Mass (mass spectrometry)
Accurate mass 
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.
Average mass 
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.
Exact mass 
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.
Isotopomer and Isotopologue 
An isotopomer (isotopic isomer) is an isomer having the same number of each isotopic atom but differing in their positions. For example, CH3CHDCH3 and CH3CH2CH2D are a pair of constitutional isotopomers.
An isotopomer should not be confused with isotopologues, which are chemical species that differ only in the isotopic composition of their molecules or ions. An example is water, where three of its hydrogen-related isotopologues are: HOH, HOD and DOD, where D stands for deuterium (2H).
Kendrick mass 
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 number 
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.
Molecular mass 
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.
Monoisotopic mass 
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 instead of the isotopic average mass. 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 term is designed for measurements in mass spectrometry primarily with smaller molecules. It is not typically useful as a concept in physics or general chemistry. Monoisotopic mass is typically expressed in unified atomic mass units (u), also called daltons (Da).
Most abundant mass 
Nominal mass 
The nominal mass of an ion or molecule is calculated using the integer mass (ignoring the mass defect) of the most abundant isotope of each element. This is the equivalent of summing the mass numbers of all constituent atoms. For example H = 1, C = 12, O = 16, etc. The nominal mass of H2O is 18, for example.
See also 
- Francis William Aston
- Atomic mass unit
- Gram-molecular weight
- List of elements by atomic mass
- Nitrogen rule (mass spectrometry)
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