The atomic mass (ma) is the mass of an atom, most often expressed in unified atomic mass units. The atomic mass may be considered to be the total mass of protons, neutrons and electrons in a single atom (when the atom is motionless). However, scientists do not include the electrons because their total mass is so insignificant that only the mass of the protons and neutrons are included in the atom's mass. The atomic mass is sometimes incorrectly used as a synonym of relative atomic mass, average atomic mass and atomic weight; however, these differ subtly from the atomic mass. The atomic mass is defined as the mass of an atom, which can only be one isotope at a time and is not an abundance-weighted average. In the case of many elements that have one dominant isotope the actual numerical difference between the atomic mass of the most common isotope and the relative atomic mass or standard atomic weights can be very small such that it does not affect most bulk calculations but such an error can be critical when considering individual atoms. For elements with more than one common isotope the difference even to the most common atomic mass can be half a mass unit or more (e.g. chlorine). The atomic mass of an uncommon isotope can differ from the relative atomic mass or standard atomic weight by several mass units.
The relative atomic mass (Ar) (also known as atomic weight and average atomic mass) is the average of the atomic masses of all the chemical element's isotopes as found in a particular environment, weighted by isotopic abundance. This is frequently used as a synonym for the standard atomic weight and it is not incorrect to do so since the standard atomic weights are relative atomic masses, although it is less specific to do so. Relative atomic mass also refers to non-terrestrial environments and highly specific terrestrial environments that deviate from the average or have different certainties (number of significant figures) than the standard atomic weights.
The standard atomic weight refers to the mean relative atomic mass of an element in the local environment of the Earth's crust and atmosphere as determined by the IUPAC Commission on Atomic Weights and Isotopic Abundances. These are what are included in a standard periodic table and is what is used in most bulk calculations. An uncertainty in brackets is included which often reflects natural variability in isotopic distribution rather than uncertainty in measurement. For synthetic elements the isotope formed depends on the means of synthesis, so the concept of natural isotope abundance has no meaning. Therefore, for synthetic elements the total nucleon count of the most stable isotope (ie, the isotope with the longest half-life) is listed in brackets in place of the standard atomic weight. Lithium represents a unique case where the natural abundances of the isotopes have been perturbed by human activities to the point of affecting the uncertainty in its standard atomic weight, even in samples obtained from natural sources such as rivers.
The relative isotopic mass is the relative mass of the isotope, scaled with carbon-12 as exactly 12. No other isotopes have whole number masses due to the different mass of neutrons and protons, as well as loss/gain of mass to binding energy. However, since mass defect due to binding energy is minimal compared to the mass of a nucleon, rounding the atomic mass of an isotope tells you the total nucleon count. Neutron count can then be derived by subtracting the atomic number.