![]() So, to be very clear, they are not the same. There must be confusion here about molar mass being the same as molecular mass. So, the molar mass (mass in gram of one mole) of glucose is 180.0 g/mol. For example, glucose has a formula mass of 180.0 a.m.u. Its numerical value is equal to the formula mass that is expressed in a.m.u. Molar mass is the mass in grams of one mole of a substance. ![]() The average atomic mass of lithium will be as Lithium atom contains a mixture of 7.5% of 6Li and 92.5% of 7Li with 6.01 a.m.u and 7.02 a.m.u masses. The average atomic mass of carbon will be as Īverage atomic mass = Σ (mass of individual isotope) (Its percentage abundance) Carbon atoms contain a mixture of 98.89% of 12C isotope with a mass of 12.00000 a.m.u, and 1.11% 13C isotope with a mass of 13.00335 a.m.u. The atomic mass of a given element can be determined by obtaining the sum of the product of the masses of isotopes and their percentage abundances. Such information is obtained by spectrometric techniques (mass spectrometer).įor example, Hydrogen has three isotopes ( 1H 1 ≈ 99.972%, 1H 2 ≈ 0.0156%, 1H 3 ≈ 10 -18%), Chlorine has two isotopes (Cl 35= 75%, Cl 37= 25%), etc. The first step in average atomic mass determination is the correct determination of the number of isotopes and their relative abundances of these isotopes. It can be calculated, as explained below. For example, the atomic mass of C is 12.011. This is why atomic masses usually appear in decimals. Therefore, the average atomic masses of atoms are is taken as the average of the masses of radioactive isotopes. This is because most elements occur in nature as a mixture of isotopes. The atomic masses are not the same as the atomic numbers in the periodic table. One atomic mass unit is equal to 1/12 of the mass of the carbon-12 atom. The relative mass of an atom is called atomic mass or atomic weight. ![]() For this reason, chemists use the relative atomic mass scale instead of atomic masses. The mass of individual atoms is very small and cannot be expressed in terms of grams or kilograms. The mass number (number of protons and neutrons) of a sample atom when related to the standard 12C, (1/12 th of the mass) gives the mass of an atom, which we call atomic mass or weight. This is accomplished via spectrometric techniques. The mass of an atom can be found no matter how small it is unless the number of protons and neutrons in its nucleus is known. This standard is met by element carbon, which has an isotopic abundance of almost 98.89% for carbon-12 ( 12C). Now, as there are isotopes in compounds as well, the standard has to be the one with the highest relative isotopic abundance, otherwise, the answers will not be as certain. It has to be related to some known mass which can be kept as a standard. Its mass cannot be measured directly by any known technique. The detailed differences between atomic mass and molar mass are as below.Īn atom is very small, in the order of picometers. Molar mass is the weight of the Avogadro number of particles of a substance which is why its unit becomes, the number of grams of a substance in one mole of a substance (g/mol). Molar mass is a macroscopic (large scale) unit of masses of atoms or molecules. The value of 1 a.m.u is 1.66 x 10 -24 g which is exactly the mass of 1 mole of carbon atoms. If we assume for an instant that C-12 is 12 a.m.u with 6 protons and 6 neutrons, and as a hydrogen atom (with one proton) is 12 times lighter than the C-12 atom, the relation between them is 1/12. For example, the mass of one hydrogen atom (protium) could not be measured directly. That is why scientists have to use relative terms to define the mass of atoms. On the other hand, the molar mass is the mass of the one mole of particles ( atoms or molecules) expressed in grams per mole (g/mol).Ītomic mass is so small that it cannot be measured individually. ![]() One a.m.u is the mass of an element relative to the 1/12 mass of a carbon atom. Atomic mass, also known as atomic weight is the mass of an individual atom expressed in atomic mass units (a.m.u). ![]()
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