Early Atomic Theories, the Bohr Model, Quantum Theory, and Magnetism

Comprehensive vocabulary review flashcards covering Dalton, Thomson, and Rutherford atomic models, the emergence of quantum theory, the Bohr model of hydrogen, electron transitions, ion configurations, and atomic magnetism.

Modern Atomic Theory (Dalton)

The theory formulated by John Dalton in 1803 proposing the solid sphere model of the atom as the foundational building block of modern chemistry.

Plum Pudding Model

The atomic model proposed by J.J. Thomson in 1904 following his discovery of the electron via cathode ray tubes, depicting negatively charged electrons embedded within a positively charged sphere.

Proton

A positively charged subatomic particle with a charge of $+1$ and a mass of approximately $1\,\text{amu}$, discovered by Ernest Rutherford in 1911 and located in the nucleus of an atom.

Gold Foil Experiment

An experiment performed by Ernest Rutherford in 1911 directing alpha particles at thin gold foil; the observation that most particles passed straight through while a few deflected or reflected led to the conclusion that the atom is mostly empty space with a dense, positively charged nucleus.

Planetary Model Collapse

A major flaw in Rutherford's model where classical physics dictated that orbiting, accelerating electrons should continuously radiate electromagnetic energy, lose speed, and spiral into the nucleus, resulting in the collapse of the atom.

Quantum of Energy

A discrete packet or unit of energy that has a specific wavelength and energy, demonstrating that energy is not continuous.

Photon

A quantum of light energy whose energy is directly proportional to its radiation frequency and inversely proportional to its wavelength.

Photoelectric Effect

A phenomenon discovered by Heinrich Hertz in which light shone onto a metal surface causes the emission of electrons from that metal, where the kinetic energy of emitted electrons depends on the frequency of the light.

Planck's Quantum Hypothesis

The hypothesis introduced in 1900 by Max Planck proposing that light is composed of discrete packets of energy (quanta) and that the emission of energy occurs only in discrete multiples rather than a continuous wave.

Planck's Constant ($h$)

The fundamental proportionality constant in Planck's equation ($E = hf$) relating photon energy to electromagnetic wave frequency, valued at $6.63 \times 10^{-37}\,\text{kJ}\cdot\text{s}$.

Stationary State

A fixed circular orbit in Bohr's model where an electron possesses constant energy and does not emit electromagnetic radiation while circulating.

Ground State

The lowest possible energy level configuration available to electrons within an atom.

Excited State

A temporary condition occurring when an electron absorbs energy from a photon and transitions to an orbit located farther from the nucleus.

Emission Spectrum

A pattern of discrete bright lines produced on a dark background when electrons transition from a higher energy state to a lower energy state, releasing energy in the form of photons.

Absorption Spectrum

A continuous spectrum interrupted by dark lines (a black-line spectrum) produced when cold gases absorb specific wavelengths of light as electrons transition from lower to higher energy levels.

Principal Quantum Number ($n$)

A positive integer designating the main energy level or orbit of an electron, which determines the maximum electron capacity of that orbit through the formula $2n^2$.

Ionization Energy

The amount of energy required to completely remove an electron from the ground state of an atom.

Bohr's First Postulate

The postulate asserting that electrons can only move within certain fixed orbits corresponding to specific energy levels without losing energy.

Bohr's Second Postulate

The postulate stating that an electron can only transition from one stationary orbit to another when it absorbs or emits a photon of energy precisely equal to the difference between the energy levels ($\Delta E = hf$).

Balmer Series

The series of visible spectral transitions in the hydrogen atom that occur when an excited electron transitions from higher energy levels ($n \ge 3$) down to the $n = 2$ level, producing four distinct lines at $410\,\text{nm}$, $434\,\text{nm}$, $486\,\text{nm}$, and $656\,\text{nm}$.

Lyman Series

The series of ultraviolet spectral lines in the hydrogen atom produced when electrons drop from higher energy levels to the $n = 1$ ground state.

Paschen Series

The series of infrared spectral lines in the hydrogen atom resulting from electron transitions ending at the $n = 3$ energy level.

Ion

An atom or molecule possessing a net electric charge caused by the loss or gain of one or more valence electrons.

Cation

A positively charged ion formed when an atom loses one or more electrons.

Anion

A negatively charged ion formed when an atom gains one or more electrons.

Isoelectronic Species

Different atoms or ions that possess the identical electron configuration, such as $\text{Cl}^-$, $\text{K}^+$, and $\text{Ar}$.

Cation Electron Removal Rule

The principle for determining cation configurations stating that electrons must be removed first from the orbital with the highest principal quantum number ($n$), removing outer $p$ electrons before $s$ electrons, and $s$ electrons before $d$ electrons.

Magnetic Domains

Regions within magnetic materials where atoms with unpaired electrons align their magnetic fields; they are randomly oriented in unmagnetized matter and line up with their north poles pointing in the same direction in magnetized matter.

Paramagnetism

A property of materials like aluminum or platinum that become magnetized when placed in an external magnetic field due to unpaired electrons, but lose their magnetism immediately when the field is removed.

Ferromagnetism

A property of materials such as iron and nickel that can retain their magnetic alignment and properties even after an external magnetic field has been removed.