Isotope relative mass table

Magnesium [7439-95-4] atomic number 12, is in Group 2 (IIA) of the Periodic Table between beryllium and calcium. It has an electronic configuration of 1T2T2 3T and a valence of two. The element occurs as three isotopes with mass numbers 24, 25, and 26 existing in the relative frequencies of 77, 11.5, and 11.1%, respectively. 

The important odorant, 2-acetyltetrahydropyridine (ACTPY), has already been mentioned in the previous section. ACTPY and ACPY play key roles in the aroma of popcorn.227 Freeze-dried maize contains relatively high amounts of proline (155), whereas ornithine is not detectable (< 5 mg kg ) Schieberle treated a low-molecu-lar-mass fraction of an aqueous extract of maize in different ways and determined ACTPY and ACPY by isotope dilution assay (Table 5.2). Steam-distillation extraction gave 130 times as much ACTPY than ACPY however, in the presence of added 2-oxo-propanal, the amount of the former was multiplied by 4, but that of the latter by 29. Dry-heating, as in popping, increased the latter further, but the former became undetectable. 

Table 1 provides the composition of the terrestrial atmosphere, which is the standard for all noble-gas analyses, and indicates the general relative abundances of the isotopes. The masses of the rare-gas planetary reservoirs are given in Table 2. 

When a pure elemental gas, such as neon, was analyzed by a mass spectrometer, multiple peaks (two in the case of neon) were observed (see Fig. 1.11). Apparently, several kinds of atoms of the same element exist, differing only by their relative masses. Experiments on radioactive decay showed no differences in the chemical properties of these different forms of each element, so they all occupy the same place in the periodic table of the elements (see Chapter 3). Thus the different forms were named isotopes. Isotopes are identified by the chemical symbol for the element with a numerical superscript on the left side to specify the measured relative mass, for example °Ne and Ne. Although the existence of isotopes of the elements had been inferred from studies of the radioactive decay paths of uranium and other heavy elements, mass spectrometry provided confirmation of their existence and their physical characterization. Later, we discuss the properties of the elementary particles that account for the mass differences of isotopes. Here, we discuss mass spectrometry as a tool for measuring atomic and molecular masses and the development of the modern atomic mass scale. 

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... 

You know from Chapter 2 that average atomic masses of the elements are given on the periodic table. For example, the average mass of one iron atom is 55.8 u, where u means atomic mass units. The atomic mass unit is defined so that the atomic mass of an atom of the most common carbon isotope is exactly 12 u, and the mass of 1 mol of the most common isotope of carbon atoms is exactly 12 g. The mass of 1 mol of a pure substance is called its molar mass. For example, the molar mass of iron is 55.847 g, and the molar mass of platinum is 195.08 g. Relative masses of elements are demonstrated in Figure 12.4. The molar mass is the mass in grams of the average atomic mass. 

Figure 15. Jupiter atmospheric Xe composition measured by the Galileo Probe mass spectrometer (Mahaffy et al. 2000), calculated as the abundance of each isotope M divided by the total abundance for all M and plotted relative to the NEA-Xe composition (Table 1) represented in the same way. The five most abundant Jovian Xe isotopes are indicated by the shaded symbols. U-Xe, SW2-Xe, and AVCC-Xe compositions in Table 1 are shown for comparison in the same representation. The data indicate deficits at the two heaviest isotopes relative to NEA-Xe and AVCC-Xe, but uncertainties are too large to rule between SW-Xe and U-Xe as the Jovian composition.

In Chap. 3 we pointed out that the formulas for both ionic and covalent compounds ignored the nature of the chemical bonds of compounds. A formula is a statement of the combining ratios or the relative numbers of atoms of each kind in the simplest unit representing the composition of the compound. We also learned in Chap. 2 that the atoms of the elements have different masses, which were expressed as relative masses, and that the quantity of any element which is numerically equal to its atomic mass must contain the same number of atoms as the corresponding quantity of any other element. The table of atomic masses assures us that 55.85 g of iron contains the same number of atoms as 12.00 g of the C isotope of carbon. Can we expand the concept of relative mass to compounds as well as elements For example, how many grams of water will contain the same number of fundamental units as 12 g of C carbon We know that the fundamental unit of water is a molecule with the composition H2O. Its relative mass must be equal to the sum of the masses of the atoms in the molecule ... 

The atomic weights listed in the periodic table are relative numbers ( C = 12.0000. ..) based upon the weighted average of naturally occurring isotopes (e.g., the atomic mass of chlorine is 35.45 reflecting the roughly 3 1 ratio of Cl to Cl). The isotope Cl has 17 protons (atomic number = 17) and 18 neutrons in its nucleus C1 has 17 protons and 20 neutrons. A more precise analysis combines the relative abundances and precise relative masses of the two stable nuclides of chlorine ( Cl 75.78 percent 34.968853. Cl 24.22 percent 36.965903) as follows Relative Atomic Mass of Chlorine 0.7578 (34.968853) + 0.2422 (36.965903) = 35.45 It is noteworthy that on rare occasions, lUPAC may introduce a very slight modification to the atomic mass provided for an element in the periodic table. The relative masses of the nuclides are known to... 

The relative mass difference of isotopes in ionic liquids is higher than in aqueous solution. This is because solvation of ions by water tends to diminish mass differences. Consequently, greater differences in electrophoretic mobility are possible in ionic liquids. Countercurrent electrophoresis in columns allowed isolation of highly enriched isotope mixtures without great difficulty. Some examples of the mass difference effect are summarized in Table 1. These results depend on the temperature as well as the composition of the ionic liquid. 

Table 1 Relative abundances on Earth of the stable isotopes used in stable isotope ratio mass spectrometry...

Based on the relative percentages, we should be able to decide if this answer makes sense. The individual isotopes have masses of roughly 35 and 37, so a 50/50 ratio would lead to an average mass of about 36. But the actual abundance of the Cl isotope is greater than that of Cl, so the average mass should be closer to 35. Thus our answer of 35.45 seems reasonable. And of course we can check the answer by consulting a periodic table. 

The relative masses of the elements as given in the periodic table are referred to as atomic masses or atomic weights. We will use the term atomic weights in this book. In those cases where the naturally occurring element exists in the form of a mixture of isotopes, the recorded atomic weight is the average value for the naturally occurring isotope mixture. This idea is discussed in Section 2.5. 

I Table 14.2 Precise Masses and Natural Abundances of Isotopes Relative to 100 Atoms of the Most Abundant Isotope... 

Atomic mass (atomic weight) Relative mass of an atom. By current international agreement, the standard for all atomic masses is the isotope carbon-12, which is arbitrarily assigned an atomic mass of exactly 12. Atomic mass was the original numerical basis for the periodic table. 

Molecular Weights of Peptides and Proteins. Traditionally, organic mass spec-troscopists have utilized peaks corresponding to monoisotopic ion species (containing only H, N, 0, etc.) to determine molecular or fragment masses (Table 10.1). While their relative abundances are low, the contribution to the molecular-ion mass from one or several and isotopes increases for peptides... 

A peak that appears two mass units higher than the mass of the molecular ion peak is called the M + 2 peak. The intensity of the M + 2 peak of ethane is only 0.01% of the intensity of the molecular ion peak. The contribution due to two deuterium atoms replacing hydrogen atoms would be (0.016 x 0.016)7100 = 0.00000256%, a negligible amount. To assist in the determination of the ratios of molecular ion, M + 1, and Af + 2 peaks. Table 3.5 lists the natural abundances of some common elements and their isotopes. In this table, the relative abundances of the isotopes of each element are calculated by setting the abundances of the most common isotopes equal to 100. 

Each element that has neither a stable isotope nor a characteristic natural isotopic composition is represented in this table by one of that element s commonly known radioisotopes identified by mass number and relative atomic mass. 

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