Atomic Structure Physical Chemistry Vocabulary

Vocabulary flashcards covering key terms, concepts, quantum numbers, and fundamental principles of Atomic Structure from Physical Chemistry lecture notes.

Plum-Pudding Model

Thomson's 1904 model of the atom proposing a uniform sphere of positive charge in which electrons are present at discrete locations.

Distance of Closest Approach ($r_0$)

The distance $r_0$ from the nucleus at which an approaching $\alpha$-particle comes to rest as its initial kinetic energy is completely converted into electrostatic potential energy.

Wavelength ($\lambda$)

The distance between two nearest crests or two nearest troughs of a wave.

Frequency ($\nu$)

The number of complete waves that pass through a given point in $1\,\text{second}$.

Wave Number ($\bar{\nu}$)

The reciprocal of wavelength ($\bar{\nu} = \frac{1}{\lambda}$), representing the number of waves present in $1\,\text{cm}$.

Amplitude ($a$)

The height of a wave's crest or the depth of its trough.

Planck's Quantum Theory

A theory postulating that radiant energy is emitted or absorbed discontinuously in discrete packets of energy called quanta, where energy $E = h\nu$.

Photon

The smallest discrete packet or quantum of light energy.

Photoelectric Effect

The ejection of electrons from a metal surface when light radiation of a suitable frequency strikes the metal surface.

Threshold Frequency ($\nu_0$)

The minimum frequency of incident radiation required to just cause the ejection of photoelectrons from a given metal surface.

Work Function ($\phi$)

The threshold energy ($\phi = h\nu_0$) representing the minimum energy of a photon required to liberate an electron from a metal surface.

Stopping Potential ($V_0$)

The minimum reverse potential applied to stop photocurrent completely, where maximum kinetic energy satisfies $K.E._{\text{max}} = e V_0$.

Stationary Orbits

Specific circular orbits in which revolving electrons maintain constant energy without radiating electromagnetic radiation.

Bohr Radius ($r_0$)

The radius of the smallest orbit ($n = 1$) for a hydrogen atom ($Z = 1$), equal to $0.529\,\text{\AA}$ ($5.29 \times 10^{-11}\,\text{m}$).

Ground State

The lowest potential energy state of an atom or ion.

Excited State

Any state of an atom or ion having higher energy than its ground state ($n > 1$).

Ionisation Energy

The minimum energy required to completely remove an electron from the ground state of an atom or ion to infinity ($n = \infty$).

Ionisation Potential

The potential difference through which a free electron must be accelerated from rest so that its kinetic energy equals the ionisation energy of the atom.

Excitation Energy

The energy required to move an electron from the ground state of an atom to any specified higher energy state.

Binding Energy

The energy required to move an electron from any specific energy state $n$ to infinity ($n = \infty$).

Emission Line Spectrum

A spectrum consisting of well-defined bright sharp lines corresponding to definite wavelengths separated by dark spaces, emitted by an excited atomic gas.

Absorption Spectrum

A spectrum obtained when white light passes through a substance, displaying dark lines or bands corresponding to absorbed wavelengths within an otherwise continuous spectrum.

Lyman Series

A spectral line series of atomic hydrogen produced by electronic transitions from $n_2 \ge 2$ to $n_1 = 1$, lying in the ultraviolet region.

Balmer Series

A spectral line series of atomic hydrogen produced by electronic transitions from $n_2 \ge 3$ to $n_1 = 2$, lying in the visible region.

Paschen Series

A spectral line series of atomic hydrogen produced by electronic transitions from $n_2 \ge 4$ to $n_1 = 3$, lying in the infrared region.

Rydberg Constant ($R_H$)

An atomic spectroscopy constant equal to $109677\,\text{cm}^{-1}$ used in calculating atomic hydrogen line wavelengths.

De Broglie Wavelength

The wavelength ($\lambda = \frac{h}{p} = \frac{h}{m v}$) associated with a moving material particle exhibiting wave-particle duality.

Heisenberg Uncertainty Principle

A quantum mechanical principle stating that it is impossible to simultaneously measure both the exact position ($\Delta x$) and exact momentum ($\Delta p$) of a subatomic particle.

Orbital

A three-dimensional region of space around the nucleus where the probability of finding an electron is maximum (about 90-95%).

Node

A surface or point around the nucleus where the probability density ($\psi^2$) of finding an electron is zero.

Radial Probability Distribution Function

The function ($4\pi r^2 \psi^2(r)\,dr$) measuring the probability of finding an electron in a spherical shell of thickness $dr$ at a distance $r$ from the nucleus.

Principal Quantum Number ($n$)

The quantum number designating the primary electron shell, specifying the size and major energy level of an orbital.

Azimuthal Quantum Number ($l$)

The quantum number specifying subshell type, orbital angular momentum, and spatial shape, ranging from $0$ to $(n - 1)$.

Magnetic Quantum Number ($m$)

The quantum number determining the spatial orientation of an orbital in a subshell, taking integral values from $-l$ to $+l$.

Spin Quantum Number ($s$)

The quantum number designating the spin direction of an electron around its own axis, taking values of $+\frac{1}{2}$ or $-\frac{1}{2}$.

Aufbau Principle

The principle stating that electrons progressively fill atomic orbitals in order of increasing energy starting from the lowest available energy orbital.

Pauli Exclusion Principle

The principle stating that no two electrons in a single atom can have identical values for all four quantum numbers.

Hund's Maximum Multiplicity Rule

The rule stating that degenerate orbitals in a subshell are each filled singly with parallel spin electrons before pairing begins.

Bohr-Bury $(n + l)$ Rule

A rule stating that the orbital with the lower $(n + l)$ value has lower energy; if $(n + l)$ values are equal, the orbital with the lower principal quantum number $n$ has lower energy.

Exchange Energy

The stabilization energy released when electrons with parallel spins exchange positions among degenerate orbitals in a subshell.

Paramagnetism

The property of substances containing unpaired electrons that causes them to be weakly attracted by an external magnetic field.

Diamagnetism

The property of substances lacking unpaired electrons that causes them to be weakly repelled by an external magnetic field.