GOB Chemistry Study Notes: Matter, Atoms, Compounds, Organic Molecules, Reactions, and Intermolecular Forces

Ch. 1: Chemistry Basics – Matter and Measurement

Elements, Compounds, and Mixtures

• Element: A pure substance consisting of only one type of atom (e.g., O2, Fe).

• Compound: A substance formed from two or more elements chemically bonded (e.g., H2O).

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

• Heterogeneous Mixture: Non-uniform composition (e.g., salad, sand in water).

Elements: Symbols, Atomic Number, Atomic Weight

• Chemical Symbol: One- or two-letter abbreviation for an element (e.g., Na for sodium).

• Atomic Number (Z): Number of protons in the nucleus.

• Atomic Weight: Weighted average mass of an element's isotopes.

• Macronutrients: Elements required in large amounts (e.g., C, H, O, N, P, S).

• Micronutrients: Elements required in trace amounts (e.g., Fe, Zn, Cu).

Physical vs. Chemical Change

• Physical Change: Alters form but not composition (e.g., melting ice).

• Chemical Change: Produces new substances (e.g., rusting iron).

Balancing Equations

• Law of Conservation of Mass: Atoms are neither created nor destroyed in chemical reactions.

• Balance equations by adjusting coefficients, not subscripts.

Measurements and Units

• Metric Prefixes: kilo (k), centi (c), milli (m), micro (μ), nano (n), etc.

• Converting Units: Use dimensional analysis (factor-label method).

• Significant Figures: Reflect precision of measurements.

• Scientific Notation: Expresses numbers as $a \times 10^n$.

Main Quantities and Units

• Mass: Measured in grams (g) or kilograms (kg).

• Volume: Measured in liters (L) or milliliters (mL).

• Density: $\text{Density} = \frac{\text{mass}}{\text{volume}}$

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

• Common Units: lb, in, cm, mL, dL, L, oz.

Derived Units

• Derived from base units (e.g., m/s for speed).

Ch. 2: Atoms and Radioactivity

Structure of the Atom

• Nucleus: Contains protons and neutrons.

• Electron Cloud: Region where electrons are likely found.

• Atomic Number (Z): Number of protons.

• Atomic Mass: Sum of protons and neutrons.

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

Ch. 3: Compounds – How Elements Combine

Electron Arrangement and the Octet Rule

• Electrons fill shells around the nucleus; stable configurations have 8 valence electrons (octet rule).

• Common oxidation states relate to how atoms gain or lose electrons to achieve an octet.

• Periodic trends affect ion formation (e.g., metals lose electrons, nonmetals gain electrons).

Ionic and Covalent Compounds

• Ionic Compounds: Formed by transfer of electrons (e.g., NaCl).

• Covalent Compounds: Formed by sharing electrons (e.g., H2O).

• Lewis structures represent bonding in molecules.

Counting Units: The Mole

• Mole: SI unit for amount of substance.

• $1 \text{ mole} = 6.022 \times 10^{23}$ entities (Avogadro's number).

• $\text{Moles} = \frac{\text{mass}}{\text{molar mass}}$

VSEPR Theory: Molecular Shape

• Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shapes based on electron pair repulsion.

Ch. 4: Introduction to Organic Compounds

Organic Molecules and Structures

• Organic molecules contain carbon and hydrogen, often with O, N, S, P, or halogens.

• Structures can be expanded, condensed, or skeletal (line-angle) formulas.

Nomenclature of Alkanes

• Alkanes: Saturated hydrocarbons (single bonds only).

• Prefix + -ane (e.g., methane, ethane).

Functional Groups

• Groups of atoms that impart characteristic properties (e.g., alcohols, carboxylic acids).

Lipids

• Saturated: No double bonds (solid at room temp).

• Unsaturated: One or more double bonds (liquid at room temp).

• Trans vs. Cis: Refers to arrangement around double bonds.

• Polysaturated: Multiple double bonds.

Isomerism

• Cis/Trans Isomers: Same formula, different spatial arrangement around double bonds.

• Chirality: Molecules with non-superimposable mirror images (important in biochemistry).

Ch. 5: Chemical Reactions

Types of Chemical Reactions

• Synthesis: Combining substances to form one product.

• Decomposition: Breaking down a compound into simpler substances.

• Exchange: Atoms or ions are exchanged between compounds.

• Combustion: Reaction with oxygen producing heat and light.

• Reversible/Irreversible Reactions: Some reactions can go both ways.

• Oxidation: Loss of electrons (increase in C–O bonds).

• Reduction: Gain of electrons (increase in C–H bonds).

• Hydrolysis: Splitting with water.

• Condensation: Joining with loss of water.

Chemical Reaction Properties

• Endothermic: Absorbs energy.

• Exothermic: Releases energy.

• Activation Energy: Minimum energy required to start a reaction.

• Kinetics: Study of reaction rates.

Ch. 6: Carbohydrates – Life's Sweet Molecules

• Carbohydrates are reviewed in class; focus on structure and function of sugars.

Ch. 7: States of Matter and Their Attractive Forces

Attractive Forces

• Dispersion Forces: Weak attractions due to temporary dipoles in all molecules.

• Dipole-Dipole: Attractions between polar molecules.

• Hydrogen Bonding: Strong dipole-dipole interaction involving H bonded to N, O, or F.

• Ionic Attraction: Strongest, between oppositely charged ions.

Attractive Forces and Solubility

• Like dissolves like: Polar solutes dissolve in polar solvents; nonpolar in nonpolar.

• Amphipathic Compounds: Contain both hydrophilic and hydrophobic regions (e.g., surfactants, micelles, lipid bilayers).

Example Table: Types of Intermolecular Forces

Type of Force	Relative Strength	Example
Dispersion	Weakest	All molecules (e.g., CH4)
Dipole-Dipole	Intermediate	HCl
Hydrogen Bonding	Strong	H2O, NH3
Ionic Attraction	Strongest	NaCl

Additional info: These notes are expanded from handwritten class notes and include academic context for clarity and completeness. Page references and some details (e.g., specific examples, page numbers) are inferred or generalized for study purposes.