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Chapter 7
7.1
The Mole
• The SI base unit used to measure the amount of a substance whose number
of particles is the same as the number of atoms of carbon in exactly 12
grams of carbon-12.
• The mole is a counting unit used to count out a given number of particles.
• Avogadro’s number- 6.022 x 1023, the number of atoms or molecules in
1.00 mol.
o 6.022 x 1023 particles = 1 mole
• Molar Mass – The mass in grams of one mole of a substance.
Converting between Moles, Particles and Molar Mass
Moles to number of particles:
How many ions are in 0.187mol of Na+ ions?
0.187 mol Na+ ions x (6.022 x 1023 ions/ 1 mole) = 1.123 x 1023 ions.
Number of Particles to moles:
3.01 x 1023 molecules SO2 x (1 mole / 6.022 x 1023 molecules) = 0.50 mol
Particles to mass
2.11 x 1024 atoms Cu x (1mol/6.022 x1023 atoms) x (63.55 grams/1 mole) =
2.23 x 102g
Mass to Particles
237 g Cu x (1 mol/63.55 g) x (6.022 x 1023/1 mol) = 2.24 x 1024 atoms
7.2
Average Atomic Masses
Most elements are mixtures of isotopes. Average atomic masses represent
weighted averages of the atomic masses of all naturally occurring isotopes of an
element.
Calculating an average atomic mass
Isotope
Percentage
Decimal Fraction Contribution
Copper-63 69.17%
0.6917
62.94 amu x 0.6917
Copper-65 30.83%
0.3083
64.93 amu x 0.3083
The average atomic mass is the sum of the contributions.
(62.94amu x 0.6917) + (64.93amu x 0.3083) = 63.55amu
Chemical Formulas and Moles
• Chemical formulas give a ratio of elemental components.
• Ionic formulas show the simplest ratio of cations and anions.
• Covalent formulas (including polyatomic ions) show both elements and the
number of each element.
Formulas can be used to calculate Molar Masses
• From formulas we can tell what elements (or ions) are present and in what
quantities.
• Molar masses of individual elements (found on the periodic table) are
summed to determine molar masses of molecular compounds.
Calculating the molar mass of a compound
• 1 mol CaCl2 = 1mol Ca + 2 mol Cl
• 1mol Ca = 1 x 40.078g/mol = 40.078g
• 2mol Cl = 2 x 35.4527g/mol = 70.9054g
• 1mol CaCl2 = 40.078g + 70.9054g = 110.983g
7.3
Percentage composition – The percentage by mass of each element in a
compound.
• Percent composition can verify a substance’s identity.
• Percent Composition can be calculated from any chemical formula.
o Find the total mass of each elemental component.
o Divide the element’s total mass by the compound's molar mass.
o Multiply by 100.
Example: CO2
Molar mass = 44.01 g/mol
Carbon
1 x 12.01 = 12.01 x 100 = 27.3%
44.01
Oxygen
2 x 16.00 = 32.00 x 100 = 72.7%
44.01
Total = 100.0%
Empirical Formula - These chemical formulas show the composition of a
compound in its simplest ratio.
Formaldehyde
= CH2O
Acetic Acid
= C 2 H4 O2
Glucose
= C6H12O6
Each has the same empirical formula = CH2O
Elemental analysis can provide this ratio. How? It’s all about the number of moles!
Example:
Expirimental analysis yields the following…
Pb - 74.51%
Cl - 24.49%
1. Convert to grams assuming 100g sample
Pb = 74.51% x 100g = 74.51g, Cl = 24.49% x 100g = 24.49g
2. Convert to moles
Pb 74.51g x (1mol/207.2) = 0.36 mol
Cl 24.49g x (1mol/35.453) = 0.69 mol
3. Convert to whole numbers by dividing each subscript by the smallest subscript.
0.36/0.36 = 1,
0.69/0.36= 1.9 (round up)
Emperical Formula = PbCl2
Molecular Formula
While formulas of ionic compounds already show the simplest whole # ratios.
Molecular formulas are multiples of empirical formulas.
• Experimental analysis can provide Molar Mass.
• By dividing the experimental molar mass by the empirical molar mass you
can determine the multiple for the molecular formula.
Example:
Empirical Formula CH2O
Empirical Molar Mass 30.03g/mole
Experimental Molar Mass 180.18g/mole
30.03g/180.18g = 6
Therefore Molecular Formula C6H12O6