Ch04_1+Combustion+Reactions

Page 2: Molecular Weight versus Molar Mass

  • Molecular Weight or Molecular Mass

    • Definition: The sum of the atomic weights of all atoms in a molecule.

    • Unit: grams/mole

    • Example: One mole of H2O has a molecular mass of 18.0 g/mol.

  • Atomic Weight or Atomic Mass

    • Definition: The atomic mass of one mole of the element.

    • Unit: grams/mole

    • Example: One mole of Cu atoms has an atomic mass of 63.5 g/mol.

Page 3: Molecular Formulas, Molecular Mass, and Stoichiometry

  • Chemical Formulas

    • Indicate the number and type of atoms in a molecule.

    • Molecular mass is calculated from the chemical formula.

  • Conversions

    • Use chemical formulas and corresponding molar mass to convert between amounts of constituents:

      • Percent composition

      • Moles

Page 4: Percent Mass

  • Composition of Pure Compounds

    • Consist of the same elements combined in the same proportions by weight.

    • Expressed as percent by weight or percent composition.

  • Example: Ethanol (C2H6O)

    • Molecular mass: 46.0 g/mol.

    • Percent composition: 52.31% C, 13.15% H, 34.72% O.

Page 5-6: Problem on Mass of Table Salt (NaCl)

  • Problem Statement

    • Determine the mass of NaCl containing 2.4 g of Na.

  • Calculations

    1. Find percentage of Na in NaCl:

      • Na: Atomic mass = 22.99 g/mol.

      • Molecular mass of NaCl = 58.44 g/mol.

      • % Na = (22.99 g / 58.44 g) × 100 = 39.34%.

    2. Set ratio to find mass of NaCl:

      • (2.4 g Na / x g NaCl) = (39.34 g Na / 100 g NaCl)

      • Result: x = 6.10 g NaCl.

Page 7-8: Problem on Finding Mass Percent of Cl in C2Cl4F2

  • Problem Statement

    • Find the mass percent of Cl in C2Cl4F2.

  • Calculations

    1. Determine molecular mass of C2Cl4F2:

      • 2 × (12.01 g/mol C) = 24.02 g,

      • 4 × (35.45 g/mol Cl) = 141.80 g,

      • 2 × (19.00 g/mol F) = 38.00 g,

      • Total = 203.8 g/mol.

    2. Calculate mass percent of Cl:

      • Cl total mass = 141.80 g.

      • Mass percent Cl = (141.80 g / 203.8 g) × 100 = 69.58%.

Page 9-10: Strategy for Determining Empirical and Molecular Formulas

  • Step 1: Determine empirical formula from percent composition.

    • Convert percent to mass, mass to moles.

    • Divide by smallest number of moles; if not whole numbers, multiply to get whole numbers.

  • Step 2: Determine molecular formula from empirical formula.

    • Calculate unit mass of empirical formula, divide by molecular mass to find factor (n). Populate molecular formula accordingly.

Page 11-14: Problem on Determining Molecular Formula of Compound with B and H

  • Problem Statement

    • A compound of B (81.10%) and H with molecular mass of 53.3 g/mol.

  • Calculations

    1. Determine empirical formula:

      • B = 81.10 g, H = 18.90 g in a 100 g sample.

      • Moles: 7.502 mol B, 18.75 mol H.

      • Ratio: 5 H to 2 B atoms (Empirical formula = B2H5).

    2. Determine molecular formula:

      • Empirical mass = 26.66 g/unit; 53.0 g/26.66 g = 2.

      • Molecular formula = B4H10.

Page 15-16: Quick Review of Organic Chemistry

  • Classifying Compounds:

    • Organic compounds were initially from living origins; inorganic from nonliving sources.

    • Organic compounds are easier to decompose compared to inorganic compounds.

  • Organic Compounds:

    • Mainly composed of C and H, sometimes O, N, P, S.

    • Carbon forms four covalent bonds.

Page 17: Simple Organic Compounds

  • Characteristics of Carbon:

    • Forms limitless chains, rings, and branched structures.

    • Compound types: Hydrocarbons (only C and H); simplest hydrocarbon is methane (CH4).

Page 19: Combustion Analysis

  • Description of Technique:

    • Burn a known mass of compound, measure production.

    • Commonly used for analyzing C, H, O compounds.

    • Calculate empirical formula using mass of products obtained.

Page 20-22: Determining Empirical and Molecular Formulas from Combustion Reaction

  • Problem Statement:

    • A hydrocarbon burned produces a known mass of CO2 and H2O. Find empirical formula.

  • Steps:

    1. Write reaction equation.

    2. Relate CO2 and H2O produced to carbon and hydrogen in the hydrocarbon.

    3. Determine masses from combustion.

Page 23: Answer to the Empirical Formula Problem

  • Empirical Formula Calculations:

    • Find moles of C and H from combustion products.

    • Calculate ratios and convert to whole numbers; results in empirical formula C3H4.

Page 24-27: Problem on Determining the Molecular Formula of CxHyOz

  • Problem Statement:

    • A compound composed of C, H, and O ignites producing CO2 and H2O. Given mass: determine its formula.

  • Steps to Solve:

    1. Write the chemical equation for combustion.

    2. Use mass relationships to identify C, H, and O contributed.

    3. Calculate empirical formula from masses determined.

    4. Find the molecular formula based on the empirical mass and the given molecular mass.

    • Resulting molecular formula = C18H20O2.