Name:______________________________________
Veritas:______________________________________
Unit 5, Worksheet 1—
Relative Mass
Relative Mass From Gases
We have established that the combining ratio of gases can be explained if two
assumptions are made:
1. Equal volumes of gases contain the same number of molecules at the same
pressure and temperature.
2. Some pure elemental gases are clustered into pairs to form diatomic
molecules.
Now, you should remember that back in Unit 1 we found that iron was more
dense than aluminum. Two possible models arose to account for this difference.
A. The masses of Al and Fe atoms are about the same,
but there are more atoms of Fe than atoms of Al in
each cm3 sample.
B. One cm3 samples of Fe and Al contain about the same
number of atoms, but the Fe atoms are more massive.
A third possibility – that both the size and the mass of the atoms of these two
elements were different – also came up. At the time, we did not have enough
evidence to make a decision about these possible models.
While the reason for density variation between particle types is difficult to
determine for liquids and solids, a conclusion can be reached more easily for
gases due to the fact that particles in a gas are widely spaced. This means that
particle size does not have an effect on the volume that a given number
of gaseous particles occupy.
This makes the determination of the relative mass of individual particles in a gas
fairly simple. If one liter of gas A weighs 5 times as much as one liter of gas B
(at the same T and P) we assume that each particle of gas A weighs 5 times as
much as each particle of gas B. Work through the following example to test your
understanding of this concept.
1. The density of oxygen gas at standard temperature and pressure is 1.43 g/liter,
whereas the density of hydrogen gas under these conditions is 0.089 g/liter. How
many times more massive is one molecule of oxygen than one molecule of
hydrogen? Explain your reasoning.
You shouldn’t conclude that chemists were able to determine the molar masses of
all the elements using this technique. Measurements of the density of the
gaseous phase of many of the elements would be difficult, if not impossible.
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However, we are going to see that chemists could use another tool – the percent
composition of compounds to determine molar masses.
Relative Mass From Compounds
Many substances combine with oxygen to form a type of compound called an
oxide. We have already seen that such combinations often occur in multiple
proportions. John Dalton made the assumption that the lowest ratio was a 1:1
combination of elements. For now, we will make a similar assumption. We may
have to re-examine this assumption later.
2. Based on the % composition of each substance in the table, calculate the mass of
each of the following elements that would combine with 100 grams of oxygen.
Mass of
element
Mass of
oxygen
11.11 g of H
88.89 g of O
42.86 g of C
57.14 g of O
46.67 g of N
53.33 g of O
77.72 g of Fe
22.28 g of O
92.59 g of Hg
7.41 g of O
93.10 g of Ag
6.90 g of O
Mass of element that
would combine with
100g of oxygen
Answers from
Question 4
1.0
3. If these elements combine in a 1:1 ratio in these compounds, we would now have
relative masses for these elements. One could conclude that each atom of oxygen
is
100
= 8.0 times as heavy as an atom of hydrogen. Water presents a problem,
12.5
however; you have seen evidence that two atoms of hydrogen combine with one
atom of oxygen in water. This makes the relative mass of oxygen twice as great
as this value. Explain.
€
4. If we choose 1.0 g of hydrogen as the standard weighable amount – 1 mole – then
the mass of a mole of oxygen atoms must be 16.0 g. Now, using proportional
reasoning, use the mass of the elements relative to 100 g of oxygen you have
calculated in the table above to determine the mass of one mole of each these
elements. Record these values in the table above. How do these molar masses
compare with the values you find in the periodic table? Is the assumption that
we made at the outset valid for all of these compounds? Explain.
Name_______________________________________
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Unit 5, Worksheet 2—
Relative Mass Activity
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)
Empty vial
Vial + Washers
Vial + Hex Nuts
Vial + Bolts
Calculations
Work must be shown and quantities labeled with units and appropriate sig figs.
1. A box of hardware contains 100 pieces. Assuming there are 4 pieces in each vial,
calculate the mass of a box of each kind of hardware. Express these values in
units of g/box.
Washers:
Nuts:
Bolts:
2. If you had 1.00 kg of each kind of hardware, how many boxes of each would you
have?
Washers:
Nuts:
Bolts:
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3. You learned that a barrel of the 1” bolts had a mass of 65.2 kg. The mass of the
barrel was 9.6 kg. How many boxes of bolts are in the barrel?
4. Someone at the Home Depot tells you that a 2” bolt is 6.75 times as heavy as a
washer. What would be the mass of a box of such bolts?
5. Suppose that you were given the job of shipping 25,000 hex nuts to a customer.
How many boxes of hex nuts would this be? All you have is a hanging scale and
a barrel of hex nuts. Describe how you could determine the proper number of
pieces without physically counting them out.
Conclusion
Do you agree or disagree with the following statement? Support your answer.
“You can count by weighing.”
Extension
Each vial contains the same number of pieces of hardware. Calculate the relative
mass of each kind of hardware. Divide each mass by the mass of the smallest.
(The smallest will be 1.00)
Relative mass:
Nuts _____ : Bolts _____ : Washers _____
Suppose that the washer represented an atom of the element carbon. From your
relative masses, determine the elements that would be represented by the nut
and the bolt.
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What is a Mole in Chemistry?
Author: Dr.Badruddin Khan, teaches Chemistry in the University of Kashmir, Srinagar, India
Molecules and atoms are extremely small objects, both in size and mass.
Consequently, working with them in the laboratory requires a large collection
of them. How large does this collection need to be? A standard needs to be
introduced. In chemistry, a mole is a certain number of particles, usually of
atoms or molecules. This number is very special in chemistry and is given the
name Avogadro's number, in honor of Italian chemist and physicist Aamadeo
Avogadro, who first suggested the concept of a molecule. In theory, one could
use any number of different terms for counting particles. For example, one
could talk about a dozen (12) particles or a gross (144) of particles. The
problem with these terms is that they describe far fewer particles than one
usually encounters in chemistry. A unit like the mole is needed because of the
way chemists work with and think about matter. A "unit" is the smallest
measurable entity in the substance, generally either an atom or a molecule. A
mole is the quantity of a substance that contains 6.02 x 1023 units. One mole
of a substance is equal to the substance's atomic weight (the average weight
of an atom of an element) or molecular weight, in grams.
When chemists work in the laboratory, they typically handle a few grams of a
substance. They might mix 15 grams of sodium with 15 grams of chlorine.
But when substances react with each other, they don't do so by weight. That
is, one gram of sodium does not react exactly with one gram of chlorine.
Instead, substances react with each other atom-by-atom or molecule-bymolecule. In the above example, one atom of sodium combines with one atom
of chlorine. This ratio is not the same as the weight ratio because one atom of
sodium weighs only half as much as one atom of chlorine. A mole (symbol
Mol) is the base unit of quantity of a substance in the metric system.
The standard mole is based upon the carbon-12 isotope. Careful
measurements yield a value for NA = 6.0221367x1023. This is an incredibly
large number, almost a trillion trillion. We may say that a convenient name
is given when there is an Avogadro's number of objects; it is called a "mole".
Now the mole concept is no more complicated than the more familiar concept
of a dozen. The mass of a mole of objects will be huge if we consider a mole of
objects of appreciable size such as pennies; however a mole of atoms or
molecules is a different story. We know that the atomic mass unit (amu) is
defined as 1/12 the mass of a carbon-12 atom. Consequently we have the
relation: NA x 12 amu = 12 g. Thus, a mole of carbon-12 atoms has a mass of
just 12 g.
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The mole unit, then, acts as a bridge between the level on which chemists
actually work in the laboratory (by weight, in grams) and the way substances
actually react with each other (by individual particles, such as atoms). One
mole of any substance—no matter what substance it is—always contains the
same number of particles: the Avogadro number of particles. Let us think of
what this means in the reaction between sodium and chlorine. If a chemist
wants this reaction to occur completely, then exactly the same number of
particles of each must be added to the mixture. That is, the same number of
moles of each must be used. One can say: 1 mole of sodium will react
completely with 1 mole of chlorine. It's easy to calculate a mole of sodium; it
is the atomic weight of sodium expressed in grams. And it's easy to calculate
a mole of chlorine; it is the molecular weight of chlorine expressed in grams.
This conversion allows the chemist to weigh out exactly the right amount of
sodium and chlorine to make sure the reaction between the two elements
goes to completion. Even the tiniest speck of sodium chloride (table salt), for
example, contains trillions and trillions of particles. The term mole, by
contrast, refers to 6.022137 × 1023 particles which when written out in the
long form, is 602,213,700,000,000,000,000,000 particles.
The term mole involves the acceptance of two dictates, the scale of atomic
masses and the magnitude of the gram. Both have been established by
international agreement. Formerly, the connotation of "mole" was "gram
molecular weight." Current usage tends to apply the term "mole" to an
amount containing Avogadro's number of whatever units are being
considered. Thus, it is possible to have a mole of atoms, ions, radicals,
electrons, or quanta. This usage makes unnecessary such terms as "gramatom," "gram-formula weight," etc. All stoichiometry essentially is based on
the evaluation of the number of moles of substance. The most common
involves the measurement of mass. The convenient on gases are pressure,
volume, and temperature. Use of the ideal gas law constant R allows direct
calculation of the number of moles: n=P V/R T. T is the absolute
temperature, R must be chosen in units appropriate for P, V, and T. The
acceptance of Avogadro's law is inherent in this calculation.
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Unit 5, Worksheet 3—
What is a mole? Reading Questions
1. What is a mole in chemistry?
2. Provide another example from personal experience of a counting term.
3. Why are chemists not able to use the term “gross” (144) for counting
particles?
4. Why are mass ratios not a good way of determining the definite ratio of
elements in a compound?
5. What is Avogadro’s number? What does it represent?
6. What is another way of expressing or measuring a mole?
7. How is one mole of aluminum the same as one mole of iron? How are
they different?
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Unit 5, Worksheet 4—
Size of a Mole
To help you better visualize the enormous size of Avogadro's number, 6.02 x 1023,
consider the following analogies:
1. If we had a mole of rice grains, all the land area of the earth would be covered
with rice to a depth of about 75 meters!
2. One mole of rice grains is more grain than the number of all grain grown since
the beginning of time.
3. One mole of marshmallows (standard 1 in3 size) would cover the United States to
a depth of 650 miles.
4.
If the Mount St. Helens eruption had released a mole of particles the size of sand
grains, the entire state of Washington would have been buried to a depth equal to
the height of a 10-story building.
5. A mole of basketballs would just about fit perfectly into a ball bag the size of the
earth.
Your turn:
Show your work including units. Use dimensional analysis and show the cancellation of units. Keep 2
sf’s in your answers.
1. Assuming that each human being has 60 trillion body cells (6 x 1013) and that the
earth's population is 7 billion (7 x 109), calculate the total number of living
human body cells on this planet. Is this number smaller or larger than a mole?
Divide the larger value by the smaller to determine the relative size of the two
values.
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2. One of the fastest supercomputers can perform about 12 teraflops (1 teraflop is
1012 calculations per second). Determine how many seconds it would take this
computer to count a mole of things. Convert this figure into years.
3. If you started counting when you first learned how to count and then counted by
ones, eight hours a day, 5 days a week for 50 weeks a year, you would be judged
a 'good counter' if you could reach 4 billion by the time you retired at age 65. If
every human on earth (about 7 x 109) were to count this way until retirement,
what fraction of a mole would they count?
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Unit 5, Worksheet 5—
Gram – Mole – Particles Conversions
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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Unit 5, Lab 1
Title
Introduction
In this experiment, a measured amount of magnesium will be allowed to
react with oxygen in the air. A product of the reaction is magnesium oxide.
You will obtain data that will enable you to determine the empirical
formula of magnesium oxide, MgxOy. “Empirical” means “based on
experimental evidence”.
Question
Safety
1. Wear safety goggles at all times.
2. Handle magnesium ribbon with forceps, not your hands
3. If magnesium begins glowing very bright, do not look directly at it. It
could damage your eyes.
4. Never turn your back to Bunsen burner flame. Long hair should be
pulled back.
5. Handle hot crucible and lid with tongs.
6. Crucible and lid will break if dropped. Be careful when handling!
Procedure
1. Clean your crucible (used crucibles will not become perfectly clean) and
then rinse with dH20 (blue bottle).
2. Heat crucible for a couple of minutes to dry. Remove from flame and
allow crucible to return to room temperature.
3. Record the mass of the clean, dry crucible and lid.
4. Polish, with steel wool, 0.15-0.20g of magnesium ribbon.
The ribbon should be a bright grey color when finished.
5. Loosely curl ribbon so it will lie in the bottom of the
crucible. Do not wad the ribbon or curl it too tightly.
6. Record the mass of the crucible, lid, and magnesium
ribbon.
7. Place the crucible in a clay triangle (on a tripod) as
shown in the picture to the right.
8. Using a Bunsen burner, slowly heat the sample. Be
sure to leave the lid slightly ajar. (See picture)
9. Occasionally, lift the lid to allow more air into the
crucible (see picture). Too much air will cause the
ribbon to glow brightly. If this happens, replace the lid
immediately.
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10. Continue to heat until no noticeable change is observed in the ash at
the bottom of the crucible.
11. Remove the lid and heat for an additional 30 seconds.
12. Stop heating and allow the crucible to return to room temperature.
13. Record the mass of the crucible, lid, and ash on the same balance you
used previously.
14. Add 2-3 drops of water (blue bottle) and reheat the crucible for 1-2
minutes. No water should be remaining in crucible.
15. Once the crucible has cooled to room temperature, record the mass of
the crucible, lid, and ash.
16. If the mass has changed, reheat the crucible for another 1-2 minutes.
Allow to cool and mass again.
17. Clean crucible and lid. Make sure all materials are put back in their
appropriate locations. Wipe down lab table. Check table and floor for
trash.
Data
Create a data table for the data recorded in lab.
Observations
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
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Calculations
Show all work. All numbers should have units and the appropriate number of
sig figs.
1. Determine the mass of magnesium reacted.
2. Determine the mass of magnesium oxide (product).
3. Determine the mass of oxygen in the magnesium oxide.
4. Determine the number of moles of magnesium
5. Determine the number of moles of oxygen.
5. Determine the mole-to-mole ratio of Mg to O:
Conclusion
1. We know that atoms of compounds combine in simple, whole number
ratios (ex. H2O). What do you think is the likely ratio for the compound
containing magnesium and oxygen?
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2. What is the empirical formula of magnesium oxide?
3. How did you decide on that formula?
4. What are some possible lab errors that could lead to an incorrect formula?
Your lab errors must be logical given your mole-to-mole ratio. How, if
possible, could those errors be avoided?
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Unit 5, Worksheet 6—
Empirical and Molecular Formulas
Show all your work when solving the following problems. Circle the final answer. Be sure to
include units and the correct number of significant figures.
1. Find the empirical formula of a compound containing 32.00 g of bromine and 4.90
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 this 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 times 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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Unit 5, Worksheet 7—
More Empirical and Molecular Formulas
Show all your work when solving the following problems. Circle the final answer. Be sure to
include units and the correct number of significant figures.
1.
The compound benzene has two formulas, CH and C6H6. Which of these is
the empirical formulas and which is the molecular formula?
_______________ empirical formula
_______________ molecular formula
2.
There are two common oxides of sulfur. One contains 32 grams sulfur and 32
grams oxygen. The other oxide contains 32 grams sulfur and 48 grams
oxygen. What are the empirical formulas for the two oxides?
3.
A form of phosphorus called red phosphorus is used in match heads. When
0.062 grams of red phosphorus burns, 0.142 grams of phosphorus oxide is
formed. What is the empirical formula of this oxide?
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4.
A certain compound is composed of 7.20 grams carbon, 1.20 grams hydrogen,
and 9.60 grams oxygen. The molecular mass of the compound is 180.0 g/mol.
What is the empirical formula and the molecular formula for this compound?
5.
Oxalic acid is a compound used in cosmetics and paints. A 0.725 gram
sample of oxalic acid was found to contain 0.194 grams carbon, 0.016 grams
hydrogen, and 0.516 grams oxygen. If the molecular mass of oxalic acid is
90.04 g/mol, what is the molecular formula for oxalic acid?
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Unit 5 — More Practice Problems
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 (include units):
a. KNO3
______________ h. UF6
b. (NH4)2CO3
______________
______________ h. UF6
______________
c. Ag2CrO4 ______________ j. H3PO4 ______________
d. oxygen gas ______________ k. (NH4)2SO4
______________
e. Ca(NO3)2 ______________ l. CH3COOH______________
f. PbSO4
______________ m. Pb(NO3)2
g. Mg(OH)2
______________ n. Ga2(SO3)3 ______________
______________
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 à moles or from moles à g. Show units cancelling.
a.
12.0 g Fe x
=
moles
b.
25.0 g of Cl2 gas x
c.
0.476 g of (NH4)2SO4 x
d.
0.15 moles NaNO3 x
e.
0.0280 moles NO2 x
=
f.
0.64 moles AlCl3 x
=
=
moles
=
moles
=
g
g
g
g. Convert 30 grams of H3PO4 to moles.
h. Convert 25 grams of HF to moles.
i.
Convert 110 grams of NaHCO3 to moles.
j.
Convery 4 moles of Cu(CN)2 to grams.
k.
Convert 5.6 moles of C6H6 to grams.
l.
Convert 21.3 moles of BaCO3 to grams.
m.
If you had 2.50 moles of oxygen gas, what mass of the gas would be in the
sample?
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5. Use Avogadro’s number to do the following conversions.
a. How many atoms are there in 0.00150 moles Zn?
b. A 4.07 g sample of NaI contains how many atoms of Na?
c. How many atoms of chlorine are there in 16.5 g of iron (III) chloride, FeCl3?
d. What is the mass of 100 million atoms of gold? Could you mass this on a
balance?
e. How many molecules are there in 24 grams of FeF3?
f.
How many molecules are there in 450 grams of Na2SO4?
g. How many grams are there in 2.3 x 1024 atoms of silver?
h. How many grams are there in 7.4 x 1023 molecules of AgNO3?
i.
How many grams are there in 7.5 x 1023 molecules of H2SO4?
j.
How many grams are there in 4.5 x 1022 molecules of Ba(NO2)2?
k. How many molecules are there in 9.34 grams of LiCl?
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6. Empirical and Molecular Formulas
a. 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
b. Calculate the empirical formula of a compound that contains 4.20 g of
nitrogen and 12.0 g of oxygen.
c. 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?
d. 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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e. What is the % by mass of oxygen in water?
f.
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?
g. A 5.438 gram sample, was found to contain 2.549 grams of iron, 1.947 grams
of oxygen, and 0.9424 grams of phosphorus. What is its empirical formula?
h. Aniline, a starting material for urethane plastic foams, consists of C, H, and
N. Combustion of such compounds yields CO2, H2O, and N2 as products. If the
combustion of 9.71 g of aniline yields 6.63 g H2O and 1.46 g N2, what is its
empirical formula?
The molar mass of aniline is 93 g/mol. What is its molecular formula?
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i.
When 2.5000 g of an oxide of mercury, (HgxOy) is decomposed into the
elements by heating, 2.405 g of mercury are produced. Calculate the
empirical formula.
j.
A 5.438 gram sample, was found to contain 2.549 grams of iron, 1.947 grams
of oxygen, and 0.9424 grams of phosphorus. What is its empirical formula?
7. Find the molecular formula of the following compounds.
a. A compound with an empirical formula of CFBrO and a molar mass of 254.7
grams per mole.
b. A compound with an empirical formula of C2H8N and a molar mass of 46
grams per mole.
c. A compound with an empirical formula of C2OH4 and a molar mass of 88
grams per mole.
d. A compound with an empirical formula of C4H4O and a molar mass of 136
grams per mole.
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Practice Problems for Semester Exam (Units 1-5)
Part I: Measurement and Calculations
1. Demonstrate understanding of the use of measurements in science. You should
be able to apply the rules of significant figures to choose the answer with the
correct number of significant figures.
a. State the number of significant figures in each of the following numbers.
1) 5.432 ________
8) 17.20
________
15) 300.10
________
2) 31.2 ________
9) 0.450
________
16) 4000.5
________
3) 304.1 ________
10) 4.560
________
17) 2.30 x106 ________
4) 20.9 ________
11) 3.4x10¯4 ________
18) 1.20 x10¯8 ________
5) 0.56 ________
12) 2.7x105 ________
19) 22.0030
________
6) 0.032 ________
13) 500
________
20) 4.00900
________
7) 34.0 ________
14) 360
________
21) 34,000
________
b. Rewrite the following numbers rounded to the indicated number of sig figs.
22) 5.67 (2 SF)
___________
28) 31.8 (2 SF)
23) 30.53 (3SF)
___________
29) 221.351 (4 SF) ___________
24) 24.25 (3 SF)
___________
30) 16.05 (3 SF)
___________
25) 0.04500 (3 SF) ___________
31) 3456 (2 SF)
___________
26) 40,001 (3SF)
___________
32) 3000.34 (2 SF) ___________
27) 77 (1 SF)
___________
34) 44 (4 SF)
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___________
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c. Express the answer in the correct number of significant figures. Label with
appropriate units.
a. 21.3 g = 16.384615
1.3 cm3
__________________
b. 6.34 cm2 x 1.2 cm
1.217 cm
__________________
= 6.251437
c. 13.21m x 61.5 m = 812.415
__________________
d. 21.50 cm
8.50 in
__________________
= 2.529411765
e. 334.54 grams + 198 grams
__________________
f.
34.1 grams / 1.1 mL
__________________
g.
2.11 x 103 joules / 34 seconds
__________________
h. 0.0010 meters – 0.11 m
__________________
i.
__________________
349 cm + 1.10 cm + 100 cm
j. 450 meters / 114 seconds
__________________
k. 298.01 kilograms + 34.112 kg
__________________
l. 4 m/s x 31.221 s
__________________
d. You should also be able to measure the length of an object or volume of a liquid
to the appropriate number of significant figures based upon the measuring
instrument.
Which of the following
best expresses the width
of the business card?
a.
5 cm
b. 5.0 cm
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d. 5.50 cm
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%
%
2.)%
%
%
%
Which of the following best expresses the
3.)% volume of the liquid in the graduated cylinder?
%
a. 40 mL
%
b. 43 mL
%
c. 43.0 mL
%
4.)% c. 44.0 mL
d. 43.01 mL
%
%
%
5.)%
%
Below each cylinder, record the volume of the
%
liquid in the graduated cylinder using
Part%2% %What%are%the%readings%on%these%graduated%cylinders?%%Be%sure%to%include%units%with%your%
appropriate significant figures?
answers.%
6.)%
%
%
%
7.)%
%
%
%
8.)%
%
%
%
9.)
%
% ______________
10.)% %
%
%
%
%
%
%
!
______________
11.)%
%
%
______________
%%%%%%%12.)%
%
%
Unit%1% %Math%&%Measurement%
___________
%%%%%%13.)%
Rulers,%graduated%cylinders,%and%thermometers%
from%math9aids.com%
%
!
______________
!
______________
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2. Demonstrate proficiency in the use of scientific notation and use of dimensional
analysis in metric conversions. Know the meaning of the following metric prefixes
and be able to make conversions utilizing them: milli-, centi-, kiloExample:
150mm ×
1m
= 0.15m
1000mm
Complete the indicated conversions:
€
a. 37 g x
=
mg
b. 4.7 kg x
=
g
c. 138 m x
=
km
d. 4021 mm x
=
m
f.
1000. cL = ? L
g. 2.66 cm = ? mm
h. 64 mm = ? cm
i.
4.32 kg = ? mg
3. Be able to convert standard (decimal) notation to scientific notation and vice
versa.
Standard:
Scientific:
Standard:
Scientific:
1300
___________
___________
24,212,000
0.00155
___________
___________
0.000665
___________
1.68 x 106
0.00332
___________
___________
2.73 x 10-2
314159
___________
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4. What is the Law of Conservation of Mass?
Study the mass and change lab.
a. Under what conditions did you observe a decrease in mass?
b. Explain how could the mass decrease if matter is conserved.
c. Under what conditions did you observe an increase in mass?
d. Explain how mass can increase if matter is conserved.
5. Determine the density of an object from a data table or from a graph of Mass v
Volume.
Volume (cm3)
1.5
3.0
4.5
6.0
7.5
9.0
Mass (g)
11.7
24.0
35.1
48.0
58.5
70.0
a. Plot the data above.
b. Determine the density of the
substance.
c. What volume would 150g of the substance occupy? Show work; use labels.
d. What mass would 5 cm3 of the substance have?
Show on the graph above how you could answer the question.
Show below how you could answer the question mathematically.
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7. Complete density calculations and conversion problems.
a.
Mercury metal is poured into a graduated cylinder that holds exactly
22.5 mL. The mercury used to fill the cylinder weighs 306.0 g. From
this information, calculate the density of mercury.
b. A block of lead has dimensions of 4.50 cm by 5.20 cm by 6.00 cm. The
block weighs 1587 g. From this information, calculate the density of
lead.
c. What is the mass of the ethanol that exactly fills a 200.0 mL
container? The density of ethanol is 0.789 g/mL.
d. Find the mass of 250.0 mL of benzene. The density of benzene is
0.8765 g/mL.
e. 28.5 g of iron shot is added to a graduated cylinder containing 45.50
mL of water. The water level rises to the 49.10 mL mark, From this
information, calculate the density of iron.
Part II: The Role of Energy in Physical Change
1. Describe the ways energy is stored in solids, liquids and gases (thermal, phase,
chemical). Also describe ways energy is transferred (working, heating,
radiating).
Thermal energy – energy of motion – related to the absolute temperature
• Hotter molecules move more rapidly than slower ones; for a given volume
the gas will have a greater pressure due to the greater number of
collisions
Phase energy – energy due to attractions between molecules; the stronger the
attractions, the lower the energy of the system of particles
• Lowest for solids, greater in liquids, greatest in gas phase
2. Heat is the transfer of energy into or out of a system due to molecular collisions.
Energy is transferred from hotter (faster) molecules to colder (slower) molecules.
a. Explain why the alcohol level in a thermometer rises when it is placed in
a warmer fluid. (3-step process)
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b. Describe what happens (at the molecular level) when a glass of cold
water warms up to room temperature.
3. Be able to draw energy bar graphs to account for energy storage and transfer
in all sorts of changes.
Complete the energy bar chart for the following scenario:
An ice cube tray of water at room temperature is placed into the freezer and the
liquid changes to solid.
Motion before:
Motion after:
Arrangement before:
Arrangement after:
4. When energy is transferred to a sample of matter, either the particles speed
up (temperature increases) or they get pulled apart (phase change), but not both
at the same time. This helps account for the shape of the warming curve you got
in the Icy Hot lab.
a. On the graph above label which phases are present in each portion of the
curve.
b. Label the sections in which the thermal energy (Eth) of the sample is
changing (indicating an increase or a decrease). Label the sections where the
phase energy (Eph) is changing (indicating an increase or a decrease).
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5. Use the equations Q = mH, where H f = 334 J
J to determine
g , H v = 2260 g
the energy transferred to or from water during a phase change. Use
to determine the energy transferred to or from water
Q = mcΔt and c = 4.18 J
g˚C
during heating
€ or cooling.
€
Draw a heating curve and mark on the curve the beginning and ending points.
Using the appropriate equations and values answer the following:
a. How much energy would be required to bring 100g of ice at 0°C to its
boiling point?
b. Suppose that during the Icy Hot lab that 65 kJ of energy were
transferred to 450 g of water at 20.˚C. What would have been the final
temperature of the water?
c. An ice cube tray full of ice (235g) at –7.0˚C is allowed to warm up to room
temperature (22˚C). How much energy must be absorbed by the contents
of the tray in order for this to happen?
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Part III: Gases and Kinetic Theory
Demonstrate knowledge of the relationships that exist among the pressure, volume,
temperature, and number of molecules of a gas. You should be able to determine the
pressure of a sample of gas in a flask connected to a manometer.
1. You should be able to identify the correct graphic representation of the
relationships between volume, temperature, and pressure.
a. Which graph describes the relationship between gas pressure and volume?
Explain.
b. Which graph describes the relationship between gas pressure and the Kelvin
temperature? Explain.
2. Solve problems given volume, pressure, or temperature of gases (PVTn charts).
a. A sample of carbon dioxide has a volume of 2.0 L at a temperature of –10˚C.
What volume will this sample have when the temperature is increased to 110˚C.
Assume that the pressure does not change and that no carbon dioxide leaks from
the sample.
b. A 12.7 L sample of gas is under a pressure of 740 mm Hg at 20°C. What will
be the volume of the gas if the pressure increases to 1.00 atm and the
temperature drops to 0.0°C?
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3. Based on the height of a column of mercury in an open U-tube manometer, be
able to compare pressures on both sides of the tube.
a. Determine the pressure in each of the flasks.
Part IV: Matter and Atomic Theory
1. Describe how matter is organized.
You should be able to identify diagrams and distinguish between pure
substances, atoms, molecules, elements, compounds, mixtures, gases, and solids.
a. Which diagram(s) show only molecules?
b. Which diagram(s) show pure substances?
c. Which diagram(s) show mixtures?
d. Which diagram(s) show only atoms?
e. Does Figure B represent a compound or an element? Explain.
f.
Which diagram(s) show only an element?
g. Which diagram(s) show only a compound?
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2. Demonstrate knowledge of separation techniques for mixtures and
compounds.
a. Distinguish between the separation techniques required for mixtures
and compounds.
b. List and describe a few examples below.
c. How would you separate a mixture of water and ethanol? Describe the
process and include a temperature-time graph.
d. How would you separate a mixture of salt and sand? Describe the
process(es) involved.
e. How would you separate the compound water into its elemental
components?
Describe the process.
Identify the apparatus used.
Label which substance is found on the
anode side and which substance is
found on the cathode side.
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3. Demonstrate knowledge of the meaning of a chemical formula in terms of atoms
and molecules. Given the formula of a compound, identify the number of atoms
present.
Example: Identify the number of atoms of each kind in each compound.
Pb(NO3)2
Na3PO4
Al2(SO4)3
Part V. The mole concept and chemical reactions
Recognize that atoms are too small to count directly. We determine how many
there are in a sample by finding their mass. We use the mole to determine the
number of atoms and molecules. Molar mass (on Periodic Table) is relative
mass, based originally on hydrogen (lightest element).
1. Be able to determine the molar mass of a compound.
Determine the molar masses (using correct SFs and giving unit of measurement!).
a. Pb(NO3)2
_____________
c. BaSO4
_____________
b. MgCl2
_____________
d. oxygen gas
_____________
2. Determine the number of atoms or moles using Avogadro’s number and the molar
mass of a compound.
a. 12 g MgCl2 x
=
b. 3 moles Cl x
=
moles MgCl2
atoms Cl
c. How many moles are 1.20 x 1025 atoms of phosphorous?
d. How many atoms are in 0.750 moles of zinc?
e. Find the grams in 1.26 x 10-4 mol of HC2H3O2.
f. Find the number of moles of argon in 452 g of argon.
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3. Determine the empirical and molecular formulas from data.
a. 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. Find the empirical and
molecular formulas for this compound.
b. Aniline, a starting material for urethane plastic foams, consists of C, H, and
N. Combustion of such compounds yields CO2, H2O, and N2 as products. If the
combustion of 9.71 g of aniline yields 6.63 g H2O and 1.46 g N2, what is its
empirical formula?
The molar mass of aniline is 93 g/mol. What is its molecular formula?
4. Percent Composition
a. What is the % by mass of oxygen in water?
b. 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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Part VI. Additional Terms and Concepts
Law of Definite Proportions
Definintion:
Examples:
Law of Multiple Proportions
Definintion:
Examples:
Empedocles
Importance:
Democritus
Importance?
Dalton’s Atomic Theory
Components:
a.
b.
c.
d.
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