General Chemistry Study Guide: Foundations, Atoms, and Compounds

Chapter 1: Introduction to Chemistry

Steps of the Scientific Method

• Observation: Gathering information about phenomena.

• Hypothesis: Proposing a tentative explanation.

• Experimentation: Testing the hypothesis through controlled experiments.

• Analysis: Interpreting data to draw conclusions.

• Theory Development: Formulating a theory if the hypothesis is consistently supported.

Matter and Its Classification

• Pure Substance: Matter with a fixed composition.

  • Element: Cannot be broken down into simpler substances.

    • Monatomic: Consists of single atoms (e.g., He, Ne).

    • Diatomic: Two atoms bonded (e.g., O2, N2).

  • Compound: Composed of two or more elements chemically combined.

    • Binary: Two elements (e.g., NaCl).

    • Ternary: Three elements (e.g., H2SO4).

• Mixtures: Physical combinations of substances.

  • Homogeneous: Uniform composition (e.g., saltwater).

  • Heterogeneous: Non-uniform composition (e.g., salad).

Colloids and Suspensions

• Colloids: Mixtures with intermediate particle size (e.g., milk).

• Suspensions: Particles settle out over time (e.g., muddy water).

Physical vs. Chemical Properties and Changes

• Physical Properties: Observed without changing composition (e.g., melting point).

• Chemical Properties: Observed during a chemical change (e.g., flammability).

• Physical Change: Does not alter composition (e.g., melting ice).

• Chemical Change: Alters composition (e.g., rusting iron).

States of Matter

• Solid: Fixed shape and volume; particles vibrate in place.

• Liquid: Fixed volume, takes shape of container; particles move more freely.

• Gas: No fixed shape or volume; particles move rapidly and are far apart.

Separation Techniques

• Filtration: Separates solids from liquids.

• Distillation: Separates based on boiling points.

• Chromatography: Separates based on movement through a medium.

• Chemical Decomposition: Breaking compounds into elements or simpler compounds.

Measurement and Units

• Quantitative Measurement: Numerical data (e.g., mass, volume).

• SI System: Standard units (meter, kilogram, second, etc.).

• Temperature Scales:

  • Celsius (°C)

  • Fahrenheit (°F)

  • Kelvin (K)

Conversion Factors and Significant Figures

• Conversion Factors: Used to convert between units.

• Significant Figures: Indicate precision of measurements.

  • Rules for multiplying/dividing and adding/subtracting.

• Percent Composition: Used as a conversion factor in stoichiometry.

Chapter 2: Atoms and Elements

Subatomic Particles

• Protons: Positive charge, found in nucleus, mass ≈ 1 amu.

• Neutrons: Neutral, found in nucleus, mass ≈ 1 amu.

• Electrons: Negative charge, found in electron cloud, mass ≈ 0.0005 amu.

Atomic Structure and Notation

• Atomic Number (Z): Number of protons.

• Mass Number (A): Protons + neutrons.

• Isotopes: Atoms with same Z but different A.

• Average Atomic Mass: Weighted average based on isotopic abundance. Formula: $\text{Average Atomic Mass} = \sum (\text{fractional abundance} \times \text{isotope mass})$

Ground and Excited States

• Ground State: Lowest energy arrangement of electrons.

• Excited State: Higher energy arrangement due to electron promotion.

Periodic Table Organization

• Groups/Families: Vertical columns; similar chemical properties.

• Periods: Horizontal rows; properties change progressively.

• Metals, Nonmetals, Metalloids: Classified by physical and chemical properties.

• Special Groups:

  • Alkali Metals (Group 1)

  • Alkaline Earth Metals (Group 2)

  • Halogens (Group 17)

  • Noble Gases (Group 18)

  • Transition Metals, Lanthanides, Actinides

Valence Electrons and Charges

• Valence Electrons: Electrons in the outermost shell; determine reactivity.

• Usual Charges: Determined by group number (e.g., Group 1: +1, Group 17: -1).

The Mole and Conversions

• Mole: SI unit for amount of substance; $1 \text{mol} = 6.022 \times 10^{23}$ particles.

• Conversions:

  • Atoms ↔ Moles

  • Grams ↔ Moles

  • Grams ↔ Atoms (two-step conversion)

History of the Atom

• Law of Conservation of Mass: Mass is neither created nor destroyed.

• Law of Constant Composition: Compounds have fixed ratios of elements.

• Dalton's Atomic Theory: Atoms are indivisible, combine in whole-number ratios.

• Cathode Ray Tube: Discovery of the electron (Thomson).

• Plum Pudding Model: Electrons embedded in positive sphere.

• Gold Foil Experiment: Discovery of nucleus (Rutherford).

• Millikan Oil Drop: Measured electron charge.

• Bohr Model: Electrons in quantized orbits.

• Quantum Mechanical Model: Electron clouds, probability distributions (Schrödinger, Heisenberg).

Chapter 3: Compounds and Chemical Formulas

Types of Compounds

• Ionic Compounds: Formed from metals and nonmetals; transfer of electrons.

• Covalent Compounds: Formed from nonmetals; sharing of electrons.

Formulas and Models

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

• Molecular Formula: Actual number of atoms in a molecule.

• Structural Formula: Shows arrangement of atoms.

• Ball-and-Stick Model: 3D representation of molecular structure.

• Space-Filling Model: Shows relative sizes and positions of atoms.

Naming Compounds

• Ionic Compounds: Name cation first, then anion (e.g., sodium chloride).

• Covalent Compounds: Use prefixes to indicate number of atoms (e.g., carbon dioxide).

Octet Rule and Charges

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

• Typical Charges: Groups 1A-7A have predictable charges based on group number.

Crossover Rule for Ionic Compounds

• Crossover Rule: Used to balance charges when writing formulas for ionic compounds.

• Polyatomic Ions: Use parentheses when more than one is present in a formula.