ACS General Chemistry Practice Cards

Comprehensive set of vocabulary flashcards for ACS General Chemistry Second Edition preparation, covering foundational concepts, structure, bonding, states of matter, kinetics, equilibrium, and thermodynamics.

Isotopes

Atoms with the same number of protons but different numbers of neutrons.

Ionic Nomenclature (Main Group)

Named by the metal first followed by the nonmetal name with the suffix -ide (e.g., sodium sulfide).

Ionic Nomenclature (Transition Metals)

Named by the metal followed by the charge in Roman numerals in parentheses (e.g., iron(III) oxide).

Binary Acids

Acids containing hydrogen and a monatomic anion; named with the prefix hydro- and the suffix -ic acid.

Oxoacids (from -ate anions)

Acids containing hydrogen and a polyatomic anion ending in -ate; suffix changes to -ic acid.

Oxoacids (from -ite anions)

Acids containing hydrogen and a polyatomic anion ending in -ite; suffix changes to -ous acid.

Density

The ratio of mass to volume, often expressed in $\text{g} \cdot \text{cm}^{-3}$ or $\text{g} \cdot \text{mL}^{-1}$.

Compound

A pure substance containing two or more elements chemically combined/bonded together.

Mixture

Two or more pure substances physically combined but not chemically bonded.

Unit Conversion: $\text{kJ}$ to $\text{J}$

$$1\,kJ = 1000\,J$$

Unit Conversion: $\mu s$ to $\text{s}$

$$1\,\mu s = 10^{-6}\,s$$

Significant Figures in Measurement

Includes all known digits plus one estimated digit (e.g., reading to the bottom of the meniscus).

Atomic Number ($Z$)

The number of protons in the nucleus of an atom.

Mass Number ($A$)

The total number of protons and neutrons in the nucleus ($A = Z + n$).

Ion

An atom that has gained or lost electrons, resulting in a net electrical charge.

Average Atomic Mass

The weighted average of the masses of all naturally occurring isotopes of an element.

Diatomic Elements

Elements that naturally exist as molecules consisting of two atoms: $H_2, N_2, O_2, F_2, Cl_2, Br_2, I_2$.

Alkali Metals

Group 1 elements on the periodic table (excluding Hydrogen).

Alkaline Earth Metals

Group 2 elements on the periodic table; commonly form $+2$ ions.

Halogens

Group 17 elements on the periodic table.

Noble Gases

Group 18 elements on the periodic table; characterized by full valence shells.

Metalloids

Elements such as Si, Ge, and As that have properties intermediate between metals and nonmetals.

Rydberg Formula

Used to calculate the energy change of an electronic transition: $\Delta E = R_H (\frac{1}{n_{i}^2} - \frac{1}{n_{f}^2})$.

Photon Energy Equation

$E_{photon} = \frac{hc}{\lambda}$ where $h$ is Planck's constant and $c$ is the speed of light.

Absorption

A process where an electron is promoted to a higher energy level by taking in a photon.

Emission

A process where an electron falls to a lower energy level and releases a photon.

Principal Quantum Number ($n$)

Indicates the main energy level or shell; values are integers $1, 2, 3 \dots$

Angular Momentum Quantum Number ($l$)

Indicates the orbital shape ($s=0, p=1, d=2, f=3$); values range from $0$ to $n-1$.

Magnetic Quantum Number ($m_l$)

Indicates the orientation of the orbital; values range from $-l$ to $+l$.

Spin Quantum Number ($m_s$)

Indicates the direction of electron spin; values are either $+\frac{1}{2}$ or $-\frac{1}{2}$.

Valence Electrons

Electrons in the outermost energy level of an atom used for bonding.

Aufbau Principle

Electrons fill lower energy orbitals first before moving to higher ones.

Hund's Rule

Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron.

Pauli Exclusion Principle

No two electrons in an atom can have the same four quantum numbers.

Paramagnetic

A species containing at least one unpaired electron; attracted to a magnetic field.

Diamagnetic

A species where all electrons are paired; weakly repelled by a magnetic field.

Effective Nuclear Charge ($Z_{eff}$)

The actual positive charge experienced by an electron in a multi-electron atom ($Z_{eff} = Z - S$).

Atomic Radius Trend

Increases down a group (more shells) and decreases across a period (higher $Z_{eff}$).

Ionic Radius (Cations)

Cations are always smaller than their parent neutral atoms due to loss of electron shell or reduced repulsion.

Ionic Radius (Anions)

Anions are always larger than their parent neutral atoms due to increased electron-electron repulsion.

First Ionization Energy

The energy required to remove the first electron from a gaseous atom.

Electronegativity

A measure of the ability of an atom in a chemical compound to attract shared electrons.

Electron Affinity

The energy change that occurs when an electron is acquired by a neutral atom.

Mole (mol)

The SI unit for amount of substance; contains $6.022 \times 10^{23}$ particles.

Avogadro's Number ($N_A$)

$$6.022 \times 10^{23}\,mol^{-1}$$

Molar Mass ($M$)

The mass in grams of one mole of a substance ($g \cdot mol^{-1}$).

Empirical Formula

The simplest whole-number ratio of atoms of each element in a compound.

Molecular Formula

The actual number of atoms of each element in a molecule of a compound.

Percent Composition by Mass

(\text{mass of element} / \text{total mass of sample}) \times 100\%

Stoichiometry

The study of quantitative relationships between reactants and products in a chemical reaction.

Limiting Reactant

The reactant that is completely consumed first in a reaction, determining the maximum product formed.

Theoretical Yield

The maximum amount of product that can be produced from a given amount of reactant based on stoichiometry.

Percent Yield

(\text{actual yield} / \text{theoretical yield}) \times 100\%

Strong Electrolyte

A solute that dissociates or ionizes completely in water, conducting electricity well.

Weak Electrolyte

A solute that only partially ionizes in water, conducting electricity poorly (e.g., weak acids/bases).

Molarity (M)

A concentration unit defined as $\text{moles of solute} / \text{liters of solution}$.

Dilution Formula

$M_1 V_1 = M_2 V_2$; moles of solute remain constant during dilution.

Precipitate

An insoluble solid that forms and settles out of a liquid mixture.

Net Ionic Equation

A chemical equation that shows only those elements and compounds directly involved in the chemical change.

Spectator Ions

Ions that do not participate in a reaction and appear unchanged on both sides of a total ionic equation.

Oxidation

The loss of electrons or an increase in oxidation state.

Reduction

The gain of electrons or a decrease in oxidation state.

Oxidizing Agent

The reactant that causes oxidation in another species and is itself reduced.

Reducing Agent

The reactant that causes reduction in another species and is itself oxidized.

First Law of Thermodynamics

$\Delta E = q + w$, where the energy change equals heat plus work.

Enthalpy ($H$)

A state function representing the heat content of a system at constant pressure.

Exothermic

A process that releases heat to its surroundings (signed \Delta H < 0).

Endothermic

A process that absorbs heat from its surroundings (signed \Delta H > 0).

Specific Heat Capacity ($c$)

The amount of heat required to raise the temperature of 1 gram of a substance by $1\,K$ (or $1\,^{\circ}C$).

Hess's Law

The total enthalpy change for a reaction is the sum of all changes, independent of the pathway.

Standard Enthalpy of Formation ($\Delta H_f^\circ$)

The enthalpy change when 1 mole of a compound is formed from its elements in their standard states.

Lattice Energy

The energy required to separate 1 mole of a solid ionic compound into its gaseous ions.

Polar Covalent Bond

A bond between nonmetals with a moderate difference in electronegativity ($0.4$ to $1.6$).

Formal Charge

$\text{valence electrons} - [\text{lone pair electrons} + \frac{1}{2}(\text{bonding electrons})]$.

Octet Rule

The tendency of main group atoms to form enough bonds to obtain eight valence electrons.

Resonance Structures

Two or more valid Lewis structures that can be drawn for a single molecule; the actual structure is an average.

Bond Enthalpy

The energy required to break 1 mole of a specific bond in the gas phase.

Bond Order

The number of chemical bonds between a pair of atoms (e.g., 1 for single, 2 for double).

VSEPR Theory

Valence Shell Electron Pair Repulsion; used to predict molecular geometry based on electron group minimization.

Trigonal Planar

Geometry with 3 electron groups and 0 lone pairs; bond angle is $120^{\circ}$.

Tetrahedral

Geometry with 4 electron groups and 0 lone pairs; bond angle is $109.5^{\circ}$.

Hybridization

The mixing of atomic orbitals to form new hybrid orbitals suitable for bonding (e.g., $sp, sp^2, sp^3$).

Sigma ($\sigma$) Bond

A covalent bond formed by the end-to-end overlap of atomic orbitals.

Pi ($\pi$) Bond

A covalent bond formed by the side-by-side overlap of unhybridized p-orbitals.

Ideal Gas Law

$$PV = nRT$$

STP (Standard Temperature and Pressure)

$0\,^{\circ}C$ ($273.15\,K$) and $1\,atm$ (or $1\,bar$).

Dalton's Law of Partial Pressures

The total pressure of a gas mixture is the sum of the partial pressures of its components.

Kinetic Molecular Theory

Describes gas particles as being in constant motion, having negligible volume, and no attractions.

Hydrogen Bonding

A strong dipole-dipole attraction when H is bonded to N, O, or F.

London Dispersion Forces

Weak, temporary intermolecular forces present in all atoms and molecules due to electron fluctuations.

Vapor Pressure

The pressure exerted by a vapor in equilibrium with its liquid or solid phase.

Normal Boiling Point

The temperature at which a liquid's vapor pressure equals $1\,atm$.

Phase Diagram

A graph showing the conditions of temperature and pressure at which a substance exists as a solid, liquid, or gas.

Critical Point

The temperature and pressure above which a substance exists as a supercritical fluid.

Triple Point

The specific temperature and pressure where all three phases coexist in equilibrium.

Molality ($m$)

A concentration unit defined as $\text{moles of solute} / \text{kilograms of solvent}$.

Mole Fraction ($X$)

The ratio of the moles of one component to the total moles of all components in a mixture.

Colligative Properties

Properties of a solution that depend only on the number of solute particles, not their identity.

Freezing Point Depression

$$\Delta T_f = i K_f m$$

Boiling Point Elevation

$$\Delta T_b = i K_b m$$