Atomic Masses and Isotope Abundances

The isotopic composition of the elements (including the exact atomic mass and the abundance of the isotopes), the atomic weights of elements, definitions and abbreviations are summarized in Appendix II. 

To calculate the average atomic mass, we begin with the precise values of the atomic masses and percent abundance of each isotope. 75.77% of Cl has a mass of 34.97 amu and 24.23% of Cl has a mass of 36.97 amu. 

As agreed by the IUPAC Commission on Atomic Weights and Isotopic Abundances in 1979 [42] the relative atomic mass (atomic weight) of an element, E, can be defined for any specified sample. It is the average mass of its atoms in the sample divided by the unified atomic mass unit1 or alternatively the molar mass of its atoms divided by the standard molar mass M = Lmu = 1 gmol-1 ... 

The variations in isotopic composition of many elements in samples of different origin limit the precision to which a relative atomic mass can be given. The standard atomic weights revised biennially by the IUPAC Commission on Atomic Weights and Isotopic Abundances are meant to be applicable for normal materials. This means that to a high level of confidence the relative atomic mass of an element in any normal sample will be within the uncertainty limits of the tabulated value. By normal it is meant here that the material is a reasonably possible source of the element or its compounds in commerce for industry and science and that it has not been subject to significant modification of isotopic composition within a geologically brief period [43]. This, of course, excludes materials studied themselves for very anomalous isotopic composition. 

Table 6.2 lists the relative atomic masses of the elements in the alphabetical order of chemical symbols. The values have been recommended by the IUPAC Commission on Atomic Weights and Isotopic Abundances in 1991 [44] and apply to elements as they exist naturally on earth. 

Given the atomic masses and natural abundances of the isotopes of an element, calculate its chemical atomic mass (Section 1.4, Problems 15-18). 

Mercury (Hg) occurs in nature as a mixture of seven stable isotopes the average atomic mass of the blend is 200.6. The atomic masses and natural abundances of the isotopes are 196.0 (0.15 %), 198.0 (10.1 %), 199.0 (17 %), 200.0 (23.1 %), 201.0 (13.2 %), 202.0 (29.65 %) and 204.0 (6.8 %). Two radioactive isotopes, and ° Hg, have been widely used in toxicological studies, radiometric analysis, and also in checks of yield of analytical procedures. The fact that mercury is an isotope mixture may be of importance in analytical work using mass spectrometry, since there are no reference samples with well-defined isotope composition. 

Section 3.4 is devoted to the radioactivity and isotopes of the rare earth elements and provides the most important data for the different rare earths. For every element, the tables display the isotopes, their atomic masses, and their abundances. 

Table 1.4 Relative atomic masses and relative abundances of some stable isotopes ...

The previous discussion has centered on how to obtain as much molecular mass and chemical structure information as possible from a given sample. However, there are many uses of mass spectrometry where precise isotope ratios are needed and total molecular mass information is unimportant. For accurate measurement of isotope ratio, the sample can be vaporized and then directed into a plasma torch. The sample can be a gas or a solution that is vaporized to form an aerosol, or it can be a solid that is vaporized to an aerosol by laser ablation. Whatever method is used to vaporize the sample, it is then swept into the flame of a plasma torch. Operating at temperatures of about 5000 K and containing large numbers of gas ions and electrons, the plasma completely fragments all substances into ionized atoms within a few milliseconds. The ionized atoms are then passed into a mass analyzer for measurement of their atomic mass and abundance of isotopes. Even intractable substances such as glass, ceramics, rock, and bone can be examined directly by this technique. 

The introductory chapter is brief but provides an ample introduction to mass spectrometry and leaves one comfortable as he/she moves on to the historical and instrumentation chapters that follow. A few of the basic equations are given as part of the review of basic concepts. In these few pages Dr Becker clearly introduces the concepts of atomic mass units relative to carbon, isotopes and isotope abundance. Figures 1.1 and 1.2 go hand in hand in providing the reader with the three major parts of a mass spectrometer (source, ion separation, detection) and show various alternatives for each of these. The subtle use of color in these and subsequent figures adds an attractive benefit for the reader. 

These values are not the same as the atomic masses in the periodic table because these are the exact masses of individual isotopes.The masses in the periodic table are average masses of the element based on the masses and natural abundances of the isotopes of which it is composed. 

Table of Isotopic Abundances, Atomic Masses and Ionization Energies of Elements... 

The mass difference a single electron makes is observable using high-accuracy mass spectrometry. Table 13.12 lists the atomic weight of a proton, neutron, and electron. Table 13.13 lists selected isotopes along with their atomic number, atomic weight, monoisotopic mass, and relative abundance. 

Only two isotopes of boron (B) occur in nature their atomic masses and abundances are given in the following table. Complete the table by computing the relative atomic mass of B to four significant figures, taking the tabulated relative atomic mass of natural boron as 10.811. 

More than half of all the atoms in naturally occurring zirconium are °Zr. The other four stable isotopes of zirconium have the following relative atomic masses and abundances ... 

Naturally occurring iron consists of four isotopes with the abundances indicated here. From the masses and relative abundances of these isotopes, calculate the atomic weight of namrally occurring iron. 

Background Most elements in nature occur as a mixture of isotopes. The weighted average atomic mass of an element can be determined from the atomic mass and the relative abundance of each isotope. In this activity, you will model the isotopes of the imaginary element "Snackium." The measurements you make will be used to calculate a weighted average mass that represents the average atomic mass of "Snackium."... 

---