General Chemistry: Electronic Structure, Chemical Bonding, and Molecular Geometry

Comprehensive vocabulary flashcards covering electronic configurations, periodic trends, ionic and covalent bonding principles, chemical nomenclature, formula determination, and VSEPR molecular geometry.

Pauli Exclusion Principle

A fundamental quantum mechanical principle stating that each electron within an atom must have a unique set of four quantum numbers: $n$, $l$, $m_l$, and $m_s$.

Aufbau Principle

The rule stating that electrons fill available atomic orbitals in order of increasing energy to establish the lowest possible ground-state energy for the atom.

Hund's Rule

The rule stating that when placing electrons into degenerate orbitals (orbitals with identical energy), each orbital receives one electron with parallel spin before any orbital receives a second electron to form a pair.

Paramagnetism

A magnetic condition exhibited by an atom, ion, or molecule containing one or more unpaired electrons, causing its electron spins to align with an external magnetic field and generate an attractive force toward the magnet.

Diamagnetism

A magnetic condition exhibited by an atom, ion, or compound in which all electrons are completely paired, causing it to experience a weak repulsion when exposed to an external magnetic field.

Ionization Energy

The thermodynamic energy change ($\Delta U$) required to remove an electron from a gaseous atom or ion in its ground state, which is always positive ($IE > 0$).

Electron Affinity

The quantity of energy released when a neutral gas-phase atom captures an electron to form a gas-phase anion ($EA = -\Delta U$), typically expressed as a positive numerical value.

Electron Attachment Enthalpy

The thermodynamic state function ($\Delta_{\text{EA}}H$) representing the actual enthalpy change when a gas-phase atom accepts an electron to form a gas-phase anion ($A(g) + e^- \rightarrow A^-(g)$), which is negative for exothermic attachments.

Ionic Bonding

A type of chemical bonding between metals and nonmetals driven by the complete transfer of one or more electrons from one atom to another, held together by electrostatic (Coulombic) attraction.

Dissociation

The process occurring when a soluble ionic solid dissolves in water and its constituent cations and anions separate from each other to disperse freely throughout the aqueous solution.

Stock Notation

A systematic nomenclature system that distinguishes the oxidation state of a cation by placing Roman numerals in parentheses immediately following the metal's name (e.g., iron(II) oxide vs. iron(III) oxide).

Formula Weight

The sum of the atomic masses of all the constituent atoms specified in the empirical formula of an ionic compound.

Molecular Weight

The sum of the atomic masses of all the constituent atoms specified in the molecular formula of a covalent compound.

Empirical Formula

A chemical formula that represents the relative ratio of elements present in a compound written using the least common multiples (lowest whole-number ratio).

Molecular Formula

A chemical formula designating the exact actual number and identity of each atom bound in a discrete molecular (covalent) substance.

Acid Mine Drainage

An environmental condition produced by the oxidation of pyrite ($FeS_2$) by oxygen and water according to the reaction $4FeS_2(aq) + 15O_2(aq) + 14H_2O(l) \rightleftharpoons 4Fe(OH)_3(s) + 8SO_4^{2-}(aq) + 16H^+(aq)$, resulting in acidic drainage and an orange iron(III) hydroxide precipitate.

Covalent Bonding

A type of chemical bonding characteristic of nonmetal elements combining with other nonmetals, involving the localized sharing of valence electron pairs between atoms.

Electronegativity

The relative propensity of an atom to attract shared electron density toward itself within a chemical covalent bond.

Pure Covalent Bond

A nonpolar covalent bond characterized by perfectly equal electron sharing, occurring when the electronegativity difference ($\Delta EN$) between the two bonded atoms is exactly $0$.

Polar Covalent Bond

A covalent bond formed between atoms possessing an intermediate difference in electronegativity ($\Delta EN$ typically $0.4$ to $2.0$), causing unequal electron sharing and generating partial charges ($\delta^+$ and $\delta^-$).

Dipole Moment

A quantitative measurement of bond or molecular polarity ($\mu$) that is directly proportional to the magnitude of separated partial charges ($q$) and the distance ($r$) between them, measured in Debyes ($\text{D}$).

Resonance Structures

Two or more valid Lewis formulas for the same compound that share the exact same skeletal atomic connectivity and electron count, differing solely in the localized placement of valence electrons.

Formal Charge

An electron bookkeeping value calculated for an atom in a Lewis structure using the formula: $\text{Formal Charge} = (\text{valence } e^-) - [(\text{nonbonding } e^-) + \frac{1}{2}(\text{bonding } e^-)]$.

Expanded Octet

A hypervalent bonding arrangement in which a central nonmetal atom belonging to Period 3 or below accommodates more than eight valence electrons by utilizing empty $d$ subshells.

Bond Energy

The thermodynamic energy required to cleave one mole of a specific chemical bond in a gaseous substance into neutral fragments.

Bond Length

The measurable internuclear distance separating two covalently bonded atoms at their lowest potential energy position.

VSEPR Theory

Valence Shell Electron Pair Repulsion theory; a structural model postulating that electron groups around a central atom orient themselves as far apart in three-dimensional space as possible to minimize electrostatic repulsions.

Electron Group

Any localized region of high electron density positioned on an atom, which includes a nonbonding lone pair, a single bond, a multiple bond (double or triple), or an unpaired single electron.