Nominal isotopic mass

Atomic mass is the mass of an atomic particle, i.e. a specific isotope. When expressed in unified atomic mass units, this is called the relative isotopic mass. The word relative is added to denote the fact that all masses are scaled to that of the isotope when set to 12 u. Nominal isotope masses are more commonly used when applying analytical techniques such as Secondary Ion Mass Spectrometry (SIMS) because this significantly simplifies matters without detracting from the information content needed. This represents the number of protons and neutrons within the nucleus, i.e. equal to the atomic mass number (A). Note It was the mass spectrograph constructed by Aston in 1919 (the first mass spectrometer from which SIMS evolved as covered in Section 1.2.1) that confirmed the existence of the isotopes, and allowed for the first time, an accurate means of measuring their relative mass (that relative to H, 0, or more recently C) and distribution. 

Nominal ion mass. The mass of an ion with a given empirical formula calculated using the integer mass numbers of the most abundant isotope of each element, e.g., C = 12, H = 1, O = 16. 

The isotopic mass is the exact mass of an isotope. It is very close to but not equal to the nominal mass of the isotope (Table 3.1). The only exception is the carbon isotope which has an isotopic mass of 12.000000 u. The unified atomic mass... 

Example the molecular ions of nitrogen, N2, carbon monoxide, CO, and ethene, C2H4, have the same nominal mass of 28 u, i.e., they are so-called iso-baric ions. The isotopic masses of the most abundant isotopes of hydrogen, carbon, nitrogen and oxygen are 1.007825 u, 12.000000 u, 14.003070 u and 15.994915 u, respectively. Using these values, the calculated ionic masses are 28.00559 u for Nz"" , 27.99437 u for CQ-", and 28.03075 u for CjH/. This means they differ by some millimass units" (mmu) from each other, and none of these isobaric ions has precisely 28.00000 u (Chap. 3.3.4 and Chap. 6.9.6). 

The term mass deficiency better describes the fact that the exact mass of an isotope or a complete molecule is lower than the corresponding nominal mass. In case of for example, the isotopic mass is 15.994915 u, being 5.085 mmu deficient as compared to the nominal value. 

Note The use of nominal mass is limited to the low mass range. Above about 500 u the first decimal of isotopic mass can be larger than. 5 causing it to be rounded up to 501 u instead of the expected value of 500 u. This will in turn lead to severe misinterpretation of a mass spectrum (Chap. 6). 

An ion containing a less abundant combination of isotopes, also included under P.I.D., is not classified separately because identification is usually more simple from the more abundant isotopic combination. The mass number and relative abundance of isotopic ions can be calculated from the accompanying table. It might be argued that classification of these could be useful where the more abundant isotopic combination is obscured by another ion of nominally identical mass. This, however, will be an unusual circumstance and can be overcome by careful use of the table or by the use of exact empirical structure determination through high resolution techniques. 

Since the advent of high-resolution mass spectrometers, it is also possible to use very precise mass determinations of molecular ion peaks to determine molecular formulas. When the atomic weights of the elements are determined very precisely, it is found that they do not have exactly integral values. Every isotopic mass is characterized by a small mass defect, which is the amount by which the mass of the isotope differs from a perfectly integral mass number. The mass defect for every isotope of every element is unique. As a result, a precise mass determination can be used to determine the molecular formula of the sample substance, since every combination of atomic weights at a given nominal mass value will be unique when mass defects are considered. For example, each of the substances shown in Table 1.4 has a nominal mass of 44 amu. As can be seen from the table, their exact masses, obtained by adding exact atomic masses, are substantially different when measured to four decimal places. 

IV.lb) High mass resolution The true Isotope mass differs slightly from its nominal Integral value ("mass defect"). If the mass resolving power of the spectrometer exceeds the ratio of the integral mass to the mass defect, then two Isotopes at the same Integral mass can be separated. For in-... 

What are the differences between nominal mass, exact mass, accurate mass, mono-isotopic mass, molecular mass, and isobaric mass The nominal mass of a compound, ion, or fragment is calculated using the masses of the elements rounded to the nearest whole number (e.g., C = 12, H = 1, O = 16). For example, cholesterol with an empirical formula C27H4gO has a nominal molecular mass of 386 Da. It is common to see molecular mass referred to as molecular weight (MW). 

Atomic Number Element Symbol Nominal Mass % Relative % Isotopic Mass Average Mass... 

The latter condition will play an important role in Section 8.6. Table 8.2 contains natural isotope distributions for the elements X of (source [288]). For these elements the lowest isotope mass equals the nominal mass nix = For masses m not mentioned in the table, lx ni) = 0. The elements of n can be partitioned into three classes according to their isotope distributions (see also [194]) ... 

Class 0 Highest isotope mass and nominal element mass coincide = wx-... 

Remember that the nominal mass of fi is the sum of nominal masses of its atoms, where the nominal mass of an element atom is the isotope mass of maximal abundance. 

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