CHEM 150 - Test 1 Review: Atomic Structure and Chemical Bonding

Vocabulary practice flashcards covering quantum atomic theory, orbital shapes, periodic trends, Lewis structures, octet exceptions, and VSEPR geometries for CHEM 150 Test 1 Review.

Photoelectric Effect

The ejection of electrons from a metal surface when irradiated with light of specific threshold wavelengths, demonstrating the particle nature of light.

de Broglie Wavelength

The wavelength associated with a moving particle of matter, calculated as $\lambda = \frac{h}{mv} = \frac{h}{p}$, demonstrating wave-particle duality.

Heisenberg Uncertainty Principle

The principle stating that it is impossible to simultaneously know both the precise position and momentum of an electron, expressed as $\Delta x \cdot m\Delta v \ge \frac{h}{4\pi}$.

Electron Density ($|\psi|^2$)

The square of the wave function from the Schrödinger equation ($\hat{H}\psi = E\psi$), representing the probability distribution of finding an electron in a 3D region of space.

Nodal Plane

A flat plane passing through an atom where the probability of finding an electron is exactly zero.

Radial Node

A specific distance or spherical surface from the atomic nucleus where the electron density is zero, occurring in higher principal energy levels.

Principal Quantum Number ($n$)

The quantum number that designates the primary energy level and overall size of an atomic orbital.

Azimuthal Quantum Number ($l$)

The orbital angular momentum quantum number that specifies the subshell and defines the three-dimensional geometric shape of an orbital.

Magnetic Quantum Number ($m_l$)

The quantum number that specifies the three-dimensional spatial orientation of an orbital within a specific subshell.

Spin Quantum Number ($m_s$)

The quantum number that specifies the intrinsic angular momentum or spin orientation of an electron, taking values of $+\frac{1}{2}$ or $-\frac{1}{2}$.

Effective Nuclear Charge ($Z_{\text{eff}}$)

The actual net positive charge experienced by an electron in a multielectron atom, determined by the equation $Z_{\text{eff}} = Z - S$, where $Z$ is the atomic number and $S$ is the core electron shielding.

Isoelectronic Series

A group of atoms or ions that possess the exact same number of electrons and identical ground-state electron configurations.

Metallic Bond

A chemical bond characterized by free, delocalized valence electrons that travel between a lattice of metallic cations, producing high conductivity and high melting points.

Lattice Energy

The energy emitted when separated gaseous ions coalesce to form an ionic solid crystal lattice, as in $\text{Na}^+(g) + \text{Cl}^-(g) \rightarrow \text{NaCl}(s)$.

Pure Covalent Bond

A bond in which electrons are shared equally between two bonding atoms, characterized by an electronegativity difference of $\Delta\text{EN} \le 0.4$.

Polar Covalent Bond

A chemical bond in which bonding electrons are shared unequally between atoms, defined by an electronegativity difference range of $0.5 \le \Delta\text{EN} \le 2.0$.

Formal Charge

The hypothetical charge assigned to an individual atom in a molecule, calculated by $\text{Formal Charge} = (\#\text{ valence } e^-) - (\#\text{ lone pair } e^-) - \frac{1}{2}(\#\text{ bonding } e^-)$.

Resonance Structures

Two or more valid Lewis structures for the same molecule that possess the same atomic skeleton but differ in electron pair distribution, representing delocalized electron bonding.

Free Radical

A highly reactive molecule or atom containing an odd total number of valence electrons, leaving at least one unpaired electron (e.g., $\text{NO}$).

Incomplete Octet

An exception to the octet rule in which a central atom, commonly from Group 2 (like $\text{Be}$) or Group 13 (like $\text{B}$ in $\text{BF}_3$), is stable with fewer than 8 valence electrons.

Hypercoordinate Molecule

A molecule containing a central atom from the third period or below that accommodates more than eight valence electrons by expanding its valence shell into available $d$-orbitals.

Trigonal Bipyramidal Geometry

An electron group geometry resulting from five electron domains around a central atom, exhibiting axial bond angles of $90^\circ$ and equatorial bond angles of $120^\circ$.

Lone Pair Spatial Repulsion

The principle that nonbonding electron pairs occupy greater angular volume than bonding pairs because they are held by only a single nucleus, exerting greater repulsion and compressing adjacent bond angles.

Equatorial Lone Pair Preference

The VSEPR structural rule dictating that lone pairs in trigonal bipyramidal arrangements preferentially occupy equatorial positions rather than axial positions to minimize unfavorable $90^\circ$ lone pair–bonding pair repulsions.

Wedge and Dash Notation

A three-dimensional representation convention where solid lines indicate bonds in the plane of the page, wedges indicate bonds pointing out toward the viewer, and dashes indicate bonds pointing into the page.

Molecular Dipole Moment

The overall vector sum of individual polar bond dipoles across the three-dimensional geometry of a molecule, determining whether the molecule possesses a permanent net polarity.