Ch 11: The Mole
Section 11.1
• Dozen=
• 12
• Ream=
• 500
• Pair=
•2
• Gross=
• 144
Measuring Matter
• What is a mole?
• It is the SI base unit used to measure the
amount of a substance
• Abbreviated mol
• Represents particles
• Is also called Avogadro’s number
6.02 x 1023 (3 significant figures)
• Remember conversion factors?
• 12 roses = 1 dozen
•
So, 3.5 dozen= ? Roses
= 42
Well, 6.02 x 1023 atoms (or any
representative paticles) = 1 mol
• Review Scientific Notation
• Review Rounding
• Example:
• 3.50 mol sucrose has how many molecules?
• Work practice problems 1-3
• Example:
• How many moles are in 3.58 X 1020 atoms of Ca?
• Work practice problems 4-7
closure
• How is a mole similar to a dozen?
• What is the relationship between
avagadro’s number & one mole?
• Why do chemists use moles?
• **worksheet: The mole & Avogadro’s number
• STOP
Section 11.2 Mass & the Mole
• Different particles (atoms) have different
masses.
• Remember atomic mass.
– Each element has its own specific mass.
– Therefore each compound has its own
specific mass.
• Molar mass (g/mol)-mass in grams of any
pure substance
• The molar mass of any element is
numerically equal to its atomic mass.
• Thus…1 mol Mn = 54.94 g/mol Mn = 6.02
x 1023 atoms Mn
• Example:
• 1 mol Zn =
65.4 g/mol
• 1 mol O2 =
32.0 g/mol
•
•
•
•
Practice
3 mol Zinc =
196 g/mol Zn
1 mol H2O =
18.0g/mol H2O
1 mol sulfuric acid =
98.1 g/mol H2SO4
Mass  Mole conversions
• Example:
• Calculate the mass in grams of 0.0450 mol
Cr.
• Work Practice problems 1-4
• Hydrate: CuSO4·5H2O
4) Calculate the mass in grams of 2.45 mol
of CaCl2·2H2O
Mole Mass conversions
• Example:
• Determine the number of moles for 25.5 g
Ag. (mass mol)
• Work practice problems 5-7.
**worksheet: moles & mass
Mass  Atom Conversion
• Example:
– How many atoms of gold (Au) are in a pure
nugget having the mass of 25.0g?
– Practice problems 1-5
Atom  Mass Conversion
• Example
• A party balloon contains 5.50 x 1022 atoms
of helium (He) gas. What is the mass in
grams of the helium?
• Practice problems 6-10.
Mole  Mole
Example
• According to the following balanced
equation, how many moles of O2 is
produced from 3.00 moles of CuO?
2CuO  2Cu + O2
Practice problems 1 & 2.
Section 11.4
Empirical & Molecular Formula
• Percent Composition is the percent by
mass of each element in a compound.
Example:
• If we had 100 g of a sample of some new
compound contains 55g of element X & 45
g of element Y. What is the % of element
X & Y?
• If we already know the _chemical formula
for a compound, you can calculate its
percent composition.
• % by mass=
•
Mass of element in 1 mol of compound
X 100
Molar mass of compound
• Ex. Determine the percent composition of
H2O.
• (If you had 350. g of water, then how much
is oxygen?)
• Practice problems 1-3.
• Empirical Formula is the smallest whole
number ratio of the elements.
• Calculating Empirical formula from percent
composition:
• Directions:
• The percent should be assumed to be 100
g & converted to moles. Then we figure
out the mole ratio by dividing each by the
smallest mole.
POEM:
1.
2.
3.
4.
Percent to Mass
Mass to Mole
Divide by small
Multiply til whole.
• Example 1
• The percent composition of an oxide of
sulfur is 40.05% S & 59.95% O.
• Example 2
• Determine the empirical formula for methyl
acetate which has the following chemical
analysis:48.64% C, 8.16% H, & 43.20% O.
• Practice Problems
• **worksheet: determining empirical
formulas