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Chapter 6 Notes Packet
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6.1 Counting By Weighing
Counting by weighing is useful when measuring ___________ quantities of a
substance.
Examples:
a)
The average mass of one M & M is 4 grams. How much would 54
M & M’s weigh?
b)
How many M & M’s are in 88 grams of the candy?
6.2 Atomic Masses: Counting Atoms by Weighing
The atomic masses found on the periodic table are __________ atomic
masses for all of the ____________ of the elements. Elements masses are
measured in ________. These masses can be used to make conversions in
chemistry.
Example:
a)
Calculate the mass in amu of 25 lead atoms.
b)
Calculate the number of Helium atoms in a sample that has a mass
of 24 amu.
6.3 The Mole
In section 6.3 we used _________ for mass, but these are extremely
___________ units. In the lab a much larger unit, the _________, is the
convenient unit for ____________.
Assume we have a sample of carbon that has a mass of 12.01. What mass of
What mass of nitrogen has exactly the same number of atoms as the sample
of carbon?
If we weigh out samples of all the elements to have a mass equal to that
element’s average atomic mass in grams, all of the elements would contain
the same number of atoms. This number would be ___________________,
which is called a _______ in chemistry. This number is called __________
number for the scientist who discovered it.
This information has given us two conversion factors that we will use. The
first conversion factor is that the atomic mass in grams , which can be found
on the __________________________ is equal to 1 mole.
_______ g = 1 mol
The mass in grams changes depending on which atom you are measuring. For
example the conversion factor for aluminum would be ______ g = 1 mole.
The other conversion factor is Avogadro’s Number. This conversion never
changes. __________________ atoms = 1 mole.
Example 1: What is the mass in grams of 3.50 mol of the element copper
(Cu)?
Example 2: A chemist produced 11.9 g of aluminum. How many moles of
aluminum have been produced?
Example 3: How many moles of silver are in 3.01 x 1023 atoms?
1. What is the mass, in grams, of each of the following?
a. 1.0 mol Li
b. 2.0 mol Al
c. 6.50 mol Cu
d. 6.02 x 1023 atoms Ag (You will need to use both conversion factors!)
e. 2.57 x 1020 atoms S
2. How many moles of atoms are represented by each of the following?
a. 40.1 g Ca
b. 5.87 g Ni
c. 2.65 g Fe
d. 2.24 x 1025 atoms Zn
e. 50 atoms Ba
3. Determine the number of atoms contained in each of the following (Use
both conversion factors!)
a. 5.40 g B
b. 8.02 g S
c. 1.50 g K
6.4 Molar Mass
A compound is composed of more than one type of atom. The ________
mass is the mass in grams of one mole of a compound. To find the molar
mass of a compound we add up the atomic masses, in ___________, of all of
the atoms in the compound. For example, the compound CO2 would have a
molar mass of 44 g because carbon has a mass of 12 g and oxygen has a mass
of 16 g. We need to add the mass of two oxygen atoms, since there are two
oxygen atoms in the formula. Molar masses give us a conversion factor that
we can use for conversions between moles and grams. The number is
different depending on the formula of the compound.
_______ g = 1 mole
Calculate the molar mass of the following compounds
1. H2O
4. Mg(OH)2
2. NH3
5. Al2(SO4)3
3. C6H12O6
6. Ca(NO3)2
Calculations Using Molar Mass
Example 1: Calculate the mass of 1.48 mol of K2O.
Example 2: Calculate the number of moles in 2.4 g of H2O.
Example 3: Calculate the number of molecules in 3.2 g of CO2.
(Conversion Factor: 1 mole = 6.02 x 1023 molecules)
Perform the following calculations
1.
Find the weight of each of the following expressed in grams.
[a] 2.00 mole of carbon tetrachloride
[b] 0.25 mole of Ca(OH)2
[c] 4.1 mole of H2O
[d] 1.50 mole of HC2H3O2
[e] 5.3 x 1024 molecules of silver nitrate
[f] 2.2 x 1020 molecules of CO2
2. Give the correct number of moles of a substance.
[a] 25.0 grams of CaCO3
[b] 22.0 grams of carbon dioxide
[c] 120.3 grams of calcium nitrate
[d] 156 grams of Li2S
[e] 3.40 x 1023 molecules of H2O2
[f] 4.80 x 1023 molecules of methane, CH4
6.5 Percent Composition of Compounds
Percent composition looks at the compounds composition in terms of the
____________ of its elements. The equation for mass percent is. . .
Mass Percent of Element = mass of element in compound
Mass of compound
x
100
Example: Determine the mass percent of each element in H2O.
Determine the mass percent of each element in the following compounds
1. H2SO4
4. Ca(OH)2
2. Hydrochloric Acid
5. Sodium Oxide
3. MgCl2
6. Nitrogen Trioxide
6.6 Formulas of Compounds
Empirical Formula –
Molecular Formula –
6.7 Calculation of Empirical Formulas
To calculate the empirical formula of a compound, we first determine the relative
masses of the various ________________ that are present.
One way to find the empirical formula is to measure the masses of elements that
react to form the compound. For example, suppose we have a 3.450 g sample of
nitrogen that reacts with 1.970 g of oxygen. Determine the empirical formula for
this compound.
Sometimes the relative numbers of moles you get when you calculate an
_____________ formula will turn out to be nonintegers. When this happens, you
must convert to the appropriate whole numbers.
Problems
1.
When a 2.000 g sample of iron metal is heated in air, it reacts with oxygen
to achieve a final mass of 2.573 g. Determine the empirical formula for this
compound.
2.
A 4.550 g sample of cobalt reacts with 5.475 g of chlorine to form a binary
compound. Determine the empirical formula for this compound.
3.
Phosphoric acid is found in some soft drinks. A sample of phosphoric acid
contains 0.3086 g of hydrogen, 3.161 g of phosphorus, and 6.531 g of oxygen.
Determine the empirical formula for phosphoric acid.
4.
The simplest amino acid, glycine, has the following mass percents: 32.00%
carbon, 6.714 % Hydrogen, 42.63% Oxygen, an 18.66% Nitrogen. Determine the
empirical formula of glycine.
5.
A compound has been analyzed and found to have the following mass percent
composition: 66.75 % copper, 10.84% phosphorus, and 22.41% oxygen. Determine
the empirical formula for this compound.
6.8 Calculations of Molecular Formulas
If we know the composition of a compound in terms of the masses (or mass
percentages) of the elements present, we can calculate the empirical formula but
not the _______________ formula. To obtain the molecular formula we must
know the __________ ___________.
Example: A compound containing carbon, hydrogen, and oxygen is found to be
40.00% carbon, 6.700% hydrogen, and 53.3% oxygen. The molar mass is 120 g/mol.
Determine the molecular formula for this compound.
Problems
1.
A compound with the empirical formula CH2O was found to have a molar mass
between 89 and 91 g/mol. What is the molecular formula of the compound.
2.
A compound with the empirical formula CH2 was found to have a molar mass
of approximately 84 g. What is the molecular formula of the compound?
3. A compound consists of 65.45% C, 5.492% H, and 29.06% O on a mass basis and
has a molar mass of approximately 110 g/mol. Determine the molecular formula of
the compound.