Laboratory Manual
Name______________________________
5.0 Chemistry
Teacher________________Period_______
Table of Contents
Experiments
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
21.
22.
23.
24.
25.
26.
27.
Page
Producing and Displaying Experimental Data
Density of Liquids Lab
Density of a Solid Lab
Physical and Chemical Changes Lab
Isotopes of Pennium Lab
Beanium Lab
Flame Test Lab
Distinguishing between Metals, Nonmetals, and Metalloids
Chemical Activity Lab
Models of Covalent Compounds Lab
Mole Lab
Hydrate Lab
Predicting Products Lab
Reactivity of Metals Lab (Single Replacement Lab)
Mole Relationship in a Chemical Reaction
Stoichiometry Lab #2
Determining the Percent Yield of Copper
Heat of Dissolving Lab
Specific Heat of a Metal Lab
Determining the Molar Mass of Butane
Solutions Lab
Penny Alchemy
Molarity and Dilutions
Determining Solution Concentration using a Spectrophotometer
A Study of Chemical Equilibrium
Acid-Base Properties Lab
Determining the Concentration of a Base
1
3
5
9
11
15
17
19
21
23
25
27
29
33
37
39
43
45
47
51
53
57
59
61
63
65
69
73
2
Producing and Displaying Experimental Data
Name ______________________________
20 Points
Date _____________ Period ___________
Objective: In this lab, you are to determine the effect of the concentration of dissolved sodium chloride
(NaCl) on the temperature of boiling water. Each time you run the experiment, record the data in the
table and then graph the temperature versus the time.
Materials:
250 mL graduated cylinder
distilled water
250 mL beaker
boiling chips
balance
sodium chloride (NaCl)
thermometer clamp
thermometer
Bunsen burner
beaker tongs
ring stand
striker
wire gauze
iron ring
Procedure:
1. Use your graduated cylinder to measure 150. mL of
distilled water. Be sure to read the volume at the bottom of the meniscus. Pour the water into the
beaker and add a boiling chip.
2. Set up a ring stand and burner as above. Be sure to clamp the thermometer to the ring stand.
Watch that the water level does not fall below the thermometer bulb during the experiment.
3. For Trial A, do NOT add any salt to the water. Begin heating the water with the burner and bring
it to a full boil.
4. As soon as the water comes to a full boil, record the temperature. This will be the temperature at
time = 0 minutes. Record the temperature after 1, 2, 3, 4, and 5 minutes. Be sure to measure the
temperature to the nearest 0.1oC. Record your results in the data table for Trial A.
5. Keep your boiling water and boiling chip for Trial B, but remeasure the water and add more
distilled water if there is less than 150. mL. Repeat steps 2 through 5 again, but this time measure
10 g of sodium chloride, NaCl, with the balance and add it to the water before heating it. Record
your data in the table for Trial B.
6. If time permits, discard the salt water from Trial B, retain the boiling chip and measure 150. mL
of new distilled water. Repeat steps 1 through 5 again adding 20 g of sodium chloride, NaCl, to
the water before heating it. Record your data in the table for Trial C.
Data Table:
Trial A
Time (min.) Temp.˚C
0
1
2
3
4
5
Trial B
Time (min.) Temp.˚C
0
1
2
3
4
5
Trial C
Time (min.) Temp. ˚C
0
1
2
3
4
5
3
Analysis and Conclusions:
1. What is the variable being tested in this experiment? _______________________________
2. The control run was completed during Trial A or Trial B or Trial C. (circle one)
3. You controlled the ________________ (time or temperature) during the experiment. Therefore,
______________ would be considered the independent variable and _______________ would be
considered the dependent variable.
4. In a graph, the independent variable is displayed on the ____ axis and the dependent variable is
displayed on the ____ axis.
5. Make a graph of all 3 trials using the time and temperature data. Use 3 different colors.
Remember to label each axis with a measurement and unit. Also, include a key and title on your
graph. The graph should be a smooth curve. Only include temperature values from 90oC+.
0
1
2
3
4
5
Based on YOUR tables and graph, write a general statement describing the relationship between
boiling temperature and the amount of sodium chloride dissolved.
_________________________________________________________________________________
_________________________________________________________________________________
4
Density of Liquids Lab
Name _____________________________
Date _______________ Period ________
Introduction:
Have you ever wondered why the oil in bottles salad dressing settles on top of the water –and-vinegar
mixture? The answer has to do with the different densities of the liquids in the dressing. Oil has a lower
density than the water-and-vinegar mixture. When two liquids of different densities are mixed, the liquid
that is less dense, in this case the oil, floats above the other liquid.
Density is the mass of a substance per unit volume. The density of any substance is a ratio and may be
calculated by dividing the mass of the sample by its volume.
mass
density =
volume
The most common units of measurement for density that you will encounter in this course are g/mL and
g/cm3.
When measured at the same temperature and pressure, all samples of a particular substance have the
same density regardless of the quantity or shape of the sample. Thus, density is a characteristic property of
matter that is often used by chemists to identify a substance.
In this investigation, you will first determine the density of distilled water by finding the density of
three separate samples and calculating an average. You will then repeat this procedure for an unknown
liquid. Finally, you will compare your measured values with the accepted values for the densities of these
liquids and compute the percent error for your results.
Prelab Questions:
1. Define density. What is the formula for density?
2. How can density help you identify an unknown substance?
3. Why is the precision of the laboratory balance that you use in this investigation important? What
effect would a less precise balance have on your results?
4. What is the advantage of taking three sets of measurements instead of just one or two when
determining the density of each liquid?
5. Why should you avoid skin contact with the unknown solution?
6. Do you expect your measured values for the density of water and the unknown to differ from the
accepted values? Explain your answer.
5
Materials:
goggles
graduated cylinder
balance
unknown liquid
apron
water
Procedure: Density of Water
1) Mass a 10-mL graduated cylinder to the correct number of places and record data in Table 1.
2) Fill the graduated cylinder with 9-10 mL of distilled water and mass it again. Record in Table 1.
3) Determine the volume of the water using the graduated cylinder. Remember to read at the bottom of
the meniscus and to the correct number of places. Record data in Table 1.
4) Repeat steps #1-3 with 2 lesser volumes of water.
5) Empty the graduated cylinder, dry it and go on to step #6.
Density of Unknown
6) Write the mass of the 10-mL graduated cylinder in Table 2.
7) Repeat steps #2-5 with your unknown liquid. Remember to record the liquid # in Table 2.
Data:
TABLE 1: Water
Mass of graduated cylinder and water (g)
Trial 1
Trial 2
Trial 3
Mass of empty graduated cylinder (g)
Mass of water (g)
Volume of water (mL)
Density of water (g/mL)
TABLE 2: Unknown # ________
Mass of graduated cylinder and unknown (g)
Trial 1
Trial 2
Trial 3
Mass of empty graduated cylinder (g)
Mass of unknown (g)
Volume of unknown (mL)
Density of unknown (g/mL)
6
Calculations:
1) Calculate the mass, volume and density of the water and unknown and write in the data table.
2) Determine the average density of water:
3) Determine the average density of the unknown:
Critical Thinking Questions:
1. Based on your data, which of the liquids in the table below could be your unknown?
________________________________________
TABLE 3: Accepted Values for Density of Various Liquids
DENSITY (g/mL) at 20 C
0.789
0.786
0.791
0.830
0.998
1.114
1.261
LIQUID
Ethanol
Isopropyl alcohol
Methanol
Mineral oil
Water
Ethylene glycol
Glycerine
2. Using the formula below, calculate the percent error for the average density of water and your
unknown liquid.
percent error 
measured value - accepted value
x 100%
accepted value
Percent error for water =
7
Percent error for unknown =
3. What changes could you make in the procedure to increase the accuracy of your results? The
precision?
4. If your unknown had a density of 0.79 g/mL and you knew that it was one of the three substances
listed in the Table 3 having a density close to your value, how would you go about determining
which liquid was your unknown?
5. Ice floats on water. Is ice more or less dense than water? ________________How do you think
the density of ice affects the survival of water-dwelling organisms in environments where
temperatures fall below the freezing point?
6. When stating density of a liquid, why is it necessary to state the temperature?
8
Density of a Solid Lab
Name________________________
Chemistry I 5.0
Period______Date______________
Purpose: To determine the density of a solid.
Procedure:
1. Mass an empty 100 mL beaker. Record the mass in the data table.
2. Choose 1 metal to analyze. Add a small amount of the metal to your beaker and determine its mass.
3. Fill half of your graduated cylinder up with water. Note the initial volume of the water.
4. Add your metal to the water. Note the final volume of the water.
5. Return the wet metal to the front lab bench for drying.
6. Repeat the procedure 4 times with a different amount of the metal each time. Only use dry metal in
this lab.
Note: Strive for a 1 mL or larger difference between all volume measurements.
.
Data: Name of metal:_______________________________
Trial 1
Trial 2
Trial 3
Trial 4
Trial 5
Mass of
empty beaker
Mass of
beaker and
metal
Mass of
metal
Initial Volume
Final Volume
Volume of
metal
Graph:
Prepare a mass vs. volume line graph using your data from the table above. Do not connect the
plotted points. Rather, draw a line of best fit.
9
Calculations
1. Determine the slope of the line. This is the density of your metal. Show all work and include
units.
Remember:
Y2 – Y1
X2 - X1
2. Determine the percent error for your experiment. Use the slope for your density!
% Error 
Experimental  Accepted
x100
Accepted
3. List 2 possible sources of error in this experiment. (Be specific)
10
Physical and Chemical Changes Lab
25 Points
Name ____________________________
Period _________ Date ____________
Introduction: As you read this it is probably fall. Summer flowers are fading and dying. Leaves are
changing from green to red, yellow and orange. All these changes involve chemistry.
As you have learned, chemistry is the study of matter and the changes that it undergoes. These changes can
be classified as physical, chemical, or nuclear. This lab will focus on physical and chemical changes. When
a physical change occurs, the physical properties of a substance – such as its size, shape, state, or density –
are altered, but its chemical composition remains the same. Examples of physical changes include melting
ice, crushing gravel, tearing paper, grinding pepper, and boiling water. No new substances are formed as a
result of these changes.
Chemical changes, also known as chemical reactions, result in the formation of one or more new substances
with different chemical properties and compositions from the original material. Examples of chemical
changes include plants dying, leaves changing color, paper burning, bananas ripening, bread baking, or iron
rusting. Some signs of chemical changes include a change in color, the formation of a precipitate (a new
solid substance that settles out of solution), the production and release of a gas, or a change in temperature.
In this investigation, you will conduct tests on several substances and then use your data to determine
whether the resulting changes were chemical or physical. As you observe each change, remember to ask
yourself, “Has the change altered the identity of the substance?”
Pre-Lab Questions:
1. What are the 5 observable signs that indicate a chemical change is taking place?
_________________________________________
___________________________________________
_________________________________________
___________________________________________
_________________________________________
1) Define precipitate as it relates to this lab. ___________________________________________________
____________________________________________________________________________________
2) Identify the following as either a chemical (C) or physical (P) change:
a. burning wood
___________
d. ripening fruit
___________
b. dry ice changing to a gas
___________
e. dissolving sugar in water
___________
c. freezing water
___________
f. breaking a pencil
___________
3) Why are you instructed to feel outside the test tube after two chemicals are mixed, as in Step 5 and Step 6
of this investigation?___________________________________________________________________
4) What are some safety precautions you need to observe during this investigation?____________________
______________________________________________________________________________________
11
Problem: How can you recognize and differentiate between physical and chemical changes in matter?
Procedure: All solids must be thrown away in metal trash can! Each step is a separate
step! Wear SAFETY GOGGLES and LAB APRON at all times!!!
1. Put on your apron and goggles. Break off a small piece of wax from the bottom of a birthday
candle and place it in a DISPOSABLE test tube. Holding the test tube with tongs, heat it gently
over the burner until the wax completely melts. Place the test tube in a beaker until it cools and
then throw test tube in broken glass container. Record results in Data Table.
2. With matches, light larger candle. Secure candle to the glass square by dripping wax onto the
square and then holding the base of the candle in to the molten wax until the wax hardens. Allow
the candle to burn for two minutes and blow out. Use scoopula to scrape wax off of glass plate.
Record results in Data Table.
3. Break off a small piece of paper and place it in a watch glass. Place watch glass on wire guaze
and ignite paper with matches. Record results in Data Table.
4. Measure 5mL of distilled water in a graduated cylinder. Pour water into a test tube a add a small
amount of sodium chloride (1 small scoop with scoopula). Stir the contents to mix. Using eye
dropper, add 10 drops of silver nitrate (0.1 M AgNO3) to the NaCl-water mixture. Record results
in Data Table. Dispose in labeled waste beaker. NOTE: SILVER NITRATE IS TOXIC AND CAN
STAIN YOUR SKIN. WASH ANY SILVER NITRATE OFF YOUR SKIN IMMEDIATELY.
5. Obtain a small piece of Magnesium (Mg) ribbon and place it in a test tube. Add a few drops of
hydrochloric acid (HCl) to the test tube. Record results in Data Table. Add enough HCl until Mg
completely reacts with HCl. NOTE: HYDROCHLORIC ACID IS HIGHLY CORROSIVE AND
CAUSES BURNS. WASH ANY HYDROCHLORIC ACID OFF YOUR SKIN IMMEDIATELY.
6. Grind 1-2 crystals of copper sulfate pentahydrate (CuSO4  5H2O) with the mortar and pestle.
Place a microspatula (small amount) of the powder into a test tube. Heat gently over a bunsen
burner for 2 minutes. Allow to cool for 5 minutes; then add a few drops of water. Touch the
bottom of your test tube with your fingertip. Record your observations in the Data Table.
7. Using two pieces of weighing paper, place a SMALL amount of sulfur in one holder and a
SMALL amount of iron filings in the other holder. Run the magnet along the bottom of the
paper. Then, mix the two substances into one of the holders and run the magnet along the side of
the holder. Record results in Data Table. Use mixture for step 8!
8. Heat the iron-sulfur sample under the fume hood for several minutes until the mixture glows.
CAUTION: Do not carry out this step unless a working hood is available. Noxious fumes are
produced that must be evacuated by a fume hood. All the sample to cool for 10 minutes. Notice
its appearance. Test with a magnet. Record results in Data Table. Throw away in broken glass
container.
9. Clean up your work area and wash your hands before leaving the laboratory.
12
Observations: Data Table
STEP
BEFORE
1
Observations
DURING
AFTER
2
3
4
5
6
7
8
Critical Thinking: Analysis and Conclusions
1. Indicate whether the following changes are physical or chemical. Explain your answer.
Physical or
Change
Explanation
Chemical
Melting candle wax
Burning candle wax
Tearing paper
Burning paper
Dissolving NaCl
Mixing NaCl and AgNO3
Adding HCl to Mg
Grinding CuSO4 · 5H2O
Heating CuSO4 · 5H2O
Mixing Fe and S
Heating Fe and S
13
2. Name two possible indications that a chemical change has occurred, using examples from this
investigation.
____________________________________________________________________________________
____________________________________________________________________________________
3. A change in color does not always indicate chemical change. Explain why it could be the result of a
physical change.
_______________________________________________________________________________
_______________________________________________________________________________
4. Answer the following questions using examples from this investigation to support your answers.
a. How can substances in a mixture be separated?
______________________________________________________________________________
______________________________________________________________________________
b. How can substances in a compound be separated?
______________________________________________________________________________
______________________________________________________________________________
5. Sodium chloride dissolves in water, leaving a clear homogenous mixture with no physical evidence of
the crystals with which you started. Design an experiment that you could perform to separate the sodium
chloride from the water.
__________________________________________________________________________________
__________________________________________________________________________________
__________________________________________________________________________________
6. How could the experiment you designed in response to Question #4 differ if some sand was mixed in
with the sodium chloride before it was added to water?
__________________________________________________________________________________
__________________________________________________________________________________
__________________________________________________________________________________
14
Isotopes of Pennium Lab
Name ___________________________
Date _______________ Period ______
Introduction:
In this investigation, you will determine the relative abundance of the isotopes of pennium and the
masses of each isotope. You will then use this information to determine the atomic mass of pennium. Recall
that the atomic mass of an element is the weighted average of the masses of the isotopes of the element. The
average is based on both the mass and the relative abundance of each isotope as it occurs in nature.
Pre-Lab Questions
1. What do the 20 pennies in this investigation represent? _______________________________
2. What do the different masses of the pennies represent? ________________________________
3. What information is needed to calculate average atomic mass?
____________________________________________________________________________
Materials: Balance
20 pennies
Procedure
1. Line the 20 pennies up in chronological order.
2. Mass each penny individually. Record its year and mass in the data table.
3. Draw a line(s) through the table indicating where the mass of the penny changed 0.1 g or greater
from the previous penny.
Data Table
Penny
Mass of 20 pennies = ________ g
Year
Mass (g)
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
19
20
15
Calculations:
1. Determine the number of isotopes of Pe that are present. _____________________
2. Calculate the percent abundance of each isotope of your sample.
3. Calculate the average atomic mass of each isotope.
4. Using percent abundance and the average atomic mass of each isotope, calculate the atomic mass
of Pe.
Analysis and Conclusions:
1. Was the mass of 20 pennies equal to the mass of 20 times one penny? Explain.
______________________________________________________________________________
__________________________________________________________________
2. In what year(s) did the mass of Pe change? How could you tell?
______________________________________________________________________________
____________________________________________________________________
3. How can you explain the fact that there are different “isotopes” of pennium?
______________________________________________________________________________
____________________________________________________________________
4. Why are the atomic masses for most elements not whole numbers?
______________________________________________________________________________
____________________________________________________________________
5. How are the three isotopes of hydrogen (hydrogen-1, hydrogen-2, hydrogen-3) alike? How are
they different?
______________________________________________________________________________
____________________________________________________________________
6. Copper has two isotopes, copper-63 and copper-65. The relative abundance of copper-63 is
69.1% and copper-65, 30.9%. Calculate the average atomic mass of copper.
16
“Beanium” Lab
Name ______________________________
Period _______ Date__________________
Introduction & Purpose:
What is an isotope? What does it mean to say that the atoms in a sample of an element are isotopes of each
other? Ordinary beans are a lot bigger than atoms, but perhaps they can give you one or two clues about
isotopes.
Isotopes: atoms of the same element that differ in mass. For example, there are actually three different kinds
of hydrogen atoms. See below for the different kinds (or isotopes) of two common atoms.
ISOTOPE
Protium
Deuterium
Tritium
Carbon-12
Carbon-14
PROTONS
1
1
1
6
6
ELECTRONS
1
1
1
6
6
NEUTRONS
0
1
2
6
8
SIGNIFICANCE
“normal” hydrogen
“heavy hydrogen”
radioactive hydrogen
“normal” carbon
carbon-14 dating
This lab is designed to show you isotopes of an element in a simulation form. You will be asked to gather
data about the “isotopes” and organize the data. If atoms were as large as beans they could be sorted,
counted, and massed. In this experiment we will sort, count, and mass three different kinds of beans and
imagine that we are observing three different isotopes of the same element (let’s call it BEANIUM). The
three different isotopes are blackium, brownium, and whitium. Finally we will calculate the isotopic mass,
the isotopic abundance, and the atomic mass of the “Beanium” element.
Definitions: • Isotopic mass – the average mass of the atoms of a specific isotope of an element
• Isotopic abundance – what percent of the element’s atoms are a specific isotope
• Atomic mass – the average mass of an element’s atoms
Materials: Plastic cup of beans (black, brown, & white); balance
Procedure:
1. Obtain a plastic cup which contains many atoms (beans) of BEANIUM.
2. Sort the atoms (beans) into a group for each isotope: black, brown, white. Record the total number of
atoms and the number of each type of isotope (blackium, brownium, and whitium) in the table.
3. Determine the isotopic mass:
a) Find the total mass of each of the three isotope groups and record on data table.
b) Find the average mass of a single atom of each isotope and record to the nearest 0.001g.
EXAMPLE: I counted 340 white beans. They have a mass of 80 grams. The average mass of
white bean is 80 / 340 = 0.235 grams.
one
4. Find the isotopic abundance (% of beans) for each isotope by dividing the number of atoms of one
isotope by the total number of atoms (black, brown, plus white) and multiplying by 100%. Record on
the data table to the nearest 0.1%.
EXAMPLE: There are a total of 500 beans. 340 are white beans. Therefore, (340 ÷ 500) x 100% =
68.0% are white beans (whitium).
17
5. Determine the atomic mass for BEANIUM based on the isotopic abundances and the isotopic
masses.
FORMULA TO CALCULATE ATOMIC MASS = (blackium %) x (mass of one blackium atom) +
(brownium %) x (mass of one brownium atom) + (whitium %) x (mass of one whitium atom).
6. Place all the beans back in the plastic cup and return to the front lab bench.
Data: Show one sample (Black Beans) for each calculation. Remember significant digits in calculations!
Average Mass of Black Beans:
Percent Abundance of Black Beans:
DATA TABLE
Isotope information: Total number of bean element atoms in cup = _________
# OF BEANS
(ATOMS)
ISOTOPE
MASS OF BEANS
(ATOMS)
AVERAGE MASS OF 1
BEAN (GRAMS)
% OF
BEANS
Blackium
Brownium
Whitium
Calculations for atomic mass of BEANIUM…(use the data table to complete)
% of blackium x avg mass
“atoms”
of blackium
+
x
+
g
g
+
% of brownium x avg mass
“atoms”
of brownium
x
g
g
+
+
% of whitium x avg mass
“atoms”
of whitium
x
+
=
g
g
Atomic Mass
of BEANIUM
=
=
g
Conclusion Questions
1. Was the average mass of the beans a whole number or a decimal? Explain.
2. What is the relationship between an element’s isotopes and the element’s atomic mass?
3. Explain how this lab simulates the various isotopes of an element.
4. Copper has two isotopes, copper-63 and copper-65. The relative abundance of copper-63 is 69.1%
and copper-65, 30.9%. Calculate the average atomic mass of copper.
18
Flame Test Lab
Name __________________________________
Date ______________ Period _______________
Introduction:
As electrons absorb energy, they jump to a higher energy level called an excited state. Electrons in
excited states are not stable and fall to lower energy levels in order to become more stable. When they fall,
they release energy in the form of electromagnetic radiation. Visible light, the only type of electromagnetic
radiation that is detectable to the human eye, is emitted when electrons in an element are excited. The light
emitted from that element becomes a fingerprint for that element because it does not include every color in
the spectrum.
The specific spectrum for an element is called a line spectrum because it is made up of only certain
colors. When using a spectroscope, a researcher can differentiate the colors emitted when that element is
excited. However, with the naked eye, only a blend of the colors is noticed. In this lab, we will be
completing flame tests to determine the characteristic spectral color of certain elements. Even though there
are many other colors in an element’s line spectrum, the color noticed during the flame test is a blend of the
colors.
Materials:
Goggles
Cotton Swabs
Bunsen Burner
BaCl2
CuCl2
SrCl2
KNO3
Ca(NO3)2
CaCl2
Apron
Crucible tongs
Ba(NO3)2
Cu(NO3)2
Sr(NO3)2
LiNO3
KCl
NaCl
NaNO3
Procedure:
1. Put on goggles and apron.
2. This lab will be performed moving from station to station. There is no set order that you will
have to go to, so each group can go to a different location. You will only need the crucible tongs
from your lab drawer. Each station has one or two salt solutions in labeled beakers, a Bunsen
burner, and a supply of cotton swabs.
3. Make sure the burner is lit and has a blue or almost clear flame. With your tongs, pick up the end
of a swab and soak the cotton end in the solution. Hold the soaked-end of the swab in the clear
part of the flame and record the flame color in the Data Table. Eventually the swab will burn
giving a characteristic yellow color. Be sure you are recording the solution’s flame color. Placed
the used swab in the waste beaker.
4. When all the groups have had the opportunity to observe all the different flame colors, the class
will be instructed to return equipment and clean up the lab area. Wash your hands before leaving
the lab.
19
Data Table
Salt Solution
Color
Salt Solution
Ba(NO3)2
KNO3
BaCl2
KCl
Cu(NO3)2
NaNO3
CuCl2
NaCl
Sr(NO3)2
Ca(NO3)2
SrCl2
CaCl2
Color
LiNO3
Critical Thinking and Application:
1. Summarize the process that produces the colors seen in the flame test.
_________________________________________________________________________________
_________________________________________________________________________________
2. What is the relationship between the colors you saw and the lines of the electromagnetic spectrum
produced by the metals?
_________________________________________________________________________________
_________________________________________________________________________________
3. According to your data, determine the characteristic color produced when the following metals are
excited:
a. Lithium _____________________
e. Sodium ______________________
b. Strontium ____________________
f. Barium ______________________
c. Calcium _____________________
g. Copper ______________________
d. Potassium____________________
4. When a glass rod is heated, a yellow flame is observed around the point of heating. What does this
yellow flame indicate? Why is it observed when glass is heated?
_________________________________________________________________________________
_________________________________________________________________________________
5. How do you think metallic salts are used in fireworks?
_________________________________________________________________________________
_________________________________________________________________________________
6. Based on your results and observations, would this method be practical to determine metals in a
mixture. Explain. _________________________________________________________________
_________________________________________________________________________________
20
Distinguishing Between Metals, Nonmetals & Metalloids
Name ___________________________
Date ______________ Period _________
Introduction:
In this activity you will investigate properties of seven elements in order to classify them as metals,
nonmetals, or metalloids. You will examine each for its physical properties of color, luster, and form (for
example, is it crystalline, like table salt?) By attempting to crush each sample, you will decide whether each
element is malleable or brittle. You will also test for the physical property of electrical conductivity.
Next, you will observe differences among these elements through chemical properties. You will find out
whether each element reacts with hydrochloric acid, HCl(aq), and with a copper(II) chloride (CuCl2) solution.
Procedure
1. Appearance: Observe and record the appearance of each element. Include physical properties such as
color, luster, and also the form.
2. Crushing: A material is malleable if it flattens without shattering when struck. A sample is brittle if it
shatters into pieces. GENTLY push on an element sample with a scoopula. Decide
whether the samples are malleable or brittle.
3. Conductivity: If a conductivity apparatus is available, use it to test each sample. Touch both electrodes
to the element sample. If the bulb lights, the sample has allowed electricity to flow
through it. Such a material is called a conductor. If the bulb fails to light, the material is a nonconductor.
Data Table
Element
Appearance
Result of
Crushing
Conductivity
Reaction
with HClaq
Reaction
with CuCl2aq
Classification
A.
B.
C.
D.
E.
F.
G.
21
CHEMICAL PROPERTIES
1. Test each sample for reactivity with acid as described below. The formation of gas bubbles indicates
that a chemical reaction has taken place.
a. Obtain a well-plate and label seven wells a thru g.
b. Place a small sample of each element in their respective wells. (The samples should fit nicely in
the center of the well B about 1/3 full.)
c. Add several drops of HCl to each sample (Enough to just about fill the well.)
d. Observe the samples for five minutes and record each result.
e. Discard the well contents into the METAL TRASHCAN. Rinse the well-plate out by gently
squirting it down with some distilled water.
2. Test each element sample for reactivity with copper (II) chloride (CuCl2) solution as described below.
Changes in the appearance of any element sample indicates that a chemical reaction has taken place.
a. Repeat steps 1a and 1b.
b. Add several drops of copper (II) chloride (CuCl2) solution to each sample (Enough to just about
fill the well.)
c. Observe the samples for five minutes and record each result.
d. Discard the well contents into the METAL TRASHCAN. Rinse the well-plate out by gently
squirting it down with some distilled water.
3. Wash your hands thoroughly before leaving the laboratory.
CONCLUSIONS - Directions: Answer the following questions on a separate piece of paper.
1. Sort the eight coded elements by letter into two groups, based on similarities in their physical and
chemical properties. Put the letters of the elements in their respective groups. If an element has
properties that could place it in both groups, then put its letter in both groups.
2. Consider the following information. Then reclassify each element as a metal, a nonmetal, or a
metalloid.

ct electricity. (These are all
physical properties.)

chemical properties.)

physical properties.)

chemical properties.)

metalloids, or
semimetals.
22
Chemical Activity Lab
Name ______________________________
Date _________________ Period _______
Introduction:
Have you ever noticed that some people are more active than others? Some people play a variety of
physical sports while others like to sit and read. You can think of elements in much the same way. In the
case of an element, its activity level is a measure of its ability to react chemically with other elements.
Many metals will react with ions of another metal in solution. You can tell that a reaction has
occurred because the metallic ions that come out of solution form a solid precipitate with that metal. At the
same time, atoms of the more active metal go into solution. Chemists use the degree of activity to predict
what changes will occur in certain reactions. For example, a more active metal will always replace (or react
with) a less active metal in a compound. This is called a single replacement reaction.
In this experiment, you will use three metals and four solutions of compounds that contain different
kinds of metallic ions. You will put each metal into a separate sample of each solution and observe what
happens. If a reaction occurs, you will notice a solid precipitate forming on the metal. If a particular metal
reacts with ions of many other metals, then that metal is a chemically active metal. If a metal reacts with few
or none of the other metals, then it is chemically inactive. From your observations in this experiment, you
will be able to arrange the four metals in the order of their chemical activity.
Pre-Lab Discussion
1. What does the term chemical activity mean?
2. What evidence of chemical activity will you be looking for in this investigation?
Purpose: To compare the chemical activity of four metals.
Materials:
goggles
copper (II) nitrate solution
magnesium nitrate solution
zinc nitrate solution
silver nitrate solution
apron
well plate
copper pieces
magnesium pieces
zinc pieces
Procedure:
1. Put on your goggles andd lab apron. Using your marking pen, label the wells in the well plate
from left to right along the top row 1, 2, 3, and 4. Label the rows down the left side A, B, and C.
2. Place 4 pieces of copper, 4 pieces of Magnesium and 4 pieces of Zn in rows A, B, and C,
respectively.
3. At your lab station, place 1 drop of copper (II) nitrate in each well in column 1. Repeat this step
for the other solutions placing them in the following columns: magnesium nitrate in column 2,
zinc nitrate in column 3, and silver nitrate in column 4. Note: Silver nitrate is poisonous if
ingested. Be careful not to get it on your skin or clothing, as it will produce a stain that is hard to
remove. If any spill occurs, notify your teacher.
23
4. Record observations in Data Table A.
5. Do NOT dispose of the metals down the drain. Place them in a paper towel and put them in the
trash can. Wash off the well plate with water. Clean up your work area and wash your hands
before leaving the laboratory.
Data Table
A
Cu
B
Mg
C
Zn
D
Ag
1
Cu(NO3)2
2
Mg(NO3)2
3
Zn(NO3)2
4
Ag(NO3)
NR
NR
NR
NR
Critical Thinking Questions
1. What similarities and differences in physical properties (e.g., hardness, color, shine) did you see
when you looked at the metals?
2. Which of the FOUR metals reacted with the greatest number of solutions?
3. Which of the FOUR metals reacted with the least number of solutions?
4. List the metals from the most active to the least active.
5. The Statue of Liberty is made of copper. Use your investigation results to explain why copper is
better material for a statue than magnesium or zinc.
6. Gold does not react with any of the solutions in this investigation. What does it tell you about gold’s
activity?
7. How does the chemical activity of gold account for its use in jewelry?
8. Lead is less active than zinc but more active than copper. Predict the results if lead metal is put into
separate solutions of zinc nitrate and copper (II) nitrate.
24
Models of Covalent Compounds Lab
Name ______________________________
Date ________________ Period _________
Pre-Lab Discussion
1. What is a covalent bond? ______________________________________________________
________________________________________________________________________
2. List some of the guidelines for completing this lab. _________________________________
________________________________________________________________________
Data Table:
Formula
H2
Electron Dot Structure
Structural Formula
Initials
HBr
H2O
PH3
CH4
HClO
N2
25
Formula
CH3NH2
Electron Dot Structure
Structural Formula
Initials
CO2
H2CO3
Hint: H
bonds to O
C2H2
CH3Cl
HCOOH
HCN
H2O2
26
Mole Lab
Name ______________________________
Date ______________ Period __________
Purpose: In this lab, you will make a series of mass measurements. You will then
convert these measurements to moles and molecules.
Procedure: Make the following measurements:
Mass of Nickel
Mass of a nickel
___________ g
Mass of Swallow of Water
Mass of cup and water (before swallow)
___________ g
Mass of cup and water (after swallow)
___________ g
Mass of water swallowed
___________ g
Mass of Sugar
Mass of sugar in sugar packet
___________ g
Mass of Signature
Mass of Chalk (before writing your name 3 times) ___________ g
Mass of Chalk (after writing your name 3 times)
___________ g
Mass of chalk used in 3 signatures
___________ g
Mass of chalk used in 1 signature
___________ g
Analysis: (Show all your work, and record your final answer with the correct number of significant
digits and unit.)
1. A nickel is composed of 25.0 % nickel and 75.0 % copper.
a. Calculate the mass of nickel in the coin.
_________________
b. Calculate the number of moles of nickel in the coin. _________________
c. Calculate the number and kind of particles of nickel in the coin. _________________
27
d. Calculate the mass of copper in the coin.
_________________
e. Calculate the number of moles of copper in the coin.
_________________
f. Calculate the number and kind of particles of copper in the coin.
_________________
2. a. Calculate the number of moles of water swallowed.
b. Calculate the number and kind of particles of water swallowed.
3. Table sugar is sucrose, C12H22O11.
a. Calculate the number of moles of sugar in the packet.
__________________
__________________
__________________
b. Calculate the number and kind of particles of sugar in the packet. _________________
c. Calculate the number and kind of particles of carbon in the packet.__________________
4. Assume your chalk was 100% calcium carbonate.
a. Calculate the number of moles of calcium carbonate used in your signature.
_________________
b. Calculate the number and kind of particles of calcium carbonate used in your signature.
_________________
c. Calculate the number of kind of particles of oxygen in your signature.
__________________
28
Hydrate Lab
Name______________________________
Date ________________ Period ________
Introduction:
Many compounds are formed as a result of reactions that occur in water solutions. These compounds
appear to be dry, but when they are heated, large amounts of water are released. The water molecules are
part of a crystalline structure and are weakly bonded to the ions or molecules that make up a compound.
Such compounds are called hydrates, meaning that they contain water. The solid that remains when the
water is removed is referred to as the anhydrous salt, or anhydrate.
Hydrate + Heat
→
Anhydrate + H20
Usually the amount of water present in the a hydrate is in a whole number ratio to the moles of the
anhydrate. An example of a hydrate is magnesium sulfate. Its formula is
MgSO4 ∙ 7H2O,
indicating that seven moles of water are combined with one mole of magnesium sulfate in the crystalline
form.
In this investigation you will be given an unknown hydrate and asked to determine the percent of
water in the compound. From this information, the molar ratio of water to anhydrous salt will be calculated.
Finally, the identity of the hydrate will be determined.
Pre-Lab Discussion:
1. What is a hydrate? An anhydrate?
_________________________________________________________________________________
_________________________________________________________________________________
_____________________________________________________
2. Why do you think it is necessary to heat the evaporating dish before finding its mass?
________________________________________________________________________
3. Why must the evaporating dish be cooled before finding the mass?
_________________________________________________________________________________
_______________________________________________________________
4. Why must the mass of the anhydrous salt be measured immediately on cooling?
_________________________________________________________________________________
_______________________________________________________________
5. Why is it necessary to handle the evaporating dish only with the crucible tongs immediately after
heating? ________________________________________________________________________
Materials:
goggles and apron
ring stand
lab burner
crucible
balance
unknown hydrate
clay triangle
iron ring
striker
crucible tongs
spatula
29
Procedure:
1. Set up Bunsen burner, clay triangle and ring stand. Gently heat crucible for three minutes. Then,
allow crucible to cool for three minutes.
2. Measure and record the mass of the crucible. Record in data table.
3. Add about 2 g of your unknown hydrate to the crucible. Measure and record the mass of the crucible
and hydrate.
4. Gently heat the unknown hydrate in the crucible until there is no more popping or spattering.
5. Examine the material in the crucible. If the edges of the solid appear to be turning brown, reduce the
heat momentarily and then begin heating again at a slower rate. Heat for 5 more minutes.
6. Remove the crucible and allow it to cool for at least one minute. Immediately measure and record the
mass of the crucible and anhydrous salt.
7. Reheat crucible and contents for a few minutes, cool and measure mass again. The value should be
within 0.02 g of the last recorded mass. If it is not, reheat and remeasure the mass until the last two
measurements are within that range. Record the final mass.
8. Dispose of your substance and clean up your work area.
Observations:
Data Table
Mass of crucible
Mass of crucible and hydrate
Mass of crucible and anhydrate (after step 6)
Mass of crucible and anhydrate (after step 7)
Calculations: (Show all work and units!)
1. Calculate Mass of Hydrate
2. Calculate mass of H2O
3. Calculate moles of H2O
4. Calculate mass of anhydrous salt
30
5. Calculate moles of anhydrous salt
6. Determine the smallest whole number ratio between moles of water and anhydrous salt
7. What is the formula of the sample based on the empirical data?
8. What is the % of H2O according to the experimental data?
9. What is the actual formula for your hydrate? _______________________ (ask teacher)
10. What is the % of H2O in the actual hydrate?
11. Calculate your percent error.
31
Questions:
1. Explain the effect the following errors would have on the value for the percent water in a hydrate.
a. The hydrate is not heated long enough to drive off all the water.
b. A damp crucible was used, and it was not dried before adding the hydrate.
c. The crucible and contents were allowed to cool overnight before finding their mass.
2. Predict what would happen if you added a few drops of distilled water to the anhydrous salt
remaining at the end of this experiment.
32
Predicting Products Lab
Name:
Chemistry 5.0 - Equations
Date:
Period:
Pre-lab Questions:
1. What constitutes a positive test for each of the following gases?
a. oxygen (O2):
______________________________
b. hydrogen (H2):
______________________________
c. water vapor (H2O):
______________________________
d. ammonia (NH3):
______________________________
e. carbon dioxide (CO2) ______________________________ and/or ____________________________
2. What is the proper way to smell a substance in the lab? __________________________________________
3. What is the role of a catalyst in a reaction? How can you tell if when a substance serves as a catalyst?
______________________________________________________________________________________
Procedure:
Read the procedure for each reaction in its entirety before doing it so that you have the necessary
materials to complete the procedure. Wear goggles and an apron throughout the entire lab. Be sure to
tie back long hair. Write detailed observations in the Chart on page 3. Follow disposal instructions.
Reaction 1: Obtain a jar of Cu powder and a crucible from the lab bench and weigh the empty, clean crucible. Add
about 1.00 g of Cu powder and determine the actual mass of copper to the nearest .001 g and record.
Place the crucible on a clay triangle and heat it strongly in a Bunsen burner flame for 5 minutes. Allow
the crucible to cool and weigh it again. Record the initial and final appearances of the copper as well as
the masses. Disposal: after the copper cools, throw it into the trash can – wipe out the crucible with a
dry paper towel.
Reaction 2: Obtain a 3 cm piece of magnesium ribbon from the lab bench. Reweigh the empty, clean crucible to the
nearest .001 g. Roll up the Mg and place it on the flat bottom of the crucible and weigh the crucible and
Mg again. Place the crucible on a clay triangle and heat it strongly in a Bunsen burner flame until it
reacts. Do not look directly at the burning Mg. After it burns, and cools, reweigh the crucible and
product and record the masses and appearances in the table. Disposal: throw ashes in trash can after it
has cooled down, wipe out the crucible and return it to the lab bench.
Reaction 3: Obtain another 3 cm piece of magnesium. Add 5-10 mL of 3.0 M hydrochloric acid (HCl) to a small test
tube in a test tube rack. Add the magnesium to the acid. Using a test tube clamp, invert a second test tube
(as shown in the figure below) over the mouth of the reaction test tube and collect the gas being produced.
Keep the test tube inverted and test the collected gas by inserting a flaming splint. Record your results.
Disposal: after the Mg has reacted completely with the HCl, the solution can go down the sink with lots of
water.
33
Reaction 4: Obtain one scoop of ammonium carbonate, (NH4)2CO3, and place it into a small test tube. Clamp the test
tube and heat it in a Bunsen burner. While the test tube is heating, touch a piece of cobalt chloride paper
to the mouth of the test tube and remove it. Smell the test tube using the proper wafting technique. Hold
a flaming splint in the test tube and record the results of each test. Record. Disposal: any solid remaining
in test tube may be rinsed down the sink with water.
Reaction 5: Measure 10 mL of hydrogen peroxide and add to a test tube. Add a very small quantity (tip of spatula)
of manganese dioxide, MnO2, (catalyst) to the test tube. Hold a glowing splint inside the mouth of the test
tube. Hold a piece of cobalt chloride paper to the mouth of the test tube. Record results. Disposal: place
both into the trash.
Reaction 6: Add 1 drop of potassium iodide, KI, and 1 drop of lead (II) nitrate, Pb(NO3)2, to a watch glass. Record the
results. Disposal: wipe the watch glass with paper towel and throw paper towel in trash. Do not put it
into the sink. Be sure to wash your hands well with soap and water.
Reaction 7: Place two scoops of copper (II) carbonate, CuCO3, into a dry large test tube. Insert a stopper with a glass
bend in the test tube. Clamp the test tube to a ring stand at an angle so that the flame will touch the
bottom of the test tube only. Fill a small test tube approximately half way with limewater. Position this
test tube so that the end of the glass bend is in the test tube and is submerged in the limewater. Light the
Bunsen burner and heat the solid in the large test tube. Observe any changes in the limewater. Disposal:
once product cools, place in the trash can, clean test tube; place limewater into the sink with lots of water.
34
Test Results
Reaction
Observations
Odor
Limewater
Initial Mass:
Final Mass:
2. Burning Mg
Appearance Before:
Appearance After:
Initial Mass:
Final Mass:
Before:
3. Mg and HCl
Flaming splint
During:
After:
Before:
Flaming splint
During:
After:
Before:
5. H2O2
Splint
Appearance Before:
Appearance After:
1. Burning Cu
4. Heating
(NH4)2CO3
Cobalt
Chloride
Glowing splint
During:
After:
6. KI and
Pb(NO3)2
7. Heating
CuCO3
Before:
After:
Before:
After:
Questions:
1. Complete and Balance each equation using the test results above. Indicate the reaction type (synthesis,
S, decomposition, D, single replacement, SR, double replacement, DR, or combustion, C) on the line
provided.
Reaction 1:
type:______
_______________________________________________________________
Reaction 2:
type:______
_______________________________________________________________
35
Reaction 3:
type:______
_______________________________________________________________
Reaction 4:
type:______
_______________________________________________________________
Reaction 5:
type:______
_______________________________________________________________
Reaction 6:
type:______
_______________________________________________________________
Reaction 7:
type:______
_______________________________________________________________
2. Write complete, balanced equations for each of the following. Identify the type (S, D, SR, DR, C):
______ a. When potassium bromate is heated, it decomposes into potassium bromide and a gas that reignites
a glowing splint.
______ b. Sodium metal reacts violently with water to produce sodium hydroxide and a gas that pops in the
presence of a flame.
______ c. When calcium hydroxide is heated, it decomposes into calcium oxide and a substance that will turn
cobalt chloride paper pink.
______ d. The Bunsen burner uses methane gas, CH4. It burns in the presence of O2 and produces a
substance that will extinguish a flaming splint and a substance that turns cobalt chloride paper
pink.
______ e. Iron metal will react with oxygen gas in the air to produce rust, iron (III) oxide.
______ f. When you mix baking soda, sodium bicarbonate, with vinegar, acetic acid, they will react
violently to form sodium acetate, and a gas that extinguishes a burning splint and a substance that
turns colbalt chloride paper pink.
36
Reactivity of Metals Lab
Name________________________________
Single Replacement Reactions
Period_______Date_____________________
Problem: Which metals will replace each other in single-replacement reactions?
Materials:
5 small pieces of copper and magnesium
iron fillings
12-well spot plate
1.0 M CuCl2
1.0 M MgCl2
1.0 M FeCl3
6.0 M HCl
Procedure:
1. Place a small amount of each of the metals in the well plate as shown in the data table.
2. Add enough of each of the solutions to completely cover the each metal.
3. Record any changes in the data table. Changes include bubbles or a colored precipitate.
“Clumping or dull” do not constitute a reaction.
4. Describe what happens in each block or put “NR” for “no reaction.”
5. List the four elements in order of their chemical reactivity.
Data Table:
Solutions →
HCl
CuCl2
MgCl2
FeCl3
H2 O
Metals
↓
Cu
Mg
Fe
Questions:
1. Which of the metals reacted with the most compounds?
_______________________________
2. Which of the metals reacted with the fewest compounds?
_______________________________
37
3. List the three metals you tested from most reactive to least reactive:
____________
____________
____________
4. Which metal(s) you tested was(were) able to replace the hydrogen in HCl? __________________
5. Where would hydrogen, as an acid, be placed in the activity series you listed in #3? Rewrite the
activity series placing hydrogen within the appropriate order:
____________
____________
____________
____________
6. Compare your results with the Activity Series handout given to you in class. Do your experimental
results agree with this list? _________
7. You tested the same three metals for their reactivity with the hydrogen in water. Can any of the
three metals we tested replace the hydrogen in water? __________ Are any metals capable of
replacing the hydrogen in water? __________ Name an example: _________________________
8. If aluminum (Al) would react with a solution of iron (III) chloride, but would not react in a solution
of magnesium chloride, where would aluminum be placed on the activity series? Rewrite the series
including aluminum in the appropriate location.
_________
_________
_________
_________
_________
9. Write the equations for the reactions that occurred from the data table:
_______________________________________________________________________________
_______________________________________________________________________________
_______________________________________________________________________________
_______________________________________________________________________________
_______________________________________________________________________________
38
Mole Relationship in a Chemical Reaction
Name_______________________________
30 Points
Period______ Date____________________
INTRODUCTION: Chemists can determine the equation for a chemical reaction by identifying the type
and amount of the substances involved. The coefficients for the balanced equation can be determined by
converting mass values to moles and expressing the moles values in terms of small whole number ratios.
It is not always practical to determine mass values for all the substances involved in the reaction, particularly
if gases are produced. Therefore, chemists will often use only the mass values for one of the
reactants and one of the products. By incorporating knowledge of the Law of Conservation of Mass, the
balanced equation can be determined.
In this experiment you will react sodium hydrogen carbonate (baking soda), NaHCO3, with
hydrochloric acid, HCl, to form sodium chloride, NaCl, and two gaseous products. You will determine
the mole relationship between NaHCO3 and NaCl, and use this data along with the Law of Conservation
of mass to predict a balanced equation for this reaction.
OBJECTIVES: In this experiment you will
 react a known mass of NaHCO3 with excess HCl
 calculate the mole relationship between NaHCO3 and NaCl
 predict a balanced equation for the reaction, and
 calculate the % yield and the % error of the actual product produced
EQUIPMENT:
goggles
Spoon
Bunsen Burner
apron
watch glass
striker
evaporating dish
HCl dropper
oven
balance
baking soda
PROCEDURE:
1. Clean and dry an evaporating dish. Do not wash it unless necessary. If you do wash it, heat it
strongly for 2-3 minutes in a Bunsen burner flame and allow it to cool. Determine the mass to the
nearest 0.001g.
2. With a spoon, add about 3.000 g of sodium hydrogen carbonate, NaHCO3, to the evaporating
dish. It does not need to be exactly 3.000 g, but should be measured precisely.
39
3. Cover the evaporating dish with a watch glass. Use the 6M HCl dropper and add about
6 mL (1 mL = 20 drops) CAUTION: HCl causes burns; avoid skin and eye contact. Rinse
spills with plenty of water. Allow the drops to enter the lip of the evaporating dish so that
they gradually flow down the side as shown in Figure 11-1.
4. Continue adding the acid slowly until the reaction has stopped. Do not add more acid than
is needed.
5. Tilt the dish from side to side to make sure the HCl has reached all of the NaHCO3, if any
unreacted NaHCO3 remains, add a few more drops of HCl to complete the reaction. Remove
the watch glass cover with forceps and rinse the underside of the watch glass with a very small
amount of distilled water from the water bottle. Be careful to wash all the material into the
evaporating dish as shown in Figure 11-2.
6. Heat the evaporating dish on a hot plate or on a ring stand over low heat. Be careful that the salt
does not spatter out of the dish. Continue to dry the solid slowly until all the moisture appears to
have evaporated.
7. Remove the dish from the heat and allow it to cool. Weigh the product to the nearest 0.000 g.
8. Reheat the dish and contents strongly, let cool, and remeasure the mass. The value should be
within 0.02 g of the last recorded mass. If it is not, reheat and remeasure the mass until the last
two measurements are within that range. Record the final measurement of the NaCl.
9. After completing the data table and questions #1 and #2 and checking with your teacher, the
contents of the evaporating dish may be rinsed down the drain with plenty of water.
DATA TABLE: Measure to the nearest 0.001g.
mass of empty evaporating dish
___________g
mass of evaporating dish and NaHCO3
___________g
mass of NaHCO3
___________g
mass of evaporating dish and NaCl
___________g
mass of NaCl
___________g
QUESTIONS:
1) Calculate the number of moles of NaHCO3 used in this reaction:
2) Calculate the number of moles of NaCl produced in this reaction:
3) Write a complete balanced equation for the reaction that occurred. (see the introduction)
40
4) From your balanced equation, what is the mole ratio between NaHCO3 and NaCl? How does this
compare with your mole ratio from questions #1 and #2. Give a reason why they are similar or different.
5) Suppose you started with 3.25 moles of sodium hydrogen carbonate, how many moles of sodium chloride
would you expect to be formed? If you started with X moles of NaHCO3 how many moles of NaCl
would you expect to be formed? Explain.
6) Using the mass of sodium bicarbonate, calculate the number of grams of sodium chloride that should
have been produced. This is the expected (theoretical) yield.
7) Calculate the % yield for NaCl produced in the experiment, by comparing it to the amount of NaCl that
should have been recovered.
% Yield 
actual
x 100% 
expected
8) Calculate the % error for NaCl produced in the experiment, by comparing it to the amount of NaCl that
should have been recovered.
% Error 
expected - actual
x 100% 
expected
41
42
Stoichiometry Lab #2
Name _____________________________
Period______ Date__________________
Introduction: In this lab you will be reacting lead (II) nitrate and potassium iodide to form potassium
nitrate and the yellow precipitate, lead (II) iodide. The following skills will be required for successful
completion of this investigation: formula writing, equation writing,
balancing equations, molar mass
determination, mole conversions and the application of the mole ratio concept.
Materials:
solid potassium iodide, lead (II) nitrate solution, balance, graduated cylinder,
small beaker, Erlenmeyer flask, funnel, filter paper, wash bottle with distilled water.
Procedure:
1.
Weigh approximately 2.000 grams of solid potassium iodide using a balance. Record the actual
mass in the data table.
2.
Using a graduated cylinder, measure 10.0 mL of the lead (II) nitrate solution and pour it into a
small, clean beaker.
3.
Add the potassium iodide to the beaker and gently swirl the beaker.
4.
Obtain a piece of filter paper. Write your initials (in pencil) in the corner of the filter paper.
Determine it’s mass and record it in the data table.
5.
Set up a filtering apparatus using a funnel and an Erlenmeyer flask. Fold the filter paper, place it
in the funnel and wet it slightly with distilled water.
6.
Begin to slowly filter the reaction mixture. Be sure not to allow the liquid level to rise above the
top of the filter paper. The yellow precipitate should remain behind on the paper.
7.
Rinse the beaker with a small amount of distilled water and filter it as well in order to get as much
of the precipitate out of the beaker and onto the filter paper.
8.
After the filtering process is complete, the liquid in the Erlenmeyer flask can be discarded into the
waste jar on the front lab bench.
9.
Using a wax pencil, write your initials on a watch glass. Carefully remove your filter paper from
the funnel using a spatula and a pair of tongs. Place the filter paper, still folded, on the watch glass
and put it in the lab oven to dry.
10.
Clean all glassware used with soap and return to the drawer. Wash your hands!
11.
When dry, reweigh your filter paper and record in the data table.
43
Data Table:
1. Mass of potassium iodide used at the beginning:
______________
2. Mass of filter paper alone:
______________
3. Mass of dry filter paper with lead (II) iodide:
______________
4. Calculate the (actual) mass of the lead (II) iodide produced:
______________
Questions:
1. What type of reaction did you complete? ___________________________________________
2. Write a complete balanced equation for the reaction that occurred.
3. Calculate the theoretical (expected) yield of lead (II) iodide in this experiment.
4. Calculate the percent yield of lead (II) iodide:
5. List a reason why your yield is either higher or lower than 100%:
6. Calculate your %error:
7. If another student repeated this experiment, but started with 12.8 g of potassium iodide and recovered
14.3 g of lead (II) iodide, calculate that student’s percent yield of lead (II) iodide.
44
Determining the Percent Yield of Copper
Name ______________________
Period _____ Date ___________
Problem: What is the percent yield of copper in the reaction between copper (II) sulfate and iron?
Definition: The Limiting Reactant is the substance used up _________________ in a chemical reaction.
Procedure:
1. Using a grease pencil, write your name and your partner’s name on the smallest beaker you have
in your drawer – make sure it is clean!
2. Weigh the beaker and record its mass in the data table.
3. Measure 12.20g of copper (II) sulfate pentahydrate crystals and add these to the beaker.
4. Measure 50.0mL of deionized water using a graduated cylinder and add the water to the crystals
in the beaker.
5. Heat the beaker over a Bunsen burner to dissolve the crystals. Do NOT allow the solution to
boil.
6. While the solution is heating, one partner needs to measure out 2.24 g of iron fillings.
7. Once the crystals have dissolved, remove the beaker from the heat and place on the lab bench.
8. Add the iron fillings very slowly (a little bit at a time) to the hot solution. Be sure to stir
continuously. When all the iron has been added, allow the beaker to cool for 10 minutes.
9. Decant the liquid into a 250mL beaker by gently pouring the liquid down the stirring rod into the
beaker. Do not disturb the solid at the bottom of the beaker.
10. Add about 10ml of deionized water to the solid and stirring vigorously. Allow the solid to settle
and decant again. Repeat the rinse 3-4 times.
11. Spread the solid over the bottom of the beaker and place it in the fume hood. The beakers will be
placed in an oven at 150oF and dried overnight. Do not heat to dryness, as the Cu will oxidize to
CuO, a black power.
10. After the solid is completely dry, find the mass of the beaker and the solid copper. Record this
mass in your data table.
11. Determine the theoretical yield and % yield of Copper.
Data Table: Day #1
Mass of empty beaker
Mass of copper (II) sulfate pentahydrate
Mass of iron fillings
45
Day #2
Mass of beaker and copper
Mass of copper
Calculations:
1. Write a complete balanced equation for the reaction taking place between iron and copper(II) sulfate
pentahydrate.
2. Convert the grams of copper(II) sulfate pentahydrate to grams of copper.
3. Convert the grams of iron to grams of copper.
4. Which reactant produces the least amount of copper? _______________________
5. Which reactant is the limiting reactant? _________________________
6. What is the most amount of copper that can be produced? ____________________
This is your theoretical yield.
7. Calculate the % yield of copper.
8. Calculate the % error.
Questions:
1. What changes indicated that all of the iron reacted?
2. Was your actual yield of copper close to your theoretical yield (was your % yield 100%). If not, list possible
sources of error.
46
Heat of Dissolving Lab
Name ____________________________
Date _________________ Period ______
Purpose: To measure the ________________________ as different _________________ dissolve in
_________________ and determine the _________________________________________.
Materials:
goggles and apron
distilled water
thermometer
stirring rod
CaCl2
graduated cylinder
foam cup as a calorimeter
balance
NH4Cl
Procedure:
1.
Put on goggles and lab apron. Measure 100.0 mL of distilled water at room temperature
and pour it into the plastic foam cup. Record the mass of the water by recalling that the
density of water is 1g/mL. Record the temperature of the water in the data table to the
nearest 0.1C. Do not remove thermometer from the cup, but be careful that it does not tip
over.
2.
Using your laboratory balance, measure out 8-10 grams of ammonium chloride on a piece
of paper. Record the mass to the nearest 0.01 g.
3.
Without removing the thermometer from the cup, shake the NH4Cl from the paper into the
water and stir gently with the stirring rod until the solid is dissolved. CAUTION: Both of
the solutions in the lab are irritating to the skin. Avoid contact with them.
4.
Make sure that the bulb of the thermometer is fully immersed in the liquid. If the
temperature rises, record the highest temperature reached. If the temperature falls, record
the lowest temperature reached.
5.
Dispose of the solution by pouring it down the drain, followed by plenty of water. Rinse
the cup, dry and return it and the thermometer to the lab bench.
6.
Repeat steps 2-5 using calcium chloride.
Observations:
Data Table:
Solute
Solute
Mass (g)
Mass of
Water (g)
Mass of
Solution
(g)
Initial T
(C)
Final T
(C)
(+/-) T
(C)
Exothermic
or
Endothermic
NH4Cl
CaCl2
47
Calculations for _________________________
1. Calculate the change in temperature. T = Tf ─ Ti
2. Calculate the heat absorbed or released by the solution. The specific heat of water is 4.184J/gC.
qsur = (mass of the solution) x (T of the water) x (specific heat of the water)
3. How much heat was released/absorbed (circle one) from/into the reaction? (qrxn = -qsur)
qrxn =
4. Using the periodic table, calculate the molar mass of the solute.
5. How many moles of the solute were used in the reaction?
6. Calculate the molar heat of solution (H) in (+/-) kJ for the solute from the formula:
Heat of the reaction
q
= rxn
∆H = Molar heat of solution =
moles solute dissolved
mole
7. Calculate the percent error of your experimental value.
% error =
accepted value - experimental value
x 100% =
accepted value
The accepted value for H for NH4Cl is +14.8 kJ/mol.
The accepted value for H for CaCl2 is –81.3 kJ/mol.
48
Calculations for _________________________
1. Calculate the change in temperature. T = Tf ─ Ti
2. Calculate the heat absorbed or released by the solution. The specific heat of water is 4.184J/gC.
qsur = (mass of the solution) x (T of the water) x (specific heat of the water)
3. How much heat was released/absorbed (circle one) from/into the reaction? (qrxn = -qsur)
qrxn =
4. Using the periodic table, calculate the molar mass of the solute.
5. How many moles of the solute were used in the reaction?
6. Calculate the molar heat of solution (H) in (+/-) kJ for the solute from the formula:
Heat of the reaction
q
= rxn
∆H = Molar heat of solution =
moles solute dissolved
mole
7. Calculate the percent error of your experimental value.
% error =
accepted value - experimental value
x 100% =
accepted value
The accepted value for H for NH4Cl is +14.8 kJ/mol.
The accepted value for H for CaCl2 is –81.3 kJ/mol.
49
Critical Thinking: Analysis and Conclusions
1. When sodium chloride dissolves in water, the ions dissociate. The equation for this reaction is
NaCl(s)  Na+(aq) + Cl-(aq) Write similar ionic equations to show the dissociation in water of each
of the solutes used in the investigation.
_______________________________________________________________________________
_______________________________________________________________________________
2. Which reaction was exothermic? Which was endothermic?
_______________________________________________________________________________
3. Rewrite each of the ionic equations from Question 1 showing the molar heat of solution as a reactant
or a product.
_______________________________________________________________________________
_______________________________________________________________________________
4. When the reactants get colder in an endothermic reaction, what has happened to the heat energy?
_____________________________________________________________________________
5. Is the change in enthalpy positive or negative for an exothermic reaction? Explain. __________
_____________________________________________________________________________
6. Suggest two uses for these solution reactions in sports injuries or camping.
_______________________________________________________________________________
7. Which solids from this investigation could be used in each of your answers in the previous question?
Explain your reasoning.
_______________________________________________________________________________
_______________________________________________________________________________
8. Why is a plastic foam cup used instead of a beaker in this experiment? For what piece of equipment
is this cup a substitute?
______________________________________________________________________________
______________________________________________________________________________
9. How could you reduce experimental errors in this investigation? Explain your reasoning.
_______________________________________________________________________________
_______________________________________________________________________________
50
Specific Heat of a Metal Lab
Name ______________________________
Date ____________________ Period _____
Purpose: To determine the specific heat of a metal using a calorimeter.
Materials:
goggles and apron
beaker (400 mL)
test tube
foam cup
thermometer
balance
Bunsen burner with ring stand, wire gauze, and ring clamp
Procedure:
1. Wear safety goggles and aprons at all times.
2. Fill a 400mL beaker half full with water. Place the beaker on the wire gauze and heat until
boiling.
3. Measure and record the mass of an empty small test tube to the nearest .01g.
4. Add sample metal pieces until the test tube is half full. Record the mass of the test tube and
metal.
5. Place the test tube containing the metal into the beaker of water and continue heating. Leave the
test tube in the boiling water bath while you complete steps 6 through 9.
6. Obtain a foam cup and carefully measure the mass and record.
7. Fill the plastic foam cup about half full with distilled water at room temperature and record mass.
8. While the test tube is still in the boiling water bath, measure the temperature of the water to the
nearest .1oC and record it. It will be assumed that the temperature of the metal is the same as the
water.
9. Measure the temperature of the water in the foam cup and record.
10. Remove the test tube from the boiling water and immediately pour the metal into the
Styrofoam cup.
11. Stir the water in the cup with your stirring rod. Record the highest temperature reached.
12. Recover the metal by pouring off the water. Spread the metal on the paper towel to dry.
Do NOT let any metal particles get into the sink.
51
Analysis: Name of Metal Assigned ________________________________
1. Data Table
Mass Data
(All numbers to .01g)
small test tube
Temperature Data
(All numbers to .1oC)
boiling water
test tube and metal
(metal i)
water in cup
metal alone
(H2O i)
water in cup after metal
was added (metal f and H2O f )
foam cup
T of water in cup
(H2O f - H2O i )
foam cup and water
T of metal
water alone
(metal f - metal i)
-------------------------
2. Calculate the heat gained by the water.
Heat gained by water (qsur) = mass of water
x
T of water
x
Cp of water
3. qsur = -qrxn heat gained by the water = - heat lost by the metal
4. Calculate the specific heat capacity of the metal.
specific heat of metal = heat lost by the metal / (mass of metal x T of metal)
5. Calculate the % error for the specific heat of your metal using one of these accepted values:
Al = .897 J/goC, Cu = .385 J/goC, Cr = .449 J/goC, Pb = .129 J/goC, Fe = .448 J/goC
Conclusions:
1. What other physical properties, other than specific heat, could you use to help you identify the samples
used in this experiment? Give two.
2. Why is water excellent to use in a calorimeter?
52
Determining the Molar Mass of Butane
Gas Laws
Name___________________________
Date ________________ Period _____
Purpose: To determine the molar mass of butane (C4H10) using Dalton’s Law of Partial Pressures and the
Ideal Gas Law.
Procedure:
1. Fill a pneumatic trough and a small Erlenmeyer flask with very warm water (you will eventually need
to hold your hand under the water for a few minutes so make sure it is not too hot). Be sure that the
flask is completely filled with water.
2. Hold your hand over the mouth of the flask and carefully invert the flask into the trough without
allowing any of the water to spill out.
3. Obtain a butane lighter and carefully determine the mass to three decimal places. At no point during
this experiment are you to light the lighter. In addition to being a safety violation it will alter the
mass of the lighter. FAILURE TO ABIDE BY THIS CONDITION WILL NOT BE
TOLERATED.
4. While holding the lighter in the water under the mouth of the flask, depress the gas release button.
You will observe the flask filling with butane as it displaces the water. You must collect enough
butane so that the water level inside the flask (while touching the bottom of the trough) is at or below
the water level in the trough.
5. Shake off any water from the lighter and dry it thoroughly with a paper towel. Stand the lighter up in
the fume hood under the heat lamp for a few minutes to dry. (Important Note: Put your lighter on a
piece of paper with your initials, so you can obtain the final mass of your lighter.)
6. Equilibrate the pressure in the flask with the atmosphere (make water levels equal). While holding
the flask at this position, slide your hand under the mouth and quickly remove the flask from the
trough turning it right side up without losing any water. Keep your hand over the mouth and carry
the flask to the fume hood to release the gas.
7. Measure and record the temperature of the water in your flask.
8. Measure and record the volume of the water in your flask.
9. Measure and record the total volume of your flask (fill completely with water and measure the water
volume with a graduated cylinder.)
10. Measure and record the mass of the dry lighter. Return the lighter to the teacher.
53
Data:
Initial mass of
butane lighter
Final mass of
butane lighter
Mass of butane
used
Volume of water
in flask (mL)
Total Volume of
flask (mL)
Volume of butane
in flask (mL)
Barometric Pressure
(mm Hg)
Vapor pressure (see
table provided)
Temperature of water
°C
Temperature of water
(convert to K)
----------------
---------------
----------------
---------------
Analysis:
Pressure (P)

Since you positioned the flask so that the water levels were equal, the total pressure inside the
flask is equal to the air pressure in the room (barometric pressure). However, in addition to the
butane, the flask also contains water vapor because some of the water evaporated. According to
Dalton’s Law of Partial Pressures, the total pressure of the gas mixture is equal to the sum of the
partial pressures of the individual gases: Pbarometric = Pbutane + Pwater vapor


Using Dalton’s Law of Partial Pressures, calculate the partial pressure of butane.
Pbutane = ________________ mm Hg
Convert this value into atmospheres.
Pbutane = _______________ atm
Volume (V)

Using the volume data, calculate the volume of the butane in the flask. Express your answer in
liters.
Vbutane = _______________ L
Temperature (T)

Assuming the gas temperature to be the same as the water temperature, convert the gas
temperature into Kelvin.
Tbutane = _______________ K
54
Moles (n)

Given the pressure, volume, and temperature of butane determined above, use the ideal gas law,
PV=nRT, and calculate the number of moles (n) of butane.
Note: R is the ideal gas constant. The value for R is 0.0821 atm·L
mol·K
nbutane = _______________mol
Molar Mass (M)

Molar mass is defined as the number of grams in 1 mol of a pure substance. Dividing the mass
of the substance by the number of moles will give you the molar mass in g/mol.

Using the mass of lighter data, calculate the mass of butane delivered into the flask.
massbutane=_____________g

Calculate the experimental value for the molar mass of butane (grams divided by moles)
Mbutane=_______________g/mol
Percent Error

The chemical formula for butane is C4H10. Use the periodic table and determine the actual
molar mass of butane and calculate the percent error.
Questions:
1. Calculate the volume that your butane sample would occupy at STP.
2. Using your experimental value for the molar mass of butane, calculate the volume that 0.659 g of
butane would occupy at 63.4 kPa and 29oC.
55
Water Vapor Pressure Table
Temperature
(°C)
0.0
5.0
10.0
12.5
15.0
15.5
16.0
16.5
17.0
17.5
18.0
18.5
19.9
Pressure
(mmHg)
4.6
6.5
9.2
10.9
12.8
13.2
13.6
14.1
14.5
15.0
15.5
16.0
16.5
Temperature
(°C)
19.5
20.0
20.5
21.0
21.5
22.0
22.5
23.0
23.5
24.0
24.5
25.0
26.0
Pressure
(mmHg)
17.0
17.5
18.1
18.6
19.2
19.8
20.4
21.1
21.7
22.4
23.1
23.8
25.2
Temperature
(°C)
27.0
28.0
29.0
30.0
35.0
40.0
50.0
60.0
70.0
80.0
90.0
95.0
100.0
Pressure
(mmHg)
26.7
28.3
30.0
31.8
42.2
55.3
92.5
149.4
233.7
355.1
525.8
633.9
760.0
56
Solutions Lab
Name________________________________
Period________Date____________________
1. Carefully weigh out 2.20 grams of anhydrous sodium sulfate and add it to a large test tube. Using a
graduated cylinder, measure and add 12.8 mL of distilled water to the test tube. Stopper the test tube
and shake.
a. Describe the contents of the test tube:
b. Is the solution saturated, unsaturated or supersaturated? (circle one)
c. How do you know?
2. Remove the stopper from the test tube. Using a test tube clamp, gently heat it back and forth in a
Bunsen burner flame until all of the solute is dissolved. Remove it from the flame and add a very
small amount of solid solute to the test tube. Allow it to cool a moment, stopper and shake the
solution.
a. Describe the contents of the test tube during heating:
b. What happened to the added solute?
c. Is the solution saturated, unsaturated or supersaturated? (circle one)
d. How do you know?
57
3. Place the test tube in a beaker of ice water for five minutes. Do not disturb the solution during this
process. Gently remove the test tube and place it in a test tube rack. Describe the contents below,
then add a very small amount of solid solute to the test tube. Set the tube back in the rack. Do not
stir or shake the tube. Observe the test tube for several minutes. For cleanup, you may need to reheat
the test tube to rinse the solute down the drain.
a. Describe the contents of the test tube before the extra solute was added.
b. What happened when the extra solute was added?
c. Describe the contents of the test tube after the extra solute was added.
d. Before adding the crystal, the solution was saturated / unsaturated / supersaturated. (circle
one)
e. How do you know?
f. After adding the crystal, the solution was saturated / unsaturated / supersaturated. (circle
one)
g. How do you know?
4. Explain one simple test that will determine whether a solution is saturated, unsaturated or
supersaturated. Explain how to interpret the test results.
58
Penny Alchemy
Name ______________________________
Period _____Date ___________________
BACKGROUND: Pennies have always been made from both copper and zinc. During WWII in 1943, as a
way to save our copper resources, pennies were made of zinc-coated steel. Unfortunately, zinc is more
reactive than copper, and the pennies have become known as “white cents” to collectors. During the early
1980’s, the amount of copper used in a penny cost more than a penny. To solve this economic problem,
“new” pennies were introduced in August of 1982. The original composition of a penny was 95% Cu and
5% Zn with a total mass of around 3.10 g. The “new” penny is a thin coating of copper over a zinc core. Its
composition is 2.50% Cu and 97.5% Zn with a mass about 20% lower, around 2.500 g.
PROCEDURE:
1. Using Steel Wool, clean 2 pennies until they are completely shiny. This is important!
2. Measure approximately 1.0 gram of granulated zinc into an evaporating dish.
3. Place the two pennies, side by side, onto the zinc. Add 20 mL of 3.0 M zinc chloride (ZnCl2) solution to
the dish.
4. Set-up a ring-stand with a wire screen and place the evaporating dish on the wire screen.
5. Using a Bunsen burner, heat the mixture to a gentle boil. Continue heating until the pennies are silver
around the edges. (Be sure that the pennies do not come to rest one on top of the other.)
6. Turn off the flame, and using beaker tongs, place the evaporating dish on the lab top. Use your crucible
tongs to turn both pennies over in the evaporating dish. Place the dish back on the ring stand and bring
the solution back to a gentle boil for about 5 minutes. Shut off burner.
7. Use beaker tongs to again place the evaporating dish on the lab top. Remove the pennies with your
crucible tongs and rinse them well with distilled water.
8. Examine the pennies to make sure that they are completely “silver”. If so, pat (do not rub them) dry with
a paper towel. If they are not completely silver, heat them in the dish for another 5 minutes.
9. Take one of the pennies and place it on a hot plate at setting 3.5 in the center for a few minutes until a
change to “gold” is observed. Do not move the pennies during heating.
10. Clean up your work area. You may keep the pennies.
QUESTIONS: (continued on the back)
1. What is the % composition of the post 1982 penny before the lab?
________________________
2. What is the “silver” color on the penny made of after step #8?
________________________
3. What is the “gold” color of the penny after step #9?
________________________
4. What are these materials (question #3 – metal solutions) called?
________________________
59
5. Which of the drawings to the left (a,b or c) show the penny after it
has been in the zinc bath? ______
6. Which of the drawings to the left (a,b or c) show the penny after it
has been heated on the hot plate? ______
7. What metals must be melted together to produce amalgam fillings?
______, ______, ______, ______, _______
8. What metals must be melted together to produce pewter?
______, ______, ______, ______,
9. What metals must be melted together to produce white gold?
______ + ______
10. What metals must be melted together to produce steel?
______ + ______
60
Molarity and Dilutions Lab
Name____________________________
Period______Date________________
The following lab is meant to give you hands-on experience with making solutions using both the
molarity formula (M=n/V), as well as the formula for dilutions (M1 x V1 = M2 x V2). In particular you will
make a 0.05000M stock solution of C6H12O6. From that stock solution you will make four dilutions resulting
in solution concentrations of 0.04M, 0.03M, 0.02M, and 0.01M.
Materials:
•25.0 mL Volumetric Flask
•4 test tubes
•test tube rack
•10.0 mL graduated cylinder
Procedure and Calculations:
1. Clean test tubes, graduated cylinders and volumetric flask.
2. Make 25.0 mL of a 0.0500M C6H12O6 stock solution in the volumetric flask. Determine the mass of sugar
needed for the stock solution. Add 1 drop of food coloring to the flask. Show all work.
3. In the 4 test tubes make the following dilutions from your stock solution. Do your calculations first:
a. From the stock solution make 10.0 mL of 0.040M C6H12O6 in the 1st tube.
b. From the stock solution make 10.0 mL of 0.030M C6H12O6 in the 2nd tube.
c. From the stock solution make 10.0 mL of 0.020M C6H12O6 in the 3rd tube.
d. From the stock solution make 10.0 mL of 0.010M C6H12O6 in the 4th tube.
4. Show your teacher the dilutions and have him/her sign the bottom of the page . Clean up and return the
volumetric flask.
61
Question:
1. Observe the solutions in each test tube. Which molarity has the deepest color? The lightest?
What does this indicate about the number of C6H12O6 molecules (or concentration) in each solution you
made?
62
Determining Solution Concentration
Using a Spectrophotometer
Name________________________________
Period______Date_____________________
Introduction
The spectrophotometer is a powerful tool which can be used for colorimetric determination of
concentration. The process is based on the fact that colored ions absorb light from the visible spectrum. The
greater the concentration of the ions, the greater the absorbance (A) of light.
Conversely, the more light that is absorbed by the ions in solution, the less light that is transmitted
through the solution. Thus, the inverse of absorbance is percent transmittance (%T). Therefore,
concentration can be measured using absorbance or percent transmittance.
To determine the concentration of a colored ion in solution, a set of carefully prepared solutions of
known concentration must be first measured for absorbance or percent transmittance.
The
spectrophotometer should have the wavelength setting at a previously determined maximum absorbance (or
%T) value for the ion in question while testing the standards. A graph plotting the absorbance (or %T)
against concentration data for the standards gives a calibration curve. Then, the absorbance (or %T) can be
measured for a solution of unknown concentration, then matched with the calibration curve to determine the
concentration.
The cobalt ion, Co2+, concentration will be determined in this experiment and should be expressed in
units of molarity. The absorption maximum (%T minimum) for Co2+ occurs at a wavelength setting of 510
nm. The percent transmittance, %T, scale will be used to produce a calibration curve.
Objectives
In this experiment you will:
* prepare a set of Co2+ standard solutions of known concentrations;
* measure the %T of the known solutions and construct a calibration graph; and,
* measure the %T for several unknowns and determine their concentration from the
calibration graph.
Materials: volumetric flask
Co(NO3)2  6 H2O
5 small test tubes
distilled water
8 disposable test tubes
10 mL graduated cylinder
Procedure:
1. Make 50.0 mL of a 0.10 M Co(NO3)2  6 H2O standard stock solution. First determine the moles of
cobalt (II) nitrate hexahydrate needed for this volume of water and then calculate the grams of
Co(NO3)2  6 H2O required. Show your work below: (3 points)
2. Label five small and 5 disposable test tubes #1 thru 5. Label 2 other disposables for Unknowns A and
B.
3. Using the stock solution and a small graduated cylinder, make 10. mL dilutions in each of the small test
tubes. Tube 2 should be 0.08 M, Tube 3 is 0.06 M, Tube 4 is 0.04 M, Tube 5 is 0.02 , and Tube 1 gets
10 mL of the 0.1M stock solution. Show your calculations for the 4 test tubes below for the amount of
stock solution required for each test tube, the rest should be diluted with distilled water up to 10.0 mL.
(4 pts.)
63
4.
Transfer each of the dilutions from the small test tubes to the identically numbered disposable test tubes
filling each tube about ¾ full.
5. Carefully use a test tube rack to transport these small test tubes to a spectrophotometer.
6.
The spectrophotometer should be warmed up and the wavelength set to 510 nm. The device needs to
be zeroed. With the sample compartment empty and closed, adjust the %T (transmittance) to 0 with the
left front knob.
7. Place a disposable test tube filled with distilled water (this is called a blank), into the sample
compartment and adjust the %T to 100 %T with the right front knob (this means all of the light travels
through the totally clear water.) Wipe all liquid and smudges off of the tubes before placing in
spectrophotometer!
8. Remove the Blank (distilled water tube) and measure and record the %T for each of the five prepared
dilutions in the disposable test tubes and the unknowns A and B.
9.
9. Construct a calibration graph plotting %T on the y-axis and molarity on the x-axis for the Standard
Solutions, then draw a smooth curve. To determine the unknown molarities for A & B, plot the
transmittance on the Calibration Curve and then drop down on the X-axis to read the molarity for each
unknown. Show this on the graph paper. (8 points)
Data Table: (8 points)
Standard Solutions
Test Tube
Blank
1
2
3
Unknowns
4
5
A
B
Molarity
%T
100
Conclusions:
1. The actual concentrations of the unknown solutions were A = 0.075 M and B = 0.050 M. Calculate the
percent error of your results for A and B. (2 points)
% Error 
expected - actual
x 100% 
expected
2. Using your calibration graph, determine the %T of a 0.036 M and 0.072 M solution of Co2+. (2 points)
3. What are the maximum and minimum concentrations (molarities) that can be measured using your graph?
(2
points)
64
A Study of Chemical Equilibrium
Name ________________________
Period ___________ Date _______
In the study of chemical reactions to this point, the assumption has been made that most reactions have
proceeded from reactants to products until at least one of the reactants was completely consumed.
In some reactions, however, the products remain in contact with each other and, if the energy requirement is
low enough, the products react to form the reactants again. For example, if hydrogen and iodine molecules are sealed
in a container, they will react to form hydrogen iodine molecules: H2(g) + I2(g)  2HI(g)
The HI molecules also collide and react to form hydrogen and iodine again.
2HI(g)  I2(g) + H2(g)
When the rates of the forward and reverse reactions are equal, a state of chemical equilibrium is said to exist.
H2(g) + I2(g)  2HI(g)
The concentrations of H2, I2, and HI become constant at a given temperature and do not change until the conditions are
changed.
According to LeChatelier’s Principle, “any equilibrium system subjected to a stress tends to change as to
relieve the stress.”
Stress
H2 increased
Concentration change
Result
H2(g) + I2(g)
 2HI(g)
increase decrease
increase
Stress
I2 decreased
Concentration change
Result
H2(g) + I2(g)  2HI(g)
increase decrease decrease
Objectives: In this experiment, you will
-apply stresses to the equilibrium systems by the addition or removal of
reactants or products.
-explain the changes in equilibrium using LeChatelier’s Principle.
Procedure: Use attached table to record data. Caution: Wear goggles and aprons at all time.
Acids and bases can cause burns.
Solubility Equilibrium for NaCl
1. In a clean test tube obtain 5 mL of saturated NaCl solution.
2. Add a few drops of concentrated HCl and record the results.
3. Discard the solution and rinse the test tube.
Copper (II) Sulfate Equilibrium
1. Place 3.0 mL of 0.1M copper (II) sulfate solution in each of two test tubes. Add an equal number of
drops of 0.1 M NaOH to each test tube until a blue-white precipitate forms.
2. Place one of the test tubes in a hot-water bath. After four minutes, observe and record any changes.
Use the other tube for comparison.
3. Rinse the test tubes out with plenty of water.
Acid – Base Equilibrium
1. In a clean test tube, obtain 4 mL of distilled water. Add 2 drops of bromothymol blue (HBB).
2. Now add 0.1M NaOH, a drop at a time, until a color change occurs, then add 2 more drops. Mix by
shaking gently.
3. Add 0.1M HCl a drop at a time until a change occurs, note the change.
4. Now add 0.1M NaOH a drop at a time until a color change occurs.
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Data and Analysis
PART A
a. Observations: _____________________________________________________________
b. Using your data and the equation below, how did the reaction shift? Left or right (circle one)
c. What stress was imposed on this system? _____________
d. Due to the stress imposed in question c, how would the reaction shift? Left or right (circle one)
e. Use up and down arrows to describe effect of this stress on each substance.
NaCl(aq)

Na+(aq)
+
______________
________________
f.
Cl-(aq)
_______________
Do your answers for questions b & d coincide? Explain why the shift occurred.
__________________________________________________________________________________
__________________________________________________________________________________
PART B
Test Tube 1 & 2
a. Observations: _____________________________________________________________
b. Using your data and the equation below, how did the reaction shift? Right or left (circle one)
c. What stress was imposed on this system? _____________
d. Due to the stress imposed in question c, how would the reaction shift? Left or right (circle one)
e. Use up and down arrows to describe effect of this stress on each substance.
Blue solution
CuSO4 (aq)
________
+ 2NaOH(aq) + Heat
________
blue/white precipitate
 Cu(OH)2(s) + Na2SO4(aq)
________
________
________
f. Do your answers for questions b & d coincide? Explain why the shift occurred.
___________________________________________________________________________
___________________________________________________________________________
Heated Test Tube
a. Observations: _____________________________________________________________
b. Using your data and the equation below, how did the reaction shift? Left or right (circle one)
c. What stress was imposed on this system? _____________
d. Due to the stress imposed in question c, how would the reaction shift? Left or right (circle one)
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e. Use up and down arrows to describe effect of this stress on each substance.
Blue solution
CuSO4 (aq)
+ 2NaOH(aq) + Heat
________
________
blue/white precipitate
 Cu(OH)2(s) + Na2SO4(aq)
________
________
________
f. Do your answers for questions b & d coincide? Explain why the shift occurred.
___________________________________________________________________________
__________________________________________________________________________
PART C
Step 2
a. Observations: _____________________________________________________________
b. Using your data and the equation below, how did the reaction shift? Left or right (circle one)
c. What stress was imposed on this system? _____________
d. Due to the stress imposed in question c, how would the reaction shift? Left or right (circle one)
e. Use up and down arrows to describe effect of this stress on each substance.
Yellow
Blue
HBB(aq)
+
___________
H2O (l)

___________
H3O+(aq)
+
____________
BB-(aq)
____________
f. Do your answers for questions b & d coincide? Explain why the shift occurred.
___________________________________________________________________________
__________________________________________________________________________
Step 3
a. Observations: _____________________________________________________________
b. Using your data and the equation below, how did the reaction shift? Left or right (circle one)
c. What stress was imposed on this system? _____________
d. Due to the stress imposed in question c, how would the reaction shift? Left or right (circle one)
e. Use up and down arrows to describe effect of this stress on each substance.
Yellow
Blue
HBB(aq)
+
___________
H2O (l)
___________

H3O+(aq)
____________
+
BB-(aq)
____________
f. Do your answers for questions b & d coincide? Explain why the shift occurred.
___________________________________________________________________________
__________________________________________________________________________
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Step 4
a. Observations: _____________________________________________________________
b. Using your data and the equation below, how did the reaction shift? Left or right (circle one)
c. What stress was imposed on this system? _____________
d. Due to the stress imposed in question c, how would the reaction shift? Left or right (circle one)
e. Use up and down arrows to describe effect of this stress on each substance.
Yellow
Blue
HBB(aq)
+
___________
H2O(l)
___________

H3O+(aq)
____________
+
BB-(aq)
____________
f. Do your answers for questions b & d coincide? Explain why the shift occurred.
___________________________________________________________________________
__________________________________________________________________________
Conclusion Questions:
1. Explain how the equilibrium in Part A would be affected if solid NaCl were added to the saturated stock
solution.
2. What would happen if you cooled test tube 1 & 2 in Part B?
3. How would the addition of distilled water in Part C affect equilibrium?
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Lab: Acid-Base Properties
Name________________________________
Period________Date____________________
Introduction:
Acids and bases are common chemicals in everyday life. Many products from shampoos to fruit
juices, from medicines to cleaning agents derive much of their usefulness from their activity as acids and
bases. Acids can be classified as substances that ionize in aqueous solutions to produce hydronium ions,
H3O+ . Acids react with metals to produce hydrogen gas and turn litmus paper red. Bases can be classified
as substances that dissociate in aqueous solutions to produce hydroxide ions, OH-. Bases turn litmus paper
blue and feel slippery. The strengths of acids and bases depend on the extent to which they ionize, or
dissociate. Strong acids or bases dissociate almost completely, while weak acids or bases dissociate to a
lesser degree.
In this investigation you will observe some reactions of acids and bases with each other, with other
compounds, and with various indicators. From your observations, you should be able to describe some of
the characteristic properties of acids and bases.
Pre-Lab Questions:
1. What is an acid? ____________________________________________________________________
__________________________________________________________________________________
2. What is a base? ______________________________________________________________________
__________________________________________________________________________________
3. What is an indicator? What indicators will you be using in this experiment?______________________
___________________________________________________________________________________
___________________________________________________________________________________
4. Write balanced chemical equations for the reactions that occur when the following solutions are mixed:
a.) HNO3(aq)
+
NaOH(aq)
→
___________________________________________
b.) 2HCl(aq)
+
Ca(OH)2(aq)
→
___________________________________________
5. What safety precautions need to be observed when handling acids and bases? ____________________
__________________________________________________________________________________
6. For what gas is the reaction product being tested in Part B? Part C? _____________________________
7. Give an example of each of these reactions from the investigation:
a.) neutralization: _________________________________________________________________
b.) double replacement: _____________________________________________________________
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Part A: Acid-Base Indicators
1. Pour 15 mL of 1.0 M HCl into a small beaker.
2. Support a piece of red litmus paper with tweezers and dip it into the beaker. Record the color.
Repeat with blue litmus paper.
3. Using tweezers, dip a piece of pH paper into the beaker. Obtain the value of the pH by comparing
the color of the paper to a reference chart.
4. Test each of the remaining indicators, phenolphthalein, bromothymol blue, and methyl orange by
pouring a small amount of the acid on a watch glass and adding 1 drop of the indicator. Record the
color (note: clear is not a color). Test one indicator at a time and be sure to rinse the watch glass
between uses.
5. Repeat the procedure using .5 M NaOH.
Please place used litmus and pH papers into the trashcan!
Part B: Reactions with metals
1. Pour 20 mL of 3.0 M HCl into a small beaker.
2. Obtain 2 pieces of the following metals (magnesium, zinc, copper and iron) and 1 wood splint and 1
dry test tube.
3. Put one piece of magnesium on a watch glass. Pour a small amount of acid on the metal and quickly
invert the test tube over the metal to collect any gas produced. Insert a flaming splint. Record
reaction and splint observations.
4. Test the other 3 metals with acids in the same manner described in procedure 3, but do not conduct
the splint tests on these metals. A positive result for each test is bubbles forming.
5. Test each of the remaining pieces of the four metals with .5 M NaOH and record observations.
Part C: Reaction of acids with carbonates
1. Fill a small test tube half way with limewater.
2. Add 3 scoops of calcium carbonate to a large test tube.
3. Obtain a stopper assembly. Place the stopper into the large test tube and the rubber tubing into the
test tube with the limewater.
4. Lift off the stopper and add enough 3.0 M HCl to cover the powder. Quickly return the stopper.
5. Record observations of the reaction and changes in the limewater.
Part D: Neutralization
1. Add 10 drops of 1.0 M HCl to a clean watch glass. Add 1 drop of phenolphthalein. Test with pH
paper.
2. Add 1 drop of .5 M NaOH to the watch glass. Gently swirl the glass to mix the contents. Continue
to add the NaOH 1 drop at a time (count the drops), swirling after each drop, until there is a
permanent color change. Record the pH of the final (neutral) solution and the total number of drops
of NaOH required to change the color.
70
Data and Conclusions:
Part A
Red Litmus
Blue Litmus
pH paper
Phenolphthalein
HCl
1.0 M
NaOH
.5 M
Questions
1. What color would a vinegar solution be in the presence of:
a. Phenolphthalein? ____________________
b. litmus? ___________________
2. What color would an ammonia solution be in the presence of
a. Phenolphthalein? ___________________ b. litmus? ________________________
3. Which would have a higher pH, vinegar or ammonia? Why?
Part B
Magnesium
Zinc
Iron
Copper
3.0 M HCl
.5 M NaOH
Questions:
1. What general conclusion can be drawn regarding acids and bases and their ability to react with
metals?
2. Did any metal not fit into your general conclusion? Why?
71
3. Explain the results of the splint test.
4. Write an equation for each reaction that occurred. (Hint: The metals have a +2 charge.)
Part C
 Observations of reaction:

Observations of limewater:
Questions:
1. Write an equation for the reaction that occurred.
2. What did the limewater test for? ______________________________
3. Explain why carbonates are common ingredients in products such as Tums and Rolaids.
Part D

pH of initial solution: _______ number of drops of NaOH: _______ pH of final solution: ______
Questions:
1. Write an equation for the reaction that occurred.
2. What conclusion can be drawn about the pH of acidic, basic and neutral solutions?
3. Explain why the number of drops of HCl and the number of drops of NaOH were not the same
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Acid Base Titration
Name ______________________________
Period _______ Date __________________
Background:
When concentrations of acids and bases are expressed in molarity, it is easy to use a solution with a
known molarity to measure the molarity of another solution. This process is called a titration. Most acidbase titrations use two burets. One buret contains an acid and one contains a base. At the end of the
titration, you read both burets to see how much of each solution you used. In this experiment, you will be
given a hydrochloric acid solution with a concentration of 0.50 M. Your teacher knows the concentrations of
the unknown sodium hydroxide solutions. You will calculate the concentration to see how carefully you
performed the titration. In this titration, hydrochloric acid reacts with sodium hydroxide according to the
equation below.
HCl +
NaOH

NaCl +
H2 O
If your acid and base solutions have the same molarity, you will use equal volumes of them in the titration.
If the acid is more concentrated than the base, the volumes of the base used will be less than the volume of
the acid.
Materials:
2 burets
Unknown solution of HCl
1 buret clamp & ring stand
125 mL Erlenmeyer Flask
________ M NaOH
phenolphthalein
Procedure:
1) Accurately record the initial readings for both burets in the Data Table on the back.
2) Be sure to rinse your Erlenmeyer flask with DI water before starting.
3) Hold the Erlenmeyer flask under the acid buret. Carefully add about 10.0 mL of the acid to the
Erlenmeyer flask.
4) Add 2 drops of Phenolphthalein to the flask. The solution should remain colorless.
5) Hold the flask under the buret containing your unknown base. Slowly add the NaOH while swirling
the solution constantly. A pink color will form and disappear as you swirl the solution. The end
point is reached when one drop of NaOH solution will cause the solution to remain pink. The color
should be a faint pink.
6) Precisely record the final readings for both burets in the Data Table.
7) Discard the solution in the flask down the drain with plenty of water. Rinse the Erlenmeyer flask with
DI water again when you are done.
8) Repeat the titration 2 more times to obtain an average molarity for your unknown base.
9) Complete the Data Table by performing the necessary calculations.
73
Data Table: Unknown # ______
Titration #1
Titration #2
Titration #3
HCl
HCl
HCl
NaOH
NaOH
NaOH
Final buret reading in mL
Initial buret reading in mL
Volume used in mL
Molarity (M)
Calculations: (13 Points)
1. At the end point:
# of moles of acid = # of moles of base
2. Calculate the molarity of the unknown base for each trial. Show work here to receive full credit.
Don’t forget significant digits and units! Record the Molarity in the Data Table.
Titration #1
Titration #2
Titration #3
3. Calculate the average molarity of HCl.
Question: A standard solution of NaOH was prepared to determine the molarity of an unknown HCl
solution similar to this experiment. If 50.0 g of NaOH were dissolved in enough distilled water to prepare
500 mL of solution, what is the molarity of the standard solution? If 15.0 mL of this standard solution was
titrated with 12.5 mL of an unknown concentration of HCl, what is the molarity of HCl?
74