Chem 1411 – Final Exam Review

Mass (SI Base Unit)

Measured in kilograms (kg); the amount of matter in an object.

Length (SI Base Unit)

Measured in meters (m); the one-dimensional distance between two points.

Volume (Derived Unit)

Space occupied by matter; common units are cubic meters (m³) or liters (1 L = 1000 cm³).

Metric Prefixes

Prefixes that scale SI units by powers of ten (e.g., kilo- 10³, centi- 10⁻², milli- 10⁻³, micro- 10⁻⁶, nano- 10⁻⁹).

Precision

How closely repeated measurements agree with one another.

Accuracy

How close a measurement is to the true or accepted value.

Systematic Error

Consistent, repeatable error associated with faulty equipment or bias; affects accuracy.

Random Error

Unpredictable variations that arise from limitations of measurement; affects precision.

Significant Figures

Digits in a number that convey measured certainty, including all certain digits plus the first uncertain digit.

Exact Number

A value known with no uncertainty (e.g., counting numbers, defined constants); regarded as having infinite significant figures.

Density (ρ)

Mass per unit volume; common units g cm⁻³ or g mL⁻¹.

Law of Conservation of Energy

Energy cannot be created or destroyed, only transformed.

Joule (J)

SI unit of energy; 1 J = 1 kg m² s⁻².

Endothermic Process

A change that absorbs heat from surroundings (ΔH > 0).

Exothermic Process

A change that releases heat to surroundings (ΔH < 0).

Dimensional Analysis

Problem-solving method that uses unit factors to convert between quantities; ensures answer has correct units.

Law of Conservation of Mass

Mass is neither created nor destroyed in a chemical reaction.

Law of Definite Proportions

A given compound always contains the same elements in the same mass ratio.

Law of Multiple Proportions

When two elements form more than one compound, mass ratios of the second element that combine with a fixed mass of the first are small whole numbers.

Dalton’s Atomic Theory

Matter is composed of tiny, indivisible atoms; atoms of an element are identical; compounds form from fixed atom ratios; reactions rearrange atoms without changing them.

Cathode-Ray Experiment

J. J. Thomson showed that atoms contain negatively charged electrons and measured charge-to-mass ratio.

Oil-Drop Experiment

R. A. Millikan determined the charge of the electron (−1.602 × 10⁻¹⁹ C).

Radioactivity

Spontaneous emission of particles or rays from atomic nuclei.

Gold Foil Experiment

Rutherford’s α-particle scattering showed atoms have a small, dense, positively charged nucleus.

Proton (p⁺)

Subatomic particle in nucleus; mass ≈ 1 amu; charge +1.

Neutron (n⁰)

Neutral nuclear particle; mass ≈ 1 amu; charge 0.

Electron (e⁻)

Negatively charged particle outside nucleus; mass ≈ 0.00055 amu; charge −1.

Atomic Number (Z)

Number of protons in nucleus; defines the element.

Mass Number (A)

Sum of protons + neutrons in a nucleus.

Isotope

Atoms of same element (same Z) with different mass numbers (different neutrons).

Ion

Charged atom or group; cation (+) loses electrons, anion (−) gains electrons.

Relative Abundance

Percent of each isotope occurring naturally for an element.

Atomic Mass

Weighted average mass of an element’s isotopes in amu.

Mole (mol)

Amount containing 6.022 × 10²³ entities (Avogadro’s number).

Molar Mass

Mass of one mole of a substance (g mol⁻¹); numerically equals formula weight in amu.

Wavelength (λ)

Distance between successive crests of a wave; units m, nm.

Frequency (ν)

Number of wave cycles per second; units s⁻¹ (Hz).

Electromagnetic Radiation

Oscillating electric and magnetic fields traveling at the speed of light.

Electromagnetic Spectrum

Continuous range of electromagnetic radiation from γ-rays to radio waves.

Photon

Quantum of electromagnetic energy; energy E = hν.

Photoelectric Effect

Ejection of electrons from metal when light above a threshold frequency shines on it.

Emission Spectrum

Discrete lines of light emitted by excited atoms or ions.

Rydberg Equation

1/λ = R_H (1/n₁² − 1/n₂²); predicts wavelengths of H-atom spectral lines.

Heisenberg Uncertainty Principle

Impossible to know exact position and momentum of a particle simultaneously (Δx Δp ≥ h/4π).

de Broglie Relationship

λ = h/p; matter exhibits wave properties.

Quantum Numbers

Set (n, ℓ, mℓ, ms) that uniquely describe an electron in an atom.

s-Orbital

Spherical orbital; ℓ = 0.

p-Orbital

Dumbbell-shaped orbital; ℓ = 1.

d-Orbital

Four-lobed (or donut-with-dumbbell) orbital; ℓ = 2.

Periodic Law

Element properties recur periodically when elements are arranged by increasing atomic number.

Electron Configuration

Notation showing distribution of electrons among orbitals (e.g., 1s² 2s² 2p⁶).

Ground State

Lowest-energy arrangement of electrons in an atom.

Excited State

Any electron configuration with higher energy than the ground state.

Pauli Exclusion Principle

No two electrons in an atom can have the same four quantum numbers; max two electrons per orbital with opposite spins.

Degenerate Orbitals

Orbitals with the same energy within a subshell.

Shielding

Reduction of nuclear attraction on an electron due to other electrons.

Penetration

Proximity of an orbital’s electron density to the nucleus; s > p > d > f for same n.

Effective Nuclear Charge (Z_eff)

Net positive charge experienced by valence electrons; Z_eff = Z − shielding.

Aufbau Principle

Electrons fill orbitals from lowest to highest energy.

Hund’s Rule

Electrons occupy degenerate orbitals singly with parallel spins before pairing.

Valence Electrons

Electrons in outermost shell involved in bonding.

Core Electrons

Inner electrons not involved in bonding.

Noble Gases

Group 18 elements with full valence shells and low reactivity.

Metals

Elements that are shiny, malleable, conductive, and tend to form cations.

Nonmetals

Elements that are brittle (if solid), nonconductive, and tend to form anions or covalent bonds.

Metalloids

Elements with mixed metal/nonmetal properties; semiconductors.

Alkali Metals

Group 1 elements; highly reactive, form 1⁺ ions.

Alkaline Earth Metals

Group 2 elements; form 2⁺ ions.

Halogens

Group 17 nonmetals; form 1⁻ ions and diatomic molecules.

Transition Metals

d-block elements; often form multiple oxidation states.

Electronegativity (EN)

Ability of an atom in a bond to attract shared electrons; increases across a period, decreases down a group.

Polar Covalent Bond

Covalent bond with unequal electron sharing due to EN difference (0.4 – 1.9).

Ionic Bond

Electrostatic attraction between cations and anions; EN difference ≳ 2.0.

Percent Ionic Character

Measure comparing actual bond dipole to a fully ionic bond dipole.

Lewis Dot Symbol

Representation of valence electrons as dots around an element symbol.

Lone Pair

Pair of valence electrons not involved in bonding.

Bonding Pair

Shared pair of electrons forming a covalent bond.

Octet Rule

Main-group atoms tend to attain eight valence electrons (duet for H, He).

Formal Charge

FC = valence − (nonbonding e⁻) − ½(bonding e⁻); helps select best Lewis structure.

Resonance

Two or more equivalent Lewis structures differing only in electron placement; actual structure is a resonance hybrid.

Bond Energy (D)

Energy required to break one mole of a bond in gas phase; always positive.

Bond Length

Average distance between nuclei of bonded atoms; decreases as bond order increases.

VSEPR Theory

Predicts molecular shapes based on repulsion between electron groups around a central atom.

Dipole Moment (μ)

Vector quantity measuring bond or molecular polarity; μ = Q × r.

Hybridization

Mixing of atomic orbitals to form equivalent hybrid orbitals (e.g., sp, sp², sp³).

σ (Sigma) Bond

Bond formed by head-on orbital overlap; electron density along internuclear axis.

π (Pi) Bond

Bond formed by side-by-side overlap of p-orbitals; electron density above and below axis.

Molecular Orbital (MO) Theory

Describes electrons delocalized over entire molecule; combines atomic orbitals into bonding/antibonding MOs.

Stoichiometry

Quantitative relationships between reactants and products in a balanced equation.

Limiting Reactant

Reactant completely consumed first; limits product amount.

Theoretical Yield

Maximum product predicted from stoichiometry, assuming complete conversion.

Percent Yield

(Actual Yield / Theoretical Yield) × 100 %; indicates reaction efficiency.

Combustion Reaction

Reaction of a substance with O₂ producing heat, CO₂, and H₂O.

Redox Reaction

Reaction involving electron transfer; oxidation and reduction occur together.

Oxidation

Loss of electrons or increase in oxidation state.

Reduction

Gain of electrons or decrease in oxidation state.

Oxidizing Agent

Species that gains electrons (is reduced) and oxidizes another.

Reducing Agent

Species that loses electrons (is oxidized) and reduces another.

Solution

Homogeneous mixture of solute(s) dissolved in solvent.

Solute

Component present in lesser amount and dissolved in solvent.