9994 amu).

In general usage if a specific isotope or isotope distribution is specified when using atomic mass, the natural percentage distribution of isotopes of that element is assumed. Periodic tables, for example usually list the atomic weights of individual elements based upon the natural distribution of isotopes of that element.

Mass is an intrinsic property of matter. Weight is a measurement of gravitational force exerted on matter.

In a series of papers published between 1803 and 1805 English physicist and chemist John Dalton (1766–1844) emphasized the importance of knowing the weights of atoms and outlined an experimental method for determining those weights.

The one problem with Dalton's suggestion was that chemists had to know the formulas of chemical compounds before they could determine the weights of atoms. But they had no way of knowing chemical formulas without a dependable table of atomic weights.

Dalton himself had assumed that compounds always had the simplest possible formula: HO for water (actually H₂O), NH for ammonia (actually NH₃), and so on. Although incorrect, this assumption allowed him to develop the concept of atomic weights, but, because his formulas were often wrong, his work inevitably resulted in incorrect values for most of the atomic weights. For example, he reported 5.5 for the atomic weight of monatomic oxygen and 4.2 for monatomic nitrogen. The correct values for those weights are closer to 16 and 14.

The first reasonably accurate table of atomic weights was produced by Swedish chemist Jöns Jacob Berzelius (1779–1848) in 1814. This table had been preceded by nearly a decade of work on the chemical composition of compounds. Once those compositions had been determined, Berzelius could use this information to calculate correct atomic weights.

In this process, Berzelius was faced with a decision that confronted anyone who tried to construct a table of atomic weights: What element should form the basis of that table and what would be the atomic weight of that standard element?

The actual weights of atoms are, of course, far too small to use in any table. The numbers that we refer to as atomic weights are all ratios. To say that the atomic weight of oxygen is 16, for example, is only to say that a single oxygen atom is 16 times as heavy as some other atom whose weight is somehow chosen as 1, or eight times as heavy as another atom whose weight has been chosen as 2, or one-half as heavy as another atom whose weight was selected to be 32, and so on.

Dalton had made the logical conclusion to use hydrogen as the standard for his first atomic table and had assigned a value of 1 for its atomic weight. Because hydrogen is the lightest element, this decision assures that all atomic weights will be greater than one.

The problem with Dalton's choice was that atomic weights are determined by measuring the way elements combine with each other, and hydrogen combines with relatively few elements. So, using Dalton's system, determining the atomic weight of another element might require a two-or three-step process.

Berzelius thought it made more sense to choose oxygen as the standard for an atomic weight table. Oxygen forms compounds with most other elements whose atomic weights can, therefore, be determined in a single step. He arbitrarily assigned a value of 100 as the atomic weight of oxygen. Other chemists agreed that oxygen should be the atomic weight standard, but used other values for its weight.

Berzelius continued working on atomic weights until, in 1828, he produced a table with values very close to those accepted today.

With the introduction of the concept of molecules (e.g., that the correct formula for water was H₂O) by Stanislao Cannizarro in 1858, it also became possible to calculate molecular weights. The molecular weight of any compound is equal to the sum of the weights of all the atoms in a molecule of that compound.

The most precise work on atomic weights during the nineteenth century was that of the Belgian chemist Jean Servais Stas (1813–1891). For over a decade, Stas recalculated Berzelius' weights, producing results that were unchallenged for nearly half a century.

An even higher level of precision was reached in the work of the American chemist Theodore William Richards (1868–1918). Richards spent more than 30 years improving methods for the calculation of atomic weights and redetermining those weights. Richards was awarded the Nobel Prize in chemistry in 1914 for these efforts.

The debate as to which element was to be used as the standard for atomic weights extended into the twentieth century, with the most popular positions being hydrogen with a weight of 1 or oxygen with a weight of 16. Between 1893 and 1903, various chemical societies finally agreed on the latter standard.

The controversy over standards was complicated by the fact that, over time, physicists and chemists began to use different standards for the atomic weight table and, thus, recognized slightly different values for the atomic weights of the elements. This dilemma was finally resolved in 1961 when chemists and physicists agreed to set the atomic weight of the carbon-12 isotope as 12.0000 as the standard for all atomic weights.

See also Atomic number; Atomic theory; Chemical bonds and physical properties; Chemical elements; Chemistry; Minerals