Element relative masses

Element Isotope Atomic weight Relative abundance (%) Elemental relative mass difference Molecular relative mass difference (%) Terrestrial %0 range ppm Technical %0 precision ppm... 

A special type of substituent effect which has proved veiy valuable in the study of reaction mechanisms is the replacement of an atom by one of its isotopes. Isotopic substitution most often involves replacing protium by deuterium (or tritium) but is applicable to nuclei other than hydrogen. The quantitative differences are largest, however, for hydrogen, because its isotopes have the largest relative mass differences. Isotopic substitution usually has no effect on the qualitative chemical reactivity of the substrate, but often has an easily measured effect on the rate at which reaction occurs. Let us consider how this modification of the rate arises. Initially, the discussion will concern primary kinetic isotope effects, those in which a bond to the isotopically substituted atom is broken in the rate-determining step. We will use C—H bonds as the specific topic of discussion, but the same concepts apply for other elements. 

Relative masses of atoms of different elements are expressed in terms of their atomic masses (often referred to as atomic weights). The atomic mass of an element indicates how heavy, on the average, one atom of that element is compared with an atom of another element... 

Even if c (Equation 5-12) is fairly large, an element cannot be precisely determined—or may even escape detection—if it is present in too small amount relative to the matrix. What amount is too small depends not only upon the relative mass (or weight-fraction) of element sought but also upon the mass absorption coefficient of the matrix, as Equations 5-8 and 5-9 imply. 

If the sample consists of atoms of one element, the mass spectrum gives the isotopic distribution of the sample. The relative molar masses of the isotopes can be determined by comparison with atoms of carbon-12. If the sample is a compound, the formula and structure of the compound can be determined by studying the fragments. For example, the + 1 ions that CH4 could produce are CH4, CH3+, CH, CFI4, C+, and H4. Some of the particles that strike the detector are those that result when the molecule simply loses an electron (for example, to produce Cl I4+ from methane) ... 

Different isotopes differ in their atomic masses. The intensities of the signals from different isotopic ions allow isotopic abundances to be determined with high accuracy. Mass spectrometry reveals that the isotopic abundances in elemental samples from different sources have slightly different values. Isotopic ratios vary because isotopes with different masses have slightly different properties for example, they move at slightly different speeds. These differences have tiny effects at the level of parts per ten thousand (0.0001). The effects are too small to appear as variations In the elemental molar masses. Nevertheless, high-precision mass spectrometry can measure relative abundances of isotopes to around 1 part in 100,000. 

Once Dalton s hypotheses had been proposed, the next logical step was to determine the relative masses of the atoms of the elements. Since there was no way at that time to determine the mass of an individual atom, the relative masses were the best information available. That is, one could tell that an atom of one element had a mass twice as great as an atom of a different element (or times as much, or 17.3 times as much, etc.). How could even these relative masses be determined They could be determined by taking equal (large) numbers of atoms of two elements and by determining the ratio of masses of these collections of atoms. 

The atomic weight of an element is the relative mass of an average atom of the element compared with 12C, which has an atomic weight of exactly 12. Thus, since a sulfur atom has a mass jj times that of a carbon atom, the atomic weight of sulfur is... 

Figure 3. Relative mass differences for elements that have two or more isotopes, cast as Am/m, where Am is the unit mass difference (Am = 1), and m is the average mass of the isotopes of that element, as a function of atomic number (Z). Note that Am/m is reported in percent, and is plotted on a log scale. Elements that are discussed in this volume shown in large black squares. Other elements that have been the major focus of isotopic studies shown in gray diamonds, and include H, C, O, and S. The relatively large mass differences for the light elements generally produce the largest isotopic fractionations, whereas the magnitude of isotopic fractionation is expected to markedly decrease with increasing mass.

For elements that have three or more isotopes, isotopic fractionations may be defined using two or more isotopic ratios. Assuming that isotopic fractionation occurs through a mass-dependent process, the extent of fractionation will be a function of the relative mass differences of the two isotope ratios. For example, assuming a simple harmonic oscillator for molecular motion, the isotopic fractionation of may be related to as ... 

Isotopic fractionations for relatively light elements, such as O, are generally higher than those of higher-mass elements, as expected based on changes in their relative mass differences (Fig. 3). CaCOj - H2O curve for fractionations... 

The significant relative mass difference (c. 16%) between the two stable isotopes of Li (approximately Li 7.5%, Li 92.5%), coupled with broad elemental dispersion in Earth and planetary materials, makes this a system of considerable interest in fingerprinting geochemical processes, determining mass balances, and in thermometry. Natural mass fractionation in this system is responsible for c. 6% variation among materials examined to date (Fig. 1). Although the modem era of Li isotope quantification has begun, there are still many questions about the Li isotopic compositions of fundamental materials and the nature of fractionation by important mechanisms that are unanswered (e.g., Hoefs 1997). 

Spectral analysis shows quite clearly that the various types of atoms are exactly the same on Earth as in the sky, in my own hand or in the hand of Orion. Stars are material objects, in the baryonic sense of the term. All astrophysical objects, apart from a noteworthy fraction of the dark-matter haloes, all stars and gaseous clouds are undoubtedly composed of atoms. However, the relative proportions of these atoms vary from one place to another. The term abundance is traditionally used to describe the quantity of a particular element relative to the quantity of hydrogen. Apart from this purely astronomical definition, the global criterion of metallicity has been defined with a view to chemical differentiation of various media. Astronomers abuse the term metaT by applying it to all elements heavier than helium. They reserve the letter Z for the mass fraction of elements above helium in a given sample, i.e. the percentage of metals by mass contained in 1 g of the matter under consideration. (Note that the same symbol is used for the atomic number, i.e. the number of protons in the nucleus. The context should distinguish which is intended.)... 

Clouds of gas in the interstellar medium are called gaseous nebulas. These nebulas are regions of the interstellar medium with above-average density. The proportions of elements in the interstellar medium conform to the abundances in the table, that is, 90% hydrogen atoms, 9% helium atoms and less than 1% heavier atoms, where these percentages now refer to relative numbers of atoms rather than relative mass. 

Atomic Weight The mass of an element relative to its atoms. 

Table 5.1 Relative masses and natural abundances for some commonly occurring elements...

In 1860, an international conference of chemists convened to discuss how the masses of atoms of different elements could be measured and compared with one another. (As we explore in Section 9.2, knowing the relative masses of atoms helps chemists understand and control chemical reactions.) At the time, there was little agreement because different chemists using different experimental procedures and assuming different theories came up with different results. Progress in chemistry was stymied by this problem. 

Element Nominal Mass (Da) Exact Mass (Da) Mass Shift (Relative to Nominal Mass) (Da)... 

During this procedure, your teacher will introduce the mole concept. Use a periodic table to find the relative masses of all the elements in the molecule CuS04 and K2Cr04, respectively Cu (copper), S (sulfur) and O (oxygen) and K (potassium), Cr (chromium), O (oxygen). The relative mass in grams for any element contains the same number of atoms. This number of atoms, 6.02 x 1023, is called a mole. In the preparation of any 0.1 M solution, 0.1 mole of molecules is needed. A 0.1 M solution, by definition, contains 0.1 mole of a substance dissolved in 1.0 liter of a solvent. 

---