Scale , atomic mass

To set up a scale of atomic masses, it is necessary to establish a standard value for one particular species. The modem atomic mass scale is based on the most common isotope of carbon, 1 C. This isotope is assigned a mass of exactly 12 atomic mass units (amu) ... 

Remember to add up all the individual atomic masses, scaling up where... 

It is interesting to note that prior to 1961. two atomic mass scales were used. Chemists preferred a scale based on the assignment of exactly 16. which experience had shown as die average mass of oxygen atoms as they are found in nature. On the oilier hand, physicists preferred to base the scale on a single isotope of oxygen, namely, loO (oxygen-16). The two... 

This small quantity is not easy to work with so, as you saw in Chapter 3, a scale called the relative atomic mass scale is used. In this scale an atom of carbon is given a relative atomic mass, An of 12.00. All other atoms of the other elements are given a relative atomic mass compared to that of carbon. 

You have already seen in Chapter 3 how we can compare the masses of all the other atoms with the mass of carbon atoms. This is the basis of the relative atomic mass scale. Chemists have found by experiment that if you take the relative atomic mass of an element in grams, it always contains 6 x 1023 or one mole of its atoms. 

Before 1961, a physical atomic mass scale was used whose basis was an assignment of the value 16.00000 to 160. What would have been the physical atomic mass of 12C on the old scale ... 

At one time there was a chemical atomic mass scale based on the assignment of the value 16.0000 to naturally occurring oxygen. What would have been the atomic mass, on such a table, of silver, if current information had been available The atomic masses of oxygen and silver on the present table are 15.9994 and 107.8682. 

The nuclidic mass of 90Sr had been determined on the old physical scale (160 = 16.0000) as 89.936. Calculate the mass of 90Sr to the atomic mass scale on which 160 is 15.9949. 

Other numerical values are exact by definition. For example, the atomic mass scale was established by fixing the mass of one atom of 12C as 12.000 Ou. As many more zeros could be added as desired. Other examples include the definition of the inch (1 in = 2.5400 cm) and the calorie (1 cal = 4.184 00 J). 

Atoms are so tiny that, until recently, the masses of individual atoms could not be measured directly (Figure 3.7). However, because mass was so important in Dalton s theory, some measure of atomic masses was necessary. Therefore, a relative scale—the atomic mass scale—is used. This scale is sometimes called the atomic weight scale. On this scale, an average of the masses of all the atoms of the naturally occurring mixture of isotopes of a given element is measured relative to the mass of an atom of a standard. 

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... 

Carbon, which composes diamond, is the basis for the atomic mass scale that is used today. 

Which atom is used today as the standard for the atomic mass scale ... 

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. 

Before 1961, a physical atomic mass scale was used whose basis was an assignment of the value... 

The assumed formulas are presented in line 1. The percent composition of each compovmd, calculated in the usual way, is presented in line 3, showing that these two compovmds, indeed, have different compositions, as required by the law of multiple proportions. Line 4 contains the ratio of the mass of mercury to the mass of oxygen, for each compound. Those ratios can be expressed as the ratio of simple whole numbers (2.25 4.5 = 1 2), fulfilling a condition required by the law of multiple proportions. Notice that Dalton s ideas do not depend upon the values assigned to the elements or the formulas for the compounds involved. Indeed, the question as to which compound, red or black, is associated with which formula cannot be answered from the data available. Thus, although Dalton was unable to establish an atomic mass scale, his general theory did provide an rmderstanding of the three mass-related laws conservation, constant composition, and multiple proportion. Other information was required to establish the relative masses of atoms. 

Mass defect is defined as the difference between the exact molecular weight and nominal molecular weight of an element. The atomic mass scale defines carbon-12 with a mass of exactly 12.0000 Da, therefore all other elements will have a uniquely different mass defect. For example, the mass defects of hydrogen and oxygen are 0.007825 and —0.005085 Da, respectively. Therefore, oxidation will introduce a mass defect of—5.1 mDa, while glucuronidation will introduce a mass defect of -i- 32 mDa. The mass defects of common nonsynthetic (Phase 1) and synthetic (Phase II) metabolites typically fall within 50 mDa. Therefore with LC/MS instruments capable of high mass accuracy, it is possible to filter out matrix-related interference ions whose mass... 

This tiny mass is exactly equal to one-twelfth of the mass of one atom of carbon (isotope mass number 12). This means that the mass of one atom of gC is exactly 12 on the atomic mass scale. This is written as m( gC) = 12 u. 

The mass of one atom of an isotope on the atomic mass scale is called its Isotopic mass m. Some isotopic masses are listed in Table 3.2. Remember that the mass of atoms is defined with reference to a standard atom , i.e. one atom of carbon-12 for which m( gC) = 12.0000 u. 

Notice that isotopic masses (in atomic mass units) are very nearly equal in size to the mass number of an isotope. This is because the mass of the electrons, protons and neutrons are approximately 0, 1 and 1, respectively, on the atomic mass scale. Note also that the units of m and A are not the same - A is unitless. 

The idea of molecular mass was introduced in Chapter 3. This is the mass of one molecule of substance on the atomic mass scale. Molecular masses are calculated using the atomic masses of the constituent atoms. A list of approximate atomic masses for atoms of elements is shown in Table 8.1. You should use these for calculations, unless you are instructed otherwise. 

SECTION Z4 The atomic mass scale is defined by assigning a mass of exactly 12 amu to a atom. The atomic weight (average atomic mass) of an element can be calculated from the relative abundances and masses of that elements isotopes. The mass spectrometer provides the most direct and accurate means of experimentally measuring atomic (and molecular) weights. 

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