Ch 4Structure of Atom — Class 11 Chemistry, concept-first.

Dalton's atomic theory successfully explained the laws of chemical combination, but it treated the atom as an indivisible, hard sphere and could not explain everyday observations such as glass or ebon…

The electron was the first subatomic particle to be identified, by J. J. Thomson in 1897, through his experiments with a cathode-ray discharge tube — a sealed glass tube held at very low pressure with…

In 1911, Ernest Rutherford investigated the structure of the atom by firing a narrow beam of positively-charged α-particles at an extremely thin sheet of gold foil, with a fluorescent screen arranged…

Rutherford's nuclear model left a puzzle: the measured mass of most atoms was roughly double what their proton count alone could account for, implying the nucleus must also contain some additional, el…

The number of protons in an atom's nucleus fixes the chemical identity of the element and is called the atomic number, symbol Z.

Some elements occur in nature as a single kind of nuclide — fluorine, for instance, occurs essentially only as .

Isobars are atoms that belong to different elements — that is, they have different atomic numbers Z — but happen to share the same mass number A.

Isotones are atoms of different elements whose nuclei happen to contain exactly the same number of neutrons, even though their atomic numbers and mass numbers are both different.

Rutherford's nuclear model pictured the atom rather like a miniature solar system, with the massive, positively-charged nucleus playing the role of the sun and the much lighter electrons orbiting it l…

By the time different atomic models were being debated, two separate lines of experimental and theoretical work — results that seemed, at first, unrelated to atomic structure — turned out to be exactl…

By the start of the 20th century, physicists faced a genuine dilemma about the nature of light. Phenomena such as diffraction (light bending around obstacles/through slits) and interference (light wav…

An electromagnetic wave can be described using a handful of standard parameters. Wavelength, symbol (the Greek letter lambda), is the distance between two consecutive crests (or two consecutive trough…

In 1900, Max Planck put forward his quantum theory to explain the spectrum of radiation given off by a heated ('black') body, a phenomenon the wave picture of light could not account for.

When a substance absorbs energy — whether from incoming radiation or from heating — its atoms, molecules or ions are said to become 'excited'.

Niels Bohr's 1913 model for the hydrogen atom rests on four postulates. First, the electron can move around the nucleus only in one of several fixed circular paths of definite radius and definite ener…

Solving out Bohr's postulates for the hydrogen atom gives several concrete, testable results. The allowed stationary states are labelled by the positive integers , called principal quantum numbers.

Bohr's theory does not just describe individual stationary states — it explains WHY hydrogen's emission spectrum shows exactly the lines it does.

Despite successfully explaining hydrogen's spectrum, Bohr's model has several well-documented limitations.

Bohr's model failed for a reason deeper than any single missing detail: it treats the electron as an ordinary classical particle travelling along a precisely-defined circular path, with both its posit…

Quantum mechanics — the branch of science built to satisfy both de Broglie's matter-wave duality and Heisenberg's uncertainty principle — was developed in 1926, independently, by Werner Heisenberg (st…

When the Schrödinger equation is solved specifically for the hydrogen atom, the outcome is a set of possible electron energies (E) together with their corresponding wave functions (), and — as a natur…

Because many different wave functions solve the Schrödinger equation for a given atom, an atom possesses many atomic orbitals, and these orbitals — really just names for particular solutions — form th…

The probability of finding an electron at a given point in an atom is proportional to at that point — a quantity Max Born interpreted as the probability DENSITY of the electron at that point.

In hydrogen and in hydrogen-like species (one-electron systems), the ONLY interaction present is the attraction between the single electron and the nucleus, so orbital energy depends purely on the pri…

'Aufbau' is German for 'building up', and the Aufbau principle describes how the electrons of a ground-state atom are filled into its orbitals, one at a time, building up the atom's full electronic co…

Wolfgang Pauli stated, in 1926, that no two electrons in the same atom can ever have an identical set of all four quantum numbers.

When electrons are being filled into a set of orbitals that all belong to the same subshell — and are therefore all equal in energy (degenerate) — Hund's rule of maximum multiplicity governs the order…

The electronic configuration of an atom is simply the distribution of its electrons among its available orbitals in the ground state, built up by applying the Aufbau principle (energy order, Pauli's p…

Two elements among the first thirty break the straightforward Aufbau prediction: chromium (Z=24) and copper (Z=29).

Atoms and ions that happen to contain exactly the same TOTAL number of electrons are called isoelectronic species, and — because electronic configuration is entirely determined by electron count for a…