Atomic Number, Mass Number and Isotopes

Key Points

Neutral atoms of each element contain an equal number of protons and electrons.
The number of protons determines an element’s atomic number and is used to distinguish one element from another.
The number of neutrons is variable, resulting in isotopes, which are different forms of the same atom that vary only in the number of neutrons they possess.
Together, the number of protons and the number of neutrons determine an element’s mass number.
Since an element’s isotopes have slightly different mass numbers, the atomic mass is calculated by obtaining the mean of the mass numbers for its isotopes.

Terms

atomic mass

The average mass of an atom, taking into account all its naturally occurring isotopes.

mass number

The sum of the number of protons and the number of neutrons in an atom.

atomic number

The number of protons in an atom.

Atomic Number

Neutral atoms of an element contain an equal number of protons and electrons. The number of protons determines an element’s atomic number (Z) and distinguishes one element from another. For example, carbon’s atomic number (Z) is 6 because it has 6 protons. The number of neutrons can vary to produce isotopes, which are atoms of the same element that have different numbers of neutrons. The number of electrons can also be different in atoms of the same element, thus producing ions (charged atoms). For instance, iron, Fe, can exist in its neutral state, or in the +2 and +3 ionic states.

Mass Number

An element’s mass number (A) is the sum of the number of protons and the number of neutrons. The small contribution of mass from electrons is disregarded in calculating the mass number. This approximation of mass can be used to easily calculate how many neutrons an element has by simply subtracting the number of protons from the mass number. Protons and neutrons both weigh about one atomic mass unit or amu. Isotopes of the same element will have the same atomic number but different mass numbers.

Atomic number, chemical symbol, and mass numberCarbon has an atomic number of six, and two stable isotopes with mass numbers of twelve and thirteen, respectively. Its average atomic mass is 12.11.

Scientists determine the atomic mass by calculating the mean of the mass numbers for its naturally-occurring isotopes. Often, the resulting number contains a decimal. For example, the atomic mass of chlorine (Cl) is 35.45 amu because chlorine is composed of several isotopes, some (the majority) with an atomic mass of 35 amu (17 protons and 18 neutrons) and some with an atomic mass of 37 amu (17 protons and 20 neutrons).

Given an atomic number (Z) and mass number (A), you can find the number of protons, neutrons, and electrons in a neutral atom. For example, a lithium atom (Z=3, A=7 amu) contains three protons (found from Z), three electrons (as the number of protons is equal to the number of electrons in an atom), and four neutrons (7 – 3 = 4).

Isotopes

Isotope, one of two or more species of atoms of a chemical element with the same atomic number and position in the periodic table and nearly identical chemical behaviour but with different atomic masses and physical properties. Every chemical element has one or more isotopes.

An atom is first identified and labeled according to the number of protons in its nucleus. This atomic number is ordinarily given the symbol Z. The great importance of the atomic number derives from the observation that all atoms with the same atomic number have nearly, if not precisely, identical chemical properties. A large collection of atoms with the same atomic number constitutes a sample of an element. A bar of pure uranium, for instance, would consist entirely of atoms with atomic number 92. The periodic table of the elements assigns one place to every atomic number, and each of these places is labeled with the common name of the element, as, for example, calcium, radon, or uranium.

Not all the atoms of an element need have the same number of neutrons in their nuclei. In fact, it is precisely the variation in the number of neutrons in the nuclei of atoms that gives rise to isotopes. Hydrogen is a case in point. It has the atomic number 1. Three nuclei with one proton are known that contain 0, 1, and 2 neutrons, respectively. The three share the place in the periodic table assigned to atomic number 1 and hence are called isotopes (from the Greek isos, meaning “same,” and topos, signifying “place”) of hydrogen.

Many important properties of an isotope depend on its mass. The total number of neutrons and protons (symbol A), or mass number, of the nucleus gives approximately the mass measured on the so-called atomic-mass-unit (amu) scale. The numerical difference between the actual measured mass of an isotope and A is called either the mass excess or the mass defect.

The Discovery Of Isotopes

Evidence for the existence of isotopes emerged from two independent lines of research, the first being the study of radioactivity. By 1910 it had become clear that certain processes associated with radioactivity, discovered some years before by French physicist Henri Becquerel, could transform one element into another. In particular, ores of the radioactive elements uranium and thorium had been found to contain small quantities of several radioactive substances never before observed. These substances were thought to be elements and accordingly received special names. Uranium ores, for example, yielded ionium, and thorium ores gave mesothorium. Painstaking work completed soon afterward revealed, however, that ionium, once mixed with ordinary thorium, could no longer be retrieved by chemical means alone. Similarly, mesothorium was shown to be chemically indistinguishable from radium. As chemists used the criterion of chemical indistinguishability as part of the definition of an element, they were forced to conclude that ionium and mesothorium were not new elements after all, but rather new forms of old ones. Generalizing from these and other data, English chemist Frederick Soddy in 1910 observed that “elements of different atomic weights [now called atomic masses] may possess identical (chemical) properties” and so belong in the same place in the periodic table. With considerable prescience, he extended the scope of his conclusion to include not only radioactive species but stable elements as well. A few years later, Soddy published a comparison of the atomic masses of the stable element lead as measured in ores rich in uranium and thorium, respectively. He expected a difference because uranium and thorium decay into different isotopes of lead. The lead from the uranium-rich ore had an average atomic mass of 206.08 compared to 207.69 for the lead from the thorium-rich ore, thus verifying Soddy’s conclusion.

The unambiguous confirmation of isotopes in stable elements not associated directly with either uranium or thorium followed a few years later with the development of the mass spectrograph (see mass spectrometry) by Francis William Aston. His work grew out of the study of positive rays (sometimes called canal rays), discovered in 1886 by Eugen Goldstein and soon thereafter recognized as beams of positive ions. As a student in the laboratory of J.J. Thomson, Aston had learned that the gaseous element neon produced two positive rays. The ions in the heavier ray had masses about two units, or 10 percent, greater than the ions in the lighter ray. To prove that the lighter neon had a mass very close to 20 and that the heavier ray was indeed neon and not a spurious signal of some kind, Aston had to construct an instrument that was considerably more precise than any other of the time. By 1919 he had done so and convincingly argued for the existence of neon-20 and neon-22. Information from his and other laboratories accumulated rapidly in the ensuing years, and by 1935 the principal isotopes and their relative proportions were known for all but a handful of elements.

Researcher:

Shenna Mae Molina