Atomic mass units, definition

How is the atomic mass unit defined How has this definition evolved ... 

Note Care has to be taken when mass values from dated literature are cited. Prior to 1961 physicists defined the atomic mass unit [amu] based on Vie of the mass of one atom of nuclide 0. The definition of chemists was based on the relative atomic mass of oxygen which is somewhat higher resulting from the nuclides and contained in natural oxygen. 

The atomic mass unit (mu) is also called the dalton (Da) - in honour of John Dalton. In response to the increase in the use of the name dalton for the unified atomic mass unit among chemists, it was suggested by IUPAC that the unified atomic mass unit (u) be renamed the dalton (Da). The definition of the unit would remain unchanged as one-twelfth the mass of a neutral 12C atom in its ground state. The International Union of Pure and Applied Chemistry (IUPAC) proposed that both units, u and Da, should be allowed in official use. 

Resolving Power (RP) A measurement of how effectively a mass analyzer can distinguish between two peaks at different, but similar m/z. Mathematically, the formula M/ AM is used, where M is the m/z value for one of the peaks and AM is the spacing, in unified atomic mass units, between the peaks. Most commonly, AM is the mass resolution, either via the 10% valley or FWHM definitions (see below). (Note that the definition used will affect the resolving power calculated.) Resolving power of 500-1000 approximately corresponds to unit resolution (e.g., at m/z 700 and FWHM resolution of 0.7, RP = 1000). 

In 1808 John Dalton proposed his atomic theory, which included the statement that when atoms of two or more elements combine to form a compound, they combine in a definite ratio by number of atoms and by mass. This is called the law of definite proportions. This provided a means to determine the mass of one atom relative to another. It was necessary to assign a mass to one element to find the mass of another element in a compound. Today we use the most common carbon isotope, assigned a mass of 12.00 atomic mass units (amu), as the basis for comparative weights of the atoms. 

In the future, it will end in the value of the Avogadro constant and the value 1 of the atomic mass unit u defined as the mass of 1/12th of the mass of the 12C atom. This will happen when the definition of the kg (now the mass of the prototype of the kilogram ) will have changed into the mass of a number of12C atoms, i.e. of the mass of NAj-m(12C). 1000/12) . 

So the mass of the 12C atom is exactly 12 atomic mass units, by virtue of the definition. 

The mass m is usually measured in Daltons (Da) or atomic mass units (amu mass of one gC12 nucleus = 12.0 Da = 12.0 amus, by definition 1 Da = 1.660538782 x 10-27 kg). Mass spectrometers can resolve isotopes rather well, and they can measure them to high precision (sometimes to 1 part in 107, but not in most routine or commercial instruments). Their mass range can be huge, from 1 amu to 100 kDa (the wider ranges have lower resolution, but always below 1 Da). The sample sizes are of the order of micrograms to nanograms. 

In Chapter 2, you learned that the mass of an atom is expressed in atomic mass units. Atomic mass units are a relative measure, defined by the mass of carbon-12. According to this definition, one atom of carbon-12 is assigned a mass of 12 u. Stated another way, 1 u = of the mass of one atom of carbon-12. 

You would never express the mass of a lump of gold, like the one in Figure 5.11, in atomic mass units. You would express its mass in grams. How does the mole relate the number of atoms to measurable quantities of a substance The definition of the mole pertains to relative atomic mass, as you learned in section 5.1. One atom of carbon-12 has a mass of exactly 12 u. Also, by definition, one mole of carbon-12 atoms (6.02 x 1023 carbon-12 atoms) has a mass of exactly 12 g. 

The atomic mass units u or Da have the same fundamental definition ... 

Use the value of Avogadro s number, the mass of one C atom (exactly 12 amu), and the definition of a mole to calculate the number of atomic mass units per gram. 

Thus the mole is defined such that a sample of a natural element with a mass equal to the element s atomic mass expressed in grams contains 1 mole of atoms. This definition also fixes the relationship between the atomic mass unit and the gram. Since 6.022 X 1023 atoms of carbon (each with a mass of 12 amu) have a mass of 12 grams, then... 

The masses for the elements listed in the table inside the back cover of this text are relative masses in terms of atomic mass units (amu) or daltons. The atomic mass unit is based on a relative scale in which the reference is the C carbon isotope, which is assigned a mass of exactly 12 amu. Thus, the amu is by definition 1/12 of the mass of one neutral c atom. The molar mass of is then... 

To generate a relationship between mass of carbon and number of carbon atoms, we need to know the weighted average mass of the carbon atoms found in nature. Experiments show that 98.90% of the carbon atoms in natural carbon are carbon-12, and 1.10% are carbon-13, with six protons, seven neutrons, and six electrons. Related experiments show that each carbon-13 atom has a mass of 13.003355 u. From the definition of atomic mass unit, we know that the mass of each carbon-12 atom is 12 u. The following setup shows how the weighted average mass of carbon atoms is calculated. 

Biomolecules are molecules that are specifically found inside of living things and have some function related to life. This includes molecules found in plants, animals, insects, bacteria, or even viruses (which are often considered not to be alive in a technical sense). Biomolecules span a wide range of sizes, some weighing only about 50 atomic mass units (amu), while others weigh millions of amu (see Atoms and Molecules to review the definition of amu). They are in our hair, skin, tissues, organs, and just about everywhere else in our bodies too. 

The mass of an atom depends on the number of electrons, protons, and neutrons it contains and all atoms of a given isotope are identical in mass. The SI unit of mass (the kilogram) is too large to function as a convenient unit for the mass of an atom, thus a smaller unit is desirable. In 1961, the International Union of Pure and Applied Chemistry (lUPAC) defined the atomic mass unit (u) to be exactly equal to one-twelfth the mass of one carbon-12 atom. Carbon-12 ( C) is the carbon isotope that has six protons, six neutrons, and six electrons. Using this definition, we have that 1 u = 1.660539 X 10 kg. The atomic mass (sometimes called atomic weight) of an atom is then defined, relative to this standard, as the mass of the atom in atomic mass units (u). For example, the two naturally occurring isotopes of hefium, He and " He, have atomic masses of 3.01602931 u and 4.00260324 u, respectively. This means that a helium-4 (" He) atom is 4.00260324/12 = 0.33355027 times as massive as a carbon-12 atom. ... 

Before 1961, two definitions of the atomic mass unit were used. In physics, the atomic mass unit was defined as one-sixteenth the mass of one 0 atom. In chemistry, the atomic mass unit was defined as one-sixteenth the average atomic mass of oxygen. These units were slightly smaller than the current carbon-12 based unit. 

One of the most useful pieces of information that can be provided by a standard mass spectrometer is the molecular mass of a compound with an accuracy of 1 dalton (Da) (1 Da= 1 unified atomic mass unit (u)= 1.660538921(73) x 10 kg). Even better, high-resolution mass spectrometry can provide us with accurate molecular masses, which are accurate to about 10 Da, depending on the total resolution of the spectrometer. At this level the atomic masses deviate substantially from multiples of 1 Da, e.g. is 12.0000 Da by definition, but the mass of isotope is 15.9949 Da (not 16), and that of is 14.0031 Da (not 14). This sensitivity can be used to work out the elemental composition or, more critically, the isotopic distribution of atoms in a molecule, which can help to identify unknown compounds. And since the basic experimental procedure involves supplying a particular quantity of energy to a molecule or ion to cause it to fragment, we can use it to deduce compound gas-phase stabilization energies, which can be further substantiated by computational modeling. This last application is rather specialized, so discussion is provided in the on-line supplementary section for Chapter 11 on stabilization. 

A carbon-12 atom contains six electrons, six protons, and six neutrons. Assuming the mass of the atom is the sum of the masses of those parts, as given in Table 5.1, calculate the mass of the atom. Why is it not exactly 12 amu, as the definition of the atomic mass unit would suggest ... 

The molar mass (M) of a substance is the mass in granns of 1 mole of the substance. By definition, the mass of a mole of carbon-12 is exactly 12 g. Note that the molar mass of carbon is numerically equal to its atomic mass. Likewise, the atomic mass of calcium is 40.08 amu and its molar mass is 40.08 g, the atomic mass of sodium is 22.99 amu and its molar mass is 22.99 g, and so on. In general, an element s molar mass in grams is numerically equal to its atomic mass in atomic mass units. The molar mass (in grams) of any compound is numerically equal to its molecular or formula mass (in amu). The molar mass of watCT, for example, is 18.02 g, and the molar mass of sodium chloride (NaCl) is 58.44 g. 

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