General Chemistry: Atomic Structure, Quantum Mechanics, and Chemical Principles

Vocabulary practice flashcards covering stoichiometric calculations, the mole concept, chemical equations, Bohr theory, atomic spectra, wave mechanics, and electron configurations.

Atomic Mass Unit (amu)

A unit of mass defined as exactly $\frac{1}{12}$ the mass of a single unbound atom of carbon-12 ($^{12}_6C$), equal to $1.660538921 \times 10^{-24}\,g$ (or $1.6605 \times 10^{-27}\,kg$); also referred to as a dalton ($Da$).

Mole (mol)

The SI base unit for the amount of substance, defined as containing exactly the same number of elementary entities (atoms, molecules, or ions) as there are atoms in $12\,g$ of carbon-12 ($^{12}_6C$).

Avogadro's Number ($N_A$)

The proportionality factor representing the number of elementary entities per mole of substance, defined as $6.0221 \times 10^{23}\,mol^{-1}$, connecting total entities $N$ to the amount in moles $n$ via the relationship $N = n \cdot N_A$.

Molecular Weight (MW)

The total mass of a single molecule, calculated as the sum of the atomic masses of each constituent atom in the molecular formula, expressed in atomic mass units ($amu$).

Molar Mass (M)

The mass of a substance divided by its amount of substance, expressed in units of $g\,mol^{-1}$, corresponding to the mass per mole of atoms for elements, molecules for molecular compounds, or formula units for ionic compounds.

Empirical Formula

A chemical formula showing the simplest whole-number ratio of the elements present in a compound, rather than the absolute number of atoms in a molecule.

Molecular Formula

A chemical formula stating the actual, exact number of atoms of each element present in a single molecule of a compound.

Mass Percentage Composition

The percentage of a compound's total mass contributed by a specific element, calculated as $\frac{\text{Mass of element in sample}}{\text{Total mass of sample}} \times 100\%$.

Law of Conservation of Mass

A fundamental principle formulated by Antoine Lavoisier in 1785 stating that mass is neither created nor destroyed during a chemical reaction, meaning the total mass of the products must equal the total mass of the reactants.

Stoichiometric Coefficient

A numerical multiplier written directly in front of a chemical formula in a balanced reaction equation to satisfy mass and charge conservation.

Non-Redox Reaction

A chemical reaction in which the oxidation numbers of all atoms remain unchanged from reactants to products, commonly exemplified by ion exchange and acid-base neutralization reactions.

Net Ionic Equation

A simplified chemical equation that displays only the ions and compounds actively participating in a chemical change, omitting spectator ions that remain dissolved in identical form on both sides.

Atomic Emission Spectrum

A characteristic, non-continuous pattern of discrete wavelengths of electromagnetic radiation emitted by atoms or molecules when electrons transition from higher energy states down to lower energy states.

Bohr Model of the Atom

A model introduced by Niels Bohr proposing that electrons orbit the nucleus in discrete, stationary circular orbits of fixed, quantized energy, emitting or absorbing light only when jumping between levels.

Rydberg Equation

An empirical formula relating the wavelengths of spectral lines for hydrogen to transitions between energy levels: $\frac{1}{\lambda} = R_H \left(\frac{1}{n_1^2} - \frac{1}{n_2^2}\right)$, where $R_H = 109737\,cm^{-1}$.

Balmer Series

The series of hydrogen emission lines in the visible region of the spectrum that arises when excited electrons fall from higher levels ($n \ge 3$) to the principal quantum level $n = 2$.

Lyman Series

The ultraviolet spectral series of hydrogen emission lines produced when electrons make transitions from higher energy levels ($n \ge 2$) down to the ground state level $n = 1$.

Photoelectric Effect

The emission of electrons from a metallic surface when exposed to electromagnetic radiation at or above a specific threshold frequency ($\nu \ge \nu_0$).

de Broglie Hypothesis

The proposal by Louis de Broglie in 1924 that all matter exhibits wave-particle duality, with a matter wavelength defined by $\lambda = \frac{h}{mv}$, where $h$ is Planck's constant, $m$ is mass, and $v$ is velocity.

Heisenberg Uncertainty Principle

A quantum mechanical theorem stating that it is impossible to simultaneously measure both the position and momentum of a subatomic particle with arbitrary precision, governed by $m\Delta v \cdot \Delta x \ge \frac{h}{4\pi}$.

Schrödinger Equation

The foundational differential wave equation of quantum mechanics, expressed as $\hat{H}\psi = E\psi$, where $\hat{H}$ is the Hamiltonian operator containing kinetic and potential energy operators, describing the wave state and allowed energy levels of a particle.

Probability Density ($\psi^2$)

The square of the wavefunction ($\psi^2$), which represents the mathematical probability of finding an electron at a particular location in space around the nucleus.

Node

A point, line, or surface in an orbital where the electron wavefunction passes through zero ($\psi = 0$), corresponding to an exact probability density of zero ($\psi^2 = 0$).

Principal Quantum Number ($n$)

A positive integer ($n = 1, 2, 3, \dots$) that primarily dictates the overall size of an atomic orbital and the principal energy level of the electron.

Angular Momentum Quantum Number ($l$)

An integer quantum number ranging from $0$ to $n - 1$ that specifies the subshell and geometric shape of an orbital ($l = 0$ is $s$, $l = 1$ is $p$, $l = 2$ is $d$, and $l = 3$ is $f$).

Magnetic Quantum Number ($m_l$)

An integer quantum number ranging from $-l$ to $+l$ that defines the three-dimensional spatial orientation of an orbital relative to Cartesian axes.

Spin Quantum Number ($m_s$)

A quantum number describing the intrinsic spin magnetic moment of an electron, taking values of either $+\frac{1}{2}$ (spin up) or $-\frac{1}{2}$ (spin down).

p Orbitals

A set of three mutually perpendicular, two-lobed atomic orbitals designated $p_x, p_y,$ and $p_z$ ($l = 1$), each featuring a nodal plane passing through the nucleus.

Degenerate Orbitals

Atomic orbitals that possess identical energy levels, such as the three $2p$ subshell orbitals in an isolated hydrogen atom.

Aufbau Principle

The rule dictating that electrons fill lower-energy atomic orbitals completely before filling higher-energy orbitals during ground-state electron configuration buildup.

Pauli Exclusion Principle

The quantum rule stating that no two electrons in an atom can share the same four quantum numbers, which limits any single orbital to at most two electrons with opposing spins.

Hund's Rule

The rule stating that when occupying degenerate orbitals, electrons enter singly with parallel spins first before pairing up, maximizing total spin multiplicity.

Valence Electrons

The electrons occupying the outermost energy level (highest $n$ value) of an atom that participate in chemical reactions, determine reactivity, and form bonds.

Core Electrons

Inner electrons filling lower energy shells corresponding to the preceding noble gas configuration, which do not participate in chemical bonding.

Lewis Dot Structure

A visual chemical representation introduced in 1910 that denotes valence electrons as dots surrounding an element's chemical symbol to illustrate potential bonding and stable octets.