Chapter 1: Introduction to Chemistry

Early and Modern Practice of Chemistry

• Early Chemistry: Chemistry began as alchemy, focusing on transforming substances and discovering new materials.

• Scientific Method: The process involves observation, hypothesis, experimentation, and conclusion.

• Modern Chemistry: Modern chemistry is systematic and relies on empirical evidence and the scientific method.

• Major Branches: The four main branches are organic chemistry, inorganic chemistry, physical chemistry, and analytical chemistry.

• Impact: Chemistry affects daily life through medicine, energy, materials, and the environment.

Chapter PSS: Precision, Significant Figures, and Scientific Notation

Measurement and Significant Figures

• Measurement Inexactness: All measurements have some degree of uncertainty due to instrument limitations.

• Instruments: Common instruments include balances, graduated cylinders, and thermometers.

• Significant Digits: Digits in a measurement that are known with certainty plus one estimated digit.

• Rules for Calculations:

  • Addition/Subtraction: Result has the same number of decimal places as the measurement with the fewest decimal places.

  • Multiplication/Division: Result has the same number of significant figures as the measurement with the fewest significant figures.

• Scientific Notation: Expresses numbers as a product of a coefficient and a power of 10.

Example: is scientific notation for 32,000.

Chapter 2: The Metric System

Units, Conversions, and Calculations

• Base Units: Gram (mass), liter (volume), meter (distance).

• Unit Equation: Statement of two equivalent quantities (e.g., 1 m = 100 cm).

• Unit Factor: Ratio of two equivalent quantities used for conversions.

• Metric Prefixes: Prefixes modify base units by powers of ten (e.g., kilo-, centi-, milli-).

• Conversions: Use unit factors to convert between units, including metric to English units.

• Volume Calculation:

• Density: , ,

• Temperature Scales: Know freezing and boiling points of water in Fahrenheit, Celsius, and Kelvin.

• Temperature Conversions:

  • Fahrenheit to Celsius:

  • Celsius to Fahrenheit:

  • Celsius to Kelvin:

  • Kelvin to Celsius:

• Temperature vs. Heat: Temperature measures average kinetic energy; heat is energy transfer due to temperature difference.

Chapter 3: Matter and Energy

States, Classification, and Properties of Matter

• States of Matter: Solid, liquid, gas.

• Physical Changes: Changes in state (e.g., melting, boiling) are physical changes.

• Classification: Matter can be classified as mixtures or pure substances. Mixtures can be separated physically; compounds can be separated chemically into elements.

• Elements: Know names and symbols of common elements.

• Metals vs. Nonmetals: Metals are typically shiny, conductive, and malleable; nonmetals are not. Semimetals have intermediate properties.

• Physical State Prediction: Use periodic table position to predict state at room temperature.

• Chemical Formulas: Indicate the number of each atom in a compound.

• Law of Definite Composition: A compound always contains the same elements in the same proportion by mass.

• Physical vs. Chemical Properties: Physical properties can be observed without changing composition; chemical properties describe reactivity.

• Physical vs. Chemical Changes: Physical changes do not alter composition; chemical changes produce new substances.

• Conservation of Mass: Mass is conserved in chemical reactions.

• Potential vs. Kinetic Energy: Potential energy is stored; kinetic energy is energy of motion. Temperature relates to average kinetic energy.

• Forms of Energy: Chemical, electrical, mechanical, nuclear, heat, and light.

Chapter 4: Models of the Atom

Atomic Theory and Structure

• Dalton's Atomic Theory: Matter is composed of atoms; atoms of the same element are identical (later found not always true); atoms combine in simple ratios; atoms are indivisible (later found not always true).

• Thomson Model: Plum pudding model; electrons embedded in a positive sphere.

• Rutherford Model: Atom has a dense, positive nucleus; electrons orbit nucleus.

• Subatomic Particles: Proton (positive, in nucleus), neutron (neutral, in nucleus), electron (negative, outside nucleus).

• Atomic Notation: where is mass number, is atomic number.

• Isotopes: Atoms with same but different ; neutrons = .

• Relative Atomic Mass: Weighted average of isotopes.

• Light and Energy: , (where is energy, is Planck's constant, is frequency, is speed of light, is wavelength).

• Electromagnetic Spectrum: Includes gamma rays, X-rays, UV, visible (ROYGBIV), IR, microwaves, radio waves.

• Quantum Concept: Energy is quantized; electrons occupy discrete energy levels.

• Bohr Model: Electrons orbit nucleus in fixed energy levels; emission spectra explained by electron transitions.

• Sublevels and Orbitals: s, p, d, f sublevels; each holds a specific number of electrons (s: 2, p: 6, d: 10, f: 14).

• Order of Filling: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.

• Electron Configuration: Distribution of electrons among orbitals.

• Orbital Shapes: s orbitals are spherical; p orbitals are dumbbell-shaped.

Chapter 5: The Periodic Table

Organization and Trends

• Mendeleev's Periodic Law: Properties of elements repeat periodically when arranged by atomic mass.

• Moseley's Modern Law: Properties repeat when elements are arranged by atomic number.

• Groups and Periods: Groups (columns) and periods (rows); American (IA–VIIIA) and IUPAC (1–18) conventions.

• Trends:

  • Metallic character increases down a group, decreases across a period.

  • Atomic size increases down a group, decreases across a period.

  • Ionization energy increases across a period, decreases down a group.

• Valence Electrons: Outermost electrons; determine chemical properties.

• Electron Dot Formulas: Show valence electrons as dots around element symbol.

• Ionic Charge Prediction: Based on group number; e.g., Group 1 forms +1 ions.

Chapter 6: Language of Chemistry

Naming and Classifying Compounds

• Compound Types: Binary ionic, ternary ionic, binary molecular, binary acids, ternary oxyacids.

• Ions: Monoatomic cations/anions, polyatomic cations/anions.

• Naming Ions: Systematic names and formulas for common ions; predict charges for representative elements.

• Polyatomic Ions: Memorize common ions (e.g., nitrate, sulfate).

• Naming Compounds: Use systematic rules for ionic and molecular compounds, acids, and oxyacids.

Chapter 7: Chemical Reactions

Types and Evidence of Chemical Change

• Evidence for Reaction: Color change, gas formation, precipitate formation, energy change.

• Diatomic Elements: H2, N2, O2, F2, Cl2, Br2, I2.

• Writing Equations: Translate descriptions into chemical equations; balance equations to conserve mass.

• Reaction Types: Combination, decomposition, single-replacement, double-replacement, neutralization.

• Predicting Reactions: Use activity series and solubility rules to predict products and solubility.

Chapter 8: The Mole Concept

Counting and Calculating with Moles

• Avogadro's Number: particles per mole.

• Mole Calculations: Convert between particles, moles, and mass using molar mass.

• Molar Mass: Sum of atomic masses in a compound's formula.

• Percent Composition:

• Empirical Formula: Simplest whole-number ratio of elements in a compound.

• Molecular Formula: Actual number of atoms in a molecule; calculated from empirical formula and molar mass.

Example: To find moles from particles:

Additional info: These objectives cover the foundational concepts and skills required for success in an introductory college chemistry course. Students should practice problems, memorize key terms, and understand the relationships between concepts for exam preparation.