Atommasse

Atoms — and the protons, neutrons, and electrons that compose them — are extremely small. For example, a carbon atom weighs less than 2 × 10⁻²³ g. When describing the properties of tiny objects such as atoms, we use appropriately small units of measure, such as the atomic mass unit (amu). The amu was originally defined based on hydrogen, the lightest element, then later in terms of oxygen. Since 1961, it has been defined with regard to the most abundant isotope of carbon, atoms of which are assigned masses of exactly 12 amu. Thus, one amu is exactly ¹/₁₂ of the mass of one carbon-12 atom: 1 amu = 1.6605 × 10⁻²⁴ g. The Dalton (Da) and the unified atomic mass unit (u) are alternative units that are equivalent to the amu. 

Because each proton and each neutron contribute approximately one amu to the mass of an atom, and each electron contributes far less, the atomic mass of a single atom is approximately equal to its mass number (a whole number sum of protons and neutrons in the atom). For example, the mass number of a single nitrogen atom is 14 (7 protons + 7 neutrons). However, the average masses of atoms of most elements are not whole numbers because most elements exist naturally as mixtures of two or more isotopes. Isotopes are atoms of the same element with the same proton number but a different number of neutrons. The mass of an element shown in a periodic table or listed in a table of atomic masses is a weighted average mass of all the isotopes present in a naturally occurring sample of that element. The average mass is equal to the sum of each individual isotope’s mass multiplied by its fractional abundance.

For example, the element chlorine (atomic number 17) is composed of two isotopes: ³⁵Cl and chlorine ³⁷Cl. About 75.78% of all chlorine atoms are ³⁵Cl with a mass of 34.969 amu, and the remaining 24.22% are ³⁷Cl with a mass of 36.966 amu. The average atomic mass for chlorine is calculated to be:

It is important to understand that no single chlorine atom weighs exactly 35.45 amu; This value is the average mass of all chlorine atoms, and individual chlorine atoms weigh either approximately 35 amu or 37 amu. Also, because naturally occurring chlorine contains more ³⁵Cl atoms than ³⁷Cl atoms, the weighted average mass of chlorine is closer to 35 amu than to 37 amu.

The occurrence and natural abundances of isotopes can be experimentally determined using an instrument called a mass spectrometer. Mass spectrometry (MS) is widely used in chemistry, forensics, medicine, environmental science, and many other fields to analyze and help identify the substances in a sample of material. In a typical mass spectrometer, the sample is vaporized and exposed to a high-energy electron beam that causes the sample’s atoms or molecules to become electrically charged, typically by losing one or more electrons. These cations then pass through a variable magnetic field that deflects each cation’s path to the extent that depends on both its mass and charge. Finally, the ions are detected, and a plot of the relative number of ions generated versus their mass-to-charge ratios, a mass spectrum, is made. The height of each vertical feature or peak in a mass spectrum is proportional to the fraction of cations with the specified mass-to-charge ratio. Since its first use during the development of modern atomic theory, MS has evolved to become a powerful tool for chemical analysis in a wide range of applications.

Text adapted from Openstax Chemistry 2e, Section 2.3: Atomic Structure and Symbolism.