CP Chemistry Semester 2 Final Review Notes

Formula Writing, Naming Compounds & Equations:

• Elements in columns (groups) on the Periodic Table have similar chemical properties.
• Elements in rows on the Periodic Table have repeating or cyclic chemical properties.
• Naming compounds:
  • KBr: potassium bromide
  • MgCl2: magnesium chloride
  • FeS: iron (II) sulfide
  • CuF3: copper (III) fluoride
  • Pb(NO3)2: lead (II) nitrate
  • As2O3: diarsenic trioxide
  • BaSO4: barium sulfate
  • N2: nitrogen gas
  • H3PO4: phosphoric acid
  • SO2: sulfur dioxide
• Neutral chemical formulas:
  • Silicon tetrachloride: SiCl4
  • Oxygen gas: O2
  • Diphosphorus pentoxide: P2O5
  • Sodium chloride: NaCl
  • Cupric iodide: CuI2
  • Stannous oxide: SnO
  • Ammonium phosphate: (NH4)3PO4
  • Manganese (IV) oxide: MnO2
  • Iron (III) hydroxide: Fe(OH)3
  • Potassium sulfate: K2SO4
• Matching:
  • I) polyatomic ion - f) PO4-3
  • II) multivalent ion - b) Fe+2 or Fe+3
  • III) simple anion - d) S-2
  • IV) simple cation - a) Ca+2
  • V) element - g) Au
  • VI) compound - c) NaOH
  • VII) diatomic element - e) Cl
• Matching names with chemical formulas:
  • a) acid - V) H2SO4
  • b) base - VIII) Mg(OH)2
  • c) oxide - IV) Al2O3
  • d) salt - I) NaNO3
  • e) metal - II) Ca
  • f) metallic carbonate - III) MgCO3
  • g) monatomic molecule - VI) Ne
  • h) diatomic molecule - VII) Cl2
• In 2Ca3(PO4)2:
  • Calcium atoms: 6
  • Polonium atoms: 0
  • Phosphate ions: 4
  • Oxygen atoms: 16
  • Phosphorus atoms: 4
  • Calcium phosphate molecules: 2
  • Charge of calcium: +2
  • Charge of phosphate ion: -3
  • Charge of calcium phosphate: 0
• Balancing equations:
  • 2 NaBr + 1 Cl2 \rightarrow 2 NaCl + 1 Br2
  • 1 Ti2O3 + 3 CO \rightarrow 2 Ti + 3 CO_2
  • 2 Ga(OH)3 + 2 H2SO4 \rightarrow 1 Ga2(SO4)2 + 6 H_2O
  • 2 C6H6 + 15 O2 \rightarrow 12 CO2 + 6 H_2O
  • 2 AlCl3 + 3 K2CO3 \rightarrow 1 Al2(CO3)3 + 6 KCl
  • 1 Ba(OH)2 + 1 CO2 \rightarrow 1 BaCO3 + 1 H2O
• Given: ^{23}_{11}Na
  • Number of protons: 11
  • Atomic mass: 23
  • Number of neutrons: 12
  • Electrons in a neutral atom: 11
  • Atomic number: 11
• Four ways to identify a chemical change:
  • Bubbles
  • Color change
  • Temperature change
  • Change in smell/taste
  • Precipitate formation
• Precipitate: solid formed in a reaction
• Seven diatomic gases: H2, O2, N2, F2, Cl2, Br2, I2
• Reaction types:
  • Combustion: CH4 + O2 \rightarrow H2O + CO2
  • Single Replacement: LiCl + F2 \rightarrow LiF + Cl2
  • Decomposition: CaCO3 \rightarrow CaO + CO2
  • Double Replacement: FeCl3 + Na2CO3 \rightarrow Fe2(CO3)3 + NaCl
  • Composition/Synthesis: H2 + Cl2 \rightarrow HCl

Moles and Stoichiometry

• Avogadro’s number: 6.02 x 10^{23}
• Standard volume: 1 mol = 22.4 L at STP (Standard Temperature and Pressure)
• For 2(NH4)2SO4:
  • Sulfur atoms: 2
  • Hydrogen atoms: 16
  • Molar mass: 132.14 g/mol
• Moles of CO2 in 98.3 g: 98.3 g CO2 * (1 mole CO2 / 44.01 g CO2) = 2.23 moles CO2
• Grams in 3.5 moles of CH4: 3.5 moles CH4 * (16.05 g CH4 / 1 mole CH4) = 56.175 g CH4
• Molecules in 2.3 g of HCl: 2.3 g HCl * (1 mole HCl / 36.46 g HCl) * (6.02 x 10^{23} molecules HCl / 1 mole HCl) = 3.80 x 10^{22} molecules HCl
• Liters of CO from 230 grams of CO: 230 g CO * (1 mole CO / 28.01 g CO) * (22.4 L CO / 1 mole CO) = 183.93 L CO
• Molecules of H2O in 47.3 liters: 47.3 L H2O * (1 mole H2O / 22.4 L H2O) * (6.02 x 10^{23} molecules HCl / 1 mole H2O) = 1.27 x 10^{24} molecules HCl
• Combustion of C6H6 with excess oxygen:
  • Equation: 2 C6H6(g) + 15 O2(g) \rightarrow 12 CO2(g) + 6 H_2O (g)
  • Grams of H2O produced from 24.3 L of C6H6: 24.3 L C6H6 * (1 mole C6H6 / 22.4 L C6H6) * (6 mole H2O / 2 mole C6H6) * (18.02 g H2O / 1 mole H2O) = 58.65 g H2O
  • Grams of CO2 produced from 260 g of oxygen gas: 260 g O2 * (1 mole O2 / 32.00 g O2) * (12 mole CO2 / 15 mole O2) * (44.01 g CO2 / 1 mole CO2) = 286.07 g CO2
• Limiting reagent:
  • Reaction: C2H4 + 2O2 \rightarrow 2CO + 2H2O
  • 2.8 moles of C2H4 reacts with 6.5 moles of O2.
  • 2.8 moles C2H4 * (2 mole O2 / 1 mole C2H4) = 5.6 moles O2
  • C2H4 is the limiting reactant.
• Moles of water produced:
  • 2.8 moles C2H4 * (2 mole H2O / 1 mole C2H4) = 5.6 moles H2O

Gas Laws

• Boyle’s Law:
  • Equation: P1V1 = P2V2
  • Relationship: Pressure and volume are inversely related.
• Gay Lussac’s Law:
  • Equation: P1/T1 = P2/T2
  • Relationship: Pressure and temperature are directly related.
• Charles Law:
  • Equation: V1/T1 = V2/T2
  • Relationship: Volume and temperature are directly related.
• Ideal Gas Law:
  • Equation: PV = nRT
• Units of pressure: Pressure, Volume, Temperature, and Moles
• Sample of hydrogen gas:
  • Initial: 25.0 ml at 2.5 atmospheres
  • Final volume: 20.0 ml
  • P2 = (P1V1) / V2 = (25.0 mL * 2.5 atm) / 20.0 mL = 3.125 atm
• Absolute zero:
  • Temperature at which all motion of particles is minimal.
  • 0 K or -273 °C
• Standard temperature: 0 °C or 273 K
• Kelvin conversion: 25 degrees Celsius = 273 + 25 = 298 K
• Gas heated from 35 °C to 56 °C:
  • Initial volume: 25.0 L
  • V2 = (V1 * T2) / T1 = (25 L * (56 + 273 K)) / (35 + 273 K) = 26.70 L
• Hot air balloon:
  • Initial volume: 2500 L at STP
  • Final conditions: -40 °C and 0.35 atm
  • V2 = (P1 * V1 * T2) / (T1 * P2) = (1 atm * 2500 L * 233 K) / (273 K * 0.35 atm) = 6096.28 L
• Pressure exerted by 0.563 moles of gas at 25°C in a 0.650 L vessel:
  • P = (nRT) / V = (0.563 moles * 0.0821 (Latm) / (Kmol) * 298 K) / 0.650 L
• Volume of gas at 0.921 atm moved from 0.868 atm container at 80ml:
  • V2 = (P1V1) / P2 = (0.868 atm * 80 mL) / 0.921 atm = 75.40 mL
• Moles of gas in 500 mL at 32°C and 0.93 atm:
  • n = (PV) / (RT) = (0.5 L * 0.93 atm) / (0.0821 (Latm) / (Kmol) * 305 K) = 0.019 moles
• Pressure definition: Collisions of molecules against the walls of the container.
• Marshmallows in a vacuum chamber: Pressure decreased, marshmallows expanded.
• Inversely related pressure and volume: Decreasing volume increases pressure.
• Increasing kinetic energy of a gas: Increase the temperature.
• Increased kinetic energy with constant volume: Pressure increases.
• Decreased kinetic energy with constant pressure: Volume decreases.

Thermodynamics

• Reaction: Energy + NH4OH ----> NH4+1 + OH-
• Endothermic or exothermic? (Requires more context to determine; assume endothermic if energy is listed as a reactant)
• Diagram showing change in enthalpy: (Diagram not provided, needs visual representation)
• Reactants or products with more energy? Products (for endothermic)
• Change positive or negative? Positive (for endothermic)
• Reactants or products more stable? Reactants (for endothermic)
• Enthalpy: The total heat content of a system.
• Activation energy: Amount of energy needed for reactants to undergo a specified chemical reaction.
• Catalysts speed up reactions by providing an alternative pathway with lower activation energy.
• Endothermic reaction: Work is done on the system.
• Exothermic reaction: Work is done on the surroundings.
• Temperature change:
  • Exothermic: Heats up
  • Endothermic: Cools down
• Specific heat of metal cylinder:
  • Given: mass = 10 g, initial temp = 96°C, final temp = 42°C, enthalpy loss = 242.46 J
  • q = m * C * (Tf – Ti)
  • C = q / (m * (Tf – Ti)) = 242.46 J / (10g * (42 - 96)) = 0.449 J/g

Acids and Bases

• Properties of acids:
  • Sour
  • pH < 7
  • Turn litmus red
  • Phenolphthalein is colorless
  • Not slippery
  • Produce H+
• Properties of bases:
  • Bitter
  • pH > 7
  • Turn litmus blue
  • Phenolphthalein is pink
  • Slippery
  • Produce OH-
• Acid-Base Definitions:
  • Acids: Donate proton (H+) OR dissociate in water and produces H+
  • Bases: Accept proton (H+) OR dissociate in water and produces OH-
• Equations:
  • 2HCl + Mg \rightarrow MgCl2 + H2 (Single Replacement)
  • H2SO4 + 2NaOH \rightarrow Na2SO4 + 2H_2O (Neutralization)
  • 2HCl + CaCO3 \rightarrow CaCl2 + H2O + CO2 (Double Replacement and Decomposition)
• Acid + Metal = SALT + HYDROGEN GAS
• Acid + Base = SALT + WATER
• Acids give off hydronium ions in water (H+ or H3O+)
• HCl ionization in water: produces H^+ and Cl^-
• Bases give off hydroxide ions in water (OH-)
• pH values:
  • Weak acid: 6
  • Strong acid: 1
  • Weak base: 8
  • Strong base: 13
• Acids and Bases - Formulas and Location:
  • Sulfuric Acid - H2SO4 - Acid Rain
  • Nitric Acid - HNO_3 - Acid Rain
  • Acetic Acid - HCH_3COO - Vinegar
  • Milk of Magnesia - Mg(OH)_2 - Laxative
  • Sodium Hydroxide - NaOH - Cleaning Products/Lye
  • Calcium Hydroxide - Ca(OH)_2 - Lime

Nuclear

• Radiation Types:
  • Alpha (α): Symbol - α, Shielding - hand, Size - large, Dangerous - Internal (Very), External (No)
  • Beta (β): Symbol - β, Shielding - aluminum, Size - small, Dangerous - Less damage to living tissue
  • Gamma (γ): Symbol - γ, Shielding - lead, Size - very small, Dangerous - Least dangerous in terms of tissue damage
• Factors Determining Radiation Danger:
  • Dose (quantity)
  • Exposure Time
  • Area Exposed
  • Tissue Type
• Half-Life: Time for half a sample to decay
• Short Half-Life = More Dangerous (heavy dose in short time)
• Alpha Decay of Uranium-235: ^{235}{92}U \rightarrow ^42He + ^{231}_{90}Th
• Beta Decay of Uranium-235: ^{235}{92}U \rightarrow ^0{-1}e + ^{235}_{93}Np
• Isotopes: Atoms with same # of protons, different # of neutrons
• Fissionable Uranium: Uranium-238
• Fission vs. Fusion:
  • Fission: Splitting large isotope into smaller isotopes
  • Fusion: Fusing smaller isotopes into a larger isotope
• Radon Gas Source: Natural breakdown of radioactive material (uranium, thorium) in soil
• Radiocarbon Dating Range: Up to 50,000 years (half-life = 5730 years)
• Fission Reaction Bullet: Neutron
• Ionizing Radiation: Energy released by atoms (EM waves or particles)
• Radiation Measurement: Geiger Counter
• Nuclear Power Plants:
  • Process: Fission
  • Pros: Cheap energy, no carbon emission
  • Cons: Environmental, meltdown risk, water demand, nuclear waste