# Atomic mass

The atomic mass of the atom is expressed in u. Indicates how many times the mass of the atom is greater than 1/12 of the mass of the carbon atom (A = 12).

When we measure a quantity, we compare it with another as a reference. To measure the mass of our body we use the kilogram (kg) as the standard unit. If the person has a mass of 80kg it means that his mass is 80 times greater than the mass of 1kg.

Chemistry, in practice, is not interested in knowing the mass of an isolated atom, but for science it is important to know the mass of atoms compared to the mass of another atom taken as standard. Carbon was then the element that has its standardized mass (A = 12).

The mass of an atom is expressed by employing a very small unit called the atomic mass unit (u). In the past, the acronym u.m.a was used for this unit.

One unit of atomic mass (u) is 1/12 of the mass of a carbon atom (A = 12). This is equivalent to setting the value 12u as the mass of a carbon atom (A = 12).

Atomic Mass is the mass of the atom expressed in u. When the mass of an element X is stated to be 24u, it means that its mass is 24 times greater than the 1/12 mass of the carbon atom (A = 12). In other words, the atomic mass of element X is twice the atomic mass of carbon.

Note the table with some chemical elements and their atomic numbers and atomic masses.

 ELEMENT SYMBOL ATOMIC NUMBER ATOMIC PASTA SULFUR s 16 32,06 OXYGEN O 8 16,00 SODIUM At 11 23,00 ALUMINUM Al 13 26,98154 CALCIUM Here 20 40,08 HELIUM He 2 4,00260 IODINE I 53 126,9045 COPPER Ass 29 63,55

## Atomic mass and its isotopes

The atomic number and mass number are always integers, but with atomic mass this does not happen.

The atomic mass of a chemical element is based on the weighted average of its isotope masses in atomic mass units (u). This means that there are several isotopes in nature and a calculation is made, a weighted average, that takes into account the relative abundances of these isotopes to be used as the atomic mass.

Thus, atomic mass is an average of the many isotopes that exist in nature, taking into account their existing quantity. Example:

In nature there are two types of copper (with different masses).
69.09% copper (A = 63), with atomic mass = 62.93u
30.91% copper (A = 65), with atomic mass = 64.93u

What mass of these coppers is taken as a reference and placed in the periodic table? We must make the weighted average of these isotopes: 