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9/23/2011
Chapter 3
Mass Relationships
In Chemical Reactions
Chapter 3
• Measuring atomic and molecular masses
– Mass spectrometry
• The mole
– Scaling molecular mass to a size we can weigh
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Chemical formulas
Experimentally determining molecular formulas
Equations representing chemical reactions
Mass relationships in chemical reactions
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Atomic Mass Units
• Units for the mass of atoms and molecules (amu)
• Defined so that mass of 126C is exactly 12
• Gives most of common elements an integral mass )or
close)
• Not all, though
Mass Spectrometer
• Most accurate way of determining masses
One type of
Mass spec
– Expose gas to beam of electrons
• Strips electrons from particles to form ions
• Breaks some molecules into fragments
– Pass beam of charged fragments through magnetic field
– Curvature of path depends on mass of fragments
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Chlorine Mass Spectrum
• Part of chlorine mass spectrum looks like below.
• X-axis units are in amu
• Why are there two peaks?
Mass Averages
• Chlorine has 2 isotopes
37 Cl: 24.2%
– 3517Cl: 75.8%
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– Different number of neutrons
• Elements have several isotopes in nature
• The atomic mass is an average of the isotopes
– For chlorine:
– (.758)(35 amu) + (.242)(37 amu) = 35.5 amu
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Atomic Mass
• Look at Periodic Table
– Which of the following probably has more than
one common isotope?
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N
F
Al
Ge
Molecular Mass
• Add atomic masses of atoms in molecule
• H2O: 2(1.008 amu) + 16.00 amu = 18.02 amu
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The Mole
• Problem: difficult to weigh one atom or molecule
– Need to scale up masses to a size we can deal with in the
laboratory
• 1 mole = amount of substance that contains as many
particles as exactly 12 g of C-12
• Number of particles in a mole = Avogadro’s Number
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6.022 x 1023 1/mole
Know this
Molar Mass
• Molar mass = mass of one mole in grams
– Units = g/mole
– H: 1.008 g/mole
Important
• Molar mass has same value as molecular mass
– Units are different
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Molar Mass
• Use Factor-Label method to convert g <-> mole
• Question: How many moles in 46.0 g water
• Question: How many g in 0.76 moles of NaCl?
• Question: How many molecules in 100.0 g water?
Chemical Formulas
• Molecular Formula: exact number of
atoms in a molecule
– Example: ethylene glycol
– C2H6O2
• Empirical Formula: simplest whole
number ratio of atoms in molecule
Black: Carbon
Red: Oxygen
Yellow: Hydrogen
– Ethylene glycol: CH3O
– Empirical formula is used in determining
molecular formula
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Empirical Formulas
• What are the empirical formulas of these
compounds?
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N 2O 4
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C6H6O3
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C6H6O
Empirical Formulas
• Why do we care about empirical formulas?
– Experiments sometimes determine empirical formulas.
– First step in figuring out a molecular formula
• Experiment -> empirical formula
• Experiment -> molecular mass
• Put them together -> molecular formula
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Empirical -> Molecular Formula
• Suppose we know the empirical formula of a
compound
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C3H7
• What is the molecular formula of the compound?
Empirical -> Molecular Formula
• Mass spectrum of the compound
• Generally the peak with largest mass is the molecular mass
– Ignore the C-13 isotope peak
• This example: Molecular mass = 86 amu
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Empirical -> Molecular Formula
• Empirical formula = C3H7.
• Molecular mass = 86 amu
• Question: What is molecular formula of compound?
• Note: Only need approximate molecular mass
Percent Composition
• Example: H2O
• %H =
• %O =
• Percent compositions can be determined experimentally
• Can use percent composition to determine empirical
formulas
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Determining Empirical Formulas
• Experimental % composition of compound is:
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50.05% S
49.95% O
• What is the empirical formula of the compound?
Determining Molecular Formulas
• Molecular mass is around 64 g/mole.
• What is the molecular formula of the compound?
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Determining Molecular Formulas
• Compound contains only C, H, and O.
– In 2.034 g of compound, have the following:
• C: 1.031 g
• H: 0.0864 g
• O: 0.917 g
– The mass spectrum of compound indicates that its
molar mass is around 140 g/mole.
– What are empirical and molecular formulas of
compound?
Chemical Equations
• Express what occurs during a
chemical reaction
• 2 Al(s) + 3 Br2(l)
Reactants
Al2Br6(s)
Product(s)
– Coefficients = # molecules or # moles
– Equation must be balanced
• Same number elements on both sides
– Physical states in parentheses
• s, l, g, aq
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Chemical Equations
• Question: Observe the video of heating mercury(II) oxide.
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What is formula of mercury(II) oxide?
One of the products is metallic mercury
Other product of reaction supports combustion. What is this gas?
What is the balanced equation for the reaction?
Chemical Equations
• Reaction of metallic magnesium in aqueous HCl.
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What is name of HCl?
The gas liberated is hydrogen H2
What is formula and name of the other product?
What is the balanced equation for the reaction?
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Stoichiometry
• Relationship between amounts of reactants and
products
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2 H2 (g) + O2 (g) ----> 2 H2O (l)
• Question: How many moles of water will be
produced from the reaction of 2.52 moles oxygen
with excess hydrogen?
Stoichiometry
• 2 H2 (g) + O2 (g) ----> 2 H2O (l)
• Question: How many grams of oxygen are needed to
react with 3.40 g of hydrogen?
• Key to doing these problems
– Convert grams to moles
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Stoichiometry
• 2 CH3OH (l) + 3 O2 (g) ----> 2 CO2 (g) + 4 H2O (l)
• How many grams of water will be produced by
burning 25.0 mL methanol in excess oxygen?
– What do you need to look up to solve this problem?
– Use Wolfram Alpha
Combustion Analysis
CmHn + O2  mCO2 + (n/2)H2O
• Burn known amount of organic compound in excess
oxygen
• Measure masses of CO2 and H2O produced
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Combustion Analysis
• Combining three topics covered in this chapter
– Stoichiometry
• Use g of CO2 to determine moles of C in compound
• Use g of H2O to determine moles of H in compound
– Empirical formula determination
• Moles C and moles H  empirical formula
– Molecular formula determination
• Mass spectrum  molecular mass
• Empirical formula and molecular mass  molecular formula
Combustion Analysis
CmHn + O2  mCO2 + (n/2)H2O
• Equation above isn’t balanced
– All we care about is C-CO2 and H-H2O relationship
• 1 mole CO2 = 1 mole C in compound
• 1 mole H2O = 2 moles H in compound
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Combustion Analysis
CmHn + O2  mCO2 + (n/2)H2O
• Burn 0.820 g of an unknown hydrocarbon
• Obtain 2.70 g CO2 and 0.73 g H2O
• Mass spectrum:
• What is compound’s
molecular formula?
Limiting Reactant
• Don’t always have exact amounts of reactants to
completely use them all up
• A + 2B  C
• Initially:
– 5 A molecules
– 6 B molecules
• How many C molecules will be produced?
• Will there be any A or B molecules remaining when
reaction is over?
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Limiting Reactant
• A + 2B  C
• B molecules are used up first
• # of C molecules produced
depends on initial # of B molecules
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6 B molecules  3 C molecules
• B = Limiting reactant
• A is in excess
Initially:
5 A molecules
6 B molecules
Limiting Reactant
C4H9OH + NaBr + H2SO4  C4H9Br + NaHSO4 + H2O
74.1 g/mole
102.9 g/mole
• What are names of NaBr, H2SO4, NaHSO4?
• Want to use up all of C4H9OH
• Start with 50.0 g of C4H9OH
• What is minimum mass of NaBr necessary to use all
of C4H9OH ? (excess H2SO4)
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Limiting Reactant
• N2 (g) + 3 H2 (g)
28.0 g/mole
2.02 g/mole
2 NH3 (g)
17.0 g/mole
• React 10.0 g N2 with 10.0 g H2. How much NH3 will
be produced?
• Need to decide which reactant is limiting
– Calculate amount of product possible for both reactants
– Smallest amount identifies limiting reactant
• Use limiting reactant to calculate amount of product
Reaction Yield
• Theoretical yield = amount of product produced if all
limiting reactant is consumed
– Best you can do
• Actual yield = amount of product you actually get
– Usually the case that actual yield is less than theoretical
• Percent yield = 100*(Actual yield)/(Theoretical yield)
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Percent Yield
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Previous problem
Theoretical yield = 18.2 g NH3
Suppose actual yield = 15.3 g NH3
What is percent yield?
Reaction Yield
Reaction of hydrazine with dinitrogen tetraoxide
2 N2H4 (l) + N2O4 (g) → 3 N2 (g) + 4 H2O (g)
32.0 g/mole
92.0 g/mole
28.0 g/mole
What mass of N2 gas theoretically would result from the
reaction of 150.0 g of hydrazine and 100.0 g of N2O4?
If you actually obtain 66.4 g N2, what is percent yield?
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Primary Productivity
• Primary productivity: How quickly and to what
extent sunlight is converted into organic material by
plants during photosynthesis.
• Photosynthesis: Simple equation describes
production of sugars
light
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6CO2 + 6H2O
C6H12O6 + 6O2
• Nutrients: Other chemicals that also are required:
– Examples: Phosphorus and nitrogen
Limiting Nutrient
• Nutrients seldom used up at same rate
– Eventually, one of them may be depleted
• What will happen to primary production once one of
nutrients is depleted?
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Limiting Nutrient
• Limiting nutrient
– The nutrient in shortest supply in a particular ecosystem
– In essence the limiting reactant in primary production
• Once a nutrient is depleted, plant growth stops in
the ecosystem
Limiting Nutrient
• Phosphorus: limiting nutrient in many lakes
• How could we design an experiment to test this?
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Eutrophication
• Large increase in primary production in a lake
– Usually because phosphorus has been added
– Mainly due to fertilizers in runoff
– Reason phosphates are illegal in detergents in many states
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