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Atomic mass unit
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Atomic mass unit

The atomic mass unit (amu), unified atomic mass unit (u), or Dalton named after the chemist John Dalton, is a small unit of mass used to express atomic masses and molecular masses. It is defined to be 1/12 of the mass of one atom of Carbon-12. Accordingly,

1 u = 1/NA gram = 1/(1000 NA) kg

where NA is Avogadro's number.

The symbol "amu" can sometimes 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".

It is not an SI unit of mass, although it is accepted for use with SI. See SI website link below.

1 u ≈ 1.6605402 x 10-27 kg

See 1 E-27 kg for a list of objects which have a mass of about 1 u.

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. This is only a rough approximation however, since it doesn't account for the mass contained in the binding energy of the nucleus (in fact, does not account for variations in that mass relative to the total mass, compared with this ratio for C-12).

Another reason the unit is used is that it is much easier to compare masses of atoms and molecules (determine relative masses) than to measure their absolute masses, because masses in kilograms are inconveniently small numbers.

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. As an equation:

1 u = 1 gram/mole
or equivalently
1 gram = NA u

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.

Table of contents
1 Measuring Relative Atomic Masses
2 External link

Measuring Relative Atomic Masses

The relative atomic mass is measured with a mass spectrometer. After placing a sample of the element to be measured in the mass spectrometer it is bombarded with electrons which turns the atoms into positive ions. An electric field is then used to accelerate these positive ions, afterwhich the ions are deflected using a magnetic field. As a result the various isotopes are separated out due to the ions of lighter isotopes being deflected more than those heavier. This produces a mass spectrum.

This spectrum provides two things:

  1. Relative isotopic masses in the sample
  2. Abundances of the isotopes

Using Mess Spectrum Data to Calculate Relative Atomic Mass

A simple calculation may be used to calculate the relative atomic mass of the sample. This is demonstrated in the following example.

Ion Relative Mass Percentage Abundance
11C+ 11 70%
13C+ 13 30%

Therefore, the relative atomic mass of the Carbon sample is:

(70/100 x 11) + (30/100 x 13)

7.7 + 3.9

= 11.6 [this is not the true atomic mass of carbon, it is merely illustrative]

External link