An atomic mass unit (abbreviated amu) is a unit of mass equal to 1/12 the mass of the most abundant isotope of carbon, carbon 12, which is assigned a mass of 12.
The unified atomic mass unit (u), or dalton (Da), is a small unit of mass used to express atomic and molecular masses. It is defined to be 1/12 of the mass of one atom of carbon-12.
1 u = 1/NA gram = 1/(1000 NA) kg (where NA is Avogadro's number) 1 u ≈ 1.66053886 × 10−27 kg ≈ 931.49 MeV/c2 See 1 E-27 kg for a list of objects which have a mass of about 1 u.
The symbol amu for atomic mass unit can sometimes still be found, particularly in older works. Atomic masses are often written without any unit and then the atomic mass unit is implied. In biochemistry and molecular biology literature (particularly in reference to proteins), the term "dalton" is used, with the symbol Da. Because proteins are large molecules, they are typically referred to in kilodaltons, or "kDa", with one kilodalton being equal to 1000 daltons. The unified atomic mass unit is not an SI unit of mass, although it is (only by that name, and only with the symbol u) accepted for use with SI.
The unit is convenient because one hydrogen atom has a mass of approximately 1 u, and more generally an atom or molecule that contains n protons and neutrons will have a mass approximately equal to n u. (The reason is that a carbon-12 atom contains 6 protons, 6 neutrons and 6 electrons, with the protons and neutrons having about the same mass and the electron mass being negligible in comparison.) This is an approximation, since it does not account for the mass contained in the binding energy of an atom's nucleus; this binding energy mass is not a fixed fraction of an atom's total mass. The differences which result from nuclear binding are generally less than 0.01 u, however. Chemical element masses, as expressed in u, would therefore all be close to whole number values (within 2% and usually within 1%) were it not for the fact that atomic weights of chemical elements are averaged values of the various stable isotope masses in the abundances which they naturally occur. [1] [For example, chlorine has a atomic weight of 35.45 u because it is composed of 76% Cl-35 (34.96 u) and 24% Cl-37 (36.97 u)].
Another reason the unit is used is that it is experimentally much easier and more precise to compare masses of atoms and molecules (determine relative masses) than to measure their absolute masses. Masses are compared with a mass spectrometer (see below).
Avogadro's number (NA) and the mole are defined so that one mole of a substance with atomic or molecular mass 1 u will have a mass of precisely 1 gram. For example, the molecular mass of water is 18.01508 u, and this means that one mole of water has a mass of 18.01508 grams, or conversely that 1 gram of water contains NA/18.01508 ≈ 3.3428 × 1022 molecules.