Name: __________________________
Period: ___________
Unit 5 Packet – Counting Particles
Packet Contents Sheet (With Objectives on Back)
Relative Mass Lab
Atomic Theory and Structure/The Mole: Its History
and Use
Cornell Notes Page/The Mole and Avogadro’s
Number
Worksheet 1- Molar Masses of Elements
Gram Formula Mass/ Moles and Mass
Worksheet 2
Worksheet 2.5- Empirical Formula Practice
Worksheet 3- Empirical and Molecular Formulas
Empirical Formula/Molecular Formula Worksheet
Unit 5 Review (2 pages)
DO NOT, under any circumstances, throw this away!
This packet MUST be saved for the final exam.
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Unit 5 – Counting Particles - Learning Goal:
Students to be able to use Avagadro’s hypothesis to solve quantitative problems
relating to number of particles/moles and mass. Students to be able to derive
formulae and % composition of compounds.
Scale
Score
Score 4
Comment
Students show mastery of score 3 without any errors plus:
 Create models to explain how to use Avagadro’s
hypothesis to solve quantitative problems relating to
number of particles/moles and mass. Students to be
able to derive formulas and % composition of
compounds.
Score 3
Without any major errors, students can independently:
 Understand and use Avagadro’s hypothesis to solve
quantitative problems relating to number of
particles/moles and mass. Students to be able to
derive formulas and % composition of compounds.
Score 2
With one or two major errors, students can independently:
 Understand how Avagadro’s hypothesis can be used to
solve quantitative problems relating to number of
particles/moles and mass. Students to be able to
derive formulas and % composition of compounds.
Score 1
With help from the teacher, students can:
 Understand and use Avagadro’s hypothesis to solve
quantitative problems relating to number of
particles/moles and mass. Students to be able to
derive formulas and % composition of compounds.
Score 0
Even with the teachers help students show no
understanding or ability to:
 Use Avagadro’s hypothesis to solve quantitative
problems relating to number of particles/moles and
mass. Students to be able to derive formulae and %
composition of compounds.
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Name
Date
Pd
Activity: Relative Mass
Purpose
The purpose is to determine the relative mass of different kinds of hardware and to learn
to count by massing.
Data
Hardware
Mass (g) Total
Mass w/o Vial
Relative Mass
(imu- item
mass units)
Empty vial
Vial + Washers
Vial + Hex Nuts
Vial + Bolts
Calculations
Show all calculations. Be sure to label your quantities!
1. Each vial contains the same number of pieces. From each mass, subtract the mass of the
empty vial to determine the adjusted mass. (Do you think the empty vial will have an
adjusted mass?) To find the relative mass, divide the mass of each item by the mass of the
smallest item. (What will the relative mass of the smallest item be?) Don’t forget units and
significant figures. Show all work for each item below.
W:
HN:
B:
2. What does relative mean in the phrase relative mass?
3. When calculating relative mass, why is it important to be sure the same number of items are
in each vial?
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4. Which item is used to determine relative masses of the other items? Why?
5. Why do you think the units were changed from grams to imu’s in the table above?
6. Which element is used to determine relative masses of the other elements? Explain the
connection between this activity and the work of Gay-Lussac and Avogadro (Avogadro’s
Hypothesis).
Conclusion
Do you agree or disagree with the following statement? Support your answer.
“You can count by weighing.”
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Atomic Theory and Structure, The Mole: Its History and Use
by Anthony Carpi, Ph.D.
Simply put, the mole represents a number. Just as the term dozen refers to the number twelve, the mole represents the number 6.02 x 10 23. (If you're
confused by the form of this number refer to our The Metric System module).
Now that's a big number! While a dozen eggs will make a nice omelet, a mole of eggs will fill all of the oceans on earth more than 30 million times
over. Think about it: It would take 10 billion chickens laying 10 eggs per day more than 10 billion years to lay a mole of eggs. So why would we ever
use such a big number? Certainly the local donut store is not going to "supersize" your dozen by giving you a mole of jelly-filled treats.
The mole is used when we're talking about numbers of atoms and molecules. Atoms and molecules are very tiny things. A drop of water the size of
the period at the end of this sentence would contain 10 trillion water molecules. Instead of talking about trillions and quadrillions of molecules (and
more), it's much simpler to use the mole.
History of the Mole
The number of objects in one mole, that is, 6.02 x 1023, is commonly referred to as Avogadro's number. Amadeo Avogadro was an Italian physics
professor who proposed in 1811 that equal volumes of different gases at the same temperature contain equal numbers of molecules. About fifty years
later, an Italian scientist named Stanislao Cannizzaro used Avogadro's hypothesis to develop a set of atomic weights for the known elements by
comparing the masses of equal volumes of gas. Building on this work, an Austrian high school teacher named Johann Josef Loschmidt calculated the
size of a molecule of air in 1865, and thus developed an estimate for the number of molecules in a given volume of air. While these early estimates
have since been refined, they led to the concept of the mole - that is, the theory that in a defined mass of an element (its atomic weight) there is a
precise number of atoms: Avogadro's number.
Molar Mass
A sample of any element with a mass equal to that element's atomic weight (in grams) will contain precisely one mole of atoms (6.02 x 1023 atoms).
For example, helium has an atomic weight of 4.00. Therefore, 4.00 grams of helium will contain one mole of helium atoms. You can also work with
fractions (or multiples) of moles:
Mole/Weight Relationship Examples Using Helium
Moles Helium
# Helium Atoms Grams Helium
1/4
1.505 x 1023
1g
1/2
3.01 x 1023
2g
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Mole/Weight Relationship Examples Using Helium
1
6.02 x 1023
4g
2
1.204 x 1024
8g
10
6.02 x 1024
40 g
Other atomic weights are listed on the periodic table (see our Periodic Table page). For each element listed, measuring out a quantity of the element
equal to its atomic weight in grams will yield 6.02 x 1023 atoms of that element.
The atomic weight of an element identifies both the mass of one mole of that element and the total number of protons and neutrons in an atom of that
element. How can that be? Let's look at hydrogen. One mole of hydrogen atoms will weigh 1.01 grams.
A Hydrogen Atom
Each hydrogen atom consists of one proton surrounded by one electron. But remember, the electron weighs so little that it does not contribute much
to an atom's weight. Ignoring the weight of hydrogen's electrons, we can say that one mole of protons (H nuclei) weighs approximately one gram.
Since protons and neutrons have about the same mass, a mole of either of these particles will weigh about one gram. For example, in one mole of
helium, there are two moles of protons and two moles of neutrons - four grams of particles.
Molecular Weight
If you stand on a scale with a friend, the scale will register the combined weight of both you and your friend. When atoms form molecules, the atoms
bond together, and the molecule's weight is the combined weight of all of its parts.
For example, every water molecule (H2O) has two atoms of hydrogen and one atom of oxygen. One mole of water molecules will contain two moles
of hydrogen and one mole of oxygen.
Mole and Weight Relationships of Water and Its Parts
2 moles H
+ 1 mole O
= 1 mole water
+
=
A bottle filled with exactly 18.02 g water will contain 6.02 x 1023 water molecules. The concept of fractions and multiples described above also
applies to molecules: 9.01 g of water would contain 1/2 mole, or 3.01 x 1023 molecules. You can calculate the molecular weight of any compound
simply by summing the weights of atoms that make up that compound.
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Cornell Notes Page
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Name
Date
Pd
Chemistry – Unit 5 Worksheet 2
1. An old (pre-1987) penny is nearly pure copper. If such a penny has a mass of 3.3 g, how many
moles of copper atoms would be in one penny?
2. Four nails have a total mass of 4.42 grams. How many moles of iron atoms do they contain?
3. A raindrop has a mass of 0.050 g. How many moles of water does a raindrop contain?
4. What mass of water would you need to have 15.0 moles of H2O?
5. One box of Morton’s Salt contains 737 grams. How many moles of sodium chloride is this?
6. A chocolate chip cookie recipe calls for 0.050 moles of baking soda (sodium bicarbonate,
NaHCO3). How many grams should the chef mass out?
7. Rust is iron(III) oxide (Fe2O3). The owner of a l959 Cadillac convertible wants to restore it
by removing the rust with oxalic acid, but he needs to know how many moles of rust will be
involved in the reaction. How many moles of iron(III) oxide are contained in 2.50 kg of rust?
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8. First-century Roman doctors believed that urine whitened teeth and
also kept them firmly in place. As gross as that sounds, it must have
worked because it was used as an active ingredient in toothpaste and
mouthwash well into the 18th century. Would you believe it’s still
used today? Thankfully, not in its original form! Modern dentists
recognized that it was the ammonia that cleaned the teeth, and they
still use that. The formula for ammonia is NH3. How many moles are
in 0.75 g of ammonia? How many molecules?
9. Lead (II) chromate, PbCrO4, was used as a pigment in paints. How many moles of lead
chromate are in 75.0 g of lead (II) chromate? How many atoms of oxygen are present?
10. The diameter of the tungsten wire in a light bulb filament is very small, less than two
thousandths of an inch, or about 1/20 mm. The mass of the filament is so very small –
0.0176 grams – that it would take 1,600 filaments to weigh an ounce! How many
tungsten atoms are in a typical light bulb filament?
11. Two popular antacids tablets are Tums and Maalox. The active ingredient in both of these
antacids is calcium carbonate, CaCO3. Tums Regular Strength tablets contain 0.747 g and
Maalox tablets contain 0.600 g of calcium carbonate. Compare the number of formula units
of calcium carbonate in both Tums and Maalox.
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Name
Date
Pd
Chemistry – Unit 5 Worksheet 2.5
Empirical Formula Practice
The empirical formula of a compound expresses the simplest whole number ratio of elements
in that compound. The empirical formula can be calculated from the percentage by mass for
each element in the compound.
Sample Problem: The percentage composition by mass of a compound is 56.6% potassium,
8.7% carbon, and 34.7% oxygen. Find its empirical formula.
Solution: Assume that there are 100 grams of the compound. Then calculate the number of
moles of atoms of each element in the sample.
Practice Problems:
From the percent composition information in each of the following, calculate empirical
formulas. Show a labeled setup below the problem, and write your answers in the space
provided.
1. 69.6% barium, 6.1% carbon, 24.3% oxygen
1. _________________
2. 40.5% zinc, 19.9% sulfur, 39.6% oxygen
2. _________________
3. 88.8% copper, 11.2% oxygen
3. _________________
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4. 79.9% copper, 20.1% oxygen
4. _________________
5. 36.7% potassium, 33.3% chlorine, 30.0% oxygen
5. _________________
6. 28.2% potassium, 25.6% chlorine, 46.2% oxygen
6. _________________
7. 40.2% potassium, 26.8% chromium, 33.0% oxygen
7. _________________
8. 26.6% potassium, 35.3% chromium, 38.1% oxygen
8. _________________
9. 56.3% oxygen, 43.7% phosphorus
9. _________________
10. 90.7% lead, 9.33% oxygen
10. _________________
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Name
Date
Pd
Chemistry – Unit 5 Worksheet 3
Empirical and Molecular Formulas
Show all your work when solving the following problems. Be sure to include units and label
your answer.
1. Find the empirical formula of a compound containing 32.0 g of bromine and 4.9 g of
magnesium.
2. What is the empirical formula of a carbon-oxygen compound, given that a 95.2 g sample of the
compound contains 40.8 g of carbon and the rest oxygen?
3. A compound was analyzed and found to contain 9.8 g of nitrogen, 0.70 g of hydrogen, and 33.6
g of oxygen. What is the empirical formula of the compound?
4. A compound composed of hydrogen and oxygen is found to contain 0.59 g of hydrogen and 9.40
g of oxygen. The molar mass of this compound is 34.0 g/mol. Find the empirical and
molecular formulas.
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5. A sample of iron oxide was found to contain 1.116 g of iron and 0.480 g of oxygen. Its molar
mass is roughly 5 x as great as that of oxygen gas. Find the empirical formula and the
molecular formula of this compound.
6. Find the percentage composition of a compound that contains 17.6 g of iron and
10.3 g of sulfur. The total mass of the compound is 27.9 g.
7. Find the percentage composition of a compound that contains 1.94 g of carbon, 0.48 g of
hydrogen, and 2.58 g of sulfur in a 5.00 g sample of the compound.
8. What is the % by mass of oxygen in Mg(NO3)2 ?
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Chemistry – Unit 5 Review
1. Definitions
a. mole
b. molar mass
c. Avogadro’s number
d. empirical formula
e. molecular formula
2. Find the molar mass of the following:
a. KNO3
d. oxygen gas
b. (NH4)2CO3
e. Ca(NO3)2
c. Ag2CrO4
f. PbSO4
3. Consider the masses of various hardware below.
Type
Mass (g)
Relative mass
Washer
1.74
Hex nut
3.16
Anchor
3.00
Bolt
7.64
a. Do the calculations necessary to complete the table.
b. Explain the connection between these calculations and the atomic masses in the Periodic
Table.
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4. Convert from g  moles or from moles  g. Show units.
a. 12.0 g Fe x
b. 25.0 g of Cl2 gas x
=
moles
=
c. 0.476 g of (NH4)2SO4 x
moles
=
moles
d. 0.15 moles NaNO3 x
=
g
e. 0.0280 moles NO2 x
=
g
f. 0.64 moles AlCl3 x
=
g
5. Use Avogadro’s number to do the following. Show work, use labels.
a. How many atoms are there in 0.00150 moles Zn?
b. If you had 2.50 moles of oxygen gas, what mass of the gas would be in the sample?
c. A 4.07 g sample of NaI contains how many atoms of Na?
d. How many atoms of chlorine are there in 16.5 g of iron (III) chloride, FeCl3?
e. What is the mass of 100 million atoms of gold? Could you mass this on a balance?
6. Calculate the empirical formula of a compound that contains 4.20 g of nitrogen and 12.0 g of
oxygen.
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7. When 20.16 g of magnesium oxide reacts with carbon, carbon monoxide forms and 12.16 g of
Mg metal remains. What is the empirical formula of magnesium oxide?
8. What is the molecular formula of each compound?
Empirical Formula
Actual Molar Mass of
Compound
CH
78 g/mole
NO2
92 g/mole
Molecular Formula
9. A compound is composed of 7.20 g of carbon, 1.20 g of hydrogen and 9.60 g of oxygen. The
molar mass of the compound is 180 g/mole. Determine the empirical and molecular formulas
of this compound.
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10. What is the % by mass of oxygen in water?
11. A compound of iron and oxygen is found to contain 28 g of Fe and 8.0 g of O. What is the %
by mass of each element in the compound?
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