Masterton and Hurley Chapter 3

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Chapter 3 Mass Relations in Chemistry
Stoichiometry
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3.1 Atomic Masses

• Atomic mass (atomic weight) The atomic mass
  of an element indicates how heavy, on average, an
  atom of an element is when compared to an atom of
  another element
• Atomic mass units (amu) the units for atomic
  masses on the periodic table

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The Carbon-12 Scale

• Mass of one 12C atom 12 amu (exactly)
• Note that 12C and C-12 mean the same thing

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Atomic Masses and Isotopic Abundances

• Mass spectrometer a device used to
  experimentally determine the atomic mass of an
  atom
• Isotopic abundances the percentage of each
  isotope that exists in nature (also, determined
  using the mass spec.)

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Figure 3.1 Mass Spectrometer

• A mass spectrometer is used to determine atomic
  masses

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Figure 3.2 Mass Spectrum of Cl

• The area under the peak in the mass spectrogram
  gives the isotopic abundance

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Atomic Mass Calculations
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Example 3.1
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Masses of Individual Atoms Avogadros Number

• Avogadros Number The number of atoms that is
  equal to the atomic mass of any element
• NA 6.02 X 1023
• For Example
• 6.02x1023 H atoms in 1.008 grams of H (atomic
  mass of H 1.008)

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Figure 3.3 One Mole of Several Substances
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Example 3.2
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3.2 The Mole

• Mole equal to Avogadros Number, equal to
  6.02x1023 particles of a substance

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Specialized units

• The correct name for a particle of a substance
  based on the type of matter
• Atom the representative particle for an element
• example Fe, S, etc.
• ion the representative particle for a charged
  particle
• example Na1, Cl-1, NH41,etc.
• Molecule the representative particle for a
  molecular compound (made up of non-metals)
• example CO2, CH4, etc.
• Formula unit the representative particle for an
  ionic compound (metal and non-metal or polyatomic
  ion)
• example KCl, MgSO4, etc.

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Molar mass

• Molar mass (MM) the mass of 1 mole of a
  substance equal to the atomic mass on the
  periodic table
• Round to the tenths place from the Periodic Table
  to simplify calculations

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Calculating molar mass

• 1. find the mass of the element on the periodic
  table
• 2. multiply by the number of atoms of that
  element in the formula (distribute parenthesis)
• 3. sum the relative masses of the individual
  elements

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Molar Masses of Some Substances
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Practice

• Example Determine the molar masses of the
  following substances.
• Aluminum CaSO4 (NH4)3P

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The Significance of the Mole

• In the laboratory, substances are weighed on
  balances, in units of grams
• The mole allows us to relate the number of grams
  of a substance to the number of atoms or
  molecules of a substance

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Mole-Gram Conversions

• Molar mass can be used like any other conversion
  factor
• Molar mass (g) 1 mole

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Example 3.3
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3.3 Mass Relations in Chemical Formulas

• Percent composition from formula
• percent composition - the number of grams of each
  element in 100 g of the compound
• Part x 100 composition
• whole

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2 types of Comp. problems

• Given data as masses of elements, etc.
• Use the masses of each element and the total mass
  of the compound

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2 types of Comp. problems

• 2. Given the chemical formula of the compound
• a. use the relative molar masses of each
  element and the molar mass of the compound

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Example 3.4
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Subscripts

1. Represent the atom ratio in a compound
2. Represent the mole ratio in a compound

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Diatomic Elements

• Elements that due to there chemical reactivity
  exist only as molecules of 2 atoms in nature.
• 7 diatomic elements
• Br I N Cl H O F
• H O N Cl Br I F

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Simplest Formula from Chemical Analysis
(Empirical Formula)

• Simplest formula (empirical formula) the
  simplest whole number ratio of the atoms in a
  compound

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Calculating the empirical formula

• 1. can be determined from masses of the
  individual elements or the composition of the
  elements in a compound
• 2. if s are given consider the sample to be of
  100 grams and so the s become the masses in
  grams
• 25.6 25.6 g
• 3. convert the mass of each element to moles
• 4. divide each number of moles by the smallest
  number of moles of all of the answers to 3
• 5. If the answers to 4 are whole numbers, these
  are the subscripts in the empirical formula.
• If any of the answers to 4 is not a
  whole number, convert all answers to a common
  fraction. Multiply each fraction by the
  denominator resulting in a whole number and these
  are the subscripts in the empirical formula.

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Common Fractions

• 0.25 0.33 0.50 0.66 0.75
• ¼ 1/3 ½ 2/3 ¾

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Example 3.5 Simplest Formula from Masses of
Elements
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Example 3.6 Simplest Formula from Mass Percents

• The compound that gives vinegar its sour taste is
  acetic acid, which contains
• the elements carbon, hydrogen, and oxygen. When
  5.00g of acetic acid is
• analyzed it is found to contain 2.00g of carbon,
  0.336g of hydrogen, and
• 2.66g of oxygen. What is the empirical formula
  of acetic acid?

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Molecular Formula

• Molecular formula the true ratio of the atoms
  in a compound as it exists naturally may be the
  same as the empirical formula will relate to the
  empirical formula in whole number ratios

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Molecular Formula

• Calculating the molecular formula
• 1. find the molar mass of the empirical formula
• 2. divide the molecular molar mass (given in the
  question) by the empirical molar mass
• MMM
• EMM
• 3. multiply each subscript in the empirical
  formula by the answer to 2, these are the
  subscripts for the molecular formula

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Example 3.7