Atomic mass and relative atomic mass both are same only then why should we calculate these separately
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because atomic mass meansThe atomic mass (ma) is the mass of an atom. Its unit is the unified atomic mass units(symbol: u, or Da) where 1 unified atomic mass unit is defined as 1⁄12 of the mass of a single carbon-12 atom, at rest.[1] For atoms, the protons and neutrons of the nucleus account for almost all of the mass, and the atomic mass measured in u has nearly the same value as the mass number.
When divided by unified atomic mass units or daltons to form a pure numeric ratio, the atomic mass of an atom becomes a dimensionless value called the relative isotopic mass (see section below). Thus, the atomic mass of a carbon-12 atom is 12 u or 12 daltons (Da), but the relative isotopic mass of a carbon-12 atom is simply 12.
The atomic mass or relative isotopic mass refers to the mass of a single particle, and therefore is tied to a certain specific isotopeof an element. The dimensionless standard atomic weight instead refers to the average (mathematical mean) of atomic mass values of a typical naturally-occurring mixture of isotopes for a sample of an element. Atomic mass values are thus commonly reported to many more significant figures than atomic weights. Standard atomic weight is related to atomic mass by the abundance ranking of isotopes for each element. It is usually aboutthe same value as the atomic mass of the most abundant isotope, other than what looks like (but is not actually) a rounding difference.
The atomic mass of atoms, ions, or atomic nuclei is slightly less than the sum of the masses of their constituent protons, neutrons, and electrons, due to binding energy mass loss (as per E=mc2).[2]
while as relative atomic mass
meansRelative atomic mass (symbol: Ar) or atomic weight is a dimensionless physical quantitydefined as the ratio of the average mass of atoms of a chemical element in a given sample to one unified atomic mass unit. The unified atomic mass unit (symbol: u or Da) is defined as being 1⁄12 of the atomic mass of a carbon-12 atom.[1][2] Since both values in the ratio are expressed in the same unit (u), the resulting value is dimensionless; hence the value is said to be relative.
For a single given sample, the relative atomic mass of a given element is the weighted arithmetic mean of the masses of the individual atoms (including their isotopes) that are present in the sample. This quantity can vary substantially between samples because the sample's origin (and therefore its radioactive history or diffusion history) may have produced unique combinations of isotopic abundances. For example, due to a different mixture of stable carbon-12 and carbon-13 isotopes, a sample of elemental carbon from volcanic methane will have a different relative atomic mass than one collected from plant or animal tissues.
The more common but distinct quantity known as standard atomic weight is an application of the relative atomic mass values obtained from multiple different samples. It is sometimes interpreted as the expected rangeof the relative atomic mass values for the atoms of a given element from all terrestrial sources, with the various sources being taken from Earth.[3] "Atomic weight" is often loosely and incorrectly used as a synonym for standard atomic weight (incorrectly because standard atomic weights are not from a single sample). Standard atomic weight is nevertheless the most widely published variant of relative atomic mass.
Additionally, the continued use of the term "atomic weight" (for any element) as opposed to "relative atomic mass" has attracted considerable controversy since at least the 1960s, mainly due to the technical difference between weight and mass in physics.[4] Still, both terms are officially sanctioned by the IUPAC. The term "relative atomic mass" now seems to be replacing "atomic weight" as the preferred term, although the term "standardatomic weight" (as opposed to the more correct "standard relative atomic mass") continues to be used.
When divided by unified atomic mass units or daltons to form a pure numeric ratio, the atomic mass of an atom becomes a dimensionless value called the relative isotopic mass (see section below). Thus, the atomic mass of a carbon-12 atom is 12 u or 12 daltons (Da), but the relative isotopic mass of a carbon-12 atom is simply 12.
The atomic mass or relative isotopic mass refers to the mass of a single particle, and therefore is tied to a certain specific isotopeof an element. The dimensionless standard atomic weight instead refers to the average (mathematical mean) of atomic mass values of a typical naturally-occurring mixture of isotopes for a sample of an element. Atomic mass values are thus commonly reported to many more significant figures than atomic weights. Standard atomic weight is related to atomic mass by the abundance ranking of isotopes for each element. It is usually aboutthe same value as the atomic mass of the most abundant isotope, other than what looks like (but is not actually) a rounding difference.
The atomic mass of atoms, ions, or atomic nuclei is slightly less than the sum of the masses of their constituent protons, neutrons, and electrons, due to binding energy mass loss (as per E=mc2).[2]
while as relative atomic mass
meansRelative atomic mass (symbol: Ar) or atomic weight is a dimensionless physical quantitydefined as the ratio of the average mass of atoms of a chemical element in a given sample to one unified atomic mass unit. The unified atomic mass unit (symbol: u or Da) is defined as being 1⁄12 of the atomic mass of a carbon-12 atom.[1][2] Since both values in the ratio are expressed in the same unit (u), the resulting value is dimensionless; hence the value is said to be relative.
For a single given sample, the relative atomic mass of a given element is the weighted arithmetic mean of the masses of the individual atoms (including their isotopes) that are present in the sample. This quantity can vary substantially between samples because the sample's origin (and therefore its radioactive history or diffusion history) may have produced unique combinations of isotopic abundances. For example, due to a different mixture of stable carbon-12 and carbon-13 isotopes, a sample of elemental carbon from volcanic methane will have a different relative atomic mass than one collected from plant or animal tissues.
The more common but distinct quantity known as standard atomic weight is an application of the relative atomic mass values obtained from multiple different samples. It is sometimes interpreted as the expected rangeof the relative atomic mass values for the atoms of a given element from all terrestrial sources, with the various sources being taken from Earth.[3] "Atomic weight" is often loosely and incorrectly used as a synonym for standard atomic weight (incorrectly because standard atomic weights are not from a single sample). Standard atomic weight is nevertheless the most widely published variant of relative atomic mass.
Additionally, the continued use of the term "atomic weight" (for any element) as opposed to "relative atomic mass" has attracted considerable controversy since at least the 1960s, mainly due to the technical difference between weight and mass in physics.[4] Still, both terms are officially sanctioned by the IUPAC. The term "relative atomic mass" now seems to be replacing "atomic weight" as the preferred term, although the term "standardatomic weight" (as opposed to the more correct "standard relative atomic mass") continues to be used.
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