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Atomic mass | Atomic mass definition, Units, & Facts

Atomic Mass

Atomic mass

Atomic mass refers to the mass of an atom (ma or m). Although the unified atomic mass unit (symbol: Da) is widely used to express atomic mass, the kilogramme (symbol: kg) is the SI measure of mass (u). The mass of an unbound carbon-12 atom in its ground state is 112 Da. Protons and neutrons make up almost all of an atom's mass, with electrons and nuclear binding energy making up the majority of the remaining mass. As a result, the mass number and the atomic mass have values that are highly comparable. The atomic mass can be used to convert between mass in kilogrammes and mass in daltons.

Atomic Mass  Formula


Atomic Mass Formula


By dividing the atomic mass ma of an isotope by the atomic mass constant mu, one can derive the relative isotopic mass, which is a dimensionless number. Consequently, a carbon-12 atom's relative isotopic mass is 12 whereas its atomic mass is 12 Da by definition. The relative molecular mass is the total of all the atoms' respective isotopic masses.

A specific isotope of an element is described by its relative isotopic mass and atomic mass. The elemental atomic mass, which is the average (mean) atomic mass of an element, weighted by the abundance of the isotopes, is useful since things are typically not isotopically pure. The weighted mean relative isotopic mass of a (average naturally occurring) mixture of isotopes is what is known as the dimensionless (standard) atomic weight.

Due to binding energy mass loss (per E = mc2), the atomic mass of atoms, ions, or atomic nuclei is somewhat less than the total of the masses of its constituent protons, neutrons, and electrons.

Atomic mass of hydrogen is 1.00784 u & Atomic mass is oxygen is 8

Relative Isotopic Mass

Not to be confused with the averaged quantity atomic weight (see above), which is an average of values for many atoms in a particular sample of a chemical element, is relative isotopic mass, a feature of a single atom.

While relative isotopic mass has no dimensions and no units, atomic mass is an absolute quantity. Relative isotopic mass refers to this scaling in relation to carbon-12, as shown by the word "relative" in the name. This loss of units is caused by the employment of a scaling ratio with respect to a carbon-12 reference.

The mass of a particular isotope (more specifically, any single nuclide) multiplied by the mass of carbon-12, where the latter must be established experimentally, yields the relative isotopic mass. The mass of an isotope or nuclide relative to 1/12 of the mass of a carbon-12 atom is equivalently known as the relative isotopic mass of that isotope or nuclide.

For instance, a carbon-12 atom's relative isotopic mass is precisely 12. For reference, a carbon-12 atom has an exact mass of 12 daltons. The atomic mass of a carbon-12 atom can also be stated in any other mass units, such as kg, where the value is 1.99264687992(60)10⁻²⁶ kg.

The relative isotopic mass numbers of nuclides other than carbon-12 are not whole numbers, but they are always close to whole numbers, just as the related atomic mass when stated in daltons.

Relationship between atomic and molecular masses

Molecules have definitions that are similar. By summing the atomic masses—not the conventional atomic weights—of a compound's constituent atoms, one can determine the molecular mass of the complex. On the other hand, the conventional atomic weights are commonly used to calculate the molar mass (not the atomic or nuclide masses). As a result, molar mass and molecular mass have somewhat different numerical values and refer to distinct ideas. The total mass of a molecule's individual atomic masses is known as the molecule's molecular mass. Molar mass is the average of all the masses of the individual molecules that make up an ensemble that is chemically pure but isotopically heterogeneous. In both scenarios, it is necessary to account for the multiplicity of the atoms (the number of times it occurs), which is often done by multiplying each unique mass by the multiplicity.

History

John Dalton, Thomas Thomson, and Jöns Jakob Berzelius were the first researchers to calculate the relative atomic masses of atoms between 1803 and 1805, and between 1808 and 1826. According to Prout's idea, which was put forth in the 1820s, all atomic masses would turn out to be precise multiples of hydrogen, relative atomic mass (also known as atomic weight) was originally defined relative to that of the lightest element, hydrogen, which was taken to be 1.00. However, Berzelius quickly demonstrated that this wasn't even close to being accurate; in fact, for some elements, like chlorine, the relative atomic mass, at 35.5, is almost exactly halfway between two integral multiples of hydrogen. Later, it was discovered that this was primarily caused by a mixture of isotopes and that, to within about 1%, the atomic masses of pure isotopes, or nuclides, are multiples of the mass of hydrogen.

By using Avogadro's law, Stanislao Cannizzaro improved relative atomic masses in the 1860s (notably at the Karlsruhe Congress of 1860). By comparing the vapour density of a group of gases with molecules containing one or more of the chemical element in question, he developed a law to determine the relative atomic masses of elements: the different amounts of the same element contained in different molecules are all whole multiples of the atomic weight.

There were two distinct atomic-mass scales used by chemists and physicists in the 20th century, up to the 1960s. The natural mixture of oxygen isotopes had an atomic mass of 16, according to the "atomic mass unit" (amu) scale used by chemists, although the same number 16 was only given to the atomic mass of the most prevalent oxygen isotope by physicists (16O, containing eight protons and eight neutrons). However, as natural oxygen also contains oxygen-17 and oxygen-18, two distinct tables of atomic mass were required. The carbon-12, or 12C,-based unified scale satisfied the physicists' requirement that the scale be based on a pure isotope and was numerically comparable to the scale used by chemists. As the "unified atomic mass unit," this was chosen. The dalton and symbol "Da" are the current International System of Units (SI) major recommendations for this unit's designation. The accepted names and symbols for the same unit are "unified atomic mass unit" and "u."

In the majority of contemporary usage, the word "atomic weight" is gradually being phased out in favour of the term "relative atomic mass." This change in terminology, which dates back to the 1960s, has been the subject of intense discussion in the scientific community. It was brought about by the adoption of the unified atomic mass unit and the recognition that the term "weight" wasn't quite accurate. The main justification for keeping the term "atomic weight" was that it was well understood by experts in the field, that it was already in use (in the sense that it is currently defined), and that "relative atomic mass" could be mistaken for relative isotopic mass (the mass of a single atom of a given nuclide expressed dimensionlessly relative to 1/12 of the mass of carbon-12; see section above).

A secondary synonym for atomic weight called "relative atomic mass" was introduced in 1979 as a compromise. Twenty years later, the term "relative atomic mass" has replaced these synonyms as the preferred word.

The word "standard atomic weights," which refers to the standardised expectation atomic weights of various samples, has not been altered because the term "standard relative atomic mass" would have been created by simply replacing "atomic weight" with "relative atomic mass."

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