Structure of the Atom - A Short Historical Journey Through Atomic Models

Charged Particles in Matter: Atoms consist of three subatomic particles: protons ($p^+$), electrons ($e^-$), and neutrons ($n^0$).

Thomson’s Model of an Atom: Often called the 'Plum Pudding' or 'Watermelon' model. He proposed that an atom consists of a positively charged sphere and the electrons are embedded in it. The total positive charge equals the total negative charge, making the atom electrically neutral.

Rutherford’s Gold Foil Experiment: Rutherford discovered the nucleus. He concluded that most of the space inside the atom is empty, there is a small, dense, positively charged center called the nucleus, and nearly all the mass of the atom resides in the nucleus.

Bohr’s Model of the Atom: To overcome the limitations of Rutherford's model, Bohr proposed that electrons revolve in discrete orbits called 'energy levels' or 'shells' ($K, L, M, N...$). Electrons do not radiate energy while revolving in these discrete orbits.

Distribution of Electrons: The Bohr-Bury scheme states that the maximum number of electrons in a shell is given by $2n^2$, where $n$ is the orbit number. The outermost shell can hold a maximum of $8$ electrons.

Valency: The combining capacity of an atom. It is determined by the number of valence electrons (electrons in the outermost shell).

Atomic Number ($Z$): The total number of protons present in the nucleus of an atom. For a neutral atom, number of protons = number of electrons.

Mass Number ($A$): The sum of the total number of protons and neutrons present in the nucleus of an atom. Protons and neutrons are collectively called nucleons.

Isotopes: Atoms of the same element having the same atomic number ($Z$) but different mass numbers ($A$). Examples include Protium ($^1_1H$), Deuterium ($^2_1H$), and Tritium ($^3_1H$).

Isobars: Atoms of different elements with different atomic numbers, which have the same mass number ($A$). Example: Calcium ($^{40}_{20}Ca$) and Argon ($^{40}_{18}Ar$).