I'Atomic Masses

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Chapter 3 Stoichiometry

• I. Atomic Masses
• Mass spectrometers and atomic masses
• Carbon is used as the standard for atomic masses
• In 1961, 12C was assigned a mass of exactly 12
  atomic mass units (amu)
• A mass spectrometer can compare other elements or
  isotopes to 12C
• 13C is found to be 1.0836129 times as massive as
  12C
• Mass of 13C (1.0836129)(12 amu) 13.003355 amu

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• Average Atomic Masses
• Why isnt the mass of Carbon listed at exactly 12
  on the periodic table?
• Natural carbon 12C(98.89), 13C(1.11), and 14C
  (negligible)
• Avg. Mass (0.9889)(12 amu) (0.0111)(13.0034
  amu) 12.01 amu
• No individual atoms have this mass
• On average, all natural carbon atoms have this
  mass
• 12.01 amu is the correct value to use for
  calculations involving carbon
• Sample Ex. 3.1 What is avg. mass of Cu?
• 69.09 63Cu (62.93 amu) and 30.91 65Cu (64.93
  amu)
• Calculate the Mass (in amu) of 75 atoms of Al
• a. Determine the mass of 1 Al atom 1 atom of Al
  26.98 amu
• b. Use the relationship as a conversion factor

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• II. The Mole
• We use a package for atoms and molecules called a
  mole
• A mole
• a. the number of Carbon atoms in 12 g of 12C
• 6.022 x 1023 units Avogadros Number
• The amount of an element equal to its atomic mass
  
• 1 mole of natural C atoms weighs 12.01 g and has
  6.02 x 1023 atoms
• 1 mole of He atoms weighs 4.003 g and has 6.02 x
  1023 atoms
• 1 mole of Al atoms weighs 26.98 g and has 6.02 x
  1023 atoms
• Sample Ex. 3.2 What is the mass of 6 Americium
  atoms?
• Calculating moles, mass, and atoms
• 1. Sample Ex. 3.3 atom / moles in 10g Al

Cu
I2
Hg
Al
Fe
S
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• 2. Sample Ex. 3.4 5.68 mg Si ? atoms Si
• Molar Mass
• 1. The molar mass is the mass in grams of one
  mole of a compound
• 2. The relative weights of molecules can be
  calculated from atomic masses
• a. Water H2O 2(1.008 g/mol) 16.00 g/mol
  18.02 g/mol
• b. 1 mole of H2O will weigh 18.02 g, therefore
  the molar mass of H2O is 18.02 g
• 1 mole of H2O will contain 16.00 g of oxygen and
  2.02 g of hydrogen
• Calculations with molar mass
• a. Sample Ex. 3.6 Juglone C10H6O3 Find Molar
  mass and moles in 0.0156 g

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• Sample Ex. 3.8 How many molecules of isopentyl
  acetate (C7H14O2) and how many carbon atoms are
  released in one bee sting (1 mg)?
• D. Mass Percent
• Composition of a substance can be given multiple
  ways
• Molecular formula gives ratio of the number of
  each element
• The Mass Percent gives the composition by mass of
  each element
• Example Ethanol, C2H6O has a molar mass of 46.07
  g

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• Sample Ex. 3.10 Determine the Mass Percent of
  C14H20N2SO4.
• III. Determining the Formula of a Compound
• Suppose we make a new compound composed of C,H,
  and N
• We probably know what it was we were trying to
  make
• That doesnt mean we succeeded
• We must have evidence that the compound we made
  is what we wanted
• Lets turn a weighed amount of the compound into
  CO2, and H2O

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• B. Analyzing the Results
• 0.1156 g of our new compound gives 0.1638 g CO2
  and 0.1676 g H2O
• (1C)(12.01g/C) (2 O)(16.00g/O) 44.01 g total
  for CO2
• (2H)(1.008g/H) (1 O)(16.00g/O) 18.02 g total
  for H2O
• Next, we find out how much carbon and hydrogen
  was in our sample
• Then, we determine the Mass Percents for our new
  compound
• Finally, we need to determine the molecular
  formula of our compound
• The easiest way to do this is to work with a
  theoretical 100.00g
• We can then change the percent masses to grams of
  each element
• 38.67 g Carbon, 16.22 g Hydrogen, 45.11 g Nitrogen

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• Molecular formulas compare moles, not grams
• Write the formula with these numbers of moles,
  then divide by the smallest number to get these
  subscripts as whole numbers
• This is called the Empirical Formula smallest
  whole number ratio of elements in the formula
• Molecular Formula is the actual number of each
  element
• Could be C2H10N2 or C3H15N3 etc
• (C1H5N1)n all possible molecular formulas for
  this empirical formula
• We can use a Mass Spectrometer on our compound to
  find its molecular mass is 31.06 g/mol. In this
  case CH5N is correct molecular formula.
• C. Sample Ex. 3.113.13 give more practice

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• Hints for determining empirical and molecular
  formulas
• Convert mass to grams of the element in 100
  total grams of sample
• Determine the Empirical Formula of an unknown if
  its Percent Composition is 47 Carbon, 47 Oxygen
  and 6.0 Hydrogen
• Change these masses to moles by using the atomic
  mass of each element
• Divide each number of moles by the smallest to
  get a small ratio
• Round off to a whole if molar s are close to
  that whole number
• Multiply the whole ratio by a factor to get all
  numbers to whole numbers
• Multiply empirical formula by factor needed to
  give the correct molar mass
• Molar Mass of unknown 204.2 g/mol but C4H6O3
  102.09 g/mol
• Correct Molecular Formula C4H6O3 x 2 C8H12O6