NAME____________________________________ PER____________ DATE DUE____________
ACTIVE LEARNING I N C HEMISTRY E DUCATION
“ALICE”
CHAPTER 3
PHYSICAL
STATES
OF
MATTER
3-1
© 1997, A.J. Girondi
NOTICE OF RIGHTS
All rights reserved. No part of this document may be reproduced or transmitted in any form by any means,
electronic, mechanical, photocopying, or otherwise, without the prior written permission of the author.
Copies of this document may be made free of charge for use in public or nonprofit private educational
institutions provided that permission is obtained from the author . Please indicate the name and address
of the institution where use is anticipated.
© 1997 A.J. Girondi, Ph.D.
505 Latshmere Drive
Harrisburg, PA 17109
[email protected]
Website: www.geocities.com/Athens/Oracle/2041
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© 1997, A.J. Girondi
SECTION 3.1 Physical Properties of Matter
You have already seen that chemistry is a science that deals with the composition of substances
and the changes in composition that they may undergo. By simply observing a chemical substance, you
cannot always determine the composition of that substance. You can usually distinguish one chemical
substance from another by experimentation. Sometimes the experiments you perform will be qualitative
activities, and at other times they will be mainly quantitative.
Chemical substances are usually distinguished by their appearance, taste, odor, feel, and other
similar properties. Such characteristics are called physical properties. We can recognize specific
substances by their physical properties. Just as you can recognize your friends by their physical
appearance, you can also recognize chemical compounds by their physical appearance and properties.
To obtain a better understanding of physical properties, we will attempt to distinguish between two
different metals on the basis of their physical properties.
ACTIVITY 3.2 Identifying Metals By Physical Properties
Get a sample of metal A and a sample of metal B from the materials shelf. Metals A and B look quite
similar. It would be difficult for the untrained eye to determine the identities of these metals based on their
appearance alone. But, there are ways of identifying these two metals based on their physical properties
listed in Table 3.1.
If you were to refer to a reference book, such as the Handbook of Chemistry and Physics, you
would probably have a difficult time trying to determine the identity of metal A and metal B. There are
simply too many metals with the same or similar properties because the physical properties that you have
described are very general.
Would you categorize the first four observations called for in Table 3.1 as quantitative or qualitative?
{1}_____________________
Why?{2}_______________________________________
There are several physical properties of substances that are more quantitative in nature. Many
substances are more easily identified by quantitative means. For example, every pure substance (element
or compound) has a specific temperature and pressure at which it boils. If we were to melt metals A and B,
we could identify them by comparing their measured melting points with the melting points of metals
found in a reference book.
Another physical property that is a good quantitative measure is the density of a substance.
Density is the mass of a given volume of a substance. You would probably agree that a block of cement is
heavier than a block of styrofoam of the same volume. As a result, we would say that the cement is more
dense than styrofoam. We will try to identify metals A and B by calculating the density of each.
Density can be measured in a very simple way. The
mathematical formula for density is given at right. The
measurements you will need to determine density are the
masses of your metals and the volumes occupied by the
metals. To determine the densities of metals A and B use the
following procedure.
density =
mass (grams)
volume (mL)
.
Procedure:
1. At this time, make the first four observations of metals A and B listed in Table 3.1. Record your
observations in the table. Following the remaining steps to determine the densities of metals A and B.
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© 1997, A.J. Girondi
2. Obtain a 50 mL graduated cylinder. Add 40.0 mL of water to the graduate. Measure the mass of the
graduate containing 40.0 mL of water. Record this mass on a piece of paper.
3. Add as many pieces of metal A as needed to the water in the graduate to make the total volume
between 47.0 and 50.0 mL. All of the pieces of metal must be totally submerged. Calculate the volume of
the metal in the graduate by subtracting the original volume of water (40.0 mL) from the final volume
reading. Record this result as the volume of metal A in Table 3.1.
4. Weigh the graduate and content (water + metal). From this final mass subtract the original mass of the
graduate + water. Record this difference as the mass of metal A in Table 3.1.
5. Dry the pieces of metal A with a towel and return them to the materials shelf.
6. Repeat steps 1 through 4 using metal B.
Table 3.1
Observed Properties of Metals A and B
Physical Property
Metal A
Metal B
1. Physical State
________________
________________
2. Color
________________
________________
3. Luster
________________
________________
________________
________________
(solid, liquid, gas)
(shiny, dull, etc.)
4. Texture
(rough, smooth)
Table 3.2 lists various common metals and their densities
including metals A and B. Calculate the density of each metal in the
space provided below. Be sure to include units with your answers.
Enter the results into Table 3.1. Compare your calculated densities
of metals A and B to the densities of the metals in Table 3.2. Find the
metals in Table 3.2 whose densities most closely match those for
metals A and B. Enter the identities of metals A and B in Table 3.3.
Using the Handbook of Chemistry and Physics or whatever
other reference resource is available to you, look up metals A and B
and list any other properties that are given. A copy of the handbook
may be available in your lab area and/or in the school library. Enter
the additional properties of these metals into Table 3.3.
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Table 3.2
Densities of Some
Common Metals
Metal
lead
zinc
mercury
tin
nickel
platinum
aluminum
Density (g/mL)
11.34
7.14
13.59
5.75
8.90
21.45
2.70
© 1997, A.J. Girondi
Metal A: Show density calculations in space below.
Metal B: Show density calculations in space below.
Table 3.3
Identities and Densities of Metals A and B
Metal A.
Mass: _____________ g; Volume: _____________ mL
Density: ____________ g/mL; Identity: _____________________________
Metal B.
Mass: _____________ g; Volume: _____________ mL
Density: ____________ g/mL; Identity: _____________________________
ACTIVITY 3.3 Comparing the Densities of Liquids
Liquids also have densities, but they are determined in a slightly different way. Use the procedure
below to determine the density of liquids.
1. Measure the mass of a clean, dry 10 mL graduated cylinder to the nearest 0.01 g. Record the mass in
Table 3.4.
2. Obtain a bottle of alcohol, oil, or vinegar from the materials shelf. Carefully measure exactly 10.0 mL of
one of the liquids into the graduate. It is extremely important for you to read the meniscus correctly in this
activity. Use a dropper to adjust the volume to exactly 10.0 mL, if needed.
3. Measure the mass of the liquid to the nearest 0.01 g, and record this mass in Table 3.4. Discard the
liquid into the sink after use.
4. Repeat steps 1 through 3 for the other liquids. When finished with the oil, wash the graduate using a
small test tube brush and some detergent. Rinse well with water.
5. Calculate the mass of each liquid using the data collected. Calculate the density of each liquid by
dividing its mass by its volume (10.0 mL). Be sure that each measurement or calculation in your table has
a unit with it.
6. Wash and return all glassware.
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© 1997, A.J. Girondi
Table 3.4
Calculated Densities of Selected Liquids
Alcohol
Oil
Vinegar
Mass of graduate + liquid
________
________
________
Mass of graduate
________
________
________
Mass of 10 mL of liquid
________
________
________
Density of liquid (g/mL)
________
________
________
Volume of liquid used
_10 mL__
_10 mL__
_10 mL__
Use the data in Table 3.4 to answer the following questions.
1. Which liquid appears to be most dense? _____________________________________________
2. Which liquid appears to be least dense? _____________________________________________
3. If oil is added to water, 2 layers form. One layer is water, while the other is oil. Which substance would
you expect to find on the bottom and which on top? (The density of water is about 1.00 g/mL) Explain:
_____________________________________________________________________________
_____________________________________________________________________________
Put about 2 mL of oil into a test tube. Carefully add about 2 mL of water. Observe the behavior of the
water as you add it. Which liquid is on top?_________________ Was your prediction correct?_______
(Use detergent and water again to clean the oily glass.) Discard the liquids into the sink.
4. Specific gravity is the ratio of the density of a liquid compared to
the density of water (1.00 g/mL). Since it is a ratio, specific gravity
has no units. For example, suppose the density of a liquid is 1.80
g/mL. The specific gravity of that liquid is calculated as shown at
right.
S.G. = 1.80 g/mL = 1.8
1.00 g/mL
Notice that the only difference between the density of a liquid and its specific gravity is that
density has units and specific gravity does not. What is the specific gravity of the oil that you
used?__________
SECTION 3.4 Phase Change Equations
Phase changes are examples of changes in physical properties. Phase changes can be
described in the form of equations. Most solids when heat sufficiently undergo a change in state from a
solid to a liquid phase. Most of us are familiar with the process called melting. This phase change can be
written as an equation:
∆
water(s) ---------->
water(l)
The ∆ symbol means “heat added” while the subscripts (s) and (l) refer to solid and liquid, respectively.
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© 1997, A.J. Girondi
The reverse process can be written as: water(l) ----> water(s) and is called freezing. Using these two
examples, write a similar equation to show the phase change involved in boiling. Use the subscript (g) to
represent a gas or vapor. Use any appropriate symbols as well.
{3}_______________________________________________
Condensation is the opposite of boiling. Write an equation to show the phase change involved in
condensation.
{4}_______________________________________________
ACTIVITY 3.5 Energy Changes Which Accompany Phase Changes
In this activity you will be observing a phase change firsthand. You are going to obtain data and
draw what is known as a cooling curve. Remember to always wear safety glasses and an apron.
1. Fill a 250 mL beaker about 1/3 full of tap water and bring the water to a boil. Obtain a stopwatch, if
possible, or have available a wrist watch or classroom clock with a second hand.
2. Obtain a test tube which already contains a small amount (about 2 cm) of paradichlorobenzene (PDB)
from the materials shelf. Keep the tube of PDB away from any flame! Your lab should have good
ventilation during this activity.
3. Turn off the heat under the beaker of water and place the tube of PDB into the hot water. Allow it to
remain there until the solid is completely melted. At that point, place a thermometer directly into the
melted PDB. Do not remove the tube of PDB from the hot water until the temperature of the melted PDB
is between 65oC and 80oC. (The higher the better within that range if you have time.)
Do not try to use a flame around or under the melted PDB. It is very flammable!
4. When the appropriate temperature has been reached, remove the tube of PDB from the hot water and
place it into a test tube rack on your lab table. Start timing immediately, and record the temperature of the
cooling PDB every 30 seconds.
5. Record the temperature as precisely as possible to the nearest 0.5oC, and use the thermometer to
carefully stir the PDB 10 seconds before each reading is taken. You will need a partner to help keep time,
read the thermometer, and record the temperatures. Enter all data into Table 3.5.
6. After you have obtained 3 or 4 consecutive readings which are within 1 degree of each other, you can
stop if your class is close to ending. However, if you have time you should continue taking readings to
complete as much of the data table as possible. If the PDB gets too solid to stir, discontinue that. Stop
the readings if the temperature reaches 40 oC.
7. When you have collected all needed data, reheat the water in your beaker and put the tube of PDB in
it. Remelt the PDB so that you can remove the thermometer. Remove the tube from the hot water and
allow it to cool. DO NOT remove the PDB from the test tube. Return all equipment to the materials shelf.
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© 1997, A.J. Girondi
Table 3.5
Cooling Temperatures
Time (sec)
Temp(oC)
Time (sec)
Temp(oC)
____0____
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
Prepare a graph of your data on the grid below. Record temperatures on the Y axis and time on
the X axis. What was the dependent variable in this activity?{5}____________________
independent variable?{6}______________________.
What was the
Compare your graph to the graph of a normal
cooling curve for a pure substance (element or compound) shown in Figure 3.1. Notice how the graph
has been divided into five sections. The data you collected reflects only a part of the overall cooling curve
for PDB.
An interesting thing to notice on the cooling curve is that there are two plateaus where
temperature remains constant. Each plateau signifies a phase change (phase changes occur at constant
temperatures).
In section B-C of the graph in Figure 3.1, condensation is occurring.
Section D-E
represents freezing.
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© 1997, A.J. Girondi
vapor
vapor + liquid
liquid
liquid + solid
solid
A
B
C
Temp
(oC)
D
E
F
Time (sec) ----->
Figure 3.1 Cooling Curve of a Pure Substance
If we were to pretend that the curve in Figure 3.1 is a cooling curve for PDB, what letter on the curve most
closely approximates the area in which you collected data in this activity?{7}__________
Before you go any further, draw a circle around that area of the graph in Figure 3.1 above which best
represents the range in which you collected your cooling curve data. Seek help if you need it.
SECTION 3.6 Kinetic and Potential Energy Changes During Phase
Changes
You may ask, “Why is the temperature constant while a phase change is occurring?” In order to
understand the answer, you must be aware that temperature is a measure of kinetic energy - the energy
due to the motion of the particles. When particles lose kinetic energy, their temperature decreases (since
temperature is a measure of kinetic energy). However, when particles move farther apart or closer
together, they gain or lose potential energy which is energy of position. Particles move closer together
when they lose potential energy, and farther apart when they gain potential energy. Potential energy
changes do not involve temperature changes.
When molecules have a lot of kinetic energy, their motion keeps them from being drawn closer
together by attractive forces that exist between them. However, if the particles lose kinetic energy, there
will come a point at which the attractive forces will pull them closer together. This involves a loss of
potential energy which occurs at a constant temperature. The result will be a phase change like
condensation or freezing.
When substances are heated, the molecules move faster and gain kinetic energy as temperature
increases. Eventually, they may gain enough kinetic energy to overcome some of the attractive forces
between them and move farther apart. When they move apart they are gaining potential energy, and they
do this at a constant temperature. This results in a phase change such as melting or vaporization. There
are plateaus on a warming curve just as there are on a cooling curve. On such a curve, particles are moving
faster when temperature is rising. They are moving farther apart when the temperature is constant during
the heating process. Pure liquids boil at constant temperatures (at constant pressure) because the
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© 1997, A.J. Girondi
particles are moving farther apart and potential energy is being stored as the liquid becomes a vapor.
Refer back to Figure 3.1 as you answer the following questions. In which sections of the cooling
curve is the temperature changing?{8}____________________________
According to the curve, does
the temperature change or remain constant during a phase change?{9}__________________________
In
which
section(s)
of
the
curve
cooler?{10}__________________
slower?{11}_____________________
is
the
substance
changing
temperature
and
getting
In which section(s) of the curve are the particles moving
In which section(s) of the curve are the particles moving at a
constant rate?{12}______________________
In which section(s) are particles moving closer
together?{13}_______________________
The freezing point of a pure substance (element or compound) is the temperature at which it
changes from the liquid to the solid phase. On the graph, the freezing point (FP) of that substance is the
temperature at which the plateau occurs. Look at the graph you plotted in Activity 3.5. What is your
estimate of the freezing point of the PDB? {14}_______________
Steam at 100oC contains more heat and can, therefore, cause a more serious burn than boiling
water at the same temperature. Explain this in terms of the kinetic and potential energy content of the
steam and of the boiling water. {15}_____________________________________________________
_____________________________________________________________________________
_____________________________________________________________________________
Figure 3.2 is a graph showing the warming curve of water. It is the opposite of a cooling curve.
Divide the graph into sections. Label with a pencil which phase or phases (solid, liquid, vapor) are present
in each section of the graph, and label those areas of the curve where phase changes (melting and
vaporization) are occurring.
According to Figure 3.2, what is the melting point of ice?{16}__________ In the same way, what is
your estimate of water’s boiling point based on Figure 3.2? {17}_____________ How do the melting and
freezing points compare? {18}_________________________________________________________
One rather interesting point about phase changes is that even though phase changes occur, you
always have the same chemical substance you began with. The solid, liquid, and gaseous phases of water
or PDB are easily converted back and forth by simply adding or removing heat energy. (See Figure 3.3)
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© 1997, A.J. Girondi
125
100
75
oC
50
25
0
-25
Time (sec) ------->
Figure 3.2
Warming Curve of Water
LIQUID WATER
add energy
add energy
WATER
VAPOR
ICE
remove energy
remove energy
LIQUID WATER
Figure 3.3
Energy and Phase Changes
Every pure substance melts and boils at a
specific temperature. Table 3.5 contains a list of
some selected pure substances. As can be seen,
each substance has its own characteristic melting
and boiling point. As a result, some substances
exist as gases at room temperature while others are
solids or liquids. Use the melting and boiling point
information provided to determine which of the
substances listed are solids, which are liquids, and
which are gases at room temperature (25oC) and
normal pressure.
3-11
Table 3.5
Melting and Boiling Points
Substance
Neon
Water
Acetone
Sodium
Ethyl Alcohol
Copper
M.P. (oC)
-248.7
0
-95.0
97.5
-117.3
1083.0
B.P. (oC)
-245.8
100.0
56.6
899.0
78.5
2582.0
© 1997, A.J. Girondi
The solids are:
{19}________________________________________________________________
The liquids are: {20}________________________________________________________________
The gases are : {21}________________________________________________________________
ACTIVITY 3.7
Physical and Chemical Changes
Now that we have learned the fine points about phase changes, we are going to look at two other
kinds of changes - physical changes and chemical changes. A physical change refers to a change in the
appearance of a substance. Like a phase change, when a physical change occurs, you also end up with
the same substance you began with - it just looks different. If you tear a sheet of paper into small pieces,
you have brought about a physical change. You still have paper, which is chemically unchanged. Is
tearing a piece of paper also a phase change?{22}_________Why or why not?{ 23}__________________
______________________________________________________________________________
When ice is changed to liquid water and then to steam, is water undergoing a physical change?________
Why or why not? {24}_______________________________________________________________
A somewhat different type of change is called a chemical change. After a chemical change
occurs, you no longer have the same chemical substance. A new and different chemical substance is
formed. If you were to set fire to paper, would you still have paper? Obviously not! Chemical changes
result in the formation of substances with new chemical compositions. There are some key signs that you
can look for to determine whether or not a chemical change has occurred. For example, if a gas is given off
or if a precipitate (solid) is formed during a change, you can be fairly certain that a chemical change has
occurred. Often (but not always) a color change can indicate a chemical change. A temperature change
may or may not indicate that a chemical change has occurred. Color changes and temperature changes
sometimes occur during physical changes, so they are not clear evidence of chemical changes.
In this activity you will witness some physical and chemical changes. You will need 4 test tubes, a
test tube rack, and the chemicals labeled “ALICE 3.7” on the materials shelf. Follow the directions below,
and record all of your observations in Table 3.6. In Table 3.6, you must place check marks in the column
that best describes what happens during each change. You may check more than one column for each
change as needed. Pay close attention to whether or not a solid substance (precipitate) forms. It will
either lie on the bottom of the tube, or will remain suspended and make the liquid contents cloudy. Clear
solutions do not contain precipitates. (Do not confuse the terms clear and colorless. Clear liquids may or
may not be colorless.)
Chemical Change 1: Rinse a small test tube with distilled water. Add about 20 drops of AgNO 3 (solution
A) to the tube. Add about 10 drops of K2CrO4 solution (solution B) to the same tube. Record any
observations in Table 3.6 by placing check marks in the proper boxes.
Chemical Change 2: Add about 10 drops of Pb(NO3)2 (solution C) to a small clean, rinsed test tube. Add
about 20 drops of NaI (solution D) to the same tube. Shake the contents gently. Record observations.
Chemical Reaction 3: Add about 2 mL of HCl (solution E) to a clean 150 mm (medium) test tube. Place a
small piece of zinc metal into the tube. Record observations.
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© 1997, A.J. Girondi
Chemical Reaction 4: Add about 2 mL of H2SO 4 (solution F) to a 150 mm test tube. Note the temperature
of the solution by feeling the outside of the tube with your hand. Next, add about 4 mL of NaOH solution
(solution G). Caution: keep these solutions off skin and out of eyes! Feel the tube again. Record.
Table 3.6
Evidence of Chemical Changes
Color Change
Gas Given Off
Precipitate Forms
Temp. Change
(yes or no)
(yes or no)
(yes or no)
(yes or no)
Change 1:
Change 2:
Change 3:
Change 4:
In chemical change 4, an energy change was the only easily noticeable evidence of a chemical
change. However, energy changes by themselves are not enough evidence to indicate for sure that a
chemical change has occurred, because many physical changes are also accompanied by energy
changes. To observe a physical change in which there is an energy change, place about 2 mL of water in
a clean test tube. Note the general temperature of the water by pressing the tube on the inside of your
wrist. Using forceps, place 2 or 3 pellets of solid sodium hydroxide (NaOH) into the tube of water. Shake
the tube gently for 30 seconds or so, and press the bottom of the tube on the inside of your wrist. Do not
let the solution or the NaOH pellets contact you! Note the relative temperature of the contents of the
tube. NaOH does not react with water, it merely dissolves. Dissolving is a physical change. What change
in temperature did you observe? {25}____________________________________________________
Discard the contents of the tube into the sink and run the water. Rinse the tube with water.
Your experience from this activity should tell you that if you note a color change, a gas being
produced, a temperature change, or the formation of a precipitate, you should at least suspect that a
chemical change has occurred. This is particularly true if more than one of these events has occurred,
because color and temperature changes themselves are not conclusive evidence.
Identify each of the following changes as either physical (P)or chemical (C).
a. getting a flat tire
{26}___________
b. a campfire burning
{27}___________
c. seeing your “breath” when you exhale on a cold day
d. digesting food
{29}___________
e. rain changing to ice
{30}___________
f. drying the laundry
{31}___________
g. a bomb explosion
{32}___________
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{28}___________
© 1997, A.J. Girondi
SECTION 3.8
Density Problems
Earlier in this chapter, you worked with the concept of density as a physical property of matter. Do
the problems below. The units used to express density are usually g/mL or g/cm3. The g/mL unit is used
more frequently for liquids, while g/cm3 is used more often for solids. This is really just a matter of
preference, since 1 mL and 1 cm3 represent exactly the same volume. That is, 1 mL = 1 cm 3. Use
dimensional analysis wherever possible. Show all work neatly, even if only one or two steps are needed.
Problem 1. What is the density of gasoline if 350 mL has a mass of 245 grams?
Problem 2. A block of metal has the dimensions 5.0 cm X 7.0 cm X 20.0 cm, and its mass is 5.0 kg.
What is the density of the metal in g/mL?
Problem 3. What is the density of hydrogen gas (in g/mL) if 100. liters have a mass of 8.93 X 10-3 kg?
Problem 4. What is the density of milk in g/mL if 1.0 quart has a mass of 1.0 kg? (1.06 qt. = 1.00 L)
ACTIVITY 3.9
Determining the Thickness of Aluminum Foil
In this activity you will use what you have learned about density to calculate the thickness of
aluminum foil. Now you can imagine how tough this would be to do using only a ruler. Aluminum foil isn’t
very thick! We will accomplish this task in an indirect way. The formulas you will use are already familiar to
you. The volume of a regular object can be found using the formula V = L x W x H, where L = length, W =
width, and H = height. If the object is a piece of aluminum foil, we can alter the formula to V = L x W x T,
where T = thickness. Since this activity involves finding the thickness, we can solve the formula for T
which gives T = V/(L x W). From this formula you can see that in order to determine the thickness, you
need to know the length and width of the piece of foil, and you also need to know the volume of it.
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© 1997, A.J. Girondi
How do you calculate the volume of a piece of aluminum foil? That’s where density becomes
useful. The density (d) of aluminum is 2.70 g/1 cm3. If you measure the mass of the foil and if you know
the density of the metal, then you can use the formula at right below to find the volume:
Use this formula to get the volume of the foil
Since:
V = gX
cm3
= cm3
g
and since:
density =
g
cm3
then:
V =gX
1
d
Do you see how these two formulas can be used to derive this one?
Once you have the volume, you can substitute that value into the
formula shown at right. Measure the length and width (in cm)of the
rectangular piece of foil, and substitute those values along with the
volume into the formula. When you solve the formula for T, the units will
cancel and leave you with centimeters. You will have calculated the
thickness using the mass and density!
T =
V
L x W
Procedure: Cut or obtain 2 square pieces (about 10.0 cm x 10.0 cm) of aluminum foil - one regular and
one heavy duty. Weigh each piece on the laboratory balance, and measure the length and width of each
piece in centimeters. Enter the data into Table 3.7 below. Return the foil squares to the appropriate
containers. Complete the calculations, showing all work neatly in the spaces provided.
Table 3.7
Aluminum Foil Data
Regular or Heavy Duty
Length (cm)
Width (cm)
Mass (g)
_________________
_________
_________
_________
_________________
_________
_________
_________
SHOW your calculations in the spaces below. Express results to two decimal places.
Step 1. Calculation of the volume of the foil.
a. regular foil
b. heavy duty foil
Step 2. Calculation of the thickness of the foil.
a. regular foil
b. heavy duty foil
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© 1997, A.J. Girondi
Step 3. According to your calculations, determine how many times thicker the heavy duty foil is compared
to the regular foil. (Divide the thickness of the heavy duty foil by the thickness of the regular foil.)
Result: the heavy duty foil is _____________ times thicker.
SECTION 3.10
Accuracy, Precision, Percentage Error and Percentage
Deviation
Let’s use your data from Activity 3.9 to determine your accuracy and your precision. Accuracy
refers to the closeness of your result to the actual or accepted value. Precision refers to the closeness of
your results to each other (if you did an experiment more than once) or to that of other lab groups who did
the same experiment. It is possible to be precise without being accurate. In other words, you can get a
result that is close to what others got, while still being inaccurate. This can happen if your class is using
poor measuring instruments or a bad technique. However, in order to be accurate you must also be
precise. In other words, if everyone in your class is getting a result which is close to the accepted value
(accurate), then those results will also be close to each other (precise). Accuracy is often expressed as
percentage error, while precision is often expressed as percentage deviation.
Step 1. Calculation of Error. Obtain the accepted value for the thickness of the two kinds of aluminum foil
from your instructor. Calculate your error and percent error as shown below. Quantities enclosed in
vertical bars such as: |O - A| refer to absolute value.
Error = |O - A|
O = your observed value
% error =
A = the accepted value
|O - A|
X 100
A
Be careful when you use the formula above to calculate % error. It involves subtraction, division, and
multiplication. When using a calculator to solve such equations, you must do all your addition and
subtraction and then get a subtotal before doing the multiplication and/or division. If you follow the rules
for significant digits, you will need to round twice. List your error results below.
% error for the regular foil: __________%
% error for the heavy foil:
__________%
In this experiment, you should be able to achieve less than 5% error. How did you do? ______________
______________________________________________________________________________
Step 2. Calculation of Deviation. Get thickness results from three other lab groups in your class. Enter
the data into Table 3.8. Calculate the average (mean) thickness for each foil. To calculate your deviation
(precision) use the formulas below.
Deviation = |O - M|
O = your observed value
% deviation =
M = average (mean) value
|O - M|
X 100
M
Note that the calculation of deviation also involves an absolute value. Use the same precautions when
using a calculator as you did when calculating % error. List your deviation results below Table 3.8.
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© 1997, A.J. Girondi
Table 3.8
Lab Group Data
Group
Thickness of
Regular Foil (cm)
Thickness of
Heavy Foil (cm)
Your Group
____________
___________
Lab Group 1
____________
___________
Lab Group 2
____________
___________
Lab Group 3
____________
___________
Average (Mean)
____________
___________
% deviation for the regular foil: __________%
% deviation for the heavy foil:
__________%
Did you manage to achieve a small deviation (5% or less)? ___________________________________
If your group had a small error but a high deviation, that means that your small error was probably
just due to luck, or maybe off-setting errors. If your group had a large error (over 5%) but a small deviation,
that means that your high error was due to the measuring instruments or the procedure – not to human
error. If your group had a small error and a small deviation, that means that the instruments and your lab
techniques and skills were good (you were both accurate and precise)! If you got a large error and also a
large deviation, plan to pursue a career in something other than science, because the error was all your
fault! Comment on the performance of your group: ________________________________________
______________________________________________________________________________
______________________________________________________________________________
SECTION 3.11
Review Problems
Problem 5. A piece of nickel sheet metal has a width of 11.5 cm, a length of 4.66 cm, and a mass of
58.18 grams. What is the thickness of this sheet of nickel metal? (Use any needed data from Table 3.2.)
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© 1997, A.J. Girondi
Problem 6.
A student experimentally tried to determine the thickness of the metal mentioned in
problem 5. Her observed (O) result was 0.126 cm. The results from three other students were 0.130 cm,
0.114 cm, and 0.134 cm.
a. Calculate her % error. (Your answer to problem 5 can be used as the accepted (A) value.)
b. Calculate her % deviation. (You will need to calculate the mean, M, of her value and the three others.)
c. Based on her % error and % deviation, comment on these results assuming an acceptable value of 5%
for both error and deviation. _________________________________________________________
______________________________________________________________________________
______________________________________________________________________________
Problem 7. The “Miracle Thaw” is a heavily advertised sheet of “space age” metal on which you can
placed frozen foods for rapid defrosting (according to the manufacturer). The Miracle Thaw is 45 cm long,
25 cm wide, and 0.50 cm thick. It has a mass of about 1519 grams. What is the “space age” metal which
composes this “miracle?” Use data from Table 3.2. (Let the consumer beware.)
Problem 8. A cork cube weighs 500.0 grams and has a density of 0.025 g/cm3. Calculate the length of
one side of the cube in cm. Hint: the volume formula is V = s3 where s is the length of a side. You will
need to take a cube root in this problem. Use the appropriate key(s) on your calculator.
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© 1997, A.J. Girondi
ACTIVITY 3.12
The Barge Problem
--> --> -->
Teacher Demonstration
<-- <-- < --
With the background you have gained thus far in this chapter, explain what would happen in the
situation described below. You must use the concept of density in your explanation. Once you have
formulated your explanation, you will carry out an experiment to illustrate what actually happens. Take the
density of water as 1.00 g/mL , and keep in mind that , in order to float, a boat must displace a volume of
water equal to its own mass.
A barge is hauling metal beams through the Panama canal. While the barge is in a closed lock in
the canal, all of the metal falls overboard into the water in the lock. Do you think the water level in the lock
will rise, fall, or remain unchanged when the metal falls in? ___________________________________
Your teacher will demonstrate an experiment to test your hypothesis. A 100 mL beaker will be
used as the barge, a 400 mL beaker about half-full of water for the lock, and pieces of lead metal for the
beams. Enough metal will be used to make a difference, but not to sink the barge! A piece of tape or a
grease pencil will be used to mark the water level before and after the metal spills. Describe the result of
the experiment:_________________________________________________________________
Do the results support your original hypothesis?____________ Whether your answer is yes or no,
explain this phenomenon: _________________________________________________________
_____________________________________________________________________________
_____________________________________________________________________________
_____________________________________________________________________________
_____________________________________________________________________________
_____________________________________________________________________________
ACTIVITY 3.13
Sublimation
--> --> -->
Teacher Demonstration
<-- <-- < --
Because toxic fumes are involved, your teacher will conduct this demonstration under an exhaust
fume hood in your lab area. You are going to witness a rather unusual type of phase change. It involves
solid iodine crystals which change directly to a vapor without going through the liquid state. This is known
as sublimation. You may be familiar with dry ice (which is solid CO2) which also undergoes sublimation.
Care must be taken because iodine vapors are very dangerous and can cause serious stains and burns.
Look at Figure 3.4. Note that an ice cube is placed on top of a watch glass which is located on top
of a beaker or evaporating dish containing some iodine crystals. As heat is applied, the iodine crystals
slowly turn into a deep violet-colored vapor which is much heavier than air. When the vapor comes into
contact with the cool underside of the watch glass it turns directly from the vapor phase back to the solid
phase in a phase change known as deposition. Crystals may also form on the cool sides of the container.
Do not allow iodine to come into contact with you or your clothing! On the line below, write a phase
change equation for the sublimation of solid iodine. Remember to use the symbol for added heat.
{33}_______________________________________
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© 1997, A.J. Girondi
ice cube
watch glass
deposited crystals
iodine vapor
iodine crystals
Figure 3.4 The Sublimation of Iodine Crystals
SECTION 3.14
Learning Outcomes
Read through the learning outcomes below. Place a check mark in front of each outcome that you
have mastered. When complete, arrange to take the test on Chapter 3, and move on to Chapter 4.
_____1. Identify general physical properties of matter (color, phase, texture, luster, etc.)
_____2. Solve problems involving the density of a solid or a liquid.
_____3. Given proper information, determine the specific gravity of a liquid.
_____4. Label, interpret, and explain the parts of heating and cooling curves for pure substances.
_____5. Distinguish between physical changes, chemical changes, and phase changes.
_____6. Distinguish between accuracy and precision, and calculate percentage error and percentage
deviation given the necessary information.
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© 1997, A.J. Girondi
SECTION 3.15
Answers to Questions and Problems
Questions:
{1} qual; {2} no measurement involved; {3} water(l) ---∆--> water(g); {4} water(g) ------> water(l) ;
{5} temperature; {6} time; {7} area around D and up to E; {8} A-B, C-D, E-F; {9} remains constant
{10} A-B, C-D, E-F; {11} A-B, C-D, E-F; {12} B-C, D-E; {13} B-C, D-E; {14} about 53oC;
{15} Steam at 100 oC has the same amount of kinetic energy but has more potential energy than water at
100 oC since it has gone through an additional phase change.
{16} 0oC; {17} 100oC; {18} the same; {19} sodium, copper; {20} water, ethyl alcohol, acetone; {21} neon;
{22} no; {23} no change in phase - still a solid; {24} yes, because it is still water in each phase
{25} increase; {26} physical; {27} chemical; {28} physical; {29} chemical; {30} physical; {31} physical;
{32} chemical; {33} I2(s) ----∆----> I2(g)
Problems:
1. 0.70 g/mL
2. 7.1 g/mL
3. 8.93 X 10-5 g/mL
4. 1.1 g/mL
5. 0.122 cm
6. a. 3.28 %
b. 0 %
c. Since both error and deviation were under 5% in this type of activity, she has done well. The fact
that there was 0% deviation indicates that the 3.28% error was probably due to the equipment or
procedure.
7. aluminum
8. 27 cm.
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© 1997, A.J. Girondi
SECTION 3.16
Student Notes
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© 1997, A.J. Girondi