Structure of Atom notes

Structure of Atom explains how scientists moved from the idea of an indivisible atom to the modern quantum mechanical picture. The chapter begins with the discovery of electrons, protons and neutrons through discharge tube, canal ray and neutron experiments. It then compares Thomson’s plum pudding model, Rutherford’s nuclear model and their limitations. Bohr’s model introduces fixed orbits, quantized energy and hydrogen spectrum calculations. The quantum mechanical model replaces definite paths with probability-based orbitals and includes de Broglie waves, Heisenberg uncertainty and Schrödinger’s equation. Quantum numbers describe each electron completely, while Aufbau principle, Pauli exclusion principle and Hund’s rule explain electronic configuration. For NEET, this chapter is highly important because it combines conceptual theory, spectra, formulas and configuration-based MCQs.

The discovery of subatomic particles proved that atoms are divisible. Cathode ray discharge tube experiments showed that negatively charged rays travel from cathode to anode, move in straight lines, produce fluorescence and get deflected by electric and magnetic fields. J. J. Thomson identified these rays as electrons and measured the charge-to-mass ratio, e/m. Millikan’s oil drop experiment later helped determine electron charge, allowing calculation of electron mass. Canal rays, observed in modified discharge tubes, were positively charged and led to the identification of the proton. James Chadwick discovered the neutron by bombarding beryllium with alpha particles and observing neutral radiation. These discoveries formed the basis of atomic structure, nuclear charge, isotopes and mass number.

Atomic models developed as experimental evidence improved. Thomson proposed the plum pudding model, where negatively charged electrons were embedded in a uniform positively charged sphere, making the atom neutral. Rutherford tested this model using alpha particles directed at a thin gold foil. Most alpha particles passed straight through, some deflected slightly and very few bounced back. This proved that most of the atom is empty space and almost all positive charge and mass are concentrated in a tiny nucleus. Rutherford’s nuclear model placed electrons around the nucleus, but it could not explain atomic stability because revolving electrons should continuously lose energy and fall into the nucleus. It also failed to explain line spectra. Atomic spectra later supported quantized energy levels.

Bohr model corrected Rutherford’s instability problem for hydrogen-like atoms by introducing quantization. Bohr proposed that electrons revolve only in certain permitted circular orbits called stationary states without radiating energy. Angular momentum is quantized as mvr = nh/2π. Radiation is emitted or absorbed only when an electron jumps between two energy levels, with energy difference ΔE = hν. For hydrogen-like species, orbit radius increases as n²/Z, velocity varies as Z/n and energy is negative, proportional to −Z²/n². The model successfully explains the line spectrum of hydrogen and spectral series such as Lyman, Balmer and Paschen. However, it fails for multi-electron atoms, fine structure, Zeeman effect and the wave nature of electrons.

The quantum mechanical model replaces Bohr’s fixed circular orbits with orbitals, which are regions of high probability of finding electrons. It is based on the dual nature of matter, de Broglie’s hypothesis and Heisenberg’s uncertainty principle. de Broglie proposed that moving particles have wavelength λ = h/mv, explaining why electrons show wave character. Heisenberg stated that position and momentum of a microscopic particle cannot both be measured exactly at the same time. Schrödinger developed a wave equation whose solutions are wave functions, ψ. The square of the wave function, ψ², gives probability density of finding an electron. Thus, modern atomic structure does not describe exact paths but predicts allowed energies, orbitals and electron distribution around the nucleus.

Atomic orbitals are three-dimensional regions around the nucleus where the probability of finding an electron is high. Each electron in an atom is described by four quantum numbers. The principal quantum number, n, gives shell, size and energy level. The azimuthal quantum number, l, gives subshell and orbital shape; l = 0, 1, 2, 3 correspond to s, p, d and f. The magnetic quantum number, m, describes orbital orientation and has values from −l to +l. The spin quantum number, s, describes electron spin as +1/2 or −1/2. Orbital shapes include spherical s orbitals, dumbbell-shaped p orbitals, cloverleaf d orbitals and complex f orbitals. Nodes are zero-probability regions and are important for orbital structure.

Electronic configuration describes the distribution of electrons among orbitals of an atom or ion. Aufbau principle states that electrons fill lower-energy orbitals before higher-energy orbitals, generally guided by the n + l rule. Pauli exclusion principle states that an orbital can contain maximum two electrons and they must have opposite spins. Hund’s rule states that electrons occupy degenerate orbitals singly with parallel spins before pairing. The filling order explains configurations such as 1s² 2s² 2p⁶. Some atoms, especially chromium and copper, show exceptions because half-filled and fully-filled subshells are extra stable due to symmetry and exchange energy. Ions are formed by adding or removing electrons; for transition metal cations, electrons are removed from ns before (n − 1)d.