Atomic mass unit

The unified atomic mass unit (u), or dalton (Da), 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/N_A gram = 1/(1000 N_A) kg   (where N_A is Avogadro's number)

1 u ≈ 1.66053886 x 10^-27 kg

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. See SI website link below.

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 only a rough approximation however, 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.

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 (N_A) 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 N_A/18.01508 ≈ 3.3428 × 10²² molecules.

Contents

1 Measuring relative atomic masses 2 History 3 See also 4 External links

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

History

The chemist John Dalton was the first to suggest the mass of one atom of hydrogen as the atomic mass unit. Francis Aston, inventor of the mass spectrometer, later used 1/16 of the mass of one atom of oxygen-16 as his unit.

Before 1961, the physical atomic mass unit was defined as 1/16 of the mass of one atom of oxygen-16, while the chemical atomic mass unit was defined as 1/16 of the average mass of an oxygen atom (taking the natural abundance of the different oxygen isotopes into account). Both units are slightly smaller than the unified atomic mass unit, which was adopted by the International Union of Pure and Applied Physics in 1960 and by the International Union of Pure and Applied Chemistry in 1961.

See also

• Dalton

External links

• SI website on acceptable non-SI units (http://www1.bipm.org/en/si/si_brochure/chapter4/table7.html)
• Accepted value of 1 u as of 2002 (http://physics.nist.gov/cgi-bin/cuu/Value?ukg)ca:Unitat de massa atòmica

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