Atomic Structure and Quantum Theory – Key Vocabulary

50 English vocabulary flashcards summarizing major terms from the lecture on atomic structure, radiation, quantum theory, and the Bohr model.

Atomic Structure

The arrangement of sub-atomic particles (protons, neutrons, electrons) within an atom.

Cathode Rays

Streams of fast-moving electrons produced in a discharge tube when high voltage is applied across electrodes.

J. J. Thomson

Physicist who discovered the electron and measured its charge-to-mass ratio (e/m).

Charge-to-Mass Ratio (e/m) of Electron

The ratio of an electron’s electric charge to its mass, determined by Thomson as 1.76 × 10¹¹ C kg⁻¹.

Plum Pudding Model

Thomson’s early atomic model in which electrons are embedded in a positively charged ‘pudding.’

Anode Rays (Canal Rays)

Positive ion beams observed in discharge tubes, leading to discovery of the proton.

E. Goldstein

Scientist who discovered anode (canal) rays using perforated cathodes.

Proton

Positively charged sub-atomic particle found in the nucleus; charge +1e and mass ≈ 1 u.

Electron

Negatively charged sub-atomic particle with charge −1e and mass 9.1 × 10⁻³¹ kg.

Rutherford’s Alpha Scattering Experiment

Gold-foil experiment that revealed atoms have a dense, positively charged nucleus.

Nuclear Model of Atom

Rutherford’s model where electrons orbit a small, massive nucleus containing protons (and later neutrons).

Nucleus

Central core of an atom containing protons and neutrons; holds most of the atom’s mass.

Alpha Particles

Helium nuclei (²⁴He²⁺) emitted in radioactive decay; used by Rutherford for scattering studies.

Atomic Number (Z)

Number of protons in an atom’s nucleus; determines element identity.

Mass Number (A)

Total number of protons and neutrons (nucleons) in an atomic nucleus.

Isotopes

Atoms of the same element (same Z) with different mass numbers (different neutrons).

Isobars

Atoms with the same mass number but different atomic numbers.

Isotones

Atoms with the same number of neutrons but different atomic and mass numbers.

Radioactivity

Spontaneous emission of particles or radiation from unstable atomic nuclei.

Alpha Decay

Radioactive process where a nucleus emits an alpha particle, reducing A by 4 and Z by 2.

Beta Decay

Radioactive emission of a beta particle (electron or positron) accompanied by a change in Z by ±1.

Gamma Radiation

High-energy electromagnetic radiation emitted from excited nuclei; carries no charge.

Electromagnetic Radiation

Energy propagated as oscillating electric and magnetic fields, including light, X-rays, etc.

Wavelength (λ)

Distance between successive crests of a wave; measured in metres or nanometres.

Frequency (ν)

Number of wave cycles passing a point per second; measured in hertz (Hz).

Speed of Light (c)

Constant speed of electromagnetic waves in vacuum, 3.00 × 10⁸ m s⁻¹.

Black Body Radiation

Continuous spectrum emitted by an ideal object that absorbs and re-radiates all incident energy.

Planck’s Quantum Theory

Concept that energy is emitted or absorbed in discrete packets called quanta; E = hν.

Quantum

Smallest discrete quantity of energy, proportional to radiation frequency.

Planck’s Constant (h)

Proportionality constant in E = hν; value 6.626 × 10⁻³⁴ J s.

Photon

A quantum of electromagnetic radiation; massless particle of light with energy hν.

Photoelectric Effect

Ejection of electrons from a metal surface when illuminated by light of sufficient frequency.

Work Function (Φ)

Minimum energy required to eject an electron from a metal surface in the photoelectric effect.

Threshold Frequency (ν₀)

Minimum light frequency needed to release electrons from a given metal.

Einstein’s Photoelectric Equation

hν = Φ + ½ m v²; relates photon energy to work function and kinetic energy of emitted electron.

Stopping Potential

Reverse voltage needed to halt photoelectrons; used to measure their maximum kinetic energy.

Bohr’s Model of Hydrogen Atom

Atomic model where electrons move in quantized circular orbits with angular momentum nh/2π.

Orbit (Bohr)

Fixed circular path of an electron around the nucleus with defined energy.

Energy Levels (n)

Discrete allowed energies of electrons in atoms, labeled by principal quantum number n.

Rydberg Constant (R_H)

Spectroscopic constant (1.097 × 10⁷ m⁻¹) used in hydrogen emission wavelength formula.

Spectral Series

Grouped sets of hydrogen emission lines corresponding to electronic transitions to a common lower level.

Lyman Series

Hydrogen spectral lines produced by transitions to n = 1; lies in the ultraviolet region.

Balmer Series

Hydrogen lines for transitions to n = 2; visible region.

Paschen Series

Infrared hydrogen lines due to transitions to n = 3.

Brackett Series

Infrared hydrogen lines from transitions to n = 4.

Pfund Series

Infrared hydrogen lines from transitions to n = 5.

Wave-Particle Duality

Concept that particles such as electrons and photons exhibit both wave and particle properties.

Electron Volt (eV)

Energy gained by an electron moving through 1 volt; 1 eV = 1.602 × 10⁻¹⁹ J.

Angular Momentum Quantization

Bohr postulate that electron angular momentum equals nh/2π (n = 1, 2, 3…).

Charge of Electron (e)

Fundamental negative electric charge, −1.602 × 10⁻¹⁹ coulombs.