Atomic mass

The atomic mass is the mass of an atom, most often expressed in unified atomic mass units. Atomic mass is sometimes incorrectly used as a synonym for relative atomic mass, average atomic mass and atomic weight; the latter differ subtly in atomic mass. *Atomic mass is defined as the mass of an atom, which can only be of one isotope at a time, and is not a weighted average of the abundances of the isotopes.* For many elements that have one dominant isotope, the actual numerical similarity/difference between the atomic mass of the most common isotope and the relative atomic mass or standard atomic weight may be very small, such that it does not affect many crude calculations, but such an error can be critical when considering individual atoms.

The standard atomic weight refers to the average of the relative atomic masses of an element in the local environment of the Earth's crust and the Earth's atmosphere, as determined by the Commission on IUPAC Atomic Weights and Isotopic Abundances. These values are those that are included in a standard periodic table, and are most often used for ordinary calculations. An uncertainty is included in parentheses that often reflects the natural variability in the isotopic distribution, rather than the uncertainty in the measurement. For synthetic elements, the isotope formed depends on the means of synthesis, so the concept of isotopic abundance naturally doesn't make sense. Consequently, for synthetic elements, the total nucleon count of the most stable isotope (ie, the isotope with the longest half-life) is listed in parentheses in place of the standard atomic weight. Lithium represents a unique case, where the natural abundance of the isotopes has been perturbed by human activities to the point of affecting the uncertainty in its standard atomic weight, even in samples obtained from natural sources, such as rivers.

The relative atomic mass is a synonym for atomic weight and is closely related to the average atomic mass (but not a synonym for atomic mass), the weighted mean of the atomic masses of all the atoms of a chemical element found in a particular sample, weighted by isotopic abundance. This is often used as a synonym for relative atomic weight, and this usage is not incorrect, since that standard atomic weights are relative atomic masses, although it is less specific. Relative atomic mass also refers to non-terrestrial environments and highly specific terrestrial environments that deviate from the mean or have different certainties (number of significant figures) than standard atomic weights.

Relative isotopic mass is the relative mass of a given isotope (more specific, any single nuclide), scaled with carbon-12 as exactly 12. No nuclides other than carbon-12 have exactly an integer number of masses on this scale. This is due to two factors:

1. the different mass of neutrons and protons that act to change the total mass in the núclides with proton/neutron relationships other than the 1:1 ratio of carbon-12; and
2. an exact number will not be found if there is a mass loss/gain different from the nuclear bond energy relative to the average nuclear bond energy of the carbon-12, however, since any mass defect due to the nuclear bond energy is a small fraction (less than 1 %) compared to the mass of a nucleon (even less than the average mass by nucleon in the carbon-12, which is moderate to a strongly attached mass), and given The neutron count can be derived by subtraction of the atomic number.

Mass defects in atomic masses

Link energy by isotopes core.

The amount by which atomic masses deviate from their mass number is as follows: the deviation starts positive at hydrogen-1, decreasing to a minimum at iron-56, iron-58, and nickel-62, then it increases to positive values in the heavier isotopes, as the atomic number increases. This corresponds to the following: nuclear fission in an element heavier than iron produces energy, and fission of anything lighter than iron requires energy. The opposite is true for nuclear fusion reactions: fusion in elements lighter than iron produces energy, and fusion in elements heavier than iron requires energy.

Measuring Atomic Masses

The process that was followed historically to determine the real masses of the atoms of the different elements was similar to that followed in the clips model, initially working with gases and comparing the masses of gases located in containers with the same pressure conditions, volume and temperature: since the masses were different, but there was the same number of particles (according to the model of matter and Avogadro's principle), it was because the particles had different real masses. Currently the direct comparison and measurement of the masses of the atoms is obtained with the use of a mass spectrometer.

Conversion factor between atomic mass unit and grams

The standard scientific unit for handling atoms in macroscopic quantities is the mole, which is arbitrarily defined as the amount of substance that has as many atoms or another unit as there are atoms in 12 grams of carbon of the C-12 isotope. The number of atoms in a mole is called Avogadro's number, whose value is approximately 6.022 x 10²³ mol⁻¹. A mole of a substance always contains exactly the relative atomic mass or molar mass of said substance, expressed in grams; however, this is not true for atomic mass. For example, the standard atomic weight of iron is 55.847 g/mol, and consequently one mole of iron as it is typically found on Earth has a mass of 55.847 grams. The atomic mass of the isotope ⁵⁶Fe is 55.935 u, and one mole of ⁵⁶Fe will theoretically weigh 55.935 g, but it has not been found such pure amounts of ⁵⁶Fe isotope on Earth.

The formula for conversion between atomic mass unit and SI mass in grams for a single atom is:

1u=MuNA=1.660539040(83)× × 10− − 24g=1.660539040(83)× × 10− − 27Kg{displaystyle 1 {rm {u}}={frac {M_{rm {u}}}{N_{rm {A}}}}=1.660,539,040(83)times 10}{-24}{rm {g}}=1.660,539,040(83rs 10^{

where Mu{displaystyle M_{u}} is the constant of molar mass and NA{displaystyle N_{A}} It's Avogadro's number.

Relation between atomic mass and molecular mass

Similar definitions apply to molecules. The molecular mass of a compound can be calculated by adding the atomic-molecular masses of its constituent atoms (nuclides). One can also calculate the undefined molar mass by adding the relative atomic masses of the elements given in the molecular formula. In both cases, the multiplicity of the atoms (the number of times it is present) must be taken into account, generally by multiplying each single mass by its inverse multiplicity.

History

In the history of chemistry, the first scientists to determine atomic weights were John Dalton between 1803 and 1808, and Jöns Jakob Berzelius between 1808 and 1826. Atomic weights were originally defined in relation to the element hydrogen, the most light, taking it as 1, and in 1820, Prout's hypothesis indicated that the atomic masses of all elements should be an integral multiple of the weight of hydrogen. However, Berzelius soon proved that this hypothesis did not always hold, and in some cases, such as chlorine, the atomic weight fell almost exactly between two multiples of the weight of hydrogen. This was later shown to be due to an isotope effect, with the atomic mass of pure isotopes, or nuclides, being a multiple of the mass of hydrogen, within 0.96%.

In the 1860s, Stanislao Cannizzaro refined atomic weights by applying Avogadro's law (at the Karlsruhe Congress of 1860). He formulated a law for determining the atomic weights of elements: different amounts of the same element contained in different molecules are all integer multiples of the atomic weight, and he determined the atomic weights and molecular weights by comparing the vapor density of a set of gases with molecules containing one or more of the chemical element in question.

In the early 20th century, until the 1960s, chemists and physicists used two atomic mass scales. Chemists used a scale such that the natural mixture of oxygen isotopes had an atomic mass of 16, while physicists assigned the same number 16 to the atomic mass of the most common oxygen isotope (containing eight protons and eight neutrons). However, because both oxygen-17 and oxygen-18 are also present in naturally occurring oxygen, this led to 2 different tables of atomic masses.^{[citation needed]} The unified scale, based on carbon-12, ¹²C, fulfilled the physicists' requirement to base the scale on a pure isotope, while being numerically close to the scale of the chemicals.