Chapters 1-7 & 9 Study Notes

• Study notes for Chapters 1-7 & 9

Module 1 (Chapter 1)

• Matter: Anything that occupies space and has mass.
• Mass: Measures the amount of matter in an object.
• Weight: Measures gravitational force acting on an object.
• Physical change examples: grinding coffee beans, boiling water for soup, melting gold, solid dry ice changing to a gas.
• Chemical change examples: burning coal.
• Diatomic and Homoatomic molecules example: Oxygen gas (O2).
• Diatomic and Heteroatomic molecules examples: Carbon monoxide (CO), Methane gas (CH4), and Water (H2O).
• Triatomic and Heteroatomic molecule example: Glucose (C6H12O6).
• Homogeneous matter: Has the same properties throughout a sample.
• Solutions: Homogenous mixtures of two or more pure substances.
• Heterogeneous matter: Properties that are not the same throughout the sample.
• Element: A pure substance consisting of only one kind of atom in homoatomic molecules or individual atoms form.
• Compound: Pure substance consisting of two or more kinds of atoms in heteroatomic molecules or individual atoms.
• 55°F is equivalent to 13°C.
• 37°C is equivalent to 310K.
• 256 K equals -17°C.
• 23°C converts to 73°F.
• Standard scientific notation: 372,000 is 3.72 x 10⁵, 0.00000384 is 3.84 x 10⁻⁶, 400 is 4 x 10², and 0.039401 is 3.9401 x 10⁻².
• Significant figures: 0.00450 has 3, 140,000,000 has 2, 0.1005030 has 7, 150.01 has 5, 0.1010 has 4.
• Aluminum Density: Aluminum has a density of 2.7 g/mL.
• Density Equation: Density (D) equals mass (g) divided by volume (mL).
• Mass of Aluminum Sample: An aluminum sample with a volume of 60.0 mL has a mass of 162 g.
• Volume of Aluminum Sample: An aluminum sample of 98.5 g has a volume of 36.5 mL.

Module 1 (Chapter 3)

• Germanium (Ge) is in Group IVA (14) and period 4.
• Aluminum (Al) is in Group IIIA (13) and period 3.
• Gold (Au) is in Group IB (11) and period 6.
• Fluorine (F) is in Group VIIA (17) and period 2.
• Antimony (Sb) is in Group VA (15) and period 5.
• Nonmetals: Hydrogen (H), Chlorine (Cl), and Iodine (I).
• Metals: Chromium (Cr) and Sodium (Na).
• Metalloids: Silicon (Si) and Boron (B).
• Larger atomic radius: Strontium (Sr), Rubidium (Rb), Tellurium (Te), Iodine (I), and Germanium (Ge).

Module 1 (Chapter 4)

• Beryllium (Be) contains 2 valence electrons and forms a 2+ ion charge.
• Phosphorus (P) contains 5 valence electrons and forms a 3- ion.
• Oxygen (O) contains 6 valence electrons and forms a 2- ion.
• Rubidium (Rb) contains 1 valence electron and forms a 1+ ion.
• Xenon (Xe) contains 8 valence electrons and forms no ion.
• Ionic compounds with Barium (Ba): BaSe, Ba₃P₂, BaI₂, and Ba₃As₂.
• Ionic Compounds - Nomenclature: SrS is Strontium sulfide, CaF₂ is Calcium floride, BaCl₂ is Barium chloride, Li₂O is Lithium oxide, MgO is Magnesium oxide.
• Covalent vs Ionic: LiBr (Ionic), SO₂ (Covalent), CsF (Ionic), MgI₂ (Ionic), SrO (Ionic).
• Binary Covalent Compounds - Nomenclature: SiO₂ Silicone dioxide, P₂O₅ Diphosphorus pentoxide, N₂O₅ Dinitrogen pentoxide, CCl₄ Carbon tetrachloride, BF₃ Boron trifluoride.

Module 2 (Chapter 2)

• Number of atoms in molecules: NH₃ (1 Nitrogen, 3 Hydrogens), C₂H₄O₂ (2 Carbons, 4 Hydrogens, 2 Oxygens), H₃BO₃ (3 Hydrogens, 1 Boron, 3 Oxygens), HN2 (1 Hydrogen, 2 Nitrogens).
• Neutral elements: C (6 protons, 6 electrons), H (1 proton, 1 electron), Ar (18 protons, 18 electrons), Mg (12 protons, 12 electrons), P (15 protons, 15 electrons).
• Isotopes: ³²S (16 protons, 16 electrons, 18 neutrons), ⁹Be (4 protons, 4 electrons, 5 neutrons), ¹³¹Xe (54 protons, 54 electrons, 77 neutrons), Silicon-28 (14 protons, 14 electrons, 14 neutrons), Argon-40 (18 protons, 18 electrons, 22 neutrons).
• Molecular weights: C₂H₂ (26.04 g/mol), C₆H₁₂O₆ (180.16 g/mol), O₃ (48.00 g/mol), CS₂ (76.14 g/mol), F₂ (38.00 g/mol).
• Grams from moles: He (17.0 g), H₂O (76.6 g), C₆H₁₂O₆ (766 g).
• Moles from grams: He (1.44 moles), H₂O (0.321 moles), C₆H₁₂O₆ (0.0321 moles).
• Mole factor relationships:
  • 23.18 g O are contained in 43.5 g of C₆H₁₂O₆.
  • 18.0 moles of H atoms are contained in 1.50 mol of C₆H₁₂O₆.
  • 4.50 x 10²³ C atoms are in 7.50 x 10²² molecules of C₆H₁₂O₆.

Module 2 (Chapter 5)

• Oxidation: Reaction increasing the number of oxygen bonds, decreasing the number of electrons and hydrogen bonds.
• Reduction: Reaction decreasing the number of oxygen bonds, increasing the number of electrons and hydrogen bonds.

Module 2 (Chapter 6)

• States of Matter: Solid (definite shape, high density), Liquid (indefinite shape, high density), Gas (indefinite shape, low density).
• Pressure conversions: 746 mmHg = 0.982 atm, 123 mmHg = 0.162 atm, 201 mmHg = 0.264 atm, 490 mmHg = 0.645 atm.
• Balanced Equations:
  • 2 Fe(s) + 3 Cl₂(g) → 2 FeCl₃(s)
  • 2 HgO(s) → 2 Hg(l) + 1 O₂(g)
  • 2 CO(g) + 1 O₂(g) → 2 CO₂(g)
  • 2 BrF₃(g) → 1 Br₂(g) + 3 F₂(g)

Module 3 (Chapter 7)

• Ideal Gas Law Applications:
  • 2.00 mol of oxygen confined to a volume of 400 mL at 20.0°C exerts a pressure of 120 atm.
  • 0.525 mol of hydrogen gas exerts a pressure of 3.00 atm at a temperature of 15.0°C with a volume of 4.14 L.
  • A nitrogen gas sample with a volume of 2.25 L and pressure of 300 torr and contains 0.100 mol at 108 K.
• Dilutions:
  • Diluting 19.6 mL of an 11.1 M hydroiodic acid solution to 100.0 mL results in a 2.18 M solution.
  • Adding 46.3 mL of water to 14.6 mL of a 0.624 M hydrochloric acid solution results in a 0.150 M solution.
• Mixture Pressures:
  • With 0.4 atm oxygen and 0.3 atm nitrogen, the total pressure is 0.7 atm.
  • With 1.2 atm total and 0.5 atm oxygen, nitrogen's atmosphere is 0.7 atm.
  • With 1.5 atm total and 0.7 atm nitrogen, oxygen's atmosphere is 0.8 atm.

Module 3 (Chapter 9)

• Solvent and Solute Identification:
  • Water (solvent), Carbon Dioxide (solute).
  • Water (solvent), Zinc Acetate (solute).
  • Water (solvent), Acetone (solute).
  • Water (solvent), Hydrogen Selenide (solute).
• Compound Miscibility:
  • Acetone and Water (miscible).
  • 1-Ethoxybutane and Water (immiscible).
  • Methyl Acetate and Water (partially miscible).
  • Acetic Acid and Water (miscible).
• Molarity Calculations:
  • 2.70 L with 0.780 mol of solute has a molarity of 0.288 M.
  • 25 mL with 0.5 moles of solute has a molarity of 20 M.
  • A 4.1 g sample of NaCl in 2.1 L water has a molarity of 0.033 M.
  • 30 mL with 0.243 mol of solute has a molarity of 8.1 M.
• Solution Concentrations Conversions:
  • A 6.48 gram sample of Zinc Fluoride in 250mL H₂O solution to 2.6%(w/v)
  • A 7.22 ml sample of Propyl Alcohol in 100mL H₂O has a concentration of 7.22%(v/v)
  • A 3.40 gram sample of Cobalt Acetate in 21.1 grams of H₂O to 16.11%(w/w)
  • A 5.30 gram sample of Zinc Fluoride in 10.52mL H₂O comes out to be 50.38%(w/v)
• Filtration properties: Water, small molecules, and ions pass through a dialyzing membrane, while colloid particles do not.
• Acids: Substances which donate a proton (hydrogen ion).
• Bases: Substances which accept a proton.
• Brønsted-Lowry Acid/Base Pairs:
  • In an indicated reaction the (Acid) is HSO₃⁻ and the (Base) is NO₂.
  • In an indicated reaction the (Acid) is HSO₃ and the (Base) is ClO.
  • In an indicated reaction the (Acid) is HCN and the (Base) is HSO₃.
  • In an indicated reaction the (Acid) is HNO₂ and the (Base) is CO₃²⁻.
• Acidic, Neutral, or Basic: pH = 3 (acidic), pH = 5 (acidic), pH = 7 (neutral), and pH = 12 (basic).
• pH Calculations:
  • [H₃O⁺] = 8.2 x 10⁻⁹ M comes out to a pH of 8.09.
  • [OH⁻] = 4.5 x 10⁻⁸ M comes out to a pH of 6.65.
  • [H₃O⁺] = 2.1 x 10⁻² comes out to a pH of 1.68.
• Hydronium Concentrations Calculations:
  • pH = 5 has a concentration of [H₃O⁺] = 1 x 10⁻⁵ M.
  • pH = 3.2 has a concentration of [H₃O⁺] = 6.31 x 10⁻⁴ M.
  • pOH = 9.3 has a concentration of [H₃O⁺] = 2.0 x 10⁻⁵ M.
  • pOH = 11.3 has a concentration of [H₃O⁺] = 2.0 x 10⁻¹² M.
• Ion Concentrations Calculations:
  • When [H₃O⁺] = 1 x 10⁻⁶ M is run it becomes: [OH⁻] = 1 x 10⁻⁸ M.
  • When [H₃O⁺] = 1 x 10⁻¹M is run it becomes: [OH⁻] = 1 x 10⁻¹³ M.
  • When [OH⁻] = 1 x 10⁻¹² M is run it becomes: [H₃O⁺] = 1 x 10⁻² M.
• Relating Aqueous pH:
  • High [H₃O⁺] equals a low pH. -Low [H₃O⁺] equals a high pH.
  • High [OH⁻] equals a high pH.
  • Low [OH⁻] equals a low pH.
• Buffers: these are solutions that have the ability to resist changes when acids or bases are added.