Isotopes and Atomic Mass

A common mistake for beginners in mass spectrometry is to confuse average atomic mass and isotopic mass. For example, the average atomic mass for chlorine is close to 35.45, but this average is of the numbers and masses of Cl and Cl isotopes. This average must be used for instruments that cannot differentiate isotopes (for example, gravimetric balances). Mass spectrometers do differentiate isotopes by mass, so it is important in mass spectrometry that isotopic masses be used... 

Refer to the "Atoms and Moles" chapter for a discussion of atomic mass and isotopes. 

For several decades, mass spectrometers were used primarily to determine atomic masses and isotopic ratios. Now they are applied to a large variety of chemical problems and low resolution mass spectrometers are used for routine chemical analysis. For exanq>le, a modem mass spectrometer can easily distinguish between species such as and... 

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. 

To solve this problem, multiply the percentage of each isotope by its atomic mass and add those products. 

This information is normally shown as a superscript (atomic mass) and a subscript (atomic number) to the symbol for the element. Hence 1%C represents the isotope of carbon (atomic number 6) with the atomic mass of 14. In practice the subscript is often omitted because the atomic number is unique to the element that is represented by the appropriate letter (e.g. C for carbon). For simplicity in the spoken form and often in the written form, isotopes are often referred to as carbon-14, phosphorus-32, etc. 

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. 

The analysis methods described here have highlighted some of the systematic features in the predictive properties of several of the commonly used atomic mass models. Additional understanding of these features and the availability of many new atomic masses for isotopes far from the stability line will serve as a basis for improving the models. The need clearly exists for a comprehensive revision and update of the mass predictions. A project, coordinated by the author, has been started to accomplish this. It is expected that new sets of mass predictions from a number of groups may be available late in 1986. 

The following table lists the atomic masses and relative percent concentrations of naturally occurring isotopes of importance in mass spectrometry.1"5... 

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. 

Naturally occurring oxygen consists of three isotopes. These have the mass numbers of 16, 17, and 18, and they are present in the percentages of 99.759, 0.037, and 0.204, respectively. Before 1961,160 was the standard of atomic mass, and physicists took it as exactly 16.0000 amu. However, chemists usually considered the natural mixture of isotopes as... 

Atoms without implied hydrogens or with any specific properties, such as charge, atomic mass for isotope specification, and so on, are given in square brackets. Thus, C is methane and [C] is elemental carbon C[13C](=0)0 is acetic acid with 13C atom in the carboxyl group, and CC(=0)[0-] is acetate anion. 

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. 

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

Since conception over 100 years ago, MS has become an important analytical and research tool with diverse applications ranging from astronomical study of the solar system to materials analysis and process monitoring in chemical, oil and pharmaceutical industries. Use of MS has led to very many scientific breakthroughs including the discovery of isotopes, accurate determination of atomic mass, and the characterization of biomolecular structure. Indeed, MS is now a fundamental technique employed in pharmacology, toxicology and other biological, environmental and biomedical sciences. 

The early pioneers of chemistry, trying to verify Dalton s atomic theory, could not measure the mass of individual atoms. The best they could do was to measure the masses of equal numbers of atoms (or other known ratios of atoms) of two (or more) elements at a time, to determine their relative masses. They established one element as a standard, gave it an arbitrary value of atomic mass, and used that value to establish the atomic mass scale. The last naturally occurring mixture of isotopes that was used as a standard was oxygen, defined as having an atomic mass of exactly 16 atomic mass units (amu). That standard has been replaced see the next subsection. The atomic mass unit is tiny it takes... 

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