Understanding Isotopes in Chemistry

Study Notes

Isotopes are variants of a single element that have the same number of protons but differ in the number of neutrons.
The term "isotope" comes from the Greek words isos (same) and topos (place), reflecting their identical position in the Periodic Table.
The history of isotopes began in 1913 when Frederick Soddy described the concept, with the term coined by Margaret Todd in 1921. Soddy received the Nobel Prize in Chemistry for his work in 1921.
Initially, atoms were believed to consist only of protons and electrons, but the differing atomic masses of isotopes suggested the presence of uncharged particles, which later became known as neutrons.
Isotopes are named using the element's name followed by the atomic mass. For example, Carbon-12, Carbon-13, and Carbon-14.
Hydrogen, the most abundant element, has three naturally occurring isotopes: protium (Hydrogen-1), deuterium (Hydrogen-2), and tritium (Hydrogen-3), and four synthetic isotopes.
Stable isotopes have stable nuclei and are not radioactive, while radioactive isotopes have unstable nuclei and undergo radioactive decay.
Radioactivity can be caused by an excessive number of neutrons, leading to the conversion of neutrons to protons and electrons. Alternatively, too many protons can also lead to instability.
Isotopes can be classified based on various criteria, including radioactivity, stability, and origin. The two major types are stable isotopes (non-radioactive) and radioactive isotopes.
Isotopic abundance refers to the percentage of each isotope found in nature and is used to calculate the average atomic mass of an element.
Processes such as radioactive decay and mass fractionation lead to variations in isotopic abundance.
Isotopes share the same chemical properties due to the same number of protons and electrons, except for hydrogen.
Physical properties of isotopes can differ due to the varying number of neutrons.