Chemistry 101: The Complete Notes

Vocabulary flashcards covering key terms and concepts from the lecture notes.

Matter

Anything that has mass and occupies space (volume).

Mass

A measure of the amount of matter in an object; SI unit is the kilogram.

Volume

The amount of space occupied by a substance; common SI unit is the liter (L) or cubic meters (m^3).

Density

Mass per unit volume (D = m/V); common units g/cm^3 or kg/m^3.

State of Matter

The physical form of a substance: solid, liquid, or gas.

Solid

A state of matter with definite shape and volume.

Liquid

A state of matter with definite volume but takes the shape of its container.

Gas

A state of matter with no definite shape or volume; fills its container.

The Components of Matter

Fundamental categories in chemistry: elements, compounds, and mixtures.

Element

A pure substance that cannot be broken down into simpler substances by chemical means.

Compound

A substance composed of two or more elements chemically bonded in a fixed ratio.

Molecule

Two or more atoms bonded together; the smallest unit of a compound with the same chemical properties.

Atom

The basic unit of matter that makes up elements; consists of a nucleus and electrons.

Nucleus

Center of the atom containing protons and neutrons.

Proton

Positively charged subatomic particle in the nucleus; mass ~1 amu.

Neutron

Electrically neutral subatomic particle in the nucleus; mass ~1 amu.

Electron

Negatively charged subatomic particle surrounding the nucleus; very small mass.

Atomic Number (Z)

Number of protons in the nucleus; defines the identity of an element.

Mass Number (A)

Total number of protons and neutrons in a given isotope.

Isotope

Atoms of the same element (same Z) with different neutron numbers (A).

Complete Atomic Symbol

Notation showing Z and A for an isotope (e.g., 14C); sometimes includes the element symbol.

Element

A substance consisting of atoms with the same atomic number.

Empirical Formula

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

Molecular Formula

The actual number and type of atoms in a molecule.

Molecule

Two or more atoms bonded together; a discrete unit of a compound.

Ionic Bond

Electrostatic attraction between oppositely charged ions formed by transfer of electrons.

Covalent Bond

Bond formed by sharing electron pairs between nonmetals.

Octet Rule

Atoms tend to have eight electrons in their valence shell (with some exceptions).

Lewis Structure

Diagram showing all valence electrons and bonds around atoms.

Valence Electrons

Electrons in the outermost shell that participate in bonding.

Electron Configuration

Arrangement of electrons in an atom’s orbitals; follows Aufbau/Hund/Pauli principles.

Orbital

Region of space around the nucleus where electrons are likely to be found.

Quantum Numbers (n, l, ml, ms)

Describe the size (n), shape (l), orientation (ml), and spin (ms) of an orbital.

Aufbau Principle

Electrons fill lowest-energy orbitals first before occupying higher ones.

Hund’s Rule

Electrons fill degenerate orbitals singly before pairing.

Pauli Exclusion Principle

No more than two electrons can occupy an orbital with opposite spins.

Hybridization

Mixing of atomic orbitals (e.g., sp, sp2, sp3) to form equal-energy hybrids.

VSEPR Theory

Valence Shell Electron Pair Repulsion; shapes of molecules determined by repulsion of electron pairs.

Molecular Geometry

Three-dimensional arrangement of atoms in a molecule.

Electronic Geometry

Arrangement of electron domains around a central atom (affects shape).

Hydrogen Bond

Strong dipole-dipole attraction involving hydrogen bonded to N, O, or F.

Dipole-Dipole Interaction

Attractive forces between polar molecules due to permanent dipoles.

London Dispersion Forces

Weak intermolecular forces arising from temporary dipoles; present in all molecules.

Solubility Rules

Guidelines for predicting whether substances will dissolve in water (solubility) based on ions and functional groups.

Molarity (M)

Concentration defined as moles of solute per liter of solution.

Stock Solution

A solution prepared at a known concentration for use in experiments.

Dilution

Process of reducing concentration; governed by C1V1 = C2V2.

Molar Mass / Formula Weight (FW)

Mass of one mole of a substance in grams per mole (g/mol) or amu per molecule.

Mole

Amount of substance containing 6.02 x 10^23 particles (Avogadro’s number).

Avogadro’s Number (NA)

NA = 6.02 x 10^23; the number of particles in one mole.

Percent Composition

Mass fraction of each element in a compound, expressed as a percentage.

Empirical vs Molecular Formulas

Empirical: simplest whole-number ratio; Molecular: actual number of atoms in a molecule.

Mass Spectrometry

Analytical technique to measure masses and abundances of isotopes and fragments.

Isotopes

Atoms of the same element with different numbers of neutrons; different mass numbers.

The Mole (Mol)

A counting unit for atoms or molecules; 1 mol = 6.02 x 10^23 objects.

Stoichiometry

Relationships among quantities of reactants and products in a chemical reaction.

Balanced Chemical Equation

Equation with equal numbers of each type of atom on both sides.

Limiting Reactant

Reactant that is consumed first, limiting the extent of the reaction.

Theoretical Yield

Maximum amount of product that could be formed from given reactants.

Actual Yield

Measured amount of product obtained from a reaction.

Percent Yield

Actual yield ÷ Theoretical yield × 100%.

Hess’s Law

Total enthalpy change for a reaction equals the sum of enthalpy changes of steps.

Standard Enthalpy of Formation (ΔHf)

Enthalpy change when 1 mole of a substance forms from its elements in standard state.

Calorimetry

Technique to measure heat transferred during chemical or physical changes.

q = Cp m ΔT

Heat transfer equation used in calorimetry (solution).

Exothermic

Process that releases heat to the surroundings; ΔH is negative.

Endothermic

Process that absorbs heat from the surroundings; ΔH is positive.

Gases: STP Volume

At STP (1 atm, 273 K), 1 mole of gas occupies 22.4 L (molar volume).

Ideal Gas Law

PV = nRT; formula relating pressure, volume, amount, temperature, and R.

Partial Pressure

Pressure contributed by a single gas in a mixture.

Mole Fraction

Moles of a component divided by total moles in a gas mixture.

Solubility Rules (Ion Chart)

Guidelines predicting solubility of ionic compounds based on ions.

Law of Conservation of Mass

Matter is neither created nor destroyed in a chemical reaction.

Redox (Reduction-Oxidation)

Chemical processes involving transfer of electrons; oxidation is loss, reduction is gain.

OIL RIG

Mnemonic: Oxidation Is Loss of electrons; Reduction Is Gain of electrons.

Oxidation State / Oxidation Number

Hypothetical charge assigned to an atom in a compound or ion.

Arrhenius Acids and Bases

Acids donate H+ in solution; bases donate OH− in solution.

Molarity vs Normality

Molarity is moles per liter; normality relates to equivalents per liter.

Solute, Solvent, Solution

Solute is the substance dissolved; solvent is the dissolving medium; solution is the mixture.

Empirical Formula from Percent Composition

Convert percentages to masses, find moles, and determine the simplest whole-number ratio.