Atomic Mass Unit – Definition, Examples and History.

Dec 29, 2022

The atomic mass of an element determines its atomic mass unit. Both atomic and molecular masses are expressed by the AMU. It is also called the unified atomic mass unit or the Dalton. The atomic mass unit definition is that it is one-twelfth of the mass of a carbon-12 atom, which is the most abundant isotope of carbon found in nature. It accounts for more than 98% of carbon found in nature, having an AMU of 12.

Atomic Mass and Atomic Mass Unit

Atoms are the constituents of every element in the periodic table, each of which has a unique atomic number and mass. The number of protons in the nucleus of an atom determines its atomic number, while atomic mass is defined as the combination of the total number of neutrons and protons. It is expressed in terms of Dalton or AMU, where one AMU is the average of the neutron and the proton rest mass. Following is the way to express it:

one atomic mass unit is equal to = 1.67376 x 10^-27 kilograms = 1.67376 x 10^-24 grams

All the atomic mass calculations are performed while considering carbon-12 as the reference. Therefore, any element’s isotope’s mass is expressed as the standard of the atomic mass unit of the carbon 12 isotope.

There are six protons and neutrons in the nucleus of a carbon 12 atom to give an atomic mass of 12 AMU. As electrons have low mass, they presumably have a negligible effect. The nucleus, consequently, constitutes the entire weight of any element’s atom, meaning that the mass of a single proton or neutron is approximately one atomic mass unit.

The word approximately is important here as the masses of the individual atoms in elements apart from carbon are not whole numbers. The reason for this is that the mass is influenced by the interactions between different particles in the nucleus. Although the mass of electrons is negligible, it is considered when the mass of one atom is calculated.

History of Atomic Mass Unit

John Dalton 1802 suggested a method to define the atomic mass units in terms of protium (hydrogen-1). Later, another scientist named Wilhelm Ostwald proposed that we can express the relative mass of an atom as 1/16th of the oxygen mass. However, when the discovery of isotopes like the isotopes of oxygen occurred, confusion was created about how to express other elements’ relative atomic mass.

As a result of all this, the method to define atomic mass unit diverged, with several scientists preferring to express it on the basis of natural oxygen. In contrast, some others tried to express it based on the oxygen-16 isotope of oxygen. The latter was popularly used to express atomic mass units until 1961 when a way to eliminate the confusion came into play.

Carbon-12 was suggested to be used as the basis to express atomic mass units instead of oxygen-16. This system was given the symbol Da and u. The symbol AMU however, did not entirely cease to exist, with scientists continuing to use it even after shifting to carbon-12.

Nowadays, atomic mass units are expressed with all three symbols— u, AMU, and Da.

Therefore, one atomic mass unit is equal to = 1 u = 1 Da

Unified Atomic Mass Unit

What is the unified atomic mass unit definition? A unified atomic mass unit is commonly considered a synonym for the atomic mass unit. As a physical constant, it is accepted for the SI measurement system. Although these days the term AMU is used more commonly, it is the same as unified AMU.

Avogadro’s number NA establishes the relationship between the unified atomic mass unit and the SI of mass. The mass of a carbon-12 atom at rest as well as in-ground state is equal to 0.012 kg according to the definition of NA.

Therefore, one atomic mass unit is equal to AMU = 1.6604 x 10^-27 kg

Differentiating Isotopes with Atomic Mass Unit

Isotopes can be effectively differentiated with the help of atomic mass units by expressing their relative masses. It involves several elements that have similar atomic numbers, but their atomic numbers are different. This is because the number of protons in them is the same, while the number of neutrons is different.

Example 1

The AMU of a uranium-235 atom is approximately 235. Another isotope, uranium-238, which has an atomic mass unit value of 238, is slightly bigger, having a larger mass. The difference in the atomic mass is due to I-238, the most abundant isotope of uranium found in nature and has three excess neutrons in its atoms compared to U-235, which is employed in nuclear reactions to produce nuclear energy through nuclear fission. It is also one of the main components of atomic bombs.

Example 2

Isotope	Number of neutrons	Number of protons	Number of electrons	Atomic mass unit ( neutrons + protons)
Carbon-12	6	6	6	12
Carbon-13	7	6	6	13
Carbon-14	8	6	6	14

What is Atomic Weight?

Atomic weight is also referred to as the relative atomic mass unit value. It is the ratio of the average mass of the atoms of an element compared to some standard. Although there is interchangeable use of the terms atomic mass and atomic weight, their meanings differ. Atomic weight refers to the force that a gravitational field exerts, while atomic mass doesn’t. More specifically, an element’s atomic weight is considered as the average of the atomic masses of its isotopes.

Example

Carbon is usually found in two isotopic forms— carbon 12 and carbon 13. The atomic mass unit of carbon 12 is 12, while that of carbon 13 is 13. The availability of the carbon 12 isotope in nature is 98.89%, while that of carbon 13 is 1.11%. The average atomic mass of carbon 13 and carbon 12 is (98.89 / 100) x 12) + ((1.11 / 100) x 13) = 12.011 AMU

The atomic weight of carbon = 12.011 AMU

The example given above indicates how an element’s atomic weight varies from its isotopes’ atomic mass. This is the reason why atomic weight is not the same as atomic mass. Rather, calling it relative atomic mass is more accurate.

Numerical relationships between atoms influence most chemical reactions; therefore, atomic weight is a basic concept in Chemistry. However, when reactants and products are measured by chemists, individual atoms are not counted. Rather, atomic weights are calculated to guide decisions.

Below is a table showing the atomic mass:

Elements	Atomic mass
Helium	4.00260
Boron	10.81
Nitrogen	14.0067
Fluorine	18.998403
Sodium	22.98977
Aluminium	26.98154
Phosphorus	30.97376
Chlorine	35.453
Argon	39.948
Scandium	44.9559
Vanadium	50.9415
Chromium	51.996
Iron	55.847
Cobalt	58.9332
Zinc	65.38
Germanium	72.59
Selenium	78.96
Krypton	83.80
Strontium	87.62

When paired with the mole concept, the concept of atomic mass is a vital part of Chemistry. When an element’s atomic mass unit is quantified in AMU (such as the atomic mass unit of carbon), it has the same value as the mass of one mole of an element in grams. Therefore, as the iron atomic mass is 55.847 amu, its one mole would weigh 55.847 grams.

The ionic compounds and molecules work on the same concept. The atomic mass of one formula unit of NaCl is about 58.44 amu. Therefore, a mole of NaCl would weigh around 58.44 grams. The weight of one water molecule would be 18.02 amu, and that of a mole of water molecules would be around 18.02 grams.

Conclusion

The concept of atomic mass, as well as an atomic mass unit, is an integral part of both Physics and Chemistry. Therefore, having a fundamental understanding of their concepts is important. Hope this topic helped you gain knowledge on different topics of the atomic mass unit, such as unified atomic mass, differentiating isotopes with the atomic mass unit, how to define atomic mass unit and the history of the atomic mass unit.

Frequently Asking Questions

Q1) Why is Atomic Mass Needed?

A) The atomic mass of an atom of a molecule is employed to determine the average mass of molecules and elements and for solving stoichiometry problems.

Q2) Can Atomic Mass Change?

A) Although the atomic number does not change, the atomic mass of an element varies among its different isotopes. As atomic mass is the sum of the neutrons and protons present in a nucleus, there are more neutrons in an isotope than protons, making the element mass larger than the stable nuclei.