The atom was originally thought of as the smallest particle in the universe and something that can not be split into smaller pieces. Time has shown both of these things to be untrue, but they are a strong basis from which to begin. The atom as we know today is constructed from three different kinds of subatomic particles, protons, neutrons, and electrons. The former two are what give each atom its structure and identity, while the third allows it to react in various ways with other elements and compounds. Protons and neutrons are housed in the nucleus and constitute the main structural elements of the atom while the electrons orbit it. Protons and neutrons together can be referred to as nucleons. Let’s take a closer look at each piece.

Protons

Protons are the positively charged particles of an atom and are located within the nucleus. Each nucleus contains one or more protons. Protons have a weight of 1.6726 x 10^-24g each, or approximately one amu, atomic mass unit. They are denoted by the symbol p or p⁺ and each proton contributes +1e, which is positive 1 to the overall charge of the atom. The number of protons present in an atom are referred to as the atomic number and give an atom its elemental identity. Protons were named and described in detail by Ernest Rutherford in 1920 but were theorized since 1815.

Neutrons

Neutrons are the other particle present within the nucleus of an atom and are slightly heavier than protons but are also classified as one amu. The atomic mass number of an atom is the sum of the nucleons. Neutrons are present to provide stability to the atom and have no net charge associated with them. They are denoted by the symbol n or n⁰. Neutrons were first theorized by Ernest Rutherford in 1920 and experimental confirmation was obtained by James Chadwick in 1932.

Electrons

Electrons are the negatively charged particles that orbit the nucleus in predictable paths known as orbitals. They are by far the smallest subatomic particle and are 1836 times smaller than a proton. Electrons are denoted by the symbol e^– and contribute -1 elementary charge to the atom. Their light weight and distance from the nucleus allow them to be exchanged with other atoms rather easily and it is the electrons that bind to form compounds with other atoms. Electrons are the main components that affect characteristics such as electrical affinity, electrochemistry, magnetism, and thermal conductivity. Electrons were first theorized by Richard Laming in 1838 and were discovered by J.J Thompson in 1897.

Isotopes

An isotope is a variant of a certain element that differs in its number of neutrons. Since the number of protons denotes the chemical identity of an element, a change in the number of neutrons makes no difference in the chemical identity, only the isotopic identity of an atom. Let’s examine chlorine as an example:

Chlorine has 17 protons but has two isotopes. Chlorine-35 and chlorine-37. Chlorine-35 has 17 protons and 18 neutrons, while chlorine-37 has 17 protons and 20 neutrons.

Isotopes have varying abundances, that is to say some are more common than others in nature. Synthetic isotopes do not have abundances as they are not found in nature. Isotopes can also be classified as stable and radioactive. Radioactive isotopes are called radioisotopes and their half-lives can vary greatly, even within the same element.

Atomic Mass and Relative Atomic Mass

The atomic mass of an atom is the mass of an atom, nothing more. This is measured by simply adding up the protons and neutrons of an atom. The unit used is the unified atomic mass units (u). So, for example, an atom of Carbon-12 has 6 protons and 6 neutrons, so its atomic mass is 12u.

This is not to be confused with relative atomic mass, also known as atomic weight, which is found under the chemical symbol on the periodic table.

Relative atomic mass is a unitless quantity that is defined as the ratio of the average mass of atoms of an element in a given sample to the atomic mass constant. The atomic mass constant is defined as 1/12 the mass of a carbon-12 atom, so basically the mass of one carbon-12 nucleon. To think about this more simply let’s define the relative atomic mass as the average weight of the isotopes of an element, as this is what these numbers are used to calculate. For example, let’s calculate the relative atomic mass of chlorine.

There are two isotopes of chlorine, chlorine-35 and chlorine-37. Chlorine-35 has a natural abundance of 75%, which leaves 25% for the abundance of chlorine-37. To calculate the relative atomic mass, we need to find the average of these weights with correlation to their abundances. So:

First, we convert the percentages to decimals, 75% = 0.75 and 25% = 0.25

Next, we multiply the atomic weights of each isotope by their respective percentages and add them together

(35 x 0.75) + (37 x 0.25) = relative atomic mass

Relative atomic mass = 35.5