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Chapter 6: Chemical Reactions: Mole and Mass Relationships

General Information

• Lecture Title: Chapter Six Chemical Reactions: Mole and Mass Relationships

• Source: Fundamentals of General, Organic, and Biological Chemistry 8th Edition, © 2017 Pearson Education, Inc.

• Authors: McMurry, Ballantine, Hoeger, Peterson, Christina A. Johnson, University of California, San Diego

Lecture Outline

• 6.1 The Mole and Avogadro’s Number

• 6.2 Gram-Mole Conversions

• 6.3 Mole Relationships and Chemical Equations

• 6.4 Mass Relationships and Chemical Equations

• 6.5 Limiting Reagent and Percent Yield

Key Concepts to Review

• Problem Solving: Unit Conversions and Estimating Answers (Section 1.10)

• Molecular Formulas and Formula Units (Sections 3.8 and 4.6)

• Balancing Chemical Equations (Section 5.2)

6.1 The Mole and Avogadro’s Number

Definition of Mole

• Mole: The amount of a substance whose mass in grams is numerically equal to its molecular or formula weight.

• Molar Mass: Mass in grams of 1 mole of a substance, numerically equal to molecular weight in amu.

Avogadro’s Number

• One mole of any substance contains 6.022 × 10^23 formula units, known as Avogadro’s number (NA).

Importance of Molar Mass

• Used to convert between mass and moles.

Molecular and Atomic Weight

• Atomic Weight: Average mass of an element’s atoms.

• Molecular Weight (MW): Average mass of a substance’s molecules. Calculated by summing atomic weights of all atoms in a molecule.

Example Calculations

• Example: The molecular weight of ethylene (C2H4) is 28.0 amu.

• Equal mass ratio in samples signifies the same number of molecules/formula units.

6.2 Gram-Mole Conversions

Molar Mass as a Conversion Factor

• Molar mass facilitates conversion between mass and number of moles.

• Example: ibuprofen (C13H18O2) with a molecular weight of 206.3 amu. If we have 0.082 mol of ibuprofen, we can convert it to grams using its molar mass.

6.3 Mole Relationships and Chemical Equations

Understanding Chemical Equations

• Coefficients in balanced chemical equations represent the moles of each reactant and product.

• Use mole ratios to simplify calculations relative to reactants and products.

Example: Synthesis of Ammonia

• Reaction: 3 H2 + 2 N2 → 2 NH3

  • Mole ratios: 3 moles H2 to 2 moles N2.

6.4 Mass Relationships and Chemical Equations

Coefficients and Weights

• Coefficients indicate mole-to-mole relationships.

• Mass of reactants is typically measured in grams.

• Example: In the reaction 3 NO2 + H2O → 2 HNO3 + NO, find mass produced based on reactants.

6.5 Limiting Reagent and Percent Yield

Limiting Reagent

• Limiting Reagent: Reactant that is completely consumed first; limits the amount of product formed.

• Theoretical Yield: Maximum amount of product possible based on the limiting reagent.

• Actual Yield: Amount of product actually formed during a reaction.

• Percent Yield = (Actual Yield / Theoretical Yield) × 100.

Example: Combustion of Acetylene

• Reaction: 2 C2H2 + 5 O2 → 4 CO2 + 2 H2O

• Theoretical yield of CO2 can be calculated from the actual yield to determine percent yield.