General Chemistry: Measurements, Chemical Bonding, and Chemical Reactions

Comprehensive vocabulary review flashcards covering units of measurement, significant figures, ionic and covalent bonding, nomenclature, Lewis structures, molecular shapes, intermolecular forces, and chemical reaction stoichiometry and energetics.

Metric and SI Units

The standard international systems of units used in scientific measurement, where length is measured in meters ($\text{m}$), mass in kilograms ($\text{kg}$) for SI and grams ($\text{g}$) for metric, volume in cubic meters ($\text{m}^3$) for SI and liters ($\text{L}$) for metric, temperature in kelvins ($\text{K}$) for SI and degrees Celsius ($^\circ\text{C}$) for metric, and time in seconds ($\text{s}$).

Exact Number

A number obtained by counting discrete items or by an established definition comparing two units within the same system (such as $1\,\text{kg} = 1000\,\text{g}$ or $1\,\text{ft} = 12\,\text{in.}$); they are not measured and do not limit the number of significant figures in a calculation.

Significant Figures

All the certain digits recorded from a measuring instrument plus one final estimated digit; zeros are significant when situated between nonzero digits, at the end of a decimal number, or in scientific notation coefficients, but non-significant as leading zeros or trailing placeholders in whole numbers without decimal points.

Cubic Centimeter ($\text{cm}^3$ or $\text{cc}$)

A derived metric unit of volume representing a cube that is $1\,\text{cm}$ on each edge, which is equivalent in volume to one milliliter ($1\,\text{cm}^3 = 1\,\text{cc} = 1\,\text{mL}$).

Density

A physical property that compares the mass of a substance to the volume it occupies, expressed by the formula $\text{Density} = \frac{\text{mass of a substance}}{\text{volume of a substance}}$ with common units of $\text{g/cm}^3$ or $\text{g/mL}$ for solids and liquids, and $\text{g/L}$ for gases.

Octet Rule

The tendency of representative atoms to lose, gain, or share valence electrons during chemical bonding to achieve a stable noble gas electron arrangement containing eight valence electrons (or two for helium).

Ionic Bond

A chemical bond formed by the strong electrical attractive forces between oppositely charged ions resulting from the complete transfer of one or more valence electrons from a metal atom to a nonmetal atom.

Covalent Bond

A chemical bond that forms when atoms of nonmetals share one or more pairs of valence electrons to achieve stability and attain a noble gas electron configuration.

Cation

A positively charged ion formed when a metal atom with low ionization energy loses one or more of its valence electrons.

Anion

A negatively charged ion formed when a nonmetal atom with high ionization energy gains one or more valence electrons.

Formula Unit

The lowest whole-number ratio of ions in an ionic compound such that the total positive charge balances the total negative charge to equal zero.

Polyatomic Ion

A covalently bonded group of atoms (typically nonmetals bonded to oxygen) that bears an overall electrical charge and behaves as a single distinct ionic unit in chemical reactions.

Lewis Structure

A structural representation of a molecule or polyatomic ion showing how valence electrons are arranged around atoms as bonding pairs (lines or pairs of dots) and nonbonding lone pairs (pairs of dots).

Electronegativity

A relative scale indicating the ability of an atom in a chemical bond to attract shared electrons toward itself, increasing across a period from left to right and decreasing down a group.

Bond Polarity and Electronegativity Differences

The classification of chemical bonds based on electronegativity difference ($\Delta\text{EN}$): nonpolar covalent ($\Delta\text{EN} = 0$ to $0.4$), polar covalent ($\Delta\text{EN} = 0.5$ to $1.8$), and ionic ($\Delta\text{EN} > 1.8$) where electrons are transferred rather than shared.

Dipole

The separation of positive and negative charge centers across a polar covalent bond, denoted by partial charge symbols ($\delta^+$ and $\delta^-$) or a crossed dipole arrow pointing from the positive end to the more electronegative negative end.

VSEPR Theory

Valence Shell Electron-Pair Repulsion theory; a model stating that electron groups (single bonds, multiple bonds, or lone pairs) arrange as far apart as possible around a central atom to minimize electrostatic repulsion, thereby establishing the electron-group geometry and molecular shape.

Ion-Dipole Interaction

A relatively strong electrostatic attractive force between an ion and the oppositely charged partial dipole ends of polar molecules, which drives the dissolution of ionic compounds in polar solvents such as water.

Hydrogen Bond

An exceptionally strong, specialized type of dipole-dipole attraction that occurs when a hydrogen atom covalently bonded to a highly electronegative atom ($\text{F}$, $\text{O}$, or $\text{N}$) is attracted to a lone pair on another $\text{F}$, $\text{O}$, or $\text{N}$ atom.

Dispersion Forces

Weak intermolecular attractive forces present in all molecules caused by temporary, momentary dipoles that develop when electron clouds fluctuate unsymmetrically as molecules collide.

Chemical Change

A process during which original substances are converted into one or more new substances possessing different formulas and properties, commonly evidenced by color changes, gas evolution, precipitate formation, or heat absorption/release.

Law of Conservation of Mass

The fundamental principle stating that matter cannot be created or destroyed during an ordinary chemical reaction; the total mass of the reactants must equal the total mass of the products, and the number of reactant atoms equals the number of product atoms.

The Mole (mol)

The standard SI counting unit for chemical quantities containing Avogadro's number ($6.02 \times 10^{23}$) of formula units, molecules, or atoms.

Avogadro's Number

The exact value $6.02 \times 10^{23}$, representing the total number of particles (atoms, molecules, or ions) present in one mole of any substance, named after the Italian physicist Amedeo Avogadro ($1776\text{--}1856$).

Molar Mass

The mass in grams of one mole of an element or compound, numerically equal to the element's atomic mass or the sum of atomic masses in a compound's formula, with units of $\text{g/mol}$.

Combination Reaction

A reaction classification in which two or more elements or simpler compounds unite chemically to yield a single product ($\text{A} + \text{B} \rightarrow \text{AB}$).

Decomposition Reaction

A reaction classification in which a single reactant compound splits or breaks apart into two or more simpler elements or compounds ($\text{AB} \rightarrow \text{A} + \text{B}$).

Single Replacement Reaction

A chemical reaction in which a free, uncombined element displaces and replaces an element within a reacting compound ($\text{A} + \text{BC} \rightarrow \text{AC} + \text{B}$).

Double Replacement Reaction

A reaction (also termed a metathesis reaction) in which positive and negative ions in two reacting ionic compounds exchange places to produce two different compounds ($\text{AB} + \text{CD} \rightarrow \text{AD} + \text{CB}$).

Combustion Reaction

A reaction where a hydrocarbon or carbon-hydrogen-oxygen compound burns in the presence of oxygen gas to produce carbon dioxide ($\text{CO}_2$), water ($\text{H}_2\text{O}$), and heat energy.

Oxidation and Reduction

Interconnected processes involving electron transfer; oxidation is the loss of electrons (often involving the addition of oxygen or loss of hydrogen), whereas reduction is the gain of electrons (often involving the loss of oxygen or gain of hydrogen).

Limiting Reactant

The reactant substance that is completely used up first in a chemical reaction, thereby restricting and establishing the theoretical maximum amount of product that can be formed.

Excess Reactant

The reactant present in a quantity greater than required to react with the limiting reactant, leaving unreacted leftover material when the reaction concludes.

Heat of Reaction ($\Delta H$)

The enthalpy change occurring at constant pressure during a chemical reaction, defined as the difference between the enthalpies of the products and the reactants: $\Delta H = H_{\text{products}} - H_{\text{reactants}}$.

Endothermic Reaction

A chemical reaction that absorbs energy from its surroundings, where the energy of the products is greater than that of the reactants, resulting in a positive heat of reaction ($\Delta H > 0$).

Exothermic Reaction

A chemical reaction that releases heat energy into the surrounding environment, where the energy of the products is less than that of the reactants, yielding a negative heat of reaction ($\Delta H < 0$).