Chem111 Chapter 3:Quantum Theory and the Electronic Structure of Atoms

A complete set of vocabulary flashcards covering wave-particle duality, atomic models, line spectra, and quantum numbers based on Chapter 3 notes.

Law of Conservation of Energy

The fundamental principle stating that energy can neither be created nor destroyed.

Wavelength ($\lambda$)

The distance between identical points on successive waves.

Frequency ($\nu$)

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

Amplitude

The vertical distance from the midline of a wave to the top of its peak.

Constructive Interference

An interference pattern created when waves combine in-phase to produce a wave with larger amplitude.

Destructive Interference

An interference pattern created when waves combine out-of-phase to cancel each other out, resulting in zero or reduced amplitude.

Planck's Constant ($h$)

A fundamental physical constant equal to $6.626 \times 10^{-34}\,\text{J}\cdot\text{s}$ that relates energy to frequency ($E = h\nu$).

Photoelectric Effect

The phenomenon where electrons ($e^-$) are ejected from a metal surface when photons of sufficient light energy hit it.

Emission Spectrum

A characteristic spectrum consisting of bright lines produced when excited electrons emit photons while transitioning to lower energy levels.

Absorption Spectrum

A characteristic spectrum consisting of dark lines produced when light passes through a sample and photons of specific wavelengths are absorbed by electrons moving to higher energy levels.

Rydberg Equation

An empirical relationship used to calculate the wavelengths of spectral lines: $\frac{1}{\lambda} = 1.09737316 \times 10^7\,\text{m}^{-1} \left(\frac{1}{n_1^2} - \frac{1}{n_2^2}\right)$.

Ground State

The lowest available energy level or orbit occupied by electrons in an atom.

Excited State

An energy level higher than the ground state occupied by electrons after absorbing thermal or light energy.

De Broglie Wavelength Equation

An equation $\lambda = \frac{h}{m u}$ describing the wave-like behavior of matter by relating wavelength ($\lambda$) to mass ($m$) and velocity ($u$).

Heisenberg Uncertainty Principle

The principle stating that it is impossible to simultaneously know both the exact position ($x$) and momentum ($p$) or velocity ($u$) of a particle, defined as $\Delta x \cdot m \Delta u \ge \frac{h}{4\pi}$.

Schrödinger Wave Equation

A fundamental quantum mechanical formula ($E\psi = H\psi$) where $\psi^2$ describes the probability density of finding an electron in a given region of an atom.

Atomic Orbitals

Three-dimensional shapes describing regions of space around the nucleus with a high probability of finding an electron, capable of holding a maximum of $2$ electrons.

Principal Quantum Number ($n$)

The quantum number ($n > 0$ integers) that defines the energy level and overall size of an atomic orbital.

Angular Momentum Quantum Number ($\ell$)

The quantum number (integers from $0$ to $n - 1$) that defines the shape of an atomic orbital.

Magnetic Quantum Number ($m_\ell$)

The quantum number (integers from $-\ell$ to $+\ell$) that defines the orientation of an atomic orbital in space.

Spin Quantum Number ($m_s$)

The quantum number with possible values of $+\frac{1}{2}$ or $-\frac{1}{2}$ that designates the spin orientation of an electron.

s Orbital

A spherical atomic orbital corresponding to $\ell = 0$, consisting of $1$ orbital per subshell and accommodating a maximum of $2$ electrons.

p Orbital

A dumbbell-shaped atomic orbital corresponding to $\ell = 1$, consisting of $3$ orbitals per subshell and accommodating a maximum of $6$ electrons.

d Orbital

An atomic orbital corresponding to $\ell = 2$, consisting of $5$ orbitals per subshell and accommodating a maximum of $10$ electrons.

f Orbital

An atomic orbital corresponding to $\ell = 3$, consisting of $7$ orbitals per subshell and accommodating a maximum of $14$ electrons.