General Chemistry: Stoichiometry, Atomic Theory, Periodic Trends, and Chemical Bonding

Comprehensive vocabulary flashcards covering general chemistry topics including isotope mass calculation, empirical and molecular formulas, limiting reagents, solution concentration, precipitation rules, quantum mechanics, atomic orbitals, periodic trends, and chemical bonding energetics.

Carbon-12 ($^{12}\text{C}$)

The carbon isotope whose atomic mass is defined as exactly $12.0000\,\text{g/mol}$ (or $\text{amu}$), consisting of $6$ protons and $6$ neutrons, and possessing a relative natural abundance of $98.89\%$.

Average Atomic Mass

The weighted average mass of an element's naturally occurring isotopes, determined using the formula $\sum (\text{isotope mass} \times \text{relative abundance})$.

Empirical Formula

The chemical formula of a compound that expresses the simplest whole-number ratio of the atoms of each element present.

Molecular Formula

The chemical formula that provides the actual number of atoms of each element in a molecule of a compound, calculated by multiplying the empirical formula by the ratio of molecular mass to empirical formula mass.

Mass Percent of an Element

The percentage of a compound's total mass contributed by an individual element, calculated as $\left(\frac{\text{mass of element in 1 mole of compound}}{\text{molar mass of compound}}\right) \times 100\%$.

Limiting Reagent

The reactant that is completely consumed first in a chemical reaction, thereby determining and limiting the maximum theoretical yield of product.

Strong Electrolyte

A solute that dissociates completely into separated ions when dissolved in water, producing a solution that conducts electricity efficiently (e.g., strong acids, strong bases, and soluble salts).

The Six Strong Acids

The group of acids that dissociate completely in aqueous solution: $\text{HCl}(aq)$, $\text{H}_2\text{SO}_4(aq)$, $\text{HBr}(aq)$, $\text{HI}(aq)$, $\text{HNO}_3(aq)$, and $\text{HClO}_4(aq)$.

Weak Electrolyte

A substance that dissolves in water but ionizes only partially to create relatively few ions in solution (e.g., $\text{HF}$, $\text{CH}_3\text{COOH}$, and $\text{NH}_3$).

Non-Electrolyte

A substance that dissolves in water without generating ions, resulting in a solution that does not conduct electricity (e.g., sucrose and ethanol).

Molarity ($M$)

The concentration of a solution expressed as moles of solute per liter of solution: $C = \frac{n}{V}$.

Dilution

The process of decreasing the concentration of a solution by adding solvent, described by the conservation of moles equation $C_1 V_1 = C_2 V_2$.

Precipitate

An insoluble solid that forms and separates from an aqueous solution during a precipitation reaction.

General Solubility Rules

Empirical guidelines stating that alkali metal cations and ammonium ($\text{NH}_4^+$) salts never precipitate, nitrate ($\text{NO}_3^-$) salts are always soluble, and halides ($\text{Cl}^-$, $\text{Br}^-$, $\text{I}^-$) are generally soluble except when paired with $\text{Ag}^+$, $\text{Pb}^{2+}$, or $\text{Hg}_2^{2+}$.

Spectator Ions

Ions present in a reaction mixture that do not participate directly in the chemical change and remain dissolved in identical form before and after the reaction.

Net Ionic Equation

A chemical equation that illustrates only the species directly participating in a chemical reaction (such as precipitate formation), omitting all spectator ions.

Electromagnetic Radiation (EMR)

A form of energy propagating through space via coupled oscillating electric and magnetic fields, governed by the relation $\lambda \nu = c$, where $c = 2.9979 \times 10^8\,\text{m/s}$.

Black-Body Radiation

The continuous thermal electromagnetic radiation emitted by an idealized non-reflective body, the study of which led Max Planck to postulate that energy exchange occurs in discrete packets.

Planck's Constant ($h$)

A fundamental physical constant with a value of $6.6261 \times 10^{-34}\,\text{J}\cdot\text{s}$, relating the energy of a quantum of radiation to its frequency via $\Delta E = n h \nu$.

Quantum of Energy

The smallest discrete quantity of energy that can be emitted or absorbed in the form of electromagnetic radiation, equal to $h\nu$.

Photon

A discrete packet or particle of electromagnetic radiation possessing energy equal to $E_{\text{photon}} = h\nu$.

Photoelectric Effect

The phenomenon in which electrons are ejected from the surface of a metal when irradiated by light of sufficient frequency, described by $E_{\text{photon}} = \Phi + \text{KE}$.

Work Function ($\Phi$)

The minimum threshold energy required for an electron to escape the surface of a specific metal.

De Broglie Wavelength

The wavelength associated with a moving material particle, given by Louis de Broglie's relation $\lambda = \frac{h}{p} = \frac{h}{mv}$.

Wave-Particle Duality

The physical principle asserting that light and matter exhibit both wave-like and particle-like characteristics, evidenced by electron diffraction through a nickel crystal.

Continuous Spectrum

A spectrum, such as that produced by white light, that contains all wavelengths of visible light without interruption.

Discrete Emission Spectrum

A series of distinct, narrow lines of specific wavelengths emitted by energized gas atoms (such as excited $\text{H}_2$), demonstrating that atomic electron energies are quantized.

Bohr Energy Equation

The formula calculating the allowed energy of an electron in a one-electron atom or ion: $E_n = -2.178 \times 10^{-18}\,\text{J} \left(\frac{Z^2}{n^2}\right)$.

Ground State

The lowest possible energy state of an atom or one-electron system, corresponding to the principal quantum level $n = 1$.

Lyman Series

The set of ultraviolet hydrogen emission spectral lines resulting from electron transitions ending at the ground state level $n_f = 1$.

Balmer Series

The set of hydrogen emission spectral lines resulting from electron transitions ending at the $n_f = 2$ level, producing lines in the visible spectrum.

Paschen Series

The set of infrared hydrogen emission spectral lines arising from electron transitions that terminate at the $n_f = 3$ level.

Ionization Energy of One-Electron Species

The energy required to completely remove an electron from its ground state ($n_i = 1$) to an infinite distance ($n_f = \infty$), measured as $+8.712 \times 10^{-18}\,\text{J}$ ($5.25\,\text{MJ/mol}$) for $\text{He}^+$.

Wave Function ($\psi$)

A mathematical solution to the Schrödinger equation ($\hat{H}\psi = E\psi$) describing the wave-like state and spatial distribution of an electron.

Probability Density ($\psi^2$)

The square of the wave function, which represents the probability of finding an electron in the vicinity of the nucleus at a given coordinate.

Heisenberg Uncertainty Principle

The fundamental quantum limit stating that the position and momentum of a particle cannot simultaneously be known with arbitrary precision: $(\Delta x) \cdot (\Delta p) \ge \frac{h}{4\pi}$.

Atomic Orbital

A quantum mechanical wave function solution describing a region of space around the nucleus where an electron has a high probability of being found.

Principal Quantum Number ($n$)

The quantum number ($n = 1, 2, 3, \dots$) that primarily dictates the size and overall energy level of an atomic orbital.

Angular Momentum Quantum Number ($l$)

The quantum number with integer values from $0$ to $n - 1$ that determines the geometric shape of an orbital ($l = 0 \rightarrow s$, $l = 1 \rightarrow p$, $l = 2 \rightarrow d$, $l = 3 \rightarrow f$).

Magnetic Quantum Number ($m_l$)

The quantum number with integer values from $-l$ to $+l$ that specifies the orientation of an orbital in three-dimensional space.

Radial Node

A spherical shell around the nucleus where the probability density $\psi^2$ is zero, calculated for any orbital using the expression $n - l - 1$.

Angular Node

A nodal plane or cone passing through the nucleus where the wave function and probability density equal zero, with the total number equal to $l$.

Degenerate Orbitals

Atomic orbitals that possess identical energy levels, such as all orbitals within the same principal shell ($n$) for one-electron species like hydrogen.

Spin Quantum Number ($m_s$)

The fourth quantum number that describes the intrinsic spin state of an electron, taking values of either $+\frac{1}{2}$ (spin up) or $-\frac{1}{2}$ (spin down).

Pauli Exclusion Principle

The rule stating that no two electrons in an atom can have the same four quantum numbers, which limits any individual atomic orbital to a maximum of two electrons with opposite spins.

Aufbau Principle

The process by which electrons are progressively added into the lowest-energy available hydrogen-like orbitals as protons are added to the nucleus.

Hund's Rule

The principle stating that for degenerate orbitals, the lowest-energy electron configuration is achieved by maximizing the number of unpaired electrons with parallel spins.

Chromium Electron Configuration Anomaly

An exception to expected orbital filling where chromium adopts an electron configuration of $[\text{Ar}]\,4s^1 3d^5$ rather than $[\text{Ar}]\,4s^2 3d^4$ to maximize subshell stability.

Copper Electron Configuration Anomaly

An exception to expected orbital filling where copper adopts an electron configuration of $[\text{Ar}]\,4s^1 3d^{10}$ rather than $[\text{Ar}]\,4s^2 3d^9$.

Orbital Penetration

The ability of an electron in an outer orbital to be found close to the nucleus due to an inner radial probability bump, causing penetration order $ns > np > nd > nf$ and splitting orbital energy levels ($E_{ns} < E_{np} < E_{nd} < E_{nf}$).

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

The net positive nuclear charge experienced by an electron in a multi-electron atom, which increases from left to right across a period because electrons in the same principal energy level are ineffective at shielding one another.

Atomic Radius Trend

The periodic trend wherein atomic radius decreases from left to right across a period due to increasing $Z_{\text{eff}}$ drawing electrons closer, and increases down a group.

Ionic Radii Comparison

The principle that cations are smaller than their parent neutral atoms due to increased $Z_{\text{eff}}$, while anions are larger than their neutral atoms due to increased electron-electron repulsion within the same shell.

Ionization Energy

The energy required to remove an electron from a gaseous atom or ion in its ground state ($X(g) \rightarrow X^+(g) + e^-$), which is always a positive (endothermic) value.

Electron Affinity

The energy change associated with adding an electron to a gaseous atom; fluorine exhibits an anomaly where electron affinity is less favorable than expected due to strong electron repulsions in its compact $2p$ orbital.

Ionic Bonding

The type of chemical bonding that occurs between a metal cation and a nonmetal anion held together by mutual electrostatic attraction resulting from electron transfer.

Covalent Bonding

A chemical bond formed when two atoms share valence electrons, stabilized by the simultaneous electrostatic attraction of both nuclei for the shared electron pair.

Electronegativity

The ability of an atom in a molecule to attract shared electrons to itself, quantified on the Linus Pauling scale with fluorine being the most electronegative element ($4.0$).

Dipole Moment ($\mu$)

The quantitative product of separated charge and distance in a polar bond or molecule: $\mu = Q \times r$, conventionally measured in Debye units ($1\,\text{Debye} = 3.336 \times 10^{-30}\,\text{C}\cdot\text{m}$).

Lattice Energy ($\Delta H_{\text{LE}}$)

The enthalpy change associated with the condensation of gaseous ions into an ionic solid ($M^+(g) + X^-(g) \rightarrow MX(s)$), maximizing electrostatic attraction and minimizing repulsion.

Born-Haber Cycle

A thermodynamic cycle that utilizes Hess's law to calculate the lattice energy of an ionic compound by relating it to standard enthalpy of formation, sublimation, ionization energy, bond dissociation energy, and electron affinity.

Enthalpy of Sublimation ($\Delta H_{\text{sub}}$)

The energy required to convert one mole of a substance directly from the solid state to the gas state (e.g., $\Delta H_{\text{sub, Na}} = 108\,\text{kJ/mol}$ in the formation of $\text{NaBr}$).