Physical Chemistry - Structure of Atom

Explore Dalton's atomic theory and its four postulates on indivisible atoms, identical atoms within an element, simple ratios, and the conservation of atoms forming compounds.

Explore how the cathode ray tube experiment revealed that gases conduct electricity under low pressure and high voltage, leading to the discovery of cathode rays.

Explore the properties of cathode rays, including straight-line travel, deflection by electric and magnetic fields, and the negatively charged electrons revealed by fluorescence and e/m ratio.

Explore how Goldstein's 1886 discovery of anode rays emerged from a perforated cathode in a high-voltage, vacuumed tube, revealing canal rays and protons.

Explore the properties of anode rays: they are positively charged and deflected by electric fields, with E/M smaller than cathode rays, and gas identity determines E/M, hydrogen having the maximum.

Coulomb's law determines the force between charges using their magnitudes and distance, not charge sign. Like charges repel and opposite charges attract; protons and electrons have equal magnitude, opposite signs.

Compute the potential energy between two charges with U = (1/4πε0)(Q1Q2)/r, and relate it to energy via E = QV, including conversions between angstroms and meters.

Examine Thomson's atomic model, the plum pudding concept with electrons embedded in a positively charged sphere, and the model's limitations in explaining the gold foil experiment.

Bohr's postulates explain how electrons orbit the positively charged nucleus with quantized angular momentum and stationary orbits, using Planck's constant. When absorbing external energy, electrons excite to higher states and emit photons as they return to the ground state, with photon energy equal to the energy difference.

Explore the hydrogen spectrum via the Bohr model, deriving energy levels and the Rydberg formula for Balmer, Lyman, and Paschen series and their visible, ultraviolet, and infrared emissions.

Bohr's equation explains electron energy transitions in stationary revolving orbits and derives hydrogen's spectrum via 1/λ = Rydberg constant (1/n1^2 - 1/n2^2).

Explore Bohr's quantisation condition that restricts electron orbits by angular momentum, explains hydrogen emission spectra, and links wave-particle concepts through de Broglie's wavelength.

Explore the limitations of Bohr's model and de Broglie's explanation for atomic structure, including hydrogen line spectra, orbit stability, and wave-particle duality.

Explain the drawbacks of Bohr's atomic model, including limited applicability to multi-electron atoms, elliptical orbits, unresolved spectral line intensities, a fixed nucleus, and inability to explain the Zeeman effect.

Rutherford's model places a central nucleus of protons and neutrons, with electrons orbiting like a solar system, based on the alpha particle scattering experiment; it notes the model's limitations.

Explore Rutherford's experiment results on atomic radius and nuclear radius, using sphere volume calculations to conclude that an atom's volume is about 10^15 times the nucleus volume.

Presents rutherford's experiment results on the structure of the atom, showing the number of deviating alpha particles relates to the scattering angle via an inverse fourth-power relation.

Explore the distance of minimum approach in Rutherford's experiment and how alpha particle repulsion by the nucleus explains their scattering.

Derive the distance of minimum approach for alpha particles scattering by applying energy conservation, equating kinetic and potential energies, and analyzing the nucleus repulsion in Rutherford's experiment.

Rutherford's gold foil experiment reveals that atoms consist mostly of empty space, with a dense, positively charged nucleus at the center that repels and occasionally rebounds an alpha particle.

Chadwick's 1932 discovery reveals a neutral neutron emitted when beryllium is bombarded by an alpha particle, solving mass balance and explaining why neutrons are not deflected by electric fields.

Explore atomic number as the count of protons, equal to electrons in a neutral atom. Define mass number as protons plus neutrons and the neutrons formula shown for element notation.

Explores isotopes, atoms with the same protons but different mass numbers, illustrated by hydrogen’s protium, deuterium, and tritium. Discuss isotones, atoms with equal neutron numbers.

Explore isobars by comparing argon and calcium with mass number 40. See differing atomic numbers and neutron counts, and note isoelectronic species like Na+, Mg2+, Al3+.

This lecture outlines the limitations of Rutherford's model, highlighting its failure to explain the hydrogen spectrum and conflicts with Maxwell's electromagnetic radiation theory, which implies electron collapse into the nucleus.

Explore why Rutherford's atomic model failed to explain electron energy, spectra, and orbital stability, and how Bohr's postulates advanced a better atomic model for future study.

Define the electron volt as the kinetic energy gained by an electron under a one-volt potential, with 1 eV = 1.6 × 10^-19 J.

Explore the dual nature of light by contrasting its wave and particle natures. Wave energy transfers continuously with no matter transfer, while particle nature is discrete, including longitudinal waves.

Explore the general characteristics of transverse waves, including wavelength (lambda) and amplitude (A), and how the distance between consecutive crests or troughs defines these quantities with their units.

Explore the general characteristics of transverse waves, focusing on frequency and wave number, with f = 1/T and k = 1/λ.

Define the wave velocity v, establish its unit in meters per second, and use the v = f λ relation to connect velocity, frequency, and wavelength.

Explore how electromagnetic waves arise from electric and magnetic fields, with both fields coexisting and remaining perpendicular to each other and to the direction of propagation.

explore electromagnetic waves, confirm they travel at the speed of light and apply c = fλ to relate frequency and wavelength, rank the spectrum from radio to cosmic rays.

Explain how Planck's quantum theory describes black body radiation as discrete energy packets called photons, with energy proportional to frequency via E = h ν, introducing Planck's constant.

Explore Planck's quantum theory by linking photon energy to frequency and wavelength, deriving E=hv and E=hc/λ, and applying formulas for multiple photons and unit conversions.

understand ionization energy as the energy required to remove an electron from the outermost shell of an isolated gaseous atom, an endothermic process, illustrated with sodium.

Investigate bond energy and kinetic energy, defining bond energy as the energy to break a bond, and examine cases where photon energy matches or exceeds bond energy.

Explore the photoelectric effect, Einstein’s Nobel Prize–winning concept, through the cathode–anode setup, photon-induced electron ejection, and the resulting photocurrent, highlighting energy conservation and the particle nature of light.

Learn how the work function sets the energy threshold for the photoelectric effect, cesium’s low ionization energy as a key example, and Einstein’s equation.

Explore threshold frequency and threshold wavelength in the photoelectric effect, and learn to calculate them from work function formulas, then decide if emission occurs based on photon frequency or wavelength.

Explore how photon energy exceeding work function yields electron kinetic energy in photoelectric effect, using the relation kinetic energy equals photon energy minus work function with frequency and wavelength forms.

Define the intensity of radiation as the photon flux hitting the metal surface, and show that photoelectron kinetic energy depends on frequency (or wavelength) and threshold frequency, not on intensity.

Explore the dual nature of matter through De Broglie’s equation lambda equals h over p, which links wave properties to momentum and reveals matter’s wave–particle character.

Explore Heisenberg's uncertainty principle, showing that position and momentum cannot be measured accurately at the same time for electrons, leading to probabilistic atomic orbital regions rather than definite paths.

Explore how the four quantum numbers—principal n, azimuthal l, magnetic m, and spin quantum number ms—define an electron’s shell, subshell, orientation, and spin.

Explore Schrödinger's wave equation and the electron wave function, where the Laplacian operator yields eigenvalues and orbital probabilities around the nucleus.

Explore atomic orbitals as three-dimensional regions with maximum electron probability, governed by the Heisenberg principle, and learn the four orbital types s, p, d, and f with distinct shapes.

Explores the shapes of s, p, d, and f orbitals and explains how many electrons each can hold, with shapes aligned to the x, y, z axes.

Explore the Aufbau principle, which guides electron filling in atomic orbitals by increasing energy, using the n + l rule to prioritize lower energy orbitals.

Apply Pauli's exclusion principle to electron configurations. Enforce that no two electrons in the same orbital share all four quantum numbers n, l, m, and s, and require opposite spins.

Hund's rule of maximum multiplicity states that electrons fill empty orbitals singly before pairing, and all unpaired electrons have the same spin, minimizing repulsion.

Explores the extra stability of atomic orbitals, explaining Hund's rule, half-filled and fully filled subshells, with chromium and copper as exceptional cases and their revised electron configurations.

Learn to write electronic configuration of elements by filling electrons into orbitals in order of increasing energy, guided by atomic numbers, with examples from lithium to nickel and square-bracket notation.

Explore the significance of the psi wave function, Schrödinger's equation, and psi squared for predicting electron probabilities and atomic orbitals in three-dimensional space.