Unit 1: Chemistry, Atoms, and Atomic Structure

Introduction to Chemistry and Science

• Science is the systematic study of the natural world through observation and experiment.

• Chemistry is the branch of science that studies the composition, structure, properties, and changes of matter.

• Technology applies scientific knowledge for practical purposes.

• Alchemy is an ancient practice that combined elements of chemistry, philosophy, and mysticism, aiming to transform matter (e.g., turning base metals into gold).

• Green chemistry and sustainable chemistry focus on designing products and processes that minimize environmental impact and resource consumption.

• Basic research seeks fundamental knowledge, while applied research aims to solve specific practical problems.

Properties and Changes of Matter

• Mass is the amount of matter in an object; weight is the force of gravity on that mass.

• Physical change alters the form or appearance of matter but not its composition (e.g., melting ice).

• Chemical change results in the formation of new substances (e.g., rusting iron).

• Physical properties can be observed without changing the substance (e.g., color, melting point).

• Chemical properties describe a substance's ability to undergo chemical changes (e.g., flammability).

Measurement and Units

• Use SI units: mass (kilogram, kg), length (meter, m), volume (liter, L), temperature (Kelvin, K; Celsius, °C).

• Unit conversions use conversion factors to change from one unit to another.

• Density is mass per unit volume:

• Temperature conversion:

Laws of Chemistry

• Law of Conservation of Mass: Mass is neither created nor destroyed in a chemical reaction.

• Law of Definite Proportions: A chemical compound always contains the same elements in the same proportions by mass.

The Mole and Avogadro's Number

• Mole is the SI unit for amount of substance; 1 mole = particles (Avogadro's number).

• Used to relate mass, number of particles, and volume in chemical calculations.

Atomic Structure and the Periodic Table

• Elements are arranged in the periodic table by increasing atomic number; arrangement reflects recurring chemical properties.

• Atoms are the smallest units of elements; molecules are combinations of atoms bonded together.

• The nucleus contains protons (positive charge, mass ≈ 1 amu) and neutrons (no charge, mass ≈ 1 amu).

• Electrons (negative charge, mass ≈ 0.0005 amu) orbit the nucleus.

• Isotopes are atoms of the same element with different numbers of neutrons.

• Electron transitions between energy levels result in absorption or emission of energy (light).

Unit 2: Chemical Bonding and Structure

Ions and Ionic Compounds

• Ions are atoms or molecules with a net electric charge due to loss or gain of electrons.

• Cations are positively charged (loss of electrons); anions are negatively charged (gain of electrons).

• Number of electrons in an ion = atomic number - charge (for cations) or + charge (for anions).

• Lewis symbols represent valence electrons as dots around the element symbol.

• Ionic bonds form from electrostatic attraction between cations and anions.

• Octet rule: Atoms tend to gain, lose, or share electrons to achieve 8 valence electrons.

• Common ions have predictable charges (e.g., Na+, Cl-).

• Binary ionic compounds: name cation first, then anion (e.g., sodium chloride).

Covalent Compounds and Bonding

• Covalent bonds involve sharing of electron pairs between atoms.

• Formulas and names for covalent compounds use prefixes (e.g., CO2: carbon dioxide).

• Bond polarity: Determined by difference in electronegativity; polar bonds have unequal sharing, nonpolar have equal sharing.

• Electronegativity increases across a period and up a group in the periodic table.

• Common nonmetals form a predictable number of bonds (e.g., oxygen forms 2, nitrogen forms 3).

• Polyatomic ions are charged groups of covalently bonded atoms (e.g., SO42-).

Unit 3: Chemical Accounting

Stoichiometry and Chemical Equations

• Balanced chemical equations have equal numbers of each atom on both sides.

• Balance equations by inspection, adjusting coefficients as needed.

• Use balanced equations to determine volumes of gases involved in reactions (at same temperature and pressure).

Mole Calculations and Molar Mass

• Formula mass (ionic compounds), molecular mass (molecules), and molar mass (g/mol) are calculated by summing atomic masses.

• Avogadro's number relates moles to number of particles.

• Conversions:

  • Mass to moles:

  • Moles to mass:

• Stoichiometry: Use mole ratios from balanced equations to relate reactants and products.

Solutions and Concentration

• Molarity (M):

• Percent by volume:

• Percent by mass:

• Given concentration and one quantity, calculate the other (e.g., moles, mass, or volume).

Unit 4: Chemical Reactions

Acids, Bases, and Neutralization

• Acids release H+ ions in water; bases release OH- ions (Arrhenius definition).

• Physical properties: acids taste sour, bases taste bitter and feel slippery.

• Write balanced equations for neutralization (acid + base → salt + water) and ionization reactions.

• Strong acids/bases dissociate completely; weak acids/bases only partially dissociate.

• Everyday uses: cleaning agents, food additives, antacids, etc.

Redox Reactions and Electrochemistry

• Oxidation: loss of electrons; reduction: gain of electrons.

• Redox reactions involve transfer of electrons between species.

• Oxidizing agent: causes oxidation (is reduced); reducing agent: causes reduction (is oxidized).

• Balance redox reactions by separating into half-reactions and ensuring electron balance.

• Electrochemical cells use redox reactions to generate electrical energy; half-reactions occur at electrodes.

Unit 5: Organic Chemistry

Introduction to Organic Chemistry

• Organic chemistry studies carbon-containing compounds.

• Organic compounds typically contain C-H bonds; inorganic compounds generally do not.

Hydrocarbons

• Hydrocarbons are compounds of only carbon and hydrogen.

• Alkanes: single bonds only (saturated); general formula CnH2n+2.

• Alkenes: at least one double bond (unsaturated); general formula CnH2n.

• Alkynes: at least one triple bond (unsaturated); general formula CnH2n-2.

• Name hydrocarbons using IUPAC rules (e.g., methane, ethene, propyne).

Aromatic Compounds and Functional Groups

• Aromatic compounds contain benzene rings (alternating double bonds in a six-membered ring).

• Name simple aromatic compounds (e.g., benzene, toluene).

• Halogenated hydrocarbons have halogen atoms (F, Cl, Br, I) substituted for hydrogen.

• Functional groups are specific groups of atoms that determine the chemical properties of organic molecules (e.g., alcohols, carboxylic acids).

• Recognize and write formulas for simple alkyl groups (e.g., methyl, ethyl).

Unit 6: Chemistry of Air

States of Matter and Phase Changes

• Particles in solids vibrate in place; in liquids, they move past each other; in gases, they move freely and rapidly.

• Phase changes: melting (solid to liquid), freezing (liquid to solid), evaporation (liquid to gas), condensation (gas to liquid), sublimation (solid to gas).

Intermolecular Forces

• Types: London dispersion forces (all molecules), dipole-dipole interactions (polar molecules), hydrogen bonding (molecules with H bonded to N, O, or F).

Kinetic-Molecular Theory and Gas Laws

• Five basic concepts:

  1. Gases consist of tiny particles in constant, random motion.

  2. Gas particles are far apart relative to their size.

  3. Collisions are elastic (no energy lost).

  4. No attractive or repulsive forces between particles.

  5. Average kinetic energy is proportional to temperature.

• Four simple gas laws:

  • Boyle's Law: (at constant T, n)

  • Charles's Law: (at constant P, n)

  • Gay-Lussac's Law: (at constant V, n)

  • Avogadro's Law: (at constant P, T)

• Ideal Gas Law:

• Use these laws to solve for unknown variables when others are given.

Atmospheric Chemistry and Environmental Impact

• Air composition: ~78% N2, ~21% O2, ~1% Ar, CO2, and trace gases.

• Pollution sources: natural (volcanoes, wildfires) and anthropogenic (vehicles, industry).

• Greenhouse gases: CO2, CH4, N2O, H2O vapor, O3.

• Greenhouse effect: Greenhouse gases trap heat in the atmosphere, maintaining Earth's temperature but contributing to global warming when concentrations rise.