The mass number, commonly written as A, is the total count of the two types of particles that make up an atomic nucleus. In other words, it equals the number of protons plus the number of neutrons present in the nucleus of an atom. Because the proton count is the defining feature of a chemical species, different values of A for the same chemical element correspond to different isotopes.

Notation and simple relations

Mass number is written as a whole number attached to an element: for example, carbon with six protons and six neutrons is carbon-12 and often shown as "12C". The basic arithmetic relations are often written as A = Z + N and N = A − Z, where Z is the atomic number (proton count) and N is the neutron count. Because A counts discrete particles, it is always an integer; fractional values do not apply to an individual nucleus.

  • Atomic number (Z): the count of protons only; it determines the chemical identity (atomic number).
  • Atomic mass: a physical mass for a single atom, commonly expressed in unified atomic mass units; not identical to the integer mass number (atomic mass).
  • Relative atomic mass / atomic weight: an average value for an element on Earth reflecting natural isotopic abundances; usually a non-integer (relative atomic mass).
  • Periodic table entries: standard periodic tables list the periodic table element symbol, atomic number and standard atomic weight, but not a single mass number because many isotopes exist for each element.

Importance and applications

Knowing A is essential in nuclear science: the outcome of nuclear reactions, decay modes, and energy releases depend on proton and neutron counts. Mass number is used when writing nuclear equations, identifying isotopes for radiometric dating, medical imaging and therapy, and in nuclear fuel descriptions (for example, uranium-235 versus uranium-238). It also guides predictions about nuclear stability—certain combinations of Z and N produce stable nuclei while others are radioactive.

Measurement, history and practical examples

Mass numbers were clarified as scientists discovered isotopes and developed techniques such as mass spectrometry to measure atomic masses and compositions. Well-known examples illustrate the concept: hydrogen isotopes include protium (1H, A=1), deuterium (2H, A=2) and tritium (3H, A=3); the most common carbon isotope is carbon-12; uranium-235 (A=235) and uranium-238 (A=238) are distinct isotopes important in nuclear engineering. Although atomic mass measurements reveal small differences from integer A due to binding energy and electron mass, the mass number itself remains a simple, indispensable integer label for nuclides.

In summary, the mass number A is a concise way to specify how many nucleons (protons plus neutrons) a nucleus contains, and it is central to classifying isotopes, describing nuclear reactions, and distinguishing the microscopic identity of atomic nuclei.