General Chemistry Study Guide: Matter, Atomic Theory, and Chemical Compounds

Chapter 1: Introduction to Chemistry and Matter

Scientific Method

The scientific method is a systematic approach to understanding natural phenomena through observation, hypothesis formation, experimentation, and analysis.

• Steps: Observation, Hypothesis, Experiment, Analysis, Conclusion, and Communication of results.

Matter and Its Classification

Matter is anything that has mass and occupies space. It can be classified based on composition and physical state.

• Pure Substance: Matter with a fixed composition.

  • Element: Cannot be broken down into simpler substances.

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

    • Diatomic: Consists of two atoms (e.g., O2).

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

    • Binary: Contains two elements (e.g., NaCl).

    • Ternary: Contains three elements (e.g., H2SO4).

• Mixture: Physical combination of two or more substances.

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

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

  • Colloids and Suspensions: Mixtures with particles dispersed in a medium (e.g., milk is a colloid).

Physical and Chemical Properties & Changes

• Physical Properties: Observed without changing the substance (e.g., color, 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

Matter exists in three primary states: solid, liquid, and gas. Each state has distinct properties:

State	Fixed Shape?	Fixed Volume?	Molecular Motion
Solid	Yes	Yes	Very slow/vibrational
Liquid	No	Yes	Moderate
Gas	No	No	Fast

Separation Techniques

• Filtration: Separates solids from liquids.

• Distillation: Separates based on boiling points.

• Chromatography: Separates based on movement through a medium.

• Decomposition (Chemical): Breaking compounds into simpler substances via chemical change.

Measuring Matter

• Qualitative: Descriptive (e.g., color, odor).

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

SI System and Units

• Mass: Kilogram (kg)

• Temperature: Kelvin (K), Celsius (°C), Fahrenheit (°F)

Temperature Conversions:

• $T_{\text{K}} = T_{\text{°C}} + 273.15$

• $T_{\text{°F}} = \frac{9}{5}T_{\text{°C}} + 32$

Conversion Factors and Significant Figures

• Conversion Factors: Used to convert between units (e.g., 1 kg = 1000 g).

• Significant Figures: Indicate precision of measurements.

  • Multiplying/Dividing: Result has as many significant figures as the least precise measurement.

  • Adding/Subtracting: Result has as many decimal places as the least precise measurement.

Scientific Notation

• Expresses numbers as $a \times 10^n$ (e.g., $3.0 \times 10^4$).

Density and Percent Composition as Conversion Factors

• Density: $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

• Percent Composition: $\%\,\text{Composition} = \frac{\text{Mass of Component}}{\text{Total Mass}} \times 100$

Chapter 2: Atomic Theory and the Periodic Table

History of the Atom

• Law of Conservation of Mass (Lavoisier): Mass is neither created nor destroyed in chemical reactions.

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

• Law of Multiple Proportions: When two elements form more than one compound, the ratios of the masses of the second element combine with a fixed mass of the first element in small whole numbers.

• Dalton's Atomic Theory: Atoms are indivisible particles; atoms of the same element are identical; compounds are combinations of different atoms.

• Cathode Ray Experiment (Thomson): Discovery of the electron; led to the "plum pudding" model.

• Oil Drop Experiment (Millikan): Determined the charge of the electron.

• Gold Foil Experiment (Rutherford): Discovered the nucleus is small, dense, and positively charged.

• Discovery of the Neutron (Chadwick): Neutrons are neutral particles in the nucleus.

• Bohr Model: Electrons orbit the nucleus in defined energy levels.

• Schrödinger: Introduced quantum mechanical model; electron position described by probability.

The Atom and Subatomic Particles

• Protons: Positive charge (+1), mass ≈ 1 amu, located in nucleus.

• Neutrons: No charge, mass ≈ 1 amu, located in nucleus.

• Electrons: Negative charge (–1), mass ≈ 0.0005 amu, located outside nucleus.

Atomic Symbols and Isotopes

• Element Symbol: One or two-letter abbreviation (e.g., H, He).

• Atomic Number (Z): Number of protons.

• Mass Number (A): Number of protons + neutrons.

• Neutron Number: $\text{Neutrons} = A - Z$

• Isotopes: Atoms of the same element with different numbers of neutrons.

• Average Atomic Mass: Weighted average based on isotopic abundance.

• Calculation: $\text{Average Atomic Mass} = \sum (\text{Fractional Abundance} \times \text{Isotope Mass})$

Ions

• Cation: Positively charged ion (loss of electrons).

• Anion: Negatively charged ion (gain of electrons).

Periodic Table Organization

• Rows/Periods: Horizontal rows (numbered 1–7).

• Columns/Groups/Families: Vertical columns (numbered 1–18).

• Metals: Left and center; conduct electricity, malleable.

• Nonmetals: Right side; poor conductors, brittle.

• Metalloids: Border metals and nonmetals; intermediate properties.

• Hydrogen: Unique, nonmetal, placed above Group 1.

Major Groups and Their Properties

Group	Valence Electrons	Usual Charge
Alkali Metals (1A)	1	+1
Alkaline Earth Metals (2A)	2	+2
Halogens (7A)	7	–1
Noble Gases (8A)	8	0

• Transition Metals: Groups 3–12; variable charges.

• Lanthanides and Actinides: Inner transition metals.

The Mole and Conversions

• Avogadro's Number: $6.022 \times 10^{23}$ particles/mol.

• Conversions:

  • Atoms ↔ Moles: $\text{Moles} = \frac{\text{Atoms}}{6.022 \times 10^{23}}$

  • Grams ↔ Moles: $\text{Moles} = \frac{\text{Mass (g)}}{\text{Molar Mass (g/mol)}}$

  • Grams ↔ Atoms: Two-step conversion via moles.

Chapter 3: Chemical Compounds and Formulas

Chemical Formulas and Models

• Subscripts: Indicate number of atoms in a molecule (e.g., H2O).

• Molar Mass: Sum of atomic masses in a compound (g/mol).

Conversion Factors with Compounds

• Grams ↔ Moles: $\text{Moles} = \frac{\text{Mass (g)}}{\text{Molar Mass}}$

• Moles ↔ Molecules: $\text{Molecules} = \text{Moles} \times 6.022 \times 10^{23}$

• Grams ↔ Molecules: Two-step conversion via moles.

• Atoms ↔ Molecules: Use subscripts and Avogadro's number.

Percent Composition

• $\%\,\text{Element} = \frac{\text{Mass of Element in 1 mol}}{\text{Molar Mass of Compound}} \times 100$

Empirical, Molecular, and Structural Formulas

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

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

• Structural Formula: Shows arrangement of atoms.

• Condensed Structural Formula: Simplified notation (e.g., CH3CH2OH).

• Ball-and-Stick Model: 3D representation of atoms and bonds.

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

Types of Compounds

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

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

Octet Rule and Typical Charges

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

• Groups 1A–7A have typical charges based on their position (see table above).

Crossover Rule for Ionic Compounds

• Used to write formulas for ionic compounds by balancing charges.

• For polyatomic ions, use parentheses when more than one is needed (e.g., Ca(NO3)2).

Naming Compounds

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

• Covalent Compounds: Use prefixes (e.g., carbon dioxide).

• Binary Acids: "Hydro-" + root + "-ic acid" (e.g., HCl: hydrochloric acid).

Example: To convert 10.0 g of H2O to molecules:

• Find moles: $\text{Moles} = \frac{10.0}{18.02} = 0.555$ mol

• Find molecules: $0.555 \times 6.022 \times 10^{23} = 3.34 \times 10^{23}$ molecules

Additional info: Some explanations and tables have been expanded for clarity and completeness based on standard general chemistry curricula.