The atomic weight is the average mass of the atoms of a certain chemical element. It is known and used interchangeably as atomic mass, although literally the meaning of each one is different.
The term 'weight' in physics implies the force exerted on a gravitational field expressed in units of force like the newton. However, since 1908 the term atomic weight has been used, which today is better known as relative atomic mass; that is, they are synonymous.
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Atoms are so small, abundant and different even for the same element, that it is not an easy task to assign them a physical magnitude such as mass. Precisely over time the choice of the unit that represents the weight or atomic mass of a chemical element has varied.
Initially, the mass of the smallest atom, which is the hydrogen atom (H), was chosen as the atomic mass unit. It was later changed to the atomic mass unit of natural oxygen 1/16, and then its lighter isotope 16 O was preferred.
Since 1961, due to the great importance of the carbon atom (C), it was chosen to refer the atomic weight to its isotope C-12. Furthermore, the C atom is the central or key chemical element of organic chemistry and biochemistry.
What is the atomic weight?
The atomic weight (AP) is the average weight of the masses of the natural isotopes that make up a chemical element. The term refers to the relative atomic mass possessed by the atoms of each of the chemical elements.
As mentioned in the opening section, the term atomic weight is traditionally used, but it is actually atomic mass. Since 1961, based on the Carbon-12 atom, its value of 12 was adopted for the scale of relative atomic weights.
But what is the atomic mass then? It is the sum of protons and neutrons that the atom has, the mass contributed by the electrons being negligible. The atomic mass of hydrogen (H), for example, is 1.00974 Da, and that of Magnesium (Mg) is 24.3050 Da.
Comparing, this means that Mg atoms are heavier than H atoms: 24 times more exactly. When the values of the weight or atomic mass of some chemical element are needed, it can be obtained by consulting the periodic table.
Units
One of the first units of atomic weight, amu, was expressed as 1/16 (0.0625) of the weight of an oxygen atom.
This unit changed with the discovery of the existence of the natural isotopes of an element from 1912; therefore, isotopes could no longer be ignored.
Currently, the standard unit of atomic mass or dalton is 1/12 of the atom weight of the 12 C isotope. This is more stable and abundant than 13 C and 14 C.
A standardized atomic mass unit is the mass of a nucleon (a proton or a neutron) and is equal to 1 g / mole. This unification or standardization was carried out with a C-12 atom to which 12 atomic mass units are assigned.
And so the relative atomic weight or atomic mass can currently be expressed in grams per one mole of atoms.
How is the atomic weight calculated?
To determine the atomic weight, first the atomic mass of the isotope must be calculated, which is the sum of the number of protons and neutrons that a certain atom has.
The number of electrons it has is not taken into account, since its mass is negligible compared to that of neutrons and protons.
The same is done with each isotope of the same element. Then, knowing their natural abundance, a weighted average atomic mass of all isotopes is calculated by adding the product m ∙ A (m = atomic mass, and A the abundance divided by 100).
For example, suppose you have a cluster of iron atoms where 93% of them are 56 Fe, while 5% are 54 Fe and the remaining 2% 57 Fe. The atomic masses are already marked in the upper left corners of chemical symbols. Calculating then:
56 (0.93) + 54 (0.05) + 57 (0.02) = 55.92 g / mol Fe atoms
In that cluster, iron has an atomic weight of 55.92. But what about the rest of the entire planet Earth or the rest of the Universe? In the cluster there are only three isotopes, whose abundances change if the Earth is taken into account, where there will be more isotopes available and the calculations become more complicated.
Considerations
To calculate the atomic weight of the elements reported in the periodic table, the following must be taken into account:
-The isotopes that exist in nature of the same chemical element. Atoms of the same chemical element that have different numbers of neutrons are the isotopes of that chemical element.
-In the samples obtained from each isotope, the atomic mass of each of them is taken into account.
-The relative abundance of each of the isotopes for a certain element in the samples found in nature is also important.
-The value of the atomic weight of a single atom alone or present in a natural sample of the element can be found. Or of a group of atoms in the case of isotopes of the same element, determining the standard or average atomic weight.
-To determine the standard atomic weight of chemical elements, one or more isotopes of the same element were considered.
-There are some chemical elements such as Francium (Fr) that do not have stable isotopes and do not yet have a standardized atomic weight.
Examples
By consulting the periodic table of chemical elements, the atomic weight of a chemical element can be found; that is, those that have been calculated considering all the stable isotopes (and for that reason they usually have many decimal places).
There it is observed that the atomic number of hydrogen (H) is equal to 1, equal to its number of protons. The atomic weight of H is the least of all the elements, having a value of 1.00794 u ± 0.00001 u.
For boron, its atomic weight was determined based on two isotopes obtained in nature and its value ranges from 10,806 to 10,821.
There is no standard atomic weight in the case of non-natural or synthetic elements that do not have isotopes in nature; such as the aforementioned case of francium (Fr), polonium (Po), radon (Ra), among other chemical elements.
In these cases, the atomic weight is limited to the sum of the number of protons and neutrons of this element.
The atomic weight value is reported in parentheses, which means that it is not a standardized atomic weight. Even the value of the standard atomic weight is liable to change if more isotopes of a certain element are discovered.
References
- Helmenstine, Anne Marie, Ph.D. (June 22, 2018). Atomic Weight Definition. Recovered from: thoughtco.com
- Jimenez, V. and Macarulla, J. (1984). Physiological Physicochemistry. (6 ta. Ed). Madrid: Interamericana
- Whitten, K., Davis, R., Peck M. and Stanley, G. (2008). Chemistry. (8 ava. Ed). CENGAGE Learning: Mexico.
- Wikipedia. (2018). Standard atomic weight. Recovered from: en.wikipedia.org
- Prof. N. De Leon. (sf). Atomic Weights. Recovered from: iun.edu