Atomic Mass Law

The ratio of atomic mass to mass number (number of nucleons) varies from 0.9988381346(51) for 56Fe to 1.007825031898(14) for 1H. Atomic mass (ma) is the mass of a single atom with the unit Da or u (the Dalton). It defines the mass of a particular isotope, which is an input value for determining the relative atomic mass. An example of three isotopes of silicon is given below. The relative isotopic mass (see section below) can be obtained by dividing the atomic mass ma of an isotope by the atomic mass constant mu, resulting in a dimensionless value. Thus, the atomic mass of a carbon-12 atom is by definition 12 Da, but the relative isotopic mass of a carbon-12 atom is simply 12. The sum of the relative isotopic masses of all atoms in a molecule is the relative molecular weight. Reliable values for atomic weights serve a very different purpose when chemical raw materials are bought and sold based on the content of one or more specified constituents. Examples include expensive metal ores such as chromium or tantalum and industrial chemical soda. The content of the specified constituent is determined by quantitative analysis. The calculated value of the material depends on the atomic weights used in the calculations.

In addition to this measurement uncertainty, there are some elements that vary between sources. That is, different sources (seawater, rocks) have a different radioactive history and therefore a different isotopic composition. To reflect this natural variability, IUPAC 2010 decided to list the relative standard atomic masses of 10 elements as an interval rather than a fixed number. [18] Relative isotopic mass is specifically the ratio between the mass of a single atom and the mass of a unit of uniform atomic mass. This value is also relative and therefore dimensionless. In the history of chemistry, the first scientists to determine atomic weights were John Dalton between 1803 and 1805 and Jöns Jakob Berzelius between 1808 and 1826. The atomic weight was originally defined in relation to that of the lightest hydrogen, which was taken as 1.00. Stanislao Cannizzaro refined atomic weights in the 1860s by applying Avogadro`s law. He formulated a law to determine the atomic weights of the elements: The different quantities of the same element contained in different molecules are all integer multiples of the atomic weight and atomic weights and molecular weights determined by comparing the vapor density of a gas collection with molecules containing one or more of the chemical elements concerned [5].

Atomic mass (ma or m) is the mass of an atom. Although the unit of mass SI is the kilogram (symbol: kg), the atomic mass is often expressed in the non-SI color blindness unit (symbol: Da) – unit of equivalent and uniform atomic mass (u). 1 It is defined as 1⁄12 of the mass of a free carbon-12 atom at rest in its ground state. [1] Protons and neutrons in the nucleus make up almost all of the total mass of atoms, with electrons and nuclear binding energy making a small contribution. Thus, the numerical value of the atomic mass, when expressed in Daltons, has almost the same value as the mass number. The conversion between mass in kilograms and mass in Dalton can be done with the atomic mass constant m u = m ( 12 C ) 12 = 1 D a {displaystyle m_{rm {u}}={{m({rm {^{12}C}})} over {12}}=1 {rm {Da}}}. The concept of atomic weight is fundamental to chemistry because most chemical reactions take place in accordance with simple numerical relationships between atoms. Since it is almost always impossible to directly count the atoms involved, chemists measure reactants and products by weighing and come to their conclusions by calculations with atomic weights.

The search for the determination of the atomic weights of the elements occupied the greatest chemists of the 19th and early 20th century. Their meticulous experimental work became the key to chemical science and technology. In addition, the continued use of the term “atomic weight” (for each element) as opposed to “relative atomic mass” has been the subject of considerable controversy since at least the 1960s, mainly due to the technical difference between weight and mass in physics. [4] Nevertheless, both conditions are officially sanctioned by IUPAC. The term “relative atomic masses” now seems to replace “atomic mass” as the preferred term, although the term “standard atomic mass” (as opposed to the more correct “relative standard atomic mass”) continues to be used. The relative isotopic mass is the relative mass of the isotope, scaled with carbon-12 as exactly 12. No other isotope has integer masses of numbers due to the different mass of neutrons and protons, as well as the mass loss/gain of binding energy. However, since the mass defect due to the binding energy is minimal compared to the mass of a nucleon, rounding the atomic mass of an isotope results in the total number of nucleons. The number of neutrons can then be derived by subtracting the atomic number. For carbon, the ratio of mass (in Dalton) to mass number is defined as 1, and after carbon it becomes less than one until a minimum is reached at iron 56 (with only slightly higher values for iron 58 and nickel-62), then it increases to positive values in heavy isotopes, with an increasing atomic number. This corresponds to the fact that nuclear fission in an element heavier than zirconium produces energy, and nuclear fission requires energy in any element lighter than niobium. On the other hand, the nuclear fusion of two atoms of an element lighter than scandium (with the exception of helium) generates energy, while fusion into elements heavier than calcium requires energy.

The fusion of two 4He atoms, resulting in beryllium-8, would require energy, and the beryllium would quickly disintegrate again.