Introduction to Mass Spectrometry

When the organic chemist is faced with the task of assigning a structure to an unidentified compound, the first challenge is determination of the molecular formula. The chemist can derive the molecular formula from the elemental analysis plus the molecular mass to the nearest integer value or from the exact molecular mass measured to four decimal places. The molecular mass to the nearest integer is determined by low resolution mass spectrometry, and the exact molecular mass to four decimal places, by high resolution mass spectrometry. The elemental analysis gives the empirical formula which consists of the atoms present in the molecular formula in their simplest integer ratio. The molecular mass to the nearest integer value together with the empirical formula mass is then used to determine if the empirical formula should be doubled, tripled, etc. to obtain the molecular formula. The exact molecular mass yields directly a unique molecular formula when the molecular mass is less than approximately 1000 amu (atomic mass units).

The following example demonstrates how to obtain a molecular formula from elemental analysis data and the mass spectral molecular mass.

Sample Problem

The compound biacetyl (butane-2,3-dione) gives the following elemental analysis data and mass spectral molecular mass: C, 55.8%, H, 7.0%; molecular mass 86 amu. Calculate the empirical formula and the molecular formula.

Empirical formula calculation: 55.8% + 7.0% = 62.8%; hence the balance of 37.2% is O. (The balance of the elemental composition is always assumed to be oxygen, see section 21.2, p. 244.). The atomic ratio of the elements is obtained by dividing each percent by the respective atomic mass and is 55.8/12 to 7.0/1 to 37.2/16 = 4.65 to 7.0 to 2.32. Dividing through by 2.32, the smallest of the numbers in the ratio, gives 2:3:1 or an empirical formula of C2H3O.

Molecular formula calculation: The empirical formula mass is 43 amu which is half the molecular mass; hence, the molecular formula is twice the empirical formula or C4H6O2.


Note: This tutorial was written by Tad Koch, Professor of Chemistry at the University of Colorado, Boulder. Editorial support, web page conversion, and graphics support were provided by Patty Feist, Coordinator for the Organic Chemistry Labs. Copyright information: Original content © University of Colorado, Boulder, Chemistry and Biochemistry Department, 2003. The information on these pages is available for academic use without restriction. We do ask that you cite the source of your information. Correspondence should be directed to Patty Feist.


Next section: Natural Abundance Atomic Isotopes

Copyright information: Original content © University of Colorado, Boulder, Chemistry and Biochemistry Department, 2011. The information on these pages is available for academic use without restriction.