Atomic Structure and Electron Arrangement: The Quantum Model of the Atom

Vocabulary flashcards covering atomic structure, the quantum mechanical model, quantum numbers, electron configuration rules, flame tests, and line spectra.

Orbital

A region in space where there is a very high probability of finding an electron; if no electrons are present in a particular energy level, the orbital does not exist.

Probable Electron Density

A region in space where it is highly likely that an electron can be found, describing electrons as probability clouds rather than fixed planetary orbits.

Principal Quantum Number ($n$)

The quantum number that represents the main energy level of an electron in an atom, taking positive integer values ($1, 2, 3, \text{...}$) and holding up to $2n^2$ electrons.

Azimuthal Quantum Number ($l$)

The second quantum number in the quantum mechanical model that describes the energy sublevel and the shape of the orbital.

Magnetic Quantum Number ($m_l$)

The third quantum number that specifies the spatial orientation of an atomic orbital in space.

Spin Magnetic Quantum Number ($m_s$)

The fourth quantum number ($+0.5$ or $-0.5$) that indicates the orientation of the magnetic field generated by an electron's spin.

$s$ Orbital

A spherically shaped atomic orbital that defines a region where there is a very high probability of finding up to two electrons.

$p$ Sublevel

An energy sublevel composed of three degenerate dumbbell-shaped orbitals ($p_x, p_y, p_z$) of equal energy that can hold a maximum of six electrons in total.

$d$ Sublevel

An energy sublevel composed of five atomic orbitals that can hold a maximum of 10 electrons in total.

$f$ Sublevel

An energy sublevel composed of seven atomic orbitals that can accommodate up to 14 electrons in total.

Aufbau Principle

The principle stating that as electrons are added to an atom, they fill the lowest-energy available orbitals first before occupying higher energy levels.

Pauli Exclusion Principle

The rule stating that any atomic orbital can hold a maximum of two electrons, and when two electrons occupy the same orbital, they must have opposite spins.

Hund's Rule of Maximum Multiplicity

The rule stating that when filling degenerate orbitals of equal energy, electrons occupy each orbital singly with parallel spins before pairing up.

Degenerate Orbitals

Atomic orbitals within the same energy sublevel that have identical energy levels (such as the three orbitals in a $p$ sublevel).

Continuous Spectrum

A spectrum showing an unbroken sequence of all wavelengths across a given range, such as the full rainbow of visible light from $400\text{ nm}$ to $700\text{ nm}$.

Line Emission Spectrum

A spectrum consisting of discrete coloured lines on a dark background, produced when excited electrons fall to lower energy levels and emit photons of specific wavelengths.

Line Absorption Spectrum

A spectrum displaying dark lines within a continuous spectrum, created when visible light passes through a cold gas that absorbs specific photon wavelengths.

Quantization

The concept that physical quantities, such as energy levels and electromagnetic radiation, exist only in discrete packets or specific allowable amounts rather than a continuous range.

Photon

A discrete quantum of electromagnetic radiation whose energy is directly proportional to its frequency, calculated by $E = h \times f$.

Ground State

The lowest, most stable allowable energy state of an atom where its electrons occupy the lowest possible energy levels (such as $n = 1$ for hydrogen).

Excited State

An unstable energy state that occurs when an electron absorbs energy and transitions to a higher principal energy level ($n \ge 2$).

Convergence

The phenomenon in atomic spectra where higher energy levels become progressively closer together, causing the resulting spectral lines to converge at higher frequencies.

Flame Test

An analytical technique where metal salts are heated in a flame to excite electrons, identifying the metal by the characteristic colour of light emitted as electrons return to lower energy states.

Condensed Electron Configuration

A shorthand notation that represents inner core electrons using the symbol of the previous noble gas in square brackets, followed by the remaining valence electrons.

Transition Metal Ionization Rule

The principle that when transition metals form positive ions, electrons are removed from the sublevel with the highest principal quantum number ($n$) first, meaning $4s$ electrons are lost before $3d$ electrons.

Lyman Series

The series of spectral emission lines in atomic hydrogen that arise from electron transitions falling to the ground state ($n = 1$), producing light in the ultraviolet (UV) region.

Balmer Series

The series of spectral emission lines in hydrogen produced when electrons drop down to the $n = 2$ energy level, releasing light in the visible region.

Paschen Series

The series of spectral lines in atomic hydrogen resulting from electron transitions ending at the $n = 3$ level, emitting lower-energy infrared (IR) radiation.

Inductive Reasoning

A bottom-up approach to scientific knowledge that builds general principles or theories from repeated experimental observations and recognizable patterns.

Deductive Reasoning

A top-down approach to scientific reasoning that infers specific, testable conclusions and observations from established general premises or theoretical models.