Chapter 7: Atomic Structure and Periodicity

Vocabulary practice flashcards covering electromagnetic radiation, quantum theory, hydrogen emission spectra, the Bohr model, and atomic orbital properties.

Electromagnetic Radiation

A form of energy transmission through space featuring electric ($E$) and magnetic ($H$) fields that simultaneously oscillate in planes mutually perpendicular to each other and to the direction of propagation.

Wavelength ($\lambda$)

The distance between two consecutive peaks or troughs in a wave, commonly measured in meters ($m$), micrometers ($1\,\mu m = 10^{-6}\,m$), nanometers ($1\,nm = 10^{-9}\,m$), or picometers ($1\,pm = 10^{-12}\,m$).

Frequency ($\nu$)

The number of waves per second that pass a given point in space, expressed in oscillations per second ($s^{-1}$) or hertz ($\text{Hz}$).

Speed of Light ($c$)

The universal speed at which all electromagnetic radiation travels in a vacuum, exactly $2.99792458 \times 10^8\,m/s$, related to wavelength and frequency by $\lambda \nu = c$.

Quantum

A small, discrete packet of energy of magnitude $h\nu$. Matter can only absorb or release energy in whole-number multiples of this quantity.

Planck's Constant ($h$)

A fundamental physical constant discovered in 1901 with a value of $6.626 \times 10^{-34}\,J \cdot s$ (or $kg\,m^2/s$), relating the energy of a quantum to its radiation frequency via $\Delta E = h\nu$.

Photoelectric Effect

The emission of electrons from the surface of a metal when light shines upon it, occurring only when the incident light frequency exceeds a specific threshold frequency ($\nu_0$).

Threshold Frequency ($\nu_0$)

The minimum frequency of radiation required to dislodge electrons from a metal surface in the photoelectric effect, below which no electrons are emitted regardless of light intensity.

Photon

A quantized stream or packet of electromagnetic radiation whose energy is given by $E_{\text{photon}} = h\nu = \frac{hc}{\lambda}$.

Dual Nature of Light

The principle that electromagnetic radiation exhibits both wave-like characteristics (periodic oscillation) and particulate characteristics (discrete packets or photons).

de Broglie Equation

The relationship proposed in 1923, $\lambda = \frac{h}{mv}$, which relates the particulate properties of matter (mass $m$ and velocity $v$) to its wave properties (wavelength $\lambda$).

Continuous Spectrum

A spectrum containing all wavelengths of visible light without interruption, produced when white light passes through a prism.

Line Spectrum

A spectrum showing only discrete, individual wavelengths of light, produced when light emitted by excited atoms passes through a prism.

Bohr Model

A quantum model of the hydrogen atom developed in 1913 where an electron moves in allowed circular orbits around the nucleus, with quantized energies given by $E = -2.178 \times 10^{-18}\,J \left(\frac{1}{n^2}\right)$.

Orbital

A spatial representation bounded by a surface that encloses $90\%$ of the total electron probability.

s Orbitals

Spherical atomic orbitals characterized by the angular momentum quantum number $l=0$, increasing in size as the principal quantum number $n$ increases.

Node

A region or surface in an atom where the probability of finding an electron drops to zero ($\psi^2 = 0$).

p Orbitals

A set of atomic orbitals characterized by the quantum number $l=1$, having two lobes separated by a nodal plane at the nucleus, oriented along coordinate axes ($2p_x$, $2p_y$, $2p_z$).

d Orbitals

A set of five orbitals characterized by the quantum number $l=2$ that first appear in principal energy level $n=3$, labeled $d_{xz}$, $d_{yz}$, $d_{xy}$, $d_{x^2-y^2}$, and $d_{z^2}$.

Degenerate Orbitals

Orbitals that share the exact same energy level, such as all orbitals possessing the same principal quantum number $n$ in a hydrogen atom.

Ground State

The lowest energy state of an atom, in which the electron resides in the lowest available orbital (such as the $1s$ orbital in hydrogen).

Excited State

An unstable, higher-energy state of an atom that occurs when an electron absorbs energy and transitions to an orbital of higher principal energy level.

Aufbau Principle

The principle stating that as protons are added to the nucleus to build elements, electrons are added to atomic orbitals in order of increasing energy to establish the most stable configuration.