Unit B Study Guide
1. WHMIS Symbols
• Flammable: Easily ignites and burns (e.g., gasoline).
• Corrosive: Causes burns or destroys materials (e.g., acids like hydrochloric acid).
• Toxic (Acute): Harmful or fatal even in small doses (e.g., cyanide).
• Biohazardous: Infectious materials causing disease (e.g., viruses, bacteria).
2. Theories Developed by Scientists/Philosophers
• Democritus: Proposed the idea of indivisible particles, “atoms.”
• Dalton: Atoms are indivisible and unique to each element. Proposed atomic theory.
• Rutherford: Gold foil experiment proved the existence of a dense nucleus.
• Bohr: Electrons orbit the nucleus in fixed energy levels (planetary model). bohr-bit
3. Classification of Matter
• Pure Substances:
• Elements: Cannot be broken down further (e.g., Oxygen – O₂).
• Compounds: Composed of two or more elements chemically bonded (e.g., H₂O).
• Mixtures:
• Mechanical Mixtures: Components remain visibly distinct (e.g., trail mix).
• Solutions: Homogeneous mixture, fully dissolved (e.g., saltwater).
• Suspensions: Particles visible but settle over time (e.g., sand in water).
• Colloids: Fine particles evenly distributed; don’t settle (e.g., milk).
4. Elements vs. Compounds
• Elements: Single type of atom (e.g., Oxygen – O₂).
• Compounds: Chemically bonded atoms of different elements (e.g., H₂O, NaCl).
5. States of Matter | Changes of State
• States of Matter: Solid, liquid, gas, plasma (e.g., lightning for plasma).
• Changes of State:
• Melting: Solid to liquid.
• Freezing: Liquid to solid.
• Condensation: Gas to liquid.
• Evaporation: Liquid to gas.
• Sublimation: Solid to gas (e.g., dry ice).
• Deposition: Gas to solid (e.g., frost).
6. Chemical vs. Physical Changes
• Chemical Change: Creates new substances with new properties (e.g., burning wood → ash, CO₂).
• Physical Change: Substance changes form, not composition (e.g., ice melting into water).
7. Periodic Table
• Families (Groups/Columns):
• Halogens (Group 17): Highly reactive nonmetals (e.g., chlorine, fluorine).
• Noble Gases (Group 18): Stable, non-reactive (e.g., helium, neon).
• Periods (Rows): Reflect number of electron shells.
• Metals: Conductive, malleable (e.g., aluminum).
• Nonmetals: Poor conductors, brittle (e.g., sulfur).
8. Atomic Structure
• Protons: Positive charge, found in nucleus.
• Neutrons: Neutral charge, in nucleus.
• Electrons: Negative charge, in electron orbits.
• Key Formula: Neutrons = Atomic Mass – Atomic Number.
9. Metals, Metalloids, Nonmetals
• Metals: Conduct electricity, shiny, malleable (e.g., copper).
• Metalloids: Semi-conductive, properties of both metals/nonmetals (e.g., silicon).
• Nonmetals: Brittle, poor conductors (e.g., nitrogen).
10. Diagram of Water Molecule
• H₂O: Two hydrogen atoms covalently bonded to one oxygen atom (V-shaped structure).
11. Ionic vs Molecular Compounds
• Ionic Compounds:
• Metal + Nonmetal (e.g., NaCl).
• Naming: Change nonmetal ending to “-ide” (e.g., sodium chloride).
• Molecular Compounds:
• Nonmetals only (e.g., CO₂).
• Naming: Use prefixes (mono-, di-, tri-) to indicate atom counts.
12. Polyatomic Ions
• Examples:
• Carbonate (CO₃²⁻).
• Sulfate (SO₄²⁻).
• Hydroxide (OH⁻).
• Phosphate (PO₄³⁻).
13. Chemical Formulas
• Example: H₂SO₄ (Sulfuric Acid): 2 Hydrogen, 1 Sulfur, 4 Oxygen.
• Writing Formulas: From chemical diagrams or names (e.g., magnesium chloride → MgCl₂).
14. Endothermic vs Exothermic Reactions
• Endothermic: Absorbs heat (e.g., photosynthesis).
• Exothermic: Releases heat (e.g., combustion, freezing water).
15. Combustion vs Corrosion
• Combustion: Burning in oxygen, produces energy (e.g., burning wood).
• Corrosion: Slow oxidation (e.g., rust forming on iron).
16. Conservation of Mass
• Law: Total mass of reactants = Total mass of products.
• Example: 2H₂ + O₂ → 2H₂O (Mass conserved).
17. Factors Affecting Chemical Reactions
• Catalysts: Speed up reactions (e.g., enzymes).
• Temperature: Higher temperature = faster reaction.
• Surface Area: More surface area = faster reaction (e.g., powdered vs solid reactant).
• Concentration: Higher concentration = faster reaction.
18. Writing Formulas for Reactions
• Format: Reactant + Reactant → Product.
• Example: CH₄ + 2O₂ → CO₂ + H₂O (Combustion of methane).