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The MoleCh 11
Measuring matter11.1
11.1 Vocabulary• Review
o molecule: two or more atoms that covalently bond together to
form a unit
• Newmole
Avogadro’s number
• Main Idea - Chemists use the mole to count atoms,
molecules, ions, and formula units.
• NOTE – YOU WILL NEED A SCIENTIFIC
CALCULATOR FOR THIS CHAPTER!
How do we measure items?
You can measure mass,
or volume,
or you can count pieces.
We measure mass in grams.
We measure volume in liters.
We count pieces in MOLES.
What is the mole?
We’re not talking about this
kind of mole!
Counting Particles
• Chemists need a convenient method for accurately
counting the number of atoms, molecules, or
formula units of a substance.
• The mole is the SI base unit used to measure the
amount of a substance.
Moles (is abbreviated: mol)
It is an amount, defined as the number of carbon atoms in exactly 12 grams of carbon-12.
1 mole = 6.022 x 1023 of the representative particles.
Treat it like a very large dozen!
6.022 x 1023 is called Avogadro’s number.
Similar Words for an amount
Pair: 1 pair of shoelaces = 2 shoelaces
Dozen: 1 dozen oranges = 12 oranges
Case: 1 case of Dr. Pepper
= 24 cans Dr. Pepper
Gross: 1 gross of pencils = 144 pencils
Ream: 1 ream of paper = 500 sheets of paper
What are Representative Particles?
The smallest pieces of a substance:
1) For a molecular compound: it is
the molecule.
2) For an ionic compound: it is the
formula unit (made of ions).
3) For an element: it is the atom.
• Remember the 7 diatomic
elements? (made of molecules)
Types of questions• How many oxygen atoms in the
following?
CaCO3
Al2(SO4)3
• How many ions in the following?
CaCl2
NaOH
Al2(SO4)3
3 atoms of oxygen
12 (3 x 4) atoms of oxygen
3 total ions (1 Ca2+ ion and 2 Cl1- ions)
2 total ions (1 Na1+ ion and 1 OH1- ion)
5 total ions (2 Al3+ + 3 SO42- ions)
Converting Between Moles and Particles
• Conversion factors must be used.
• Moles to particles
Example:
Number of molecules in 3.50 mol of sucrose
Converting Between Moles and Particles (cont.)
• Particles to moleso Use the inverse of Avogadro’s number as the conversion factor.
Practice problems (round to 3 sig. figs.)
How many molecules of CO2 are in 4.56 moles of CO2?
How many moles of water is 5.87 x 1022
molecules?
How many atoms of carbon are in 1.23 moles of C6H12O6?
How many moles is 7.78 x 1024 formula units of MgCl2?
2.75 x 1024 molecules
0.0975 mol (or 9.75 x 10-2)
4.44 x 1024 atoms C
12.9 moles
10.1 CheckWhat does the mole measure?
A. mass of a substance
B. amount of a substance
C. volume of a gas
D. density of a gas
10.1 CheckWhat is the conversion factor for determining the
number of moles of a substance from a known
number of particles?
A.
B.
C. 1 particle 6.02 1023
D. 1 mol 6.02 1023 particles
Mass and the mole11.2
11.2 Vocabulary• Review
o conversion factor: a ratio of equivalent values used to express the same
quantity in different units
• Newo molar mass
• Main Idea - A mole always contains the same
number of particles; however, moles of different
substances have different masses.
The Mass of a Mole• 1 mol of copper and 1 mol of carbon have different
masses.
• One copper atom has a different mass than 1
carbon atom.
Measuring Moles Remember relative atomic mass?
- The amu was one twelfth the mass of
a carbon-12 atom.
Since the mole is the number of atoms
in 12 grams of carbon-12,
the decimal number on the periodic table
is also the mass of 1 mole of those atoms in
grams.
The Mass of a Mole (cont.)• Molar mass (MM) is the mass in grams of one mole of
any pure substance.
• Also called Formula Weight (FW)
Molar Mass
Equals the mass of 1 mole of an element in
grams (from periodic table)
12.011 grams of C has the same number of
pieces as
• 1.008 grams of H
• 55.85 grams of iron.
We can write this as: 12.011 g C = 1 mole C
We can count things by weighing them.
Using Molar Mass• Moles to mass
Using Molar Mass (cont.)• Convert mass to moles with the inverse molar mass
conversion factor.
Examples How much would 2.34 moles of carbon
weigh?
How many moles of magnesium is 24.31 g of Mg?
How many atoms of lithium is 1.00 g of Li?
How much would 3.45 x 1022 atoms of U weigh?
28.1 grams C
1 mol Mg
8.72 x 1022 atoms Li
13.6 grams U
Using Molar Mass (cont.)• This figure shows the steps to complete conversions
between mass and atoms.
11.2 CheckThe mass in grams of 1 mol of any pure substance is:
A. molar mass
B. Avogadro’s number
C. atomic mass
D. 1 g/mol
11.2 CheckMolar mass is used to convert what?
A. mass to moles
B. moles to mass
C. atomic weight
D. particles
Moles of compounds11.3
Vocabulary• Review
o representative particle: an atom,
molecule, formula unit, or ion
• Main Idea -The molar mass of a compound can be
calculated from its chemical formula and can be
used to convert from mass to moles of that
compound.
Chemical Formulas and the Mole
• Chemical formulas indicate the numbers and types
of atoms contained in one unit of the compound.
• One mole of CCl2F2 contains one mole of C atoms,
two moles of Cl atoms, and two moles of F atoms.
The Molar Mass of Compounds
• The molar mass of a compound equals the molar
mass of each element, multiplied by the moles of
that element in the chemical formula, added
together.
• The molar mass of a compound demonstrates the
law of conservation of mass.
The Molar Mass of Compounds in 1 mole of H2O molecules there are two
moles of H atoms and 1 mole of O atoms (think of a compound as a molar ratio)
To find the mass of one mole of a compound
odetermine the number of moles of the elements present
oMultiply the number times their mass(from the periodic table)
oadd them up for the total mass
Calculating Molar Mass
Calculate the molar mass of
magnesium carbonate, MgCO3.
24.3050 g + 12.0107 g + 3 x (15.9994 g)
= 84.3139
Thus, 84.33139 grams is the formula
mass for MgCO3.
Examples Calculate the molar mass of the
following and tell what type it is:
Na2S
N2O4
C
Ca(NO3)2
C6H12O6
(NH4)3PO4
= 78.045 g/mol
= 92.011 g/mol
= 12.011 g/mol
= 164.088 g/mol
= 180.156 g/mol
= 149.087 g/mol
Moles to Mass Conversion for Compounds
• For elements, the conversion factor is the molar
mass of the elements.
• The procedure is the same for compounds, except
that you must first calculate the molar mass of the
compound.
Mass to Moles Conversion for Compounds
• The conversion factor is the inverse of the molar
mass of the compound.
37
For exampleHow many moles is 5.69 g of NaOH?
38
For example
How many moles is 5.69 g of NaOH?
5 69. g
39
For example
How many moles is 5.69 g of NaOH?
5 69. g mole
g
We need to change 5.69 grams NaOH to moles
40
For example
How many moles is 5.69 g of NaOH?
5 69. g mole
g
We need to change 5.69 grams NaOH to moles
1mole Na = 22.990 g
1 mol O = 15.999 g
1 mole of H = 1.008 g
41
For example
How many moles is 5.69 g of NaOH?
5 69. g mole
g
We need to change 5.69 grams NaOH to moles
1mole Na = 22.990 g
1 mol O = 15.999 g
1 mole of H = 1.008 g
1 mole NaOH = 39.997 g
42
For example
1 mole NaOH = 39.997 g NaOH
= 0.1422606696 mol NaOH
= 0.142 mol NaOH 3 Sig Figs
Mass to Particles Conversion for Compounds
• Convert mass to moles of compound with the
inverse of molar mass.
• Convert moles to particles with Avogadro’s number.
• This figure summarizes the conversions between
mass, moles, and particles.
11.3 Check
How many moles of OH— ions are in 2.50 moles of
Ca(OH)2?
A.2.00
B. 2.50
C.4.00
D.5.00
11.3 Check
How many particles of Mg are in 10 moles of MgBr2?
A.6.02 1023
B. 6.02 1024
C.1.20 1024
D.1.20 1025
Empirical and Molecular formulas
11.4
Vocabulary
• Review
o percent by mass: the ratio of the mass of each element to
the total mass of the compound expressed as a percent
• New
percent composition
empirical formula
molecular formula
• Main Idea - A molecular formula of a compound is
a whole-number multiple of its empirical formula.
Percent Composition
• The percent by mass of any element in a
compound can be found by dividing the mass of
the element by the mass of the compound and
multiplying by 100.
Percent Composition
• The percent by mass of each element in a
compound is the percent composition of a
compound.
• Percent composition of a compound can also be
determined from its chemical formula.
50
Calculating Percent Composition
of a Compound
Like all percent problems:
part
whole
1) Find the mass of each of the
components (the elements),
2) Next, divide by the total mass of
the compound; then x 100%
x 100 % = percent
51
Example
Calculate the percent composition of a
compound that is made of 29.0 grams of
Ag with 4.30 grams of S
(Assume you have one mol of substance)
29.0 g Ag
33.3 g totalX 100 = 87.1 % Ag
4.30 g S
33.3 g totalX 100 = 12.9 % S
Total = 100 %
52
Examples
Calculate the percent
composition of C2H4?
How about Aluminum
carbonate?
85.7% C, 14.3 % H
23.1% Al, 15.4% C, and 61.5 % O
Empirical Formula• The empirical formula for a compound is the smallest
whole-number mole ratio of the elements.
• You can calculate the empirical formula from percent by mass by assuming you have 100.00 g of the compound. Then, convert the mass of each element to moles.
o SO3 has a percent composition of 40.05% S and 59.95% O.
o To find mole ratio assume 100 g of Sulfur Trioxide; this means 40.05 g S and 59.95 g O. Multiply each by molar mass to determine moles.
o 1.249 mol S and 3.747 mol O. Turn this into a whole number ratio by dividing by the smallest. S becomes 1 and O becomes 3
o [If you still don’t get whole number, multiply by smallest number that will produce a whole number. If you had 1.5 mol C, 3 mol H, 1 mol O, multiply all by 2 to get 3 mol C, 6 mol H, 3 mol O = C3H6O2
• The empirical formula may or may not be the same as the molecular formula.
Molecular formula of hydrogen peroxide = H2O2
Empirical formula of hydrogen peroxide = HO
Molecular Formula• The molecular formula specifies the actual number
of atoms of each element in one molecule or
formula unit of the substance.
• Molecular formula is always a whole-number
multiple of the empirical formula.
• To distinguish the molecular formula from empirical
formula, chemists must experimentally determine
the molar mass of the compound. Then, this
relationship is used: o Molecular formula = (empirical formula)n
• n is the factor by which molecular formula is obtained.
o Unknown substance X has formula weight of 26.04 g/mol and its empirical
formula is CH. Determine the molecular formula.
o (26.04g/mol)/ (13.02 g/mol) = 2
o ANSWER – C2H2
Empirical vs. Molecular FormulaWhat is the empirical
formula for the
compound C6H12O6?
A.CHO
B. C2H3O2
C.CH2O
D.CH3O
Which is the empirical
formula for hydrogen
peroxide?
A.H2O2
B. H2O
C.HO
D.none of the above
Formulas of hydrates11.5
Vocabulary• Review
crystal lattice: a three-dimensional geometric
arrangement of particles
• New
hydrate
Main Idea -Hydrates are solid ionic compounds in
which water molecules are trapped.
Naming Hydrates• A hydrate is a compound that has a specific
number of water molecules bound to its atoms.
• The number of water molecules associated with
each formula unit of the compound is written
following a dot.
• Sodium carbonate decahydrate =
Na2CO3 • 10H2O
Naming Hydrates
Analyzing Hydrates• When heated, water molecules are released from a
hydrate leaving an anhydrous compound.
• To determine the formula of a hydrate, find the
number of moles of water associated with 1 mole of
hydrate.
Analyzing Hydrates (cont.)
1. Weigh hydrate.
2. Heat to drive off the water.
3. Weigh the anhydrous compound.
4. Subtract and convert the difference to moles.
5. The ratio of moles of water to moles of anhydrous
compound is the coefficient for water in the
hydrate.
Use of Hydrates• Anhydrous forms of hydrates are often used to
absorb water, particularly during shipment of
electronic and optical equipment.
• In chemistry labs, anhydrous forms of hydrates are
used to remove moisture from the air and keep
other substances dry.
11.5 CheckHeating a hydrate causes what to happen?
A.Water is driven from the hydrate.
B. The hydrate melts.
C.The hydrate conducts
electricity.
D.There is no change in the
hydrate.
11.5 CheckA hydrate that has been heated and the water driven
off is called:
A.dehydrated compound
B. antihydrated compound
C.anhydrous compound
D.hydrous compound
11.5 CheckTwo substances have the same percent by mass
composition, but very different properties. They must
have the same ____.
A.density
B. empirical formula
C.molecular formula
D.molar mass