element: The Atom

Atomic Number and Mass Number

Regardless of how many atoms the element is composed of, each atom has the same number of protons in its nucleus, and this is different from the number in the nucleus of any other element. Thus this number, called the atomic number (at. no.), defines the element. For example, the element carbon consists of atoms all with at. no. 6, i.e., all having 6 protons in the nucleus; any atom with at. no. 6 is a carbon atom. By 2016, 118 elements were known, ranging from hydrogen with an at. no. of 1 to oganesson with an at. no. of 118. (See the table entitled Elements for an alphabetical list of all the elements, including their symbols, atomic numbers, atomic weights, and melting and boiling points.) The nuclei of most atoms also contain neutrons. The total number of protons and neutrons in the nucleus of an atom is called the mass number. For example, the mass number of a carbon atom with 6 protons and 6 neutrons in its nucleus is 12.

Isotopes

Although all atoms of an element have the same number of protons in their nuclei, they may not all have the same number of neutrons. Atoms of an element with the same mass number make up an isotope of the element. All known elements have isotopes; some have more than others. Hydrogen, for example, has only 3 isotopes, while xenon has 16. Approximately 300 naturally occurring isotopes are known, and more than 2,500 radioactive isotopes have been artificially produced (see synthetic elements ). There are 13 isotopes of carbon, having from 2 to 14 neutrons in the nucleus and therefore mass numbers from 8 to 20.

Not all of the elements have stable isotopes. Some have only radioactive isotopes, which decay to form other isotopes, usually of other elements (see radioactivity ). In some cases all the isotopes of an element are very unstable, and the element is therefore not found in nature. Only 94 of the elements are known to occur naturally on earth. Of these, 6 occur in minute amounts produced by the decay of other elements. These 6 extremely scarce elements and those that do not occur at all naturally were discovered when they were produced in the laboratory; they are often called the artificially produced or synthetic elements .

Atomic Mass and Atomic Weight

Atoms are not very massive; a carbon atom weighs about 2 × 10 ⁻²³ grams. Because atoms have so little mass, a unit much smaller than the gram is used. In the current system (adopted in 1960–61) the unit of atomic mass, called atomic mass unit (amu), is defined as exactly   1⁄12 the mass of an atom of carbon-12. The atomic weight of an element is the mean (weighted average) of the atomic masses of all the naturally occurring isotopes. Carbon has two principal naturally occurring isotopes, carbon-12 and carbon-13. Carbon-12, whose mass is defined as exactly 12 amu, constitutes 98.89% of naturally occurring carbon; carbon-13, whose mass is 13.00335 amu, constitutes 1.11%. (There are also small traces of the radioactive isotope carbon-14.) The atomic weight of the element is determined by multiplying the percent abundance of each isotope by the atomic mass of the isotope, adding these products, and dividing by 100. However, isotope abundance is often determined by the medium of the source, solid, liquid, or gas, and the average atomic weight may fluctuate. Thus, for carbon, [(98.89 × 12.000) + (1.11 × 13.00335)]/100 = 12.01115, which is the atomic weight of the element carbon in amu, but because the proportions of the isotopes vary depending on where the carbon is found, carbon's atomic weight is now expressed as an interval defined by the lower and upper bounds within which the atomic weight ranges: [12.0096; 12.0116]. Certain synthetic elements exist only momentarily in the form of a few short-lived isotopes; in such cases the concept of atomic weight cannot be applied.