Atomic mass is a fundamental property in chemistry and physics, representing the mass of a single atom. This value allows scientists to determine the quantity of substances in chemical reactions and analyze the composition of matter. While related to the count of particles in the nucleus, atomic mass is a distinct measurement. The mass number is a simple integer count of an atom’s protons and neutrons, but the actual atomic mass is a more precise value that is typically not a whole number. Determining atomic mass involves a combination of arbitrary standards, weighted mathematical averages, and sophisticated instrumentation.
Defining the Atomic Mass Unit
Measuring the mass of a single atom is impractical. To create a workable scale for the atomic world, scientists established a relative unit of measurement. This unit is the unified atomic mass unit, symbolized as “u” or sometimes the Dalton (Da).
The standardization of this unit evolved over time, initially using hydrogen as a reference point. Later, oxygen became the preferred standard because it readily formed compounds with many other elements. This system created confusion, however, as chemists used the average mass of naturally occurring oxygen isotopes, while physicists used the mass of the single oxygen-16 isotope. In 1961, the scientific community agreed on a unified standard to resolve this discrepancy.
The current international standard defines one unified atomic mass unit as exactly one-twelfth of the mass of a single, unbound atom of Carbon-12. Setting the mass of the Carbon-12 isotope at precisely 12 u established a consistent and highly accurate reference point for all other elements. This choice provided a stable and abundant isotope that satisfied both the chemical and physical communities.
Understanding the Weighted Average of Isotopes
The atomic mass value listed on the periodic table is almost never a whole number, despite protons and neutrons each having a mass of approximately one atomic mass unit. This fractional value results from the existence of isotopes. Isotopes are atoms of the same element that have the identical number of protons but differ in the number of neutrons, leading to different masses.
For example, a typical sample of the element chlorine is not composed of atoms all having the same mass. Chlorine primarily exists as two stable isotopes: Chlorine-35 and Chlorine-37. If the atomic mass were a simple arithmetic average of 35 and 37, the result would be 36, but the published atomic mass for chlorine is approximately 35.45 u.
The published atomic mass is a “weighted average” that accounts for the natural abundance of each isotope. Natural abundance refers to the percentage of each isotope found in a naturally occurring sample of the element. Chlorine-35 is far more common, making up about 75.77% of all chlorine atoms, while Chlorine-37 accounts for the remaining 24.23%.
To determine the elemental atomic mass, the precise mass of each isotope is multiplied by its fractional abundance, and the products are summed. Because the lighter Chlorine-35 isotope is significantly more abundant, its mass contributes disproportionately to the average, pulling the final calculated mass value closer to 35 than to 37. This weighted calculation ensures that the mass listed on the periodic table accurately reflects the element’s composition as it is encountered in nature.
Experimental Determination using Mass Spectrometry
The precise masses and natural abundances of individual isotopes, which are the necessary inputs for the weighted average calculation, are determined experimentally using a mass spectrometer. This powerful analytical technique separates atoms based on their mass-to-charge ratio (\(m/z\)) in a four-stage process.
Ionization
The first stage is Ionization, where the sample is vaporized and then bombarded with a stream of high-energy electrons. This collision knocks electrons off the atoms, converting them into positively charged ions, which is necessary for manipulation by electric and magnetic fields. In most cases, only one electron is removed, giving the ions a net charge of +1.
Acceleration
Next is Acceleration, where the positive ions are propelled through a series of plates with progressively decreasing negative electrical potential. This electric field accelerates all the ions to a high, uniform kinetic energy. Because the ions now have the same energy, their velocities differ, with lighter ions traveling faster than heavier ones.
Deflection
The third stage is Deflection, where the accelerated ion beam passes through a strong magnetic field. The magnetic field exerts a force that causes the ions to bend from their straight path. The extent of this deflection is inversely proportional to the ion’s mass-to-charge ratio (\(m/z\)). Lighter ions and ions with a greater charge are deflected more sharply, while heavier ions are deflected less.
Detection
Finally, in the Detection stage, the separated ion beams hit a detector plate. When an ion strikes the plate, it generates a small electrical current. The location where the ion hits determines its mass, and the magnitude of the current is directly proportional to the number of ions hitting that spot, which measures the relative abundance of that isotope. The resulting mass spectrum plots the relative abundance against the mass-to-charge ratio, providing the data needed to calculate the elemental atomic mass.