Mass (mass spectrometry)
- 1 Units
- 2 Molecular mass
- 3 Average mass
- 4 Mass number
- 5 Nominal mass
- 6 Accurate mass
- 7 Exact mass
- 8 Monoisotopic mass
- 9 Most abundant mass
- 10 Isotopomer and isotopologue
- 11 Kendrick mass
- 12 Mass defect (mass spectrometry)
- 13 Packing fraction (mass spectrometry)
- 14 Nitrogen rule
- 15 Prout's hypothesis and the whole number rule
- 16 See also
- 17 References
- 18 External links
The unified atomic mass unit (symbol: u) is the standard unit that is used for indicating mass on an atomic or molecular scale (atomic mass). The dalton (symbol: Da) is equivalent to the unified atomic mass unit. One unified atomic mass unit is approximately the mass of one a single proton or neutron. The unified atomic mass unit has a value of 538921(73)×10−27 kg. 1.660 The amu without the "unified" prefix is an obsolete unit based on oxygen, which was replaced in 1961.
The molecular mass (abbreviated Mr) of a substance, formerly also called molecular weight and abbreviated as MW, is the mass of one molecule of that substance, relative to the unified atomic mass unit u (equal to 1/12 the mass of one atom of 12C). Due to this relativity, the molecular mass of a substance is commonly referred to as the relative molecular mass, and abbreviated to Mr.
The average mass of a molecule is obtained by summing the average atomic masses of the constituent elements. For example, the average mass of natural water with formula H2O is 1.00794 + 1.00794 + 15.9994 = 18.01528.
The mass number, also called the nucleon number, is the number of protons and neutrons in an atomic nucleus. The mass number is unique for each isotope of an element and is written either after the element name or as a superscript to the left of an element's symbol. For example, carbon-12 (12C) has 6 protons and 6 neutrons.
The nominal mass for an element is the mass number of its most abundant naturally occurring stable isotope, and for an ion or molecule, the nominal mass is the sum of the nominal masses of the constituent atoms. Isotope abundances are tabulated by IUPAC: for example carbon has two stable isotopes 12C at 98.9% natural abundance and 13C at 1.1% natural abundance, thus the nominal mass of carbon is 12. The nominal mass is not always the lowest mass number, for example iron has isotopes 54Fe, 56Fe, 57Fe, and 58Fe with abundances 6%, 92%, 10%, and 2%, respectively, and a nominal mass of 56. For a molecule, the nominal mass is obtained by summing the nominal masses of the constituent elements, for example water has two hydrogen atoms with nominal mass 1 and one oxygen atom with nominal mass 16, therefore the nominal mass of H2O is 18.
In mass spectrometry, the difference between the nominal mass and the monoisotopic mass is the mass defect. This differs from the definition of mass defect used in physics which is the difference between the mass of a composite particle and the sum of the masses of its constituent parts.
The accurate mass (more appropriately, the measured accurate mass) is an experimentally determined mass that allows the elemental composition to be determined. For molecules with mass below 200 u, 5 ppm accuracy is often sufficient to uniquely determine the elemental composition.
The exact mass of an isotopic species (more appropriately, the calculated exact mass) is obtained by summing the masses of the individual isotopes of the molecule. For example, the exact mass of water containing two hydrogen-1 (1H) and one oxygen-16 (16O) is 1.0078 + 1.0078 + 15.9949 = 18.0105. The exact mass of heavy water, containing two hydrogen-2 (deuterium or 2H) and one oxygen-16 (16O) is 2.0141 + 2.0141 + 15.9949 = 20.0229.
When an exact mass value is given without specifying an isotopic species, it normally refers to the most abundant isotopic species.
The monoisotopic mass is the sum of the masses of the atoms in a molecule using the unbound, ground-state, rest mass of the principal (most abundant) isotope for each element. The monoisotopic mass of a molecule or ion is the exact mass obtained using the principal isotopes. Monoisotopic mass is typically expressed in unified atomic mass units.
For typical organic compounds, where the monoisotopic mass is most commonly used, this also results in the lightest isotope being selected. For some heavier atoms such as iron and argon the principal isotope is not the lightest isotope. The mass spectrum peak corresponding to the monoisotopic mass is often not observed for large molecules, but can be determined from the isotopic distribution.
Most abundant mass
Isotopomer and isotopologue
Isotopomers (isotopic isomers) are isomers having the same number of each isotopic atom, but differing in the positions of the isotopic atoms. For example, CH3CHDCH3 and CH3CH2CH2D are a pair of structural isotopomers.
Isotopomers should not be confused with isotopologues, which are chemical species that differ in the isotopic composition of their molecules or ions. For example, three isotopologues of the water molecule with different isotopic composition of hydrogen are: HOH, HOD and DOD, where D stands for deuterium (2H).
The Kendrick mass is a mass obtained by multiplying the measured mass by a numeric factor. The Kendrick mass is used to aid in the identification of molecules of similar chemical structure from peaks in mass spectra. The method of stating mass was suggested in 1963 by the chemist Edward Kendrick.
The Kendrick mass for a family of compounds F is given by
For hydrocarbon analysis, F=CH2.
Mass defect (mass spectrometry)
The mass defect used in nuclear physics is different from its use in mass spectrometry. In nuclear physics, the mass defect is the difference in the mass of a composite particle and the sum of the masses of its component parts. In mass spectrometry the mass defect is defined as the difference between the exact mass and the nearest integer mass.
The Kendrick mass defect is the exact Kendrick mass subtracted from the nearest integer Kendrick mass.
Mass defect filtering can be used to selectively detect compounds with a mass spectrometer based on their chemical composition.
Packing fraction (mass spectrometry)
The term packing fraction was defined by Aston as the difference of the measured mass M and the nearest integer mass I (based on the oxygen-16 mass scale) divided by the quantity comprising the mass number multiplied by ten thousand:
Aston's early model of nuclear structure (prior to the discovery of the neutron) postulated that the electromagnetic fields of closely packed protons and electrons in the nucleus would interfere and a fraction of the mass would be destroyed. A low packing fraction is indicative of a stable nucleus.
The nitrogen rule states that organic compounds containing exclusively hydrogen, carbon, nitrogen, oxygen, silicon, phosphorus, sulfur, and the halogens either have an odd nominal mass that indicates an odd number of nitrogen atoms are present or an even nominal mass that indicates an even number of nitrogen atoms are present in the molecular ion.
Prout's hypothesis and the whole number rule
The whole number rule states that the masses of the isotopes are integer multiples of the mass of the hydrogen atom. The rule is a modified version of Prout's hypothesis proposed in 1815, to the effect that atomic weights are multiples of the weight of the hydrogen atom.
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