General Chemistry (Chang) – Chapters 1–8 Study Notes

Comprehensive vocabulary flashcards covering the essential concepts, laws, formulas, and structural diagrams across Chapters 1 through 8 of General Chemistry.

Matter

Anything that occupies space and has mass.

Scientific Law

A concise statement of a relationship between phenomena that is always the same under the same conditions.

Theory

A unifying principle that explains a body of facts and the laws that are based on them.

Substance

A form of matter that has a definite (constant) composition and distinct properties.

Element

A substance that cannot be separated into simpler substances by chemical means.

Compound

A substance composed of atoms of two or more elements chemically united in fixed proportions that can be separated only by chemical means.

Mixture

A combination of two or more substances in which the substances retain their distinct identities and can be separated by physical methods.

Classification of Matter

The organizational hierarchy dividing matter into mixtures (homogeneous and heterogeneous, separable by physical methods) and substances (compounds and elements, separable by chemical methods).

Extensive Property

A property that depends on how much matter is being considered, such as mass or volume.

Intensive Property

A property that does not depend on the amount of matter being considered, such as density or temperature.

Temperature Scales Comparison

The calibration relationships between the Kelvin, Celsius, and Fahrenheit scales, showing water freezing at $273\,\text{K}$, $0\,^\circ\text{C}$, and $32\,^\circ\text{F}$, and boiling at $373\,\text{K}$, $100\,^\circ\text{C}$, and $212\,^\circ\text{F}$.

Accuracy

The closeness of a measurement to the true value of the quantity that was measured.

Precision

The closeness of agreement between two or more repeated measurements of the same quantity.

Law of Definite Proportions

Different samples of the same compound always contain its constituent elements in the same proportion by mass.

Law of Multiple Proportions

If two elements can combine to form more than one compound, the masses of one element that combine with a fixed mass of the other element are in ratios of small whole numbers.

Rutherford Gold Foil Experiment

An experiment measuring the deflection of $\alpha$ particles fired at a thin gold foil, demonstrating that the atom is mostly empty space with almost all its mass and positive charge concentrated in a dense core called the nucleus.

Atomic Number ($Z$)

The number of protons in the nucleus of each atom of an element, which defines the elemental identity.

Mass Number ($A$)

The total number of neutrons and protons present in the nucleus of an atom of an element.

Isotopes

Atoms that have the same atomic number ($Z$) but different mass numbers ($A$), exhibiting identical chemical properties.

Allotrope

One of two or more distinct forms of the same element, such as oxygen ($\text{O}_2$) and ozone ($\text{O}_3$), or diamond and graphite.

Empirical Formula

A chemical formula that shows which elements are present and the simplest whole-number ratio of their atoms.

Molecular Formula

A chemical formula that shows the exact number of atoms of each element in the smallest unit of a substance.

Compound Nomenclature System

The systematic decision-tree rules for naming ionic compounds using metal cations and mono- or polyatomic anions (with Roman numerals for variable-charge cations) and naming molecular compounds using Greek prefixes.

Oxoacid

An acid that contains hydrogen, oxygen, and another central element.

Hydrate

A compound that has a specific number of water molecules bound to its atoms, such as $\text{CuSO}_4 \cdot 5\text{H}_2\text{O}$.

Atomic Mass Unit (amu)

A mass unit defined as exactly $1/12$ the mass of one carbon-12 atom, equal to $1.661 \times 10^{-24}\,\text{g}$.

Mole (mol)

The amount of a substance that contains as many elementary entities as there are atoms in exactly $12\,\text{g}$ of carbon-12.

Avogadro's Number ($N_\text{A}$)

The number of elementary entities in one mole of any substance, equal to $6.022 \times 10^{23}$.

Molar Mass ($\mathcal{M}$)

The mass in grams of one mole of a substance ($\text{g/mol}$), numerically equal to its atomic or molecular mass in amu.

The Mole Method

The stoichiometric procedure that converts mass of a reactant to moles, applies the stoichiometric mole ratio from a balanced equation to find moles of product, and converts moles of product into mass.

Limiting Reagent

The reactant that is completely consumed first in a reaction, thereby determining and limiting the maximum amount of product that can form.

Theoretical Yield

The amount of product that would result if all the limiting reagent reacted completely according to the balanced chemical equation.

Electrolyte

A substance that, when dissolved in water, yields a solution that can conduct electricity due to dissociation into ions.

Hydration

The process in which an ion is surrounded by water molecules arranged in a specific orientation.

Metathesis Reaction

A double-displacement reaction that involves the exchange of parts between two reacting compounds.

Brønsted Acid

A substance capable of donating a proton ($\text{H}^+$) to another substance.

Brønsted Base

A substance capable of accepting a proton ($\text{H}^+$) from another substance.

Neutralization Reaction

A reaction between an acid and a base that produces a salt and water.

Oxidation

The chemical process in which an atom, molecule, or ion loses one or more electrons.

Reduction

The chemical process in which an atom, molecule, or ion gains one or more electrons.

Oxidation Number

The number of charges an atom would have in a molecule or ionic compound if electrons were transferred completely.

Activity Series

A ranking of metals according to their reducing strength and ability to displace hydrogen or other metal cations from aqueous solution or acids.

Gravimetric Analysis

An analytical technique based on the precise measurement of mass, typically by isolating and weighing a formed precipitate.

Titration

A procedure in which a solution of accurately known concentration (standard solution) is gradually added to an unknown solution until the chemical reaction is complete.

Equivalence Point

The point in a titration at which the added titrant has completely and stoichiometrically reacted with the unknown solute.

Atmospheric Pressure

The pressure exerted by Earth's atmosphere, standardly defined as $1\,\text{atm} = 760\,\text{mmHg} = 760\,\text{torr} = 101{,}325\,\text{Pa}$.

Boyle's Law

The volume of a fixed amount of gas maintained at constant temperature is inversely proportional to the gas pressure ($P \propto \frac{1}{V}$ or $P_1V_1 = P_2V_2$).

Charles's Law

The volume of a fixed amount of gas maintained at constant pressure is directly proportional to its absolute temperature ($V \propto T$ or $\frac{V_1}{T_1} = \frac{V_2}{T_2}$).

Avogadro's Law

At constant pressure and temperature, the volume of a gas is directly proportional to the number of moles of the gas present ($V \propto n$ or $\frac{V_1}{n_1} = \frac{V_2}{n_2}$).

Empirical Gas Laws

Schematic representations demonstrating the direct and inverse relationships of gas volume to pressure (Boyle), temperature (Charles), and mole quantity (Avogadro).

Dalton's Law of Partial Pressures

The total pressure of a mixture of gases is equal to the sum of the partial pressures that each individual gas would exert if it were present alone ($P_\text{total} = \sum P_i$).

Diffusion

The gradual mixing of molecules of one gas with molecules of another by virtue of their random kinetic properties, moving from higher to lower concentration.

Effusion

The process by which a gas under pressure escapes from one compartment through a microscopic hole into an evacuated space.

Graham's Law of Effusion

The rate of effusion or diffusion of a gas is inversely proportional to the square root of its molar mass ($\frac{r_1}{r_2} = \sqrt{\frac{\mathcal{M}_2}{\mathcal{M}_1}}$).

van der Waals Equation

An equation of state modifying the ideal gas law to account for real gas behavior, given by $\left(P + \frac{an^2}{V^2}\right)(V - nb) = nRT$, where $a$ corrects for attractive forces and $b$ corrects for finite molecular volume.

State Function

A thermodynamic property that is determined entirely by the system's present macroscopic state, regardless of how that state was reached (such as $U$, $P$, $V$, $T$, and $H$).

First Law of Thermodynamics

The principle stating that energy can be converted from one form to another but cannot be created or destroyed, expressed mathematically as $\Delta U = q + w$.

Enthalpy ($H$)

A thermodynamic quantity defined by the equation $H = U + PV$, whose change under constant-pressure conditions equals the heat exchanged ($\Delta H = q_p$).

Specific Heat ($s$)

The amount of heat energy required to raise the temperature of one gram of a substance by one degree Celsius ($1\,^\circ\text{C}$).

Standard Enthalpy of Formation ($\Delta H^\circ_\text{f}$)

The heat change that results when one mole of a compound is synthesized from its constituent elements in their standard states at a pressure of $1\,\text{atm}$.

Hess's Law

When reactants are converted to products, the total change in enthalpy is identical whether the reaction takes place in a single step or in a series of steps.

Electromagnetic Radiation Spectrum

The distribution of radiant energy categorized by wavelength and frequency, spanning from short-wavelength, high-frequency gamma rays to long-wavelength, low-frequency radio waves, including the narrow visible band ($400\text{--}700\,\text{nm}$).

Quantum

The smallest discrete quantity of energy that can be emitted or absorbed in the form of electromagnetic radiation, defined by $E = h\nu$.

Photoelectric Effect

A phenomenon in which electrons are ejected from the surface of a metal when exposed to light having a frequency higher than a specific threshold value.

Hydrogen Emission Series

The quantized spectral emission transitions of the hydrogen atom, including the Lyman ($n_\text{f} = 1$, ultraviolet), Balmer ($n_\text{f} = 2$, visible), Paschen ($n_\text{f} = 3$, infrared), and Brackett ($n_\text{f} = 4$, infrared) series.

Heisenberg Uncertainty Principle

It is fundamentally impossible to know simultaneously both the exact momentum and exact position of a subatomic particle.

Atomic Orbital

The wave function ($\psi$) of an electron in an atom, whose square ($\psi^2$) represents the electron density or probability of finding the electron in a specific region of space.

$2p$ Orbitals

A subshell containing three degenerate, mutually perpendicular dumbbell-shaped orbitals designated $2p_x$, $2p_y$, and $2p_z$, enclosing approximately $90\%$ of the electron density.

$3d$ Orbitals

A subshell containing five degenerate boundary surface diagrams, consisting of four cloverleaf shapes ($3d_{x^2-y^2}$, $3d_{xy}$, $3d_{xz}$, $3d_{yz}$) and one donut-encircled dumbbell ($3d_{z^2}$).

Aufbau Principle

The rule dictating that electrons enter and fill atomic subshells in order of increasing orbital energy, following the diagonal order beginning with $1s < 2s < 2p < 3s$.

Pauli Exclusion Principle

No two electrons in the same atom can have the exact same set of four quantum numbers ($n$, $\ell$, $m_\ell$, $m_\text{s}$).

Hund's Rule

The most stable arrangement of electrons in subshells is the one that achieves the maximum number of parallel electron spins by filling degenerate orbitals singly first.

Paramagnetic

A substance that contains one or more unpaired electrons and is consequently attracted into an external magnetic field.

Diamagnetic

A substance in which all electron spins are paired, causing it to be slightly repelled by an external magnetic field.

Representative Elements

The main group elements in Groups 1A through 7A of the periodic table, characterized by incompletely filled $s$ or $p$ subshells of the highest principal quantum level.

Valence Electrons

The outermost electrons of an atom that occupy the highest energy level and are involved in chemical bonding.

Isoelectronic

Chemical species (atoms or ions) that possess the exact same number of electrons and identical ground-state electron configurations.

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

The actual nuclear charge experienced by an electron, equal to the total positive nuclear charge ($Z$) minus the shielding constant ($\sigma$) caused by inner core electrons.

Atomic Radius Trend

Periodic variation defined as one-half the distance between nuclei of adjacent bonded atoms, which systematically decreases across a period from left to right and increases down a group.

First Ionization Energy Trend

The minimum energy in $\text{kJ/mol}$ needed to remove an electron from an isolated gaseous atom in its ground state, showing peaks for noble gases and troughs for alkali metals.

Electron Affinity

The negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion.

Diagonal Relationships

Similarities in chemical behavior between pairs of elements situated diagonally to each other in adjacent periods and groups of the periodic table, specifically $\text{Li}$--$\text{Mg}$, $\text{Be}$--$\text{Al}$, and $\text{B}$--$\text{Si}$.

Amphoteric Oxide

An oxide that exhibits both acidic and basic properties depending on the reaction conditions, exemplified by aluminum oxide ($\text{Al}_2\text{O}_3$).