CHEM 201: Electronic Structure, Periodic Trends, Molecular Representations, and Molecular Orbital Theory

Comprehensive vocabulary flashcards covering electronic configurations, light and atomic emission spectra, periodic trends, structural representations of organic molecules, and molecular orbital theory.

Valence Shell

The occupied electron shell of an atom with the highest principal quantum number ($n$), which possesses the highest energy level.

Valence Electrons

Electrons occupying the outermost shell (and $(n-1)d$ subshell for transition metals) that possess the highest energy, are chemically active, and participate in chemical reactions.

Core Electrons

Inner shell electrons that are buried deep inside an atom and do not actively participate in chemical bonding or reactions.

Cation

A positively charged ion formed when a neutral atom loses one or more valence electrons, resulting in fewer electrons than protons.

Anion

A negatively charged ion formed when a neutral atom gains one or more valence electrons, resulting in more electrons than protons.

Isoelectronic Species

Atoms or ions from different elements that have identical numbers of electrons and the exact same electronic configuration.

Octet Rule

The tendency of main-group atoms to lose, gain, or share electrons in order to achieve the stable valence shell electron configuration ($ns^2 np^6$, totaling $8$ electrons) of the nearest noble gas.

Continuous Spectrum

An unbroken distribution of light across all wavelengths and colors produced when polychromatic white light is refracted through a prism.

Atomic Emission Spectrum

A characteristic series of discrete, individual colored lines emitted when light produced by excited gaseous atoms is refracted through a prism.

Electromagnetic Radiation

A form of energy consisting of perpendicular oscillating electric and magnetic fields traveling through space as a transverse wave at the speed of light ($c = 2.998 \times 10^8\,\text{m/s}$ in a vacuum).

Wavelength ($\lambda$)

The physical distance between two successive peaks or troughs of a wave cycle, typically measured in meters ($\text{m}$).

Frequency ($\nu$)

The number of complete wave cycles or oscillations that pass a fixed point per second, expressed in units of $\text{s}^{-1}$ or Hertz ($\text{Hz}$).

Photon

A discrete, massless packet or quantum of electromagnetic energy whose magnitude is directly proportional to the frequency of the radiation ($E = h \times \nu$).

Ground State

The lowest energy and most stable electronic state of an atom in which all electrons occupy orbitals of the lowest possible energy.

Excited State

An energized, unstable state of an atom where an electron has absorbed energy and transitioned to a higher energy atomic orbital.

Relaxation

The return of an excited electron from a higher energy level back to a lower energy orbital, accompanied by the emission of a photon of light matching the energy difference ($\Delta E = h \times \nu$).

Shielding

The electrostatic screening effect whereby inner core electrons reduce the full positive nuclear attraction felt by valence electrons.

Effective Nuclear Charge ($Z^*_{\text{eff}}$)

The actual net positive charge experienced by an electron in a multielectron atom, approximated by the formula $Z^*_{\text{eff}} = Z - S$, where $Z$ is the atomic number and $S$ is the number of shielding core electrons.

Atomic Radius

Half of the experimentally measured distance between the nuclei of two identical adjacent atoms bonded or packed together in the solid state.

First Ionization Energy ($\text{IE}_1$)

The minimum energy required to remove the most loosely bound electron from one mole of isolated gaseous atoms in their ground state ($\text{X}(g) \rightarrow \text{X}^+(g) + e^-$).

Successive Ionization Energy

The energy required to sequentially remove subsequent electrons from increasingly positive gaseous ions ($\text{IE}_1 < \text{IE}_2 < \text{IE}_3$).

Electron Affinity ($\text{EA}$)

The enthalpy change ($\Delta_{\text{EA}} H$) associated with the addition of one mole of electrons to one mole of isolated gaseous atoms to form negative ions ($\text{X}(g) + e^- \rightarrow \text{X}^-(g)$).

Curved Arrows (Curly Arrows)

Visual tools in reaction mechanisms used to depict the precise movement of electron pairs from regions of high electron density to regions of low electron density during bond breaking or bond making.

Condensed Structure

A compact structural formula representing organic molecules where some or all single covalent bonds and lone pairs are omitted, and groups attached to a carbon atom are listed beside it.

Heteroatom

Any atom present in an organic molecular framework that is neither carbon nor hydrogen, such as oxygen, nitrogen, sulfur, or halogens.

Functional Group

A specific atom or group of bonded atoms within a molecule exhibiting characteristic polarity and reactivity regardless of the larger structure to which it is attached.

Constitutional Isomers

Compounds that share the identical molecular formula but possess different atom-to-atom connectivity in their chemical structures.

Stereoisomers

Compounds that share both the same molecular formula and the same atom-to-atom connectivity, but differ in the three-dimensional spatial orientation of their bonds.

Molecular Orbital (MO) Theory

A quantum mechanical bonding model wherein atomic orbital wave functions ($\psi$) combine constructively or destructively to generate delocalized molecular orbitals across a whole molecule.

Bonding Molecular Orbital

A stabilizing, lower-energy molecular orbital produced by constructive in-phase overlap of atomic orbitals, resulting in high electron density between the two nuclei along the internuclear axis.

Antibonding Molecular Orbital

A destabilizing, higher-energy molecular orbital produced by destructive out-of-phase overlap of atomic orbitals, characterized by a nodal plane with zero electron density between the two nuclei.

Bond Order

A parameter defining the stability and strength of a covalent bond, calculated as half the difference between the number of bonding electrons and antibonding electrons: $\text{Bond Order} = \frac{(\text{Bonding } e^-) - (\text{Antibonding } e^-)}{2}$.