• Explain the mole concept, and convert between grams, moles, and atoms and molecules.
• Determine mass percent composition of a sample from experimental data.
• Determine mass percent composition of a compound from its formula.
• Determine empirical formula of a compound from its mass percent composition.
• Use chemical equations to predict amounts of reactant consumed or product produced in a chemical change.
• Define limiting reagent and identify the limiting reagent in a chemical equation.
• Define and distinguish theoretical yield, actual yield, and percent yield. Compute percent yield from actual and theoretical yields.
• Define molarity. Relate molarity, moles of solute, and liters of solution.
• Explain how a solution of a desired concentration can be prepared from solid solute or from a more concentrated stock solution.

• Why use moles?
  • many properties depend on the number of molecules in the sample, not on the mass of the sample. How can the number of molecules in the sample be measured?
  • solution: convert sample mass to daltons to molecules, using the molecular weight
  • now you can count molecules in a pure sample just by weighing the sample.
  • problem: daltons are very small units, and inconvenient; can we convert grams to molecules without using them?
  • solution: define a new unit, the mole (SI unit for amount of substance)

• Defining moles
  • 1 mole of molecules has a mass equal to the molecular weight in grams.
  • examples
    • 1 mole H₂O is the number of molecules in 18.015 g H₂O
    • 1 mole H₂ is the number of molecules in 2.016 g H₂.
  • 1 mole of atoms has a mass equal to the atomic weight in grams.
  • 1 mole of particles = 6.02214 x 10²³ particles for any substance!
    • Prove this by converting 1 mole to grams to daltons to particles for any substance.
• vocabulary
  • molar mass is the mass of one mole of a substance
    • bridge between moles and grams in unit conversion problems
  • Avogadro's number is the number of molecules in one mole for any substance
    • bridge between molecules and moles in unit conversion problems

Using chemical formulas

• molecular formulas give atom-to-atom and mole-to-mole ratios
  • example: molecular formula C₆H₁₂O₆

    atom-to-atom ratios	atom-to-molecule ratios	mole-to-mole ratios (elements)	mole-to-mole ratio (compound)
    6 atoms C 12 atoms H 6 atoms C 6 atoms O 12 atoms H 6 atoms O	6 atoms C 1 molecule 12 atoms H 1 molecule 6 atoms O 1 molecule	6 mol C 12 mol H 6 mol C 6 mol O 12 mol H 6 mol O	6 mol C 1 mol C6H12O6 12 mol H 1 mol C6H12O6 6 mol O 1 mol C6H12O6

• problems that ask you to relate one substance to another require mole-to-mole ratios
  • examples
    • How many grams of H₂ can be obtained from the electrolysis of 10.0 g of H₂O?
    • How many grams of CuO can be made from a piece of copper wire weighing 0.2134 g?
    • 2.04 g of carbon reacts with 5.44 g of O₂ to form 7.48 g of compound. How many atoms of O per atom of C are in this compound?

Chemical formulas and elemental composition

• predicting the amount of product produced (theyield) from amounts of more than one reactant
  • if reactants aren't mixed in the right mole ratio, yield is limited by one of the reactants
  • strategy
    1. compute yield for each reactant
    2. the lowest yield is the theoretical yield
    3. the reactant giving the lowest yield is the limiting reactant.
  • examples
    • Ethanol (C₂H₅OH) is synthesized for industrial use by the following reaction, carried out at very high pressure:
      
      

      C2H4(g) + H2O(g)

      C2H5OH

      What is the maximum amount of ethanol that can be produced when 1.0 kg of ethylene (C₂H₄) and 1.0 kg of steam are placed in the reaction vessel?

    • Some of the acid in acid rain is produced by the following reaction:
      
      3 NO₂(g)+ H₂O(l) 2HNO₃(aq) + NO(g)
      A falling raindrop weighing 0.05 g comes into contact with 1 mg of NO₂. How much HNO₃ can be produced?
• why doesn't actual yield = theoretical yield?
  • impurities in reactants
  • mechanical losses
  • side reactions
    synthesis of 2, 4, 5-T
  • reversible reactions

    H2(g) + Cl2(g) 2 HCl(g)	actual yield close to theoretical yield
    H2(g) + I2(g) l 2 HI(g)	actual yield is much less than theoretical yield!

• calculating percentage yield
  1. the actual yield must be given
  2. compute the theoretical yield
  3. percentage yield = (actual yield/theoretical yield) x 100%
  • examples
    • 0.12 g of ozone is produced per gram of oxygen gas exposed to an electric arc. What is the percent yield of ozone?
    • The fertilizer ammonium nitrate is synthesized from ammonia and nitric acid. If 17 tons of ammonia produced 63 tons of fertilizer, what is the percent yield?
• predicting the amount of excess reactants
  • strategy
    1. find how many moles of each reactant is present before the reaction
    2. identify the limiting reactant
    3. find how much of each reactant is consumed by converting moles of limiting reactant to moles of each of the nonlimiting reactants
    4. final moles of reactant = initial moles - moles consumed
    • example
      • The following reaction is used to extract gold from pretreated gold ore:
        

        2 Au(CN)2-(aq) + Zn(s)

        2 Au(s) + Zn(CN)42-(aq)

        If 0.654 g of powdered Zn is added to a solution that contains 2.52 g of Au(CN)₂-, what masses of reactants will be left over after the reaction is complete?

  • caveats about interpreting these calculations
    • chemical equations give only mole ratios between reactants and products
    • equations are macroscopic - they don't represent actual molecular level events!
      • example: reactions with intermediates
    • equations can depict impractical or impossible reactions
      • example: formation reactions for organic molecules
    • some reactions don't run to completion: watch for the double harpoon

Copyright © 1997-2005 by Fred Senese
Comments & questions to fsenese@frostburg.edu
Last Revised 02/23/18.URL: http://antoine.frostburg.edu/chem/senese/101/moles/index.shtml