Periodic Table
1
18
1
H
1.008
3
Li
6.941
11
Na
22.99
19
K
39.10
37
Rb
85.47
55
Cs
132.9
87
Fr
(223)
2
He
4.003
10
Ne
20.18
18
Ar
39.95
36
Kr
83.80
54
Xe
131.3
86
Rn
(222)
2
4
Be
9.012
12
Mg
24.31
20
Ca
40.08
38
Sr
87.62
56
Ba
137.3
88
Ra
(226)
3
21
Sc
44.96
39
Y
88.91
57
La
138.9
89
Ac
(227)
4
22
Ti
47.88
40
Zr
91.22
72
Hf
178.5
104
Rf
(257)
5
23
V
50.94
41
Nb
92.91
73
Ta
180.9
105
Ha
(260)
6
24
Cr
52.00
42
Mo
95.94
74
W
183.9
106
Sg
(263)
7
25
Mn
54.94
43
Tc
(98)
75
Re
186.2
107
Ns
(262)
8
26
Fe
55.85
44
Ru
101.1
76
Os
190.2
108
Hs
(265)
9
27
Co
58.93
45
Rh
102.9
77
Ir
192.2
109
Mt
(266)
10
28
Ni
58.69
46
Pd
106.4
78
Pt
195.1
11
29
Cu
63.55
47
Ag
107.9
79
Au
197.0
12
30
Zn
65.39
48
Cd
112.4
80
Hg
200.6
13
5
B
10.81
13
Al
26.98
31
Ga
69.72
49
In
114.8
81
Tl
204.4
58
Ce
140.1
90
Th
232.0
59
Pr
140.9
91
Pa
(231)
60
Nd
144.2
92
U
238.0
61
Pm
(147)
93
Np
(237)
62
Sm
150.4
94
Pu
(242)
63
Eu
152.0
95
Am
(243)
64
Gd
157.3
96
Cm
(247)
65
Tb
158.9
97
Bk
(247)
66
Dy
162.5
98
Cf
(249)
67
Ho
164.9
99
Es
(254)
14
6
C
12.01
14
Si
28.09
32
Ge
72.59
50
Sn
118.7
82
Pb
207.2
15
7
N
14.01
15
P
30.97
33
As
74.92
51
Sb
121.8
83
Bi
209.0
16
8
O
16.00
16
S
32.07
34
Se
78.96
52
Te
127.6
84
Po
(210)
17
9
F
19.00
17
Cl
35.45
35
Br
79.90
53
I
126.9
85
At
(210)
68
Er
167.3
100
Fm
(253)
69
Tm
168.9
101
Md
(256)
70
Yb
173.0
102
No
(254)
71
Lu
175.0
103
Lr
(257)
Constants and Conversions
Na 
6.02 E 23
mole
F = 96486 C/mol e-
R  0.0821
R  8.314
L  atm
mol  K
J
mol  K
1 lbs = 453.6 g
1 inch = 2.54 cm
1 mile = 1.609 km
1 cal = 4.284 J
K = oC + 273.15
oC
= 5/9(oF-32)
h = 6.63 E –34 Js
c = 2.997 E 8 m/s
1 atm = 760mm
Hg
1 atm = 101,325
Pa
1 oz. = 29.6 mL
1 tera (T) = 1 E 12
1 deci (d) = 0.1
1 giga (G) = 1 E 9
1 centi (c) = 0.01
1 mega (M) = 1 E 6
1 milli (m) = 0.001
1 kilo (k) = 1000
1 micro () = 1 E –6
1 hecto (h) = 100
1 nano (n) = 1 E –9
1 deka (da) = 10
1 pico (p) = 1 E –12
2
Table of Contents
II
III
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
28
29
30
31
32
33
34
Laboratory Safety Information ...................................................................................................4
Lab Journal Guidelines .................................................................................................................5
Measurement and SI Basics - Part A ...........................................................................................7
Measurement and SI Basics – Part B.........................................................................................10
Chemical Names and Calculations ............................................................................................16
Chemical Formula of Magnesium Oxide ...................................................................................20
Determination of Calcium in Milk .............................................................................................22
Investigations of the Ideal Gas Law ..........................................................................................24
Finding the Calories in Food Products .....................................................................................27
Percent of Oxygen in the Air ......................................................................................................29
A Chemical Formula by Microscale Titration...........................................................................31
Atomic and Molecular Spectroscopy ........................................................................................33
Molecular Geometry ...................................................................................................................36
Paper Chromatography..............................................................................................................38
Freezing Point Depression ........................................................................................................40
Determination of a Chemical Rate Law ....................................................................................42
Determining Order of Reaction .................................................................................................45
Determination of a Keq ...............................................................................................................47
CO2 in Breath: An Acid-Base Reaction ......................................................................................50
Basic Acid-Base Titrations .........................................................................................................52
Determination of Ka for Acetic Acid ..........................................................................................54
Determination of H3PO4 in a Cola Beverage .............................................................................56
Ca(IO3)2 : Solubility Product - Ksp..............................................................................................57
Determination of Water Hardness ............................................................................................59
Paper Chromatography – Metal Ions ........................................................................................61
Separation of Group I Ions: Silver, Mercury & Lead ................................................................62
Separation of Group II Ions: Iron, Aluminum & Zinc ...............................................................64
Separation of Group III ions: Barium, Calcium & Ammonium ................................................66
Analysis for Common Anions ....................................................................................................68
Cumulative Qualitative Analysis ...............................................................................................69
Nuclear Chemistry ......................................................................................................................70
Qualitative Eo Table ....................................................................................................................72
Bleach: A Redox Titration ..........................................................................................................74
Determination of F and NA .........................................................................................................76
Preparation of Aspirin ...............................................................................................................78
Polymerization Reactions ..........................................................................................................80
3
II
Laboratory Safety Information
For each lab you should have your goggles, lock, calculator,
lab manual, lab journal, text book, and a flash drive
RULES: The following rules are provided to help ensure your personal safety, the safety of others in the
lab, and that everyone will have access to uncontaminated supplies and working equipment to perform the
best possible analysis.
I. AVOIDING AND CARING FOR INJURY
1. Lab goggles MUST be worn at ALL times - ESPECIALLY if you wear contact lenses.
2. Open toed shoes are STRONGLY discouraged in the lab.
3. Long hair should be restrained behind your head.
4. It is preferable for you to wear long sleeve shirts and long pants to lab.
5. Wear old clothes to lab. It is your responsibility if clothes are damaged during lab.
6. If you choose to sit on the lab bench do so carefully. Someone may have left chemicals spilled on
the bench, which may cause you to leave the seat of your pants on the bench.
7. NO eating, drinking, open toed shoes, or horseplay in the lab.
8. Report ALL accidents and injuries promptly - regardless of severity.
9. Familiarize yourself with the location of fire extinguishers, safety showers and eyewash stations.
10. BURNS: Avoid picking up hot items, i.e. heated glass, rings or ringstands used with burners. If
burned, USE COLD WATER IMMEDIATELY.
11. CUTS: Avoid putting stress on glass tubing, especially when inserting into a stopper, putting on
hose, ect. Report cuts immediately, they can be more serious than they appear.
12. CHEMICALS ON THE SKIN OR EYES: Flood with copious amounts of water both for chemicals on the
skin and in the eyes.
II. SAFE DISPOSAL TECHNIQUES
1. Never throw insoluble solids into the sink. Ask your instructor for the appropriate disposal
procedure.
2. Use the hood whenever instructed to avoid exposure to toxic or irritating gases.
3. Report ALL chemical spills for instructions in proper cleanup and disposal.
4. Keep your lab bench clean at all times. A damp paper towel should be used for wiping the surface
that should then be dried. Your area must be clean when you leave the lab.
III. SOME SAFETY PRECAUTIONS
1. Carefully read the labels on all reagent bottles before you use them. Using the wrong chemical can
give serious flawed results and cause injury. Use only the quantities and concentrations needed.
2. Perform ONLY the assigned experiments for this course.
3. Never work alone in the lab.
4. Never taste any chemical.
5. When asked to smell a chemical, gently fan the vapor toward you, NEVER smell it directly.
6. Always point a test tube being heated away from yourself and your neighbors.
7. Before leaving the Lab, be sure the water and gas jets are completely shut off.
IV. PRACTICES TO AVOID CONTAMINATION AND WASTE
1. Insert only clean spatulas into reagent bottles. Replace stock bottles to their storage shelf. Do not
waste chemicals or you may be charged for the excess.
2. Weigh solids on weighing paper or in a glass container when using the balances. Never place
chemicals directly on the balance pan.
4. ONLY return a chemical to its stock bottle if you are SPECIFICALLY instructed to do so.
4
III Lab Journal Guidelines
The purpose of a lab journal is to keep an accurate chronological record of lab activities. The journal
should always be kept up-to-date, neat, and in ink (ball point). Before any information is recorded in the
journal, each page is to be consecutively numbered. Pages should NEVER be torn out of the lab journal.
You may write on both the front and back of the pages. Below are guidelines as to how to divide the lab
journal. Each section should be CLEARLY labeled.
FRONT OF THE LAB JOURNAL
1. Inside Cover Page – On the inside front cover of the journal include your name and as much personal
information as you would like to show. You must include at least be your name, your instructor’s name,
the course number and title, and the semester.
2. Table of Contents (TOC) – The first and second pages of the journal are reserved for the TOC. Each
page will have 5 columns in this order: Experiment Number, Experiment Name, Page Numbers in
the journal, Date the experiment was performed, and Grade. Leave the grade area blank. KEEP an upto-date TOC.
#
Name
Pages
Date
Grade
INDIVIDUAL EXPERIMENTS
1. Title – Each new experiment will begin on a new page in the journal. At the top of each page will be
written both the number and name of the experiment.
2. Notes – The notes section is to be used during lab for taking notes on procedures, recording initial data
that is collected, and, perhaps, sample calculations. The notes section is also a place for directly
recording information into your lab journal so that it doesn’t get lost.
3. Introduction – Each experiment will include a short introductory paragraph describing the objective of
the experiment. The purpose need only be 3 – 4 sentences.
4. Procedure – The procedure is a thorough and accurate account of how each experiment was
performed. The procedure should be written legibly. Any corrections given during the pre-lab lecture
must be noted in the procedure. The procedure may NOT be photocopied and placed into the lab
journal.
5. Results – This section will include data, observations, calculations, and graphs. It is to include a neat
and orderly set of the data collected during a particular experiment. The lab manual usually has a nice
set of data tables that may be copied (not photocopied) into the journal. Collected data should
ALWAYS include the appropriate units. Some experiments involve recording only observations. For
such experiments, those observations are included in this section. ALL calculations should be included
within this section. If calculations are required for a data table to be completed, the calculations should
be included within the table. The calculations should be neat, in a logical order, and include units.
Always leave enough room to effectively show the calculations. Numerical conclusions should be
clearly identified by placing a box around them.
5
If an experiment requires the construction of a graph, it must be done using a spreadsheet (preferably
EXCEL). All graphs must be properly labeled and of appropriate size (4 in. x 4 in. to 5 in. x 5 in.). An
example graph is below. Notice the size of the graph and the fact that BOTH axes are labeled AND
include the appropriate unit. All graphs should be neatly taped (using clear tape NOT MASKING TAPE,
DUCT TAPE OR STAPLES) into the lab journal. DO NOT title the graph or include a legend.
1.40
Current / uA
1.20
1.00
0.80
0.60
0.40
0.20
0.00
-0.80
-0.70
-0.60
Voltage / mV
-0.50
-0.40
6. Conclusions – This section will include a short summary of the results of the experiment. It should be
explicitly stated whether or not the objectives of the lab were met. Final numerical results and/or
overall observations will be restated here. If there was significant error (> 10%) in numerical results
or the objectives of the experiment were not otherwise met, justification as to possible sources of error
should be included. Any post-lab questions should be restated and answered in this section.
Points will be deducted for incomplete sections, sloppy work, loose-leaf pages of graphs or other data. If
errors are made in writing the report, do NOT scribble through them. Mark through them with a single
line. Do NOT use Whiteout on the pages.
6
1
Measurement and SI Basics - Part A
Prelab Questions
These are to be completed on a separate sheet of paper and turned it at the beginning of lab.
Show ALL work for problems requiring calculations.
Convert the following:
1.
5’9” into meters
2.
292 pounds into kg
3.
28 mi/gallon to km/L
4.
435 mg to g
5.
2 m/s to mi/hr
Introduction and Objectives
In this experiment, basic measurement methods will be reviewed and practiced. All measurements in
chemistry, as well as any other science classes will be made using the SI (or metric) units. It is important to
always include the units with a measurement of a physical property. In addition to making measurements,
practice will be gained in the construction and use of graphs, especially using spreadsheet software, in the
analysis of lab data. An introduction to two basic statistical tools will also be explored.
Procedure
Complete the following experimental measurements.
Data/Calculations
1. Reading graduated cylinders and burets. Make sure your instructor initials your journal.
a. 10 mL graduated cylinder
Volume = ___________ mL
__________ initials
b. 50 mL graduated cylinder
Volume = ___________ mL
__________ intials
2. a. Measure the dimensions of your group’s container and calculate the volume in cubic centimeters.
b. Determine the volume (in mL) of water your group’s container will hold within it COMPLETELY full.
a.
b.
container
c. Divide the results in (a) by the corresponding results in (b). What does the outcome of the
divisions indicate about the relationship between cm3 and mL?
7
3. Temperature measurements
a. First, make a sketch below of the graph of Temp (Y-axis) versus Time (X-axis) for heating water
from room temperature to boiling and letting it continue to boil as long as liquid water is present.
Have your instructor initial your sketch BEFORE you perform the experiment. State your reasons
for the shape of the line or curve in your proposed graph.
initials
b. Heat 250 mL of water in a beaker. Use a thermometer to record the temperature of the water in 2minute intervals until your instructor tells you to stop. Keep the thermometer in the water, but
do not let it touch the bottom or the side of the container. Record the temperatures to the
nearest 0.1oC.
c. Make a graph in MS Excel of Temp (in oC) (Y-axis) vs. Time (in minutes) (X-axis). Paste the graph
below.
d. Mark on the graph the temperature at which bubbles begin to form in the water. What is the
composition of the bubbles?
e. What was the boiling point of the water? Was it different than you expected? Why or why not?
8
Density
4. Determine the density of two of the following elements as indicated by your instructor: aluminum (Al),
zinc (Zn), copper (Cu), silicon (Si), iron, (Fe) and lead (Pb). Use water displacement to determine the
volumes of the materials. Use at least 70 grams of the Zn, Cu, Fe and Pb. Use at least 20 grams of the Al
and Si.
Note: You must carefully place the metals in the container so that (1) all of the particles are
completely submerged, (2) there is no air trapped in the metal pieces, and (3) no water is spilled
or splashed from the graduated cylinder.
Metal
Mass of metal
g
Initial Volume of grad. cylinder
mL
Final Volume of grad. cylinder
mL
Metal Volume
mL
Density
g/mL
Accepted Density Value
(source:
% Difference
)
 Experimental Value - Accepted Value

Accepted Value

Class Data
Cu
Al
g/mL

 100

%
Zn
Pb
Si
Fe
5. Density of rubbing alcohol.
Suggest a method to determine the density of rubbing alcohol. Have your procedure initialed below
before you begin.
Procedure:
Density (g/mL)
Accepted Density Value (g/mL)
(source:
% Difference
 Experimental Value - Accepted Value

Accepted Value

g/mL
g/mL
)

 100

%
9
2
Measurement and SI Basics – Part B
Prelab Questions
1. Convert 55 mi/hr to m/sec
2. If 575 g of a liquid fills a cylinder that is 10 cm tall with a diameter of 12 cm, what is the liquid’s
density? Is this liquid likely to be water? Explain.
3. Convert 750 nm to cm.
4. Convert 450 uL to mL.
5. Which has the greater density a solid gold coin the size of a penny or a quarter? Explain.
Introduction and Objectives
This experiment is a continuation of Part A having the same objectives and principles.
Procedure
Complete the following experimental measurements.
Data/Calculations
1. Density vs. Sample Size
a) Record the data for each group’s block of wood on the table.
b) Make a graph of volume (x-axis) vs. mass (y-axis). Include the best-fit
linear trendline with the equation and R2 value. (See page 14-15.) Tape it in
the space below.
#
Volume
(cm3)
Mass
(g)
1
2
3
4
5
6
c) What does the graph demonstrate about the density of an object?
d) What is the density of the wood?
10
2. Penny Density and Composition. In the last 30 years the composition of the pennies changed.
a. Record the year for each penny in your set and its mass.
Year
Mass (g)
Year
Mass (g)
Year
Mass (g)
1978
1982
1986
1979
1983
1987
1980
1984
1981
1985
b. Make a graph using MS Excel (Mass = Y-axis and Year = X-axis) and determine when the
composition of the pennies changed. (Place graph below) DO NOT ADD A TRENDLINE TO THIS
GRAPH.
c. Using 30 pennies from the period before the change and 30 pennies after the change, CAREFULLY
determine the density of the two sets of pennies. (Determine both the mass and volume of the
pennies in each group ALL together.)
30 Pennies before change
30 Pennies after change
Mass (g)
Volume (cm3)
Density (g/cm3)
Mass (g)
Volume (cm3)
density (g/cm3)
d. What is the major elemental component of pennies before and after the change? Choose from those
in Exp. 1.
Before
After
11
3. Determination of the Percent of Sugar in Soda
Procedure
a. Obtain one of the hydrometers from your instructor. DO NOT INVERT THE
HYDROMETER.
b. Place the hydrometer in a 50 mL graduated cylinder containing about 40 mL of
distilled water. The top of the hydrometer should float so that the stem is
slightly out of the graduated cylinder.
c. Practice forcing the hydrometer to float vertically in the graduated cylinder
WITHOUT affecting the height at which it floats. This will be important in your
measurements to come.
d. While the hydrometer is in the graduated cylinder with distilled water, measure
the height of the stem sticking OUT of the water (in mm). (See the diagram.)
e. Remove the hydrometer from the graduated cylinder. Carefully dry off the excess
water from the hydrometer and thoroughly dry the graduated cylinder.
f. Place about 40 mL of the 4% sugar solution into it. Carefully place the hydrometer
into the solution. Twirl the hydrometer to ensure that it is floating at the appropriate
level and, again, ensure that it is floating vertically in the cylinder. Measure the height
of the stem sticking out of the water (in mm). Record the height in the table.
g. Remove the hydrometer and rinse it with distilled water. Carefully dry off the
excess water. Also thoroughly rinse AND dry the graduated cylinder.
h. Repeat steps (f) and (g) with the 8%, 12%, 16% sugar solutions and the degassed soda
solution.
i. Using Excel, make a graph of the percent sugar (x-axis) versus height of the stem (y-axis).
Include the best-fit linear trendline with the equation and R2 value on your graph. DO NOT
include the data from the soda. This is your calibration curve.
% Sugar
0
4
8
12
16
Soda
stem height
12
j.
Once you have obtained your graph with the trendline and equation, you can determine the
percent of sugar in the soda by using the best-fit equation. You have measured the height of
the stem out of the water from the soda. As with the other samples, this is a “y value”.
Substitute that y value into your best-fit equation and calculate the “x value” for that height
which corresponds to percent of sugar in the soda.
Calculations for % sugar in soda
Experimental %
k. How can you calculate the percent of sugar in the can of soda? (Hint: Check out the
information on the can.) What other information will you need?
l.
Calculate the accepted value for the percentage and the percent error (difference) in what
you obtained from your experiment. (Show your calculations.)
Accepted %
Percent Difference
m. Why does the soda have to be degassed before using it in this experiment?
13
Graphing and adding a linear trendline and equation for a best-fit line:
Enter the X data
immediately to the left of the
Y data. Highlight all of the
data and click on the “Insert”
tab on the command line.
Click on “Scatter” and
choose the upper left graph.
This will give a screen
similar to the one below.
Once you see the graph,
right click on one of the
colored points. You will get
the menu shown. Choose
“Add Trendline”.
The choice for a linear
trendline is the default
choice. Check the boxes
"Display equation…" and
"Display R-squared…" as
shown.
Choose "close" to have the
line and equation added to
the graph
14
To finish making the graph, both axes should have the proper label with title and units. There is NO need to
add a Chart Title. You should also right click on the Legend (right side of the graph) and delete it. It adds
nothing to your graph.
To make a label, under “Chart Tools” choose
“Layout” and “Axis Titles”. In “Primary
Horizontal Axis Title”, choose “Title under
axis.” In “Primary Vertical Axis Title”, choose
“Rotated Title”.
The graph is now ready to be printed and placed into your lab journal.
15
3
Chemical Names and Calculations
Prelab Questions:
1. Convert 8 oz. to mL
2. What is volume of 8.6 g of a metal having a density of 20.5 g/cm3?
3. What is the name of the following compound Ba(MnO4)2?
4. What is the formula for the following copper (I) phosphate?
5. What is the formula for nitrous acid?
Introduction and Objectives
The main objective of this experiment is to become familiar with the names of ionic compounds
including salts, oxoacids, and bases and to gain practice in solving problems related to chemical reactions
and density.
Procedure
Complete each part of the experiment.
Data/Calculations
I.
Estimate the density of a person. Think about the general shapes of the human body to develop a
method for calculation density. CLEARLY SHOW YOUR CALCULATIONS.
a. Find the mass in grams of one of your group members.
b. Find the approximate volume in cm3 of one of your group members.
(Vsphere = 4/3r3 , Vbox = lwh , Vcylinder = r2h)
Part
Volume (cm3)
Head
Torso
Arms (x2)
Hands (x2)
Legs (x2)
Density =
Feet (x2)
Total V
(c) Does this method provide a reasonable approximation for the density of a person? Explain.
16
II. Chemical Reactions and Names
Complete the following reactions as seen in the first problem. In your answers, include the names and
formulas for the products. (Note: Some compounds represent an acid and some a base.) You DO NOT
have to balance the equations.
1.
BaCl2 + AgNO3
2.
AgSCN + Al2(SO4)3
3.
CaCO3 + Pb(OH)2
4.
Co(NO3)2 + KH2PO4
5.
Fe(HCO3)2 + Cs2O
6.
LiNO2 + Mn(ClO3)2
7.
Ba(OH)2 + H2SO3
8.
Ca(OH)2 + HClO4
9.
NH4OH + HNO
10.
CuBr2 + H3PO3
17
III. Exercises
1. Determination of the water content and mole ratio of water in hydrated copper sulfate.
Weigh out about 1.0 g of the hydrated copper sulfate into a crucible. Heat the sample on HIGH for 5
minutes. Allow it to cool on a thermal plate for 5 minutes before each mass determination. It should be
cool to the touch.
a.
Mass of the crucible
g
b.
Initial Mass of CuSO4XH2O
g
c.
Mass after heating on high for 5 minutes
g
d.
Mass after heating on high for another 2 minutes
g
e.
Mass after heating on high for another 2 minutes
g
f.
Mass of the CuSO4 remaining after all of the H2O is evaporated
g
g.
Mass of the H2O removed from the CuSO4
g
h.
Moles of CuSO4
mol
i.
Moles of H2O
mol
j.
𝑖. 𝑚𝑜𝑙𝑒𝑠 𝐻2 𝑂
(Round to the nearest whole number.)
ℎ. 𝑚𝑜𝑙𝑒𝑠 𝐶𝑢𝑆𝑂4
k.
Theoretical percent of water in the hydrated copper sulfate
(Note: Use the value of X from the bottle.)
(𝑋)(𝑚𝑎𝑠𝑠 𝑜𝑓 1 𝑚𝑜𝑙𝑒 𝐻2 𝑂)
𝑥100
( 𝑚𝑎𝑠𝑠 𝑜𝑓 1 𝑚𝑜𝑙𝑒 𝐶𝑢𝑆𝑂4 ) + (𝑋)(𝑚𝑎𝑠𝑠 𝑜𝑓 1 𝑚𝑜𝑙𝑒 𝐻2 𝑂)
l.
Experimental percent of water in the hydrated copper sulfate
𝑔. 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑤𝑎𝑡𝑒𝑟
𝑥100
𝑏. 𝑚𝑎𝑠𝑠 𝑜𝑓 ℎ𝑦𝑑𝑟𝑎𝑡𝑒𝑑 𝑐𝑜𝑝𝑝𝑒𝑟 𝑠𝑢𝑙𝑓𝑎𝑡𝑒
m.
%
Percent difference
(𝑘)−(𝑙)
(𝑘)
x 100
%
%
18
2. Pennies one more time…
From the previous week’s activity, 1982 pennies were found to be composed of 95% copper and 5%
zinc. Determine the mass of one of those pennies and calculate the number of copper and zinc atoms
there are in the penny.
Year
Penny mass
g
Mass of copper
Moles copper
Atoms of copper
Mass of zinc
Moles of zinc
Atoms of zinc
3. Density of a liquid
a. Determine the mass of 5.70 x 1022 molecules of acetone, C3H6O.
Mass =
g
b. Place a 10-mL graduated cylinder on a balance. Tear out the mass of the graduated cylinder. Add
acetone dropwise until you reach the mass recorded in 3 (a). Record the volume.
Volume =
mL
c. Using the data from parts (a) and (b), calculate the density of isopropyl acetone.
Density =
d. Accepted density value
g/mL
source
19
4
Chemical Formula of Magnesium Oxide
Prelab Questions
1. What is the name for CaSO45H2O?
2. What is the molar mass of CaSO45H2O?
3. How many grams of water are in 185g of CaSO45H2O?
PbI2 + KCl  PbCl2 + KI
4. For the reaction above how many grams of KI are produced from 15.2 g PbI2 with excess KCl?
5. For the reaction above how many grams of KI are produced from 10.5 g PbI2 mixed with 18.0 g of
KCl? Which reactant is the limiting reagent? Explain.
Introduction and Objectives
The goal of this experiment is to determine the chemical formula for the compound formed from the
reaction of metallic magnesium with elemental atmospheric oxygen. The experiment makes use of the fact
that by determining the mole ratio of the components in a chemical compound one can determine the
empirical formula for that compound. Additionally, if the true molecular mass of the compound is known
the chemical formula of the compound can also be determined. Finally, the experimental results of the
experiment will be used to confirm the law of conservation of matter.
Procedure
1. Heat a clean, dry porcelain crucible with cover on a clay triangle (supported on a ring stand), using a
direct flame, for about 5 minutes. This is to make certain that the crucible is dry. The bottom of the
crucible should glow red hot toward the end of the heating.
2. Turn off the burner, and let the crucible and cover cool. Leave the crucible and cover resting on the clay
triangle.
3. Weigh the empty, dry crucible and cover using an electronic balance. Record the mass to 3 decimal
places.
4. Cut a strip of magnesium ribbon that is about 35 cm long. If it is not shiny, polish it briefly with steel
wool.
5. Being careful not to touch the ribbon too much, fold or coil the magnesium ribbon so that it will fit
inside the crucible. Don't coil it so tightly that it will be difficult for the magnesium to react completely.
Wrapping the ribbon around a pencil, pen or glass stir rod works well.
6. Carefully reweigh the crucible and cover with the magnesium ribbon inside.
7. Put the cover aside for a moment, and start heating the crucible with the magnesium in it. Have the
cover close by, with some crucible tongs ready to handle it with. The instant that the magnesium starts
to burn, put the cover on the crucible (using the crucible tongs) to put out the fire.
8. Continue heating the covered crucible for a minute or so, and then take off the cover again. Wait for the
magnesium to catch fire again, and then quickly re-cover the crucible.
9. Repeat steps 7 and 8 until the magnesium no longer catches fire when the cover is removed.
10. When the magnesium no longer catches fire, heat the sample strongly for 5 minutes, with the cover on
the crucible. Make sure the bottom of the crucible becomes red hot. Let the crucible cool until it is
warm to the touch.
11. To the cooled crucible, add about 10 drops of distilled water. Make sure to wet the entire surface of
the sample, not just one spot. This process will help decompose any magnesium nitride that was
produced in the process. A faint smell of ammonia is not unusual.
12. Warm the crucible using a gentle flame for about a minute to drive off the water. Heat it moderately
strongly for about 10 minutes. The cover can be left off of the crucible, and it does not have to become
red hot.
13. Let the crucible and contents cool until the crucible is barely warm to the touch.
14. Weigh the cooled crucible, cover, and contents of the crucible, recording this total mass.
20
Data/Calculations
1. From the data collected, calculate the mass of magnesium and oxygen consumed in the reaction.
2. Calculate the moles of magnesium and oxygen used in the reaction. (Show calculations.)
3. Determine the empirical formula for magnesium oxide. (Show calculations.)
4. Given that the true molar mass of magnesium oxide is 40.3 g/mol, what is the chemical formula for
magnesium oxide? (Show calculations.)
5. Which reactant is the limiting reagent? Explain your choice.
6. Based on your choice, calculate the maximum amount of magnesium oxide that could theoretically be
produced.
7. Calculate the percent yield of your reaction. If your percent yield is less than 90%, explain possible
sources of error.
𝑎𝑐𝑡𝑢𝑎𝑙 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑝𝑟𝑜𝑑𝑢𝑐𝑡
(
) 𝑥100
𝑡ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑝𝑟𝑜𝑑𝑢𝑐𝑡
Conclusion
Report the chemical formula for magnesium oxide and the percent yield for the reaction. Also include
the explanation if the percent yield is less than 90%.
21
5
Determination of Calcium in Milk
Prelab Questions:
1. What is the formula for dicarbon dihydride?
2. Write the balanced equation for dicarbon diydride reacting with oxygen to form carbon dioxide and
water.
3. Calculate the grams of carbon dioxide produced from the reaction of 25 g of each of the two
reactants. Which of the reactants is the limiting reagent? Explain.
4. If 45.5 g of KMnO4 is dissolved in 150 mL of distilled water, what is the molarity of the KMnO4?
5. If 5.5 mL of the above solution is removed, how many grams of KMnO4 are present?
Introduction and Objectives
Calcium is a necessary element for maintaining the structure and biological function of many living
species. It is necessary to be able to verify the proper concentration of such species within food products
and dietary supplements. One particular method for determining concentrations of species in aqueous
solutions is the time tested method of a titration. By adding a specific amount of a titrant that reacts with
the analyte in a well defined chemical reaction, it is possible to determine the amount of analyte present.
In this particular experiment a rather large titrant molecule, EDTA – ethylenediaminetetraacetic acid, is
used to determine the amount of calcium in a solution of milk. An indicator must also be used to determine
when a sufficient amount of EDTA has been added because all of the species of interest are colorless. A
relatively large amount of base is added to the reaction to promote the best environment in which the
EDTA can function. The reaction of interest is given below:
EDTA + Ca+2  EDTA-Ca complex
The most important point to realize in the reaction is that the EDTA and calcium react in a 1:1 mole
ratio. Without an understanding of the reaction stoichiometry, the entire titration process is meaningless.
Procedure
1. Determine the mass of a 10-mL graduated cylinder before and after adding about 10 mL of milk to a
125 or
250-mL flask. Calculate the exact mass of milk used in the experiment.
2. Place the milk and 50 mL of deionized water in a 125 ml flask.
3. Add 5 mL of pH 10 buffer to the milk in the flask and wait 5 minutes.
4. Add about 0.10 g of the powdered indicator, Calmagite. The mixture should appear red-pink.
5. Fill a pipet buret with 0.050M EDTA solution and record the initial volume of the pipet. (RECORD
THE VOLUME AS IT APPEARS ON THE MICROBURET.) If the volume is between 4 and 5, it is read
as 4.XX NOT 5.XX
6. Titrate the milk sample until the red color changed to blue. This signals the endpoint. (Some practice
may be required to detect this change in color of the skim milk sample.)
7. Record the final volume of EDTA solution as it appears ON THE MICROBURET.
8. Repeat the experiment.
Data/Calculations
1. Use the mass and density of milk (1.02 g/mL) to determine the exact volume of milk used in each trial.
Convert mL to oz.
2. Using the exact molarity of the EDTA solution (FROM THE CONTAINER) and the volume of EDTA
required for the titration (found by subtracting the initial volume of the microburet from the final
volume), calculate the moles of Ca in the milk sample in each trial. Remember the 1:1 mole ratio
between EDTA and Ca.
22
3. Calculate the milligrams of calcium present in the skim milk sample in each trial.
4. Using the results from (3) and the exact volume of milk in (1), calculate the
mg Ca
oz milk
. Finally, calculate
the milligrams of Ca present in an 8 oz serving of the milk? Calculate the average Ca in 8 oz from your
two trials.
5. The calcium values stated on the milk container label are given as a percent Daily Value (%DV) of Ca in
an 8 oz serving (based on 1000 mg being 100 %DV of Ca). How many milligrams of Ca does this
represent?
6. What is the % difference (see Exp 1) between your average titration value and the value from the milk
container?
Conclusion
Report your average milligrams of calcium per 8 oz serving and the percent difference with the container
value.
23
6
Investigations of the Ideal Gas Law
Prelab Questions:
1. What is the mass (in grams) of 1 atom of P?
2. What is the name of H3PO3?
3. Complete and balance the reaction for H3PO3 with NaOH.
4. If 9.25 g of NaOH is dissolved in 75 mL of water, how many mL of 0.350M H3PO3 would be required
to neutralize the NaOH?
5. What volume (in mL) of 4.25M H3PO3 is required to produce 150 mL of 0.350M H3PO3?
Introduction and Objectives
The goals of this experiment are (1) to determine the molar mass of a gas, butane, (2) to discover the
relationships between pressure and volume and pressure and temperature, and (3) to determine the value
of absolute zero in degrees Celsius.
In this experiment, the ideal gas law will be used to determine the molar mass of butane gas obtained
from a disposable lighter. The mass of the lighter will be determined before and after a particular volume
of the gas is dispensed into a graduated cylinder. The ideal gas law given below will be used to determine
the moles of gas collected in the graduated cylinder from the volume, pressure, and temperature of the gas.
PV = nRT
In order to explore the relationships of temperature and volume with pressure, a Texas Instruments TI83 calculator will be connected to Vernier pressure and temperature sensors via a TI calculator-based
laboratory (TI-CBL) interface. Two graphs will be constructed with the data in Parts 2 and 3: (1) Pressure
vs. Volume, (2) Temperature vs. Pressure.
Part 1. Molar Mass of Butane
Procedure
1. Weigh the lighter. Record the mass.
2. Fill a large beaker about three-quarters full with tap water.
3. Attach a large three-prong clamp to a 100 mL graduated cylinder.
4. Fill the graduated cylinder to the top with water.
5. Cover the cylinder COMPLETELY with your hand, invert and immerse under the water surface in a large
beaker or other appropriate vessel.
6. Position the cylinder near the wall of the beaker. Position the lip 1-2 cm below the water surface.
7. Hold the modified lighter under the opening of the cylinder; press the release valve to release gas.
Collect between 85 and 95mL of gas.
8. Remove the lighter; shake off excess water; pat dry with a paper towel.
9. Adjust the internal and external water levels until they are the same; record the volume of gas in the
cylinder.
10. Measure and record the temperature of the water. Record the barometric pressure
11. Dry the lighter thoroughly. Canned air may also be used to dry the lighter. Weigh the thoroughly dried
lighter. Record the mass.
Data/Calculations for Butane
1. What mass of gas was collected?
2. What volume of gas was collected?
3. What was the temperature of the gas collected? (Same as the water in the cylinder.)
4. What was the pressure of the gas collected? (Patm = Pgas + Pwater vapor) (The water vapor pressure is found
on the chart on the wall. It is based on the temperature of the water. The value of Patm is the barometric
pressure.)
24
5. From the ideal gas law, how many moles of gas were collected? Show calc.
6. What is the molar mass of the gas? Show calc.
7. If the chemical formula for butane is C4H10, calculate the % diff between your value and the real value in
(6).
Part 2. Procedure for P vs. V Experiment
Procedure
1. Below sketch the expected graph of pressure (Y-axis) versus volume (X-axis). Explain the shape of the
line in your graph. Be sure to include an initialed graph in your lab journal.
2. Setup the CBL system and connect the pressure sensor to the CBL unit. Launch the DataMate software
to collect data using the “event with entry” mode. Your instructor will help you.
3. Connect the syringe to the pressure sensor with the plunger pulled out to 2-mL. Activate the CBL and
begin collecting data start with the initial value with the plunger at the 2-mL mark. Pull the plunger
back in 2-mL increments recording data at each value until you reach the 12-mL mark. Add 0.2 mL to
each volume to account for the extra tubing volume.
Data/Calculations for P vs. V
1. Using Excel make a graph of Pressure (Y-axis) vs. Volume (X-axis). Make sure the graph axes are
labeled. Add the appropriate best-fit line (linear or power) to the data with the equation and R2 value.
25
Part 3. Procedure for P vs. T & Absolute Zero
Procedure
1. Below sketch the expected graph of pressure (X-axis) versus temperature (Y-axis). Explain the shape of
the line in your graph. Be sure to include an initialed graph in your lab journal.
2. Have your instructor help you set up the CBL system to collect temperature and pressure data
simultaneously.
3. Collect data every 5 degrees until the temperature reaches about 70oC.
Data/Calculations for P vs. T & Absolute zero
1. Using Excel make a graph of Pressure (X-axis) vs. Temperature (Y-axis). Make sure the graph axes are
labeled. Include the appropriate best-fit line (linear or power) through the data, the equation of the
line, and the value of the correlation coefficient.
2. Determine the value for absolute zero (in oC) as obtained from your data.
Conclusion
Report your value for the molar mass of the gas in Part 1, the value of the exponent on X in equation given
for the best-fit line in Part 2 and how this compares to the value that you would expect to obtain and your
value of absolute zero in part 3.
26
7
Finding the Calories in Food Products
Prelab Questions
1. What is the formula for sodium dihydrogen phosphate monohydrate?
2. What volume of 5.68M HCl is required to make 450 mL of .750M HCl? What volume of water is
required?
3. What is the molarity of HCl if 56.0 mL of 0.350M HCl is mixed with 43.2 mL of 0.610M HCl?
4. What is the pressure if 18.2 g of O2 is heated to 30oC in a 500 mL sealed container?
C4H10 + O2  CO2 + H2O
5. In the reaction above, what is the pressure of CO2 produced in a 750mL flask at 25oC if 15.1 g of
C4H10 is reacted with excess oxygen?
Introduction and Objectives
The goal of this experiment is to determine the amount of energy stored as heat in various food
products. The amount of heat will be determined in joules and then converted to the typical food calorie.
This will be an indirect determination of the heat energy of the food because measurements of temperature
change will be recorded for water. The following equation will be used to determine the heat energy of the
food:
Heat released by the food = - (Heat absorbed by the water + Heat absorbed by the
calorimeter)
qfood = -(qwater + qcalorimeter)
qfood = -(mwaterswaterTwater + mcalscalTcal)
where q is heat energy in joules (J), swater = 1.00 cal/goC, and scal = 0.215 cal/goC.
For the conversion of the units of joules to the food unit of Calories: 1 Food Calorie (C) = 1000 calories (c).
Procedure
1. Determine the mass of the aluminum soda can. (Ensure that it is clean and dry.) Measure out about
100 mL of water and add it to the soda can. Determine the mass of the soda can plus the water.
2. Place a thermometer into the soda can so that the bulb is submerged in the water but that it is NOT
touching the sides or the bottom of the can. Let the water stand in the can about 5 minutes and then
record the temperature of the water.
3. While you are waiting for the temperature of the water to equilibrate, carefully place a piece of the
walnut onto a wire stand and determine the initial mass of the walnut plus the stand. Make sure that
the walnut on the stand is about 1 cm from the bottom of the soda can BEFORE you ignite it.
4. Ignite the walnut while it is underneath the soda can. CAREFULLY use the thermometer to stir the
water in the can. Continue stirring and record the highest temperature of the water.
5. After the water has attained its highest temperature, record the final mass of the walnut with the stand.
6. Repeat this process with a marshmallow.
Data/Calculations
1. What was the mass of the food product that reacted?
2. What was the mass of the water in the can?
3. What was the change in temperature for the system?
27
4. How much heat energy was absorbed by the water? (See the equation above.)
5. How much heat was absorbed by the calorimeter? (See the equation above.)
6. What was the total amount of heat energy released by the food in food calories (C)?
7. What was the total amount of heat energy released by the food in food calories per gram of food (C/g)?
8. How many food calories per gram are obtained from info on the food package?
9. Calculate
C / g wa ln ut (exp)
C / g marshmallow (exp)
and
C/g
C/g
wa ln ut ( actual)
marshmallow ( actual)
10. Calculate the percent difference for the ratio values in #9.
Conclusion
Report your values for # 9 and 10 above.
28
8
Percent of Oxygen in the Air
Prelab Questions
1. In the reaction below, what is the pressure of CO2 produced in a 750mL flask at 25oC if 15.1 g of
C4H10 is reacted with 42.5 g O2 at 35oC?
C4H10 + O2  CO2 + H2O
2. What is the name for Ni(NO2)3?
3. What is the empirical formula for a compound composed of 2.1%H, 52.8% As and 45.1% O?
4. What volume of liquid X is required to make 65 mL of 0.785M solution of X? (The molar mass of X
is 60.5 g/mol, and the density of X is 1.23 g/mL.)
5. If the ΔHrxn for the reaction shown in problem #1 is -325 kJ, how much heat energy (kJ) is released
when 25.0 g of C4H10 is reacted?
Introduction and Objectives
In this experiment a very simple apparatus will be employed to determine the percent of oxygen in the
atmosphere. The reaction vessel will be a test tube containing steel wool inverted into a container of water.
The oxygen reacts will react with the steel wool according to the following reaction:
4Fe + 3O2 + 2xH2O  2Fe2O3xH2O
As the steel wool reacts with the oxygen, hydrated iron (III) oxide (rust) is produced and water will rise
into the test tube to replay the reacted oxygen. This amount of water must be carefully collected. The
percent of oxygen in the air is the same inside and outside of the test tube. To determine the percentage of
oxygen inside the test tube, a ratio will be taken of the volume of water collected in the tube after the
reaction to the volume of the entire tube.
Procedure
PART A – Water Method
1. Fill a medium size beaker about ¾ full with tap water. Record the temperature of the water.
2. Determine the mass of the pieces of steel wool you were given. Remember which piece is used for
which part.
3. Spray one piece of the steel wool with vinegar 2-3 times on the front and back side to cover the wool
but NOT soak it in vinegar.
4. THOROUGHLY dry it with a paper towel. QUICKLY, place the steel wool into the test tube ensuring that
it is not forced all the way to the bottom of the test tube. It should be as spread out as possible to
ensure the greatest surface area that will lead to a shorter reaction time.
5. QUICKLY invert the tube placing the mouth of the tube under the surface of the water. Leave it
undisturbed until the water level in the tube ceases to rise.
6. When the water level in the tube has stopped rising, raise or lower the tube in the beaker until the level
of the water inside and outside of the tube are equal and measure the height of the column of water
inside the tube
7. From its height, calculate the volume of the water. (Hint: What is the geometric shape of the water in
the tube?)
8. Determine the volume of the empty test tube.
Data/Calculations
1. Why does the water level rise in the tube?
2. Using the volume of water and volume of the tube, calculate the percent of O2 in the atmosphere.
3. Using the equation at the right, determine the actual percent of oxygen. (PT is the barometric pressure
from the barometer. Pwater is the partial pressure of water vapor from table on the board.) DO NOT
change the percent into a decimal.
𝟐𝟎. 𝟗𝟒% (
𝑷𝑻 − 𝑷𝒘𝒂𝒕𝒆𝒓 𝒗𝒂𝒑𝒐𝒓
)
𝑷𝑻
4. What is the % difference between the experimental and actual values?
29
Procedure
PART B – CBL Method
1. With the help of your instructor set up a TI-CBL system with a pressure sensor to be ready to receive
the test tube filled with steel wool once it is properly prepared. Collect data points every 15 seconds
for 15 minutes.
2. Repeat the preparation of the steel wool as in Part A. Place the steel wool into a clean test tube and,
with the help of your instructor, connect it to the pressure sensor and begin collecting data.
3. Once the data collection is finished, determine the pressure at the start of the experiment and at the
conclusion of the experiment. This information will be used to determine the percent of oxygen in the
air. Show how you calculated the % O2 from the graph and include a graph of pressure (y-axis) versus
time (x-axis) in your lab journal. Reference the journal article: J. Chem. Educ. 82(2), 2005, p.286-87 for
help if necessary.
Data/Calculations
1. Use the beginning and ending pressure to determine the percent of oxygen in the tube. (Consider the
following: If the change in pressure went from 1 atm to 0.60 atm, what would the % difference be? Think
about the process that you used to determine your answer and how it can apply to this situation.)
2. What is the % difference between the experimental and actual values?
Conclusion
Report the values for the percent of oxygen in both parts along with each percent difference from the
accepted value and the answers to the questions below.
Questions – Use data from Part A to answers these questions.
1. Use the ideal gas law (PV=nRT) to determine the moles of oxygen present in the test tube. (You must use
the partial pressure of oxygen as P. A good approximation of this is found by multiplying the barometric
pressure by the percent of oxygen found in the air. The volume, V, is the volume of the empty test tube, and
the temperature, T, is the water temperature.) Using the balanced equation, convert the moles of oxygen
to grams of Fe necessary for complete reaction. From the result determine if there was sufficient steel
wool in the test tube to react with all of the oxygen. REMEMBER UNITS. Show your calculations.
2. Considering the density of the steel wool (7.87 g/cm3 - iron) and the amount of steel wool used, explain
why the volume of iron can be ignored in calculating the total volume of air inside the test tube.
(Consider how the actual volume of iron compares to the volume of the test tube.)
30
9
A Chemical Formula by Microscale Titration
Prelab Questions
1. Name Mn2O7.
2. Balance the reaction: AgBr + Na2S2O3  Na3[Ag(S2O3)2] + NaBr
3. If 16.5 g of Na2S2O3 is reacted with excess AgBr and 10.5g Na3[Ag(S2O3)2] are produced, what is the
percent yield?
4. If for the reaction above the value of Hrxn = -248 kJ, what is the value of Hrxn for the following
reaction?
½ [ Na3[Ag(S2O3)2] + NaBr  AgBr + Na2S2O3 ]
5. What volume (in mL) of 0.35M HNO3 is required to completely neutralize 5.85 mL 0.268M Al(OH)3?
Introduction and Objectives
There are two main objectives in this experiment. The first is to become familiar with the classical wet
chemistry technique of titration. However, this titration will be done on a microscale to decrease the
amount of chemical waste produced in the process. A diagram of the microburet made from a pipet is
given at the right. Regardless of the scale the principles of titration are the same. The second objective is to
gain further experience in determining the empirical and chemical formula for a compound. Remember for
the titration that there is a one-to-one ratio between the HCl added and the hydroxide ion produced by the
calcium reaction with water.
Procedure
1. Weigh out a piece of calcium metal between 0.10 to 0.20 g and place it into a 250-mL
Erlenmeyer flask containing about 100 mL of DI (deionized) water. Swirl the flask
until all of the calcium has dissolved. One of the products of this reaction is
hydrogen gas that is observed as bubbles produced as the calcium dissolves.
2. Assemble the micropipetting equipment as shown by your instructor and fill it with
1.0M HCl. Note that when reading the pipet or a typical buret it is always read from
the top down and using the bottom of the meniscus as the reference point. If the
bottom of the meniscus falls between the 3 and 4 mL mark, the volume is read as
3.xx. One should also take care in making sure that when reading the meniscus the
eyes are level with it to avoid volume measurement error.
syringe
pipet
3. When the reaction of the calcium and water is complete, five drops of 0.1% thymol
blue indicator should be added. The solution should turn a pale blue color. At this
point, the solution is ready to be titrated. Be sure to record the initial volume of the
pipet. The titration should proceed SLOWLY and with constant swirling until the solution turns yellow.
The yellow color indicates that all of the hydroxide produced in the reaction has been neutralized by
the HCl. The solution should remain yellow for about 30 seconds. If the titration is conducted to
quickly, the solution could turn red which indicates that a large excess of the HCl has been added. The
solution should be discarded. The entire experiment should be repeated to duplicate the results.
Data/Calculations
1.
Mass of the calcium metal used
g
2.
Initial pipet reading
mL
3.
Final pipet reading
mL
4.
Volume of HCl used
mL
5.
Concentration of HCl
M
31
Calculations
6.
Moles of calcium metal used
mol
7.
Moles of HCl used
mol
8.
Moles of OH- produced by the calcium
mol
Divide (6) and (8) by the smallest value between (6) and (8) to determine the empirical
formula
9.
Empirical formula for calcium hydroxide
10.
If the true molar mass is 74.1 g/mol, the chemical formula is
Conclusion
Report the empirical formula and chemical formula determined for the calcium hydroxide.
32
10 Atomic and Molecular Spectroscopy
Prelab Questions
1. What is the name for P4O10?
2. Using only the density of silver (10.5 g/cm3) and its molar mass, approximate the volume of an
atom of silver.
3. If a car gets 35 mi/gal, what is its mileage in km/L?
4. What is the Lewis structure for SO2?
5. For the reaction below, what is the maximum amount of POCl3 produced from 42.3 g P4O10 and
15.3mL of 0.875M HCl? Which reactant is the limiting reagent?
HCl + P4O10  POCl3 + H2O
Introduction and Objectives
The main objective of this experiment is to examine the basic principles of absorption and emission of
light from atoms and molecules. Different atoms and molecules appear different colors when they are
excited to emit light or absorb visible light. This is due in part to the differences in energy levels to which
electrons transition during absorption or excitation.
Procedure
Fill out the charts and answer the questions in the data/calculation section as you are instructed in lab.
Data/Calculations
PART 1A
Calibration Curve from Helium (He)
Carefully record the position of each color line seen in the
spectrometer.
Use Excel to make a graph of wavelength (x-axis) vs. distance (yaxis) and place the best-fit trend line through your data with the
equation for the line and correlation coefficient.
 (nm)
cm
Color
447
468
492
501
588
668
707
33
Part 1B
Wavelength Determination for Other Gases
Carefully record the position of each color line seen in the spectrometer and use the equation from part
1A to determine the wavelength for each of the colors of visible light seen in the two lamps.
Spectral data for Mercury (Hg)
Color
ruler
cm
Wavelength
(nm)
Frequency
(Hz)
Energy
(J)
Frequency
(Hz)
Energy
(J)
Frequency
(Hz)
Energy
(J)
Spectral data for Hydrogen (H2)
Color
ruler
cm
Wavelength
(nm)
Spectral data for Xenon (Xe)
Color
ruler
cm
Wavelength
(nm)
34
PART 3
Fluorescence Emission
Use the spectrometer box that you used in Part 1 to determine which, if any, of the elements from the gas
discharge lamps are present in BOTH an old and new fluorescence lamp. Remember that the emission
spectrum for an element is a finger print. ALL of the particular wavelengths for an element MUST be
present for the element to be present. There may well be other wavelengths present due to other elements
for which you are not looking, but all wavelengths for a particular element must be present if the element is
present. You must show that all of the wavelengths are represented for an element in the lights to be able
to justify its presence.
Color
ruler
cm
Wavelength
(nm)
Which of the gases, if any, are confirmed to be present in the fluorescent lights?
Conclusion
No specific conclusion is expected for this experiment.
35
11 Molecular Geometry
Prelab Questions
1. Is either SF4 or OF4 plausible molecules? Explain.
2. What is the formula for iron (III) hydrogen carbonate?
3. What mass (in g) of the solid is need to make 750 mL of 0.890M iron (III) hydrogen carbonate ?
4. What is the wavelength (in nm) for a photon having a frequency of 5.76 E 14 Hz? Does this photon
fall into the visible spectrum?
5. Using the information in #4, what is the energy (in J) per photon? What is the energy per mole of
photons?
Introduction and Objectives
The objective of this experiment is to explore a variety of molecular configurations. It will be necessary
to bring and use the molecular model kits used in class to construct each of the molecules in this lab. The
instructors will check your accuracy of each model and initial that you have successfully completed its
construction.
Procedure
Complete the chart seen below.
Data/Calculations
Formula
Lewis Structure
Molecular
Geometry
Polar or
nonpolar
Hybrid. of
central atom
Initials
H2S
SF5Br
SeO2
NH3
CCl3H
BrF3
36
Formula
Lewis Structure
Molecular
Geometry
Polar or
nonpolar
Hybrid. of
central atom
Initials
XeI3Cl
BF3
PCl5
KrFCl
SF2H2
IF5
CH3CH3
CH2CH2
CHCH
CH3SH
NH2CH3
37
12 Paper Chromatography
Prelab Questions
1. What volume (in mL) of 0.890M KMnO4 contains 2.000 g Mn?
2. If a chemical reaction is endothermic requiring photons of 250 kJ/mol to produce the reaction.
What is the minimum wavelength of radiation possessing sufficient energy?
3. Use the thermochemical equations shown below to determine
the enthalpy for the reaction: CH2O(g) + N2(g) + 3H2(g)N2H4(l) + CH4O(l)
N2H4(l) + H2(g)  2NH3(g)
H=-90KJ
NH3(g)  1/2N2(g) + 3/2H2(g)
H=150KJ
CH4O(l)  CH2O(g) + H2(g)
H=-325KJ
4. In outer space, there are only about 5 atoms per cubic centimeter. Given that the temperature is
3K, what is the pressure (in atm) in outer space?
5. Complete and balance the following: TiI4 + H3PO2 What is the name of the product containing P?
Introduction and Objectives
All chromatography is based the principle of “like interacts with like”. In any chromatographic
technique, the separation occurs between the stationary phase and the mobile phase. Here the stationary
phase is a coffee filter and the mobile phases are water, ethanol and acetone. In this experiment, the paper
chromatography will be used to examine the affect of mobile phases with differing polarity on separation of
pigments in black ink.
Procedure
Part 1
1. Obtain a piece of chromatography paper. Draw a straight line across the
bottom of the piece of paper with the black marker given to you. See the
diagram at the right.
2. Place the piece of paper into a 100-mL beaker into which you have
carefully added about 20 mL of water. Make sure that there is NO water
on the inside walls of the beaker above the water. Carefully position the
paper so that the tip with the line is in the water BUT SO THAT THE
LINE IS NOT IN THE WATER.
3. Wait for the chromatograph to develop. When it is finished, record the
colors that you observe from the top of the paper down to the line.
Part 2
4. Obtain several pieces of chromatography paper. Treating them one at
time, draw a straight line across the bottom of the piece of paper with
one of the black markers as in Part 1.
5. Place the piece of paper into a 100-mL beaker into which you have
carefully added about 20 mL of water as in Part 1. Wait for the
chromatograph to develop. When it is finished, record the colors that
you observe from the top of the paper down to the line.
6. Repeat the process substituting 20 mL of water with 20 mL of methanol
and finally with 20 mL of acetone.
38
Data/Calculations
1. What is the order of colors for each chromatogram? Sketch each chromatography.
2. What is the one color that is significantly different in each?
3. If you are told that water is the most polar and acetone is the least polar of the mobile phases used,
what can you conclude about the polarity of the molecules producing the color that you listed in #2?
39
13 Freezing Point Depression
Prelab Questions
1. What is the electron configuration for Fe?
2. Which of the quantum numbers ultimately defines the size of an atom – n, l, ml or ms?
3. What is the percent of mercury in Hg2Cl2?
4. How many grams of Na2HPO4 are required to produce 100 mL of 0.565M Na2HPO4?
5. Complete and the following reaction: NaH2PO4 + NaOH. How many grams of NaOH are required to
react with 76.5mL of 0.325M NaH2PO4.
Introduction and Objectives
Colligative properties are those that depend on number of molecules dissolved in a solvent. The
presence of these foreign molecules incr eases the boiling point of the solvent (boiling point elevation) and
decreases it’s freezing point (freezing point depression). The expression used for freezing point depression
is:
ΔTf = i∙Kf∙m
where Tf is the change in freezing point, i is the number of particles the solute molecule dissociates into as
it dissolves in the solvent, Kf is the freezing point depression constant for the solvent, and m is the molality
of the solution. Remember that molality is the measure of the moles of solute per kilogram of solvent.
Procedure
Part 1
1. Weigh the 25 x 150 mm test tube in a small beaker.
2. Add 10 mL of the solvent, cyclohexane, to the test tube and reweigh to determine the precise mass of
the solvent.
3. Accurately weigh 0.30 g of the solute, naphthalene, onto a piece of weighing paper and save for later
use.
4. Continuously agitate the pure cyclohexane with the temperature probe until the solution is frozen.
Remove it from the ice bath and allow it to begin to melt and record the temperature of the
disappearance of the VERY LAST crystals. This will be the freezing point of the pure cyclohexane (Tf).
5. Refreeze the solution, and repeat the process to check the freezing point temperature. If the two trial
results do not agree to within  0.10 oC, repeat the process until they do.
Part 2
1. Using the same solution as in part 1, CAREFULLY add the naphthalene (solid) to the cyclohexane
(liquid). Use temperature probe rod to COMPLETELY dissolve the naphthalene BEFORE proceeding.
2. Pour off most of the water from the ice bath and add about 20 mL of salt. Fill the bath mostly with ice
and then just enough water to cover the ice. Stir for about 2 minutes.
3. Place the test tube into the ice bath and determine the freezing point of this new solution as in Part 1.
Again, looking for the disappearance of the last crystals.
Data and Calculations
Mass of cyclohexane (g)
Mass of cyclohexane (kg)
Freezing point of pure cyclohexane (oC) (Part 1)
avg. f.p.
Mass of nathphalene (g)
Freezing point of mixture (oC) (Part 2)
avg. f.p.
40
1. What is the value for ΔTf for the cyclochexane solutions?
2. Given that kf for cyclohexane is 20.0 oC/m and naphthalene does not dissociate in cyclohexane.
Calculate the moles of Naphthalene used in the experiment.
3. What is the molar mass of naphthalene?
4. If the chemical formula for naphthalene is C10H8, what is the percent difference between your
experimental and theoretical values of the molar mass?
41
14 Determination of a Chemical Rate Law
Prelab Questions
1. Balance the following equation: CCl4 + HF  CCl3F + CF2Cl2 + HCl
2. What is the name for the two reactants? (Name the one containing F in its aqueous form.)
3. With a pressure of CCl4 of 1.25 atm at 27oC in a 450 mL flask and excess HF, what mass of gaseous
CCl3F (in grams) could be produced?
4. Suppose that in problem #3 instead of excess HF there was 15.5 mL of 0.658M HF, what mass of
gaseous CCl3F (in grams) could be produced?
5. What is the electron configuration for Mo+3?
Introduction and Objectives
The field of Kinetics deals with reaction rates and Thermodynamics deals with reaction concentrations.
Determining the rate law of a reaction will give information as to how fast a particular reaction will occur
but will not yield information as to how much of a product will be formed. In addition, the rate law can be
used to help determine the reaction mechanisms – the set of elementary steps that occur during the
conversion of reactants to products. The following reaction will be studied in this experiment (the main
reaction is in bold).
blue complex
Starch
+
H2O2
+
2I-
+
S4O6-2
+
2H+
I2
+
+
2S2O3-2
2H2O
Recall that the generic rate law for the reaction can be written as rate = k[H2O2]x[I-]y[H+]z . A series of
reactions will be run to determine the value of x, y, z, and k for this experiment.
A set of four experimental runs will be made according to the table found in the procedure section.
Steps are made to ensure that the concentration of only one species at a time is being changed in the
experiment. The time for the experiment must be carefully measured to determine the rate under each set
of conditions. The time will be measured from the addition of the last component (H2O2) until the solution
changes color to a very deep blue. The color change is a result of the depletion of the S2O3-2, thiosulfate, and
the subsequent production of excess iodine that reacts with the starch to produce an intensely blue colored
complex.
Procedure
Add the reagents found in the table below IN ORDER from left to right and into CLEAN containers. Be
sure to constantly AND consistently stir the reaction and begin timing when the final component is
added. Record the time in seconds.
Water
(mL)
1
2
3
4
75
80
50
60
0.05M
HAc/NaAc
Buffer
30
30
30
30
Vinegar
0.8 M Hac
0
0
0
15
0.05 M
KI
(mL)
25
25
50
25
0.1%
Starch
(mL)
5
5
5
5
0.05 M
Na2S2O3
(mL)
5
5
5
5
0.8 M
H2O2
(mL)
10
5
10
10
pH
[H+] = 10-pH
4.75
3.75
42
Data/Calculations
1. Calculating Rate:
rate 
8.33 E -4
Time ( sec onds)
Recall from class that the relative rate is
[ H 2O2 ] 1 [ S2O32 ]
.

t
2
t
From the reagent table the total change in concentration for the thiosulfate is
0.00167M because it is the limiting reagent and will be completely used in the
process.
Mixture
Time
(sec)
Rate
(M/sec)
1
2
3
4
2. Finding x: use experiments 1 and 2.
rate 2  k[H 2 O 2 ] 2x [I - ] 2y [H  ] z2
rate 1  k[H 2 O 2 ]1x [I - ]1y [H  ]1z
rate 2  [H 2 O ]
rate 1  [H 2 O ]
x
2 2
x
2 1
The terms containing I- and H+ cancel because
they are constant in those two experiments.
 rate 2 

log 
 rate1 
So, x is found from x 
 [H O ]
log  2 2 2
 [H 2 O 2 ]1



Calculating [H2O2]: MiVi = MfVf - The dilution equation: using the initial volume of H2O2 (Vi) [from
the table]; the total solution volume (Vf) [sum of all volumes from the table], and the initial molarity
[from the table] of the H2O2 (Mi).
[H2O2]1
[H2O2]2
Value of x
43
3. Finding y and calculating [I-] from exp. 1 and 3: Repeat the process using KI information to find the
values of y. Use the same equations as in #2 above to calculate the rate and concentrations for iodide.
[I- ]1
[I-]3
Value of y
4. Finding z and calculating [H+] from exp. 1 and 4: To find the [H+] values for “z”, use the equation,
[H+]=10-pH (Use the 10x key on your calculator.) Put the [H+] directly into the log equation. DO NOT use
the dilution equations.
[H+ ]1
[H+ ] = 10-pH
[H+]4
[H+ ] = 10-pH
Value of z
5. Calculating k: Use all of the concentrations and rate from reaction #1 to calculate k.
Value of k
Conclusion
Write a brief conclusion including the complete rate law expression with the experimental values of x,
y, z, and k: rate = k[H2O2]x[I-]y[H+]z
44
15 Determining Order of Reaction
Prlab Questions
1. What is the molar mass of C25H30N3Cl?
2. What is the % of N in the compound in #1?
3. The average bond length for C=C is 1.33E-8 cm. What is this length in pm?
4. If 15.5 mL of 0.356M NaOH is required to react with 8.75 g of H2X, what is the molar mass of X?
5. What is the hybridization of the central atom in KrI3+?
Introduction and Objectives
The purpose of this experiment is to determine how the rate of a reaction depends on the concentration
of one of the components. The reaction of interest is between crystal violet (CV+) and sodium hydroxide:
C25H30N3+ + OH(CV+)

purple
C25H30N3OH
(CV-OH) colorless
As the reaction proceeds, the sodium hydroxide will react with the crystal violet to fade the color of the
solution. The concentration of the crystal violet will be tracked verses time using a TI-CBL data acquisition
system. The absorbance of the solution (at 565 nm) will be used in place of concentration in the
calculations because absorbance is directly proportional to concentration.
Various concentration verses time graphs will be made to determine the order of the reaction with
respect to the blue food coloring. One of the following graphs will determine the order of the reaction. As a
reminder:
If zero order:
If first order:
If second order:
A plot of Absorbance vs. time
will yield a straight line with a
negative slope.
A plot of ln Absorbance vs.
time will yield a straight line
with a negative slope according
to the following equation:
A plot of 1/Absorbance vs.
time will yield a straight line
with a positive slope according
to the following equation:
At = - kt + Ao
ln At = -kt + ln Ao
1/At = kt + 1/Ao
Procedure
1. Mix 5 mL of 2.5 E –5 M crystal violet with 5 mL of 0.1M NaOH.
2. Fill a cuvette containing a stir bar about ¾ full with this solution.
3. Place the cuvette in the colorimeter and begin collecting the data in 30 second intervals for about 8
minutes.
4. Have a lab assistant check your data. Once approved, copy the data from the calculator lists into your
journal.
45
Date/Calculations
1. In Excel, construct the following table that can be taped into your lab book. See the example below.

The Excel command to calculate the value of the natural log (ln) of absorbance is to place the
cursor in cell C2 and type: =ln B2

The Excel command to calculate the value of the inverse (1/Abs) of absorbance is to place the
cursor in cell D2 and type: =1/B2
2. From the data, make 3 graphs: Absorbance vs. time, ln Absorbance vs. time, and 1/Absorbance vs. time
(with time always being the x-axis). Once all three graphs are made they should be printed and neatly
taped into the lab journal. Each graph should have a best-fit "LINEAR" curve, equation, and R2 value.
Conclusion
Write a brief conclusion paragraph indicating the order of the reaction with respect to the crystal violet.
Be sure to justify your answer using the graphs.
46
16 Determination of a Keq
Prelab Questions
1. A 0.9% NaCl solution has a density of 1.0 g/mL. What is the molarity of this solution?
2. What is the molality of the solution in #1?
3. Which of the following concentration units is not sensitive to temperature: molarity, molality,
percent by mass?
4. From the table of information below determine the rate law for A and B in the generic reaction
(include the proper units on k):
A + B  C + D
Initial [A] (M)
Initial [B] (M)
Rate (M/s)
0.25
0.40
1.32 E -2
0.50
0.40
2.55 E -2
0.25
1.20
0.120
5. If A and B react in a 1:1 mole ratio, would it be possible for the reaction above to occur in one step
as written?
Introduction and Objectives
The equilibrium constant is a thermodynamic quantity that allows the determination of the extent of a
chemical reaction. Using the equilibrium constant it is possible to calculate the concentrations of reactants
and products present at equilibrium for a particular chemical reaction at a particular temperature. The
reaction of interest here is:
xFe3+ + ySCNFex(SCN)yn
The product (a complex ion) formed in this reaction is intensely colored red. Its concentration will be
determined using a spectrophotometer, and the other concentrations will be determined using simple
subtraction. The desired equilibrium expression is:
K 
[Fe x (SCN) y n ]
[Fe 3 ] x [SCN  ] y
Part 1 – Product Ion Formula
The experiment has two parts. The first part entails the determination of the values of x and y. These
values are necessary before the equilibrium constant, K, can be determined in part two. The values of x and
y will be determined by a spectrophotometric method called Job’s method. In Job’s Method, a graph is
made of the mole fraction of the metal verses the absorbance of the reaction solution. Remember that mole
fraction is defined as:
X Fe 
moles of Fe
moles of Fe  moles of SCN
From the graph of absorbance (Y-axis) vs. mole fraction of iron (X-axis), the mole fraction yielding the
maximum absorbance will be determined. The ratio of ions in the complex can be determined from this
mole fraction value. The ratio of maximum absorbance represents the ideal mix of ions to produce the
maximum amount of complex. Suppose the maximum absorbance appears for the mole fraction of 0.333.
The fraction equivalent to 0.333 is 1/3. From the mole fraction equation, this must represent:
1
1 mole of Fe

3 1 mole of Fe  2 moles of SCN
The formula for the complex ion would be Fe(SCN)2 and the values of x and y would be 1 and 2,
respectively. This general procedure can be used to determine the values of x and y.
Procedure
1. Use 7 clean and dry small beakers and mix the following amounts of each solution from the table below
into the specific beakers. Wait 10 minutes to allow for complete reaction.
2. Record the absorbance for each solution at 470 nm and make a graph of absorbance vs. mole fraction.
47
Data/Calculations
beaker
KSCN
0.0020 M
(mL)
Fe(NO3)3
0.0020 M
(mL)
1
5
0
2
10
2
3
8
4
4
6
6
5
4
8
6
2
10
7
0
5
Mole Fraction of
Fe
Absorbance
Write the balanced chemical equation and explain how the stoichiometric coefficients were determined.
Part 2 - Equilibrium Constant
All three of the concentrations found in the equilibrium expression below must be determined in order
to calculate the value of the equilibrium constant, K. Only the absorbance information from beakers 3, 4,
and 5 will be used for this part of the experiment. In order to determine the concentrations, use the
relationships below.
K 
[Fe x (SCN) y n ] eq
x
y
[Fe 3 ] eq
[SCN  ] eq
Data/Calculations - SHOW ALL CALCULATIONS
[Fex(SCN)yn]eq The concentration of the complex ion will be determined from Beer’s law: A = bC
where A is the measured absorbance; , is the molar absorptivity and has a value of 4500 M1cm-1 @ 470 nm; b is the cell pathlength and has a value of 1.00 cm; and C is the
concentration. Beer’s law must be solved for C.
Beaker 3
Beaker 4
Beaker 5
48
[Fe+3]eq
The free iron(III) concentration can be found by subtracting the concentration of the iron in
the product ion from the total concentration of iron added to the solution.
[Fe 3 ] eq 
3


 0.0020 M  mL Fe in the bea ker  - [Fe x (SCN)y  n ]


12 .0 mL


Beaker 3
Beaker 4
Beaker 5
[SCN-]eq
The free thiocyanide concentration can be found by subtracting the concentration of the
thiocyanide in the product ion from the total concentration of thiocyanide added to the
solution.
[SCN  ] eq 


 0.0020 M  mL SCN in the bea ker  - [Fe x (SCN)y  n ]


12 .0 mL


Beaker 3
Beaker 4
Beaker 5
Value of K
beaker
3
4
5
Avg. value of K
Conclusion
Write a short conclusion paragraph including the balance chemical equation for the iron thiocyanate
reaction and the average value of K for the reaction.
49
17 CO2 in Breath: An Acid-Base Reaction
Prelab Questions
1. For reaction 2 in the experiment below, what is the maximum amount (in grams) of Na2CO3 that can
be produced if 18.5 g of each reactant is mixed together?
2. What is the Lewis structure for SF5-1? What is the name of this ions geometry?
3. What type of intermolecular force operates between molecules of water and I2?
4. If 3.5 E 10 atoms of Pb are found in 1.0 mL of tap water, what is the molarity of Pb in the water?
5. Write the expression for Keq for reaction 1 (rxn. 1) below.
Introduction and Objectives
This experiment will allow the exploration of an acid-base reaction oddly enough by using carbon
dioxide in exhaled breath. Industrially, among other things carbon dioxide is used to carbonate beverage.
In doing so, it increases the acidity of those solutions as it dissolves in water to produce carbonic acid.
CO2(g) + H2O(l)  H2CO3(aq) (rxn. 1)
In this experiment, the carbonic acid produced is reacted with sodium hydroxide:
H2CO3(aq) + 2NaOH(aq)  Na2CO3(aq)_ + H2O(l) (rxn. 2)
The net result of the reaction is that the pressure inside the reaction vessel will decrease proportionally
to the amount of carbon dioxide that is within it. It is this change in pressure produced by the acid-base
reaction which can be used to calculate the percentage of CO2 in exhaled air.
Procedure
1. A simple diagram of the apparatus for the experiment can be
seen in Figure 1. Three pieces of glass tubing (about 10 cm
each) are placed into a 2-hole #8 rubber stopper fitting into the
mouth of a 500-mL flask.
2. Fill a small plastic bottle with approximately 10 mL of 6M NaOH
and carefully lower it into the flask with tongs WITHOUT
SPILLING. CAREFULLY place a plastic pop bottle cap on top of
the container to cover it.
3. CAREFULLY slide a straw into a piece of rubber tubing on the
glass tubing.
4. Inhale a normal breath and then exhaled through the straw and
into the flask. This procedure was repeated twice more. It was
assumed at this point that the gas in the flask was fully
composed of exhaled air because the average tidal volume of an
adult male or female is approximately 500 mL.
5. Quickly BUT CAREFULLY clamp the straw’s rubber tubing and
the rubber tubing on the second glass tube.
6. The third glass tube should be connected to a Vernier gas
pressure sensor through a short rubber tube and Luer lock
connector.
7. The CBL should be setup to collect data in the “time graph”
mode. Collect data for 500 seconds at a rate of 1 point every 10
seconds.
8. A few data points should be taken at the start of the experiment to measure a baseline value. Then the
glass weighing bottle holding the NaOH should be tipped over inside the flask releasing the NaOH.
9. Swirl the flask occasionally until the data collection is finished.
50
Data/Calculations
1. Devise a method using the initial and final pressure to calculate the percentage of carbon dioxide in
exhaled breath. (Think about what is contributing to the initial pressure values and what is not
contributing to the pressure values at the end of the experiment.)
2. Compare your value with the accepted value of 5%. If your value is < 3% or > 7%, repeat the
experiment.
Conclusion
Report the percent of carbon dioxide your group determined in exhaled breath.
51
18 Basic Acid-Base Titrations
Prelab Questions
1. What is the name for LiHCO3∙4H2O?
2. If 15.6 g of NaOH are reacted with excess CH3CO2H (monoprotic), how many grams of H2O is
produced?
3. Is the following reaction endothermic or exothermic? HCl + NaOH  NaCl + H2O + 85 kJ
4. What is the electron configuration for Cu?
5. What is the hybridization of the O in H2CO? (Note: C is the central atom in the molecule.)
Introduction and Objectives
Additional experience with acids and bases will be gained in this experiment by performing titrations.
The first titration will be a direct determination of the amount of acetic acid present in vinegar. The
titration will be performed using sodium hydroxide as the titrant. The reaction is:
NaOH + CH3CO2H
CH3CO2Na + H2O
The second titration will involve an indirect method for the determination of the amount of calcium
carbonate in an acid tablet. Calcium carbonate is the active ingredient in antacid tablets used in this
experiment. The calcium carbonate will be reacted with a known excess amount of HCl. The unreacted HCl
will be titrated with NaOH and the grams of calcium carbonate will be calculated. The reactions for the
second part of the experiment are:
2HCl + CaCO3
CaCl2 + H2CO3
HCl + NaOH
NaCl + H2O
Procedure
Part 1
1. Fill a cleaned buret with 0.300 M NaOH. Ensure that the tip of the buret is filled before performing the
titration.
2. Pipet 3 mL of vinegar into a 250-mL flask followed by 20 mL of water and 1 drop of phenolphthalein.
3. Place a piece of white paper under the flask. Record the initial volume on the pipet buret and proceed
with the titration with SLOW addition of the NaOH and constant swirling. Stop the titration AT THE
FIRST sign of the ENTIRE solution becoming a very faint shade of pink. Record the final volume on the
pipet buret. If you add too much NaOH, you must repeat the experiment.
4. Repeat the experiment to get two sets of accurate data.
Part 2
1. Place an antacid tablet in two different 250-mL flask and add 20 mL of DI water followed by 50 mL of
0.30 M HCl. NOTE: The acid must be PRECISELY added and recorded.
2. Place the flask on a hot plate and heat until the effervescence stops. Boil for 1-2 more minutes.
3. Add 2 drops of phenolphthalein indicator. The solution should turn colorless.
4. Again fill the buret with 0.100 M NaOH and proceed with the titration until the entire solution becomes
faintly pink.
5. Repeat the titration on the second tablet.
52
Data/Calculations
Part 1 - Percent of Acetic Acid in Vinegar
1. What volume of vinegar was used, and, considering that the density of vinegar is 1.05 g/mL, how many
grams were used?
2. From the volume and molarity of NaOH was used to reach the titration’s end point, calculate the moles
of NaOH required for the titration. (Recall that the titration volume is the initial pipet volume minus the
final pipet volume as recorded directly off of the pipet.)
3. Calculate the grams of acetic acid present in the vinegar.
4. Calculate the percent of acetic acid in the vinegar.
5. Calculate the % diff. in the experimental and commercial percentages of acetic acid.
Part 2 – Mass of Calcium Carbonate in an Antacid Tablet
1. Calculate the moles of HCl initially added to the flask.
2. From the volume and molarity of the NaOH, calculate the moles of NaOH required for the titration.
3. The NaOH reacted with excess HCl – not required for reaction with the antacid tablet. Calculate the
moles of excess HCl.
4. From the information in (1) and (3), calculate the moles of HCl that reacted with the calcium carbonate
in the antacid tablet.
5. Calculate the moles, grams, and milligrams of CaCO3 present in the antacid tablet based on the moles of
HCl required that reacted with the CaCO3 in the tablet.
6. Calculate the % diff. between the experimental and commercial values of the milligrams of calcium
carbonate present in the antacid tablets.
Conclusion
Write a brief conclusion paragraph including the average values of the percentage of acid in vinegar
and the average milligrams of calcium carbonate in the antacid tablets.
53
19 Determination of Ka for Acetic Acid
Prelab Question
1. What is the pH of a solution made by mixing 25.0 mL of 0.450M HCl and 35.5 mL of 0.650M HNO3?
2. Propane, C3H8, reacts with O2 to produce CO2 and H2O. If a typical gas grill tank holds 20 pounds of
propane, what volume of CO2 gas is produced at 25oC and 1 atm pressure?
3. In the picture below, suppose the container on the left is 12 L and holds He at a pressure of 10 atm
and that the container on the right is 20 L and holds N2 at a pressure of 15 atm. The temperature of
the whole system is 25oC. What will be the pressure when the valve in the middle is opened?
4. What is the Lewis structure for CNS-1?
5. Which of the following is/are polar molecules (first atom is central atom)? BrF5, CO2, CH2Cl2
Introduction and Objectives
In this experiment, two procedures will be explored for determining the dissociation constant, Ka, for a
weak acid – acetic acid. The reaction of interest is:
CH3CO2H  CH3CO2- + H+
so that the equilibrium constant expression is:
Ka 


[CH3CO2 ][H ] [Ac ]eq[H ]eq

[CH3CO2H]
[HAc]eq
Therefore, in order to calculate the value of Ka, it would be necessary to determine each of the three
concentrations appearing in the expression at equilibrium. Alternatively, a graph from the HendersonHasselbalch can be used to determine Ka.
 [Ac  ]eq
pH  pK a  log
 [HAc] eq





A graph of pH versus the log term will be made to determine the pKa for acetic acid. In the second part
of the experiment, the value of Ka will be determined from a titration curve.
A pH meter will be used in each solution to determine the concentration of hydrogen ion. The key
component in a pH meter is the sensing electrode. The electrode is filled with a 1.0 M HCl solution so that
the only difference between the solution inside the electrode and outside is the concentration of acid. The
difference in acid concentration produces a difference in electrical potential or voltage across the very
fragile glass membrane at the bottom of the electrode. It is this potential difference that is measured and
converted into pH. EXTREME care should be taken not to break the glass bulb.
Part 1. Ka Determination
Procedure
Obtain the specified amounts of the following solutions. Carefully place the electrode down into the
solutions and gently swirl until the pH reading stabilizes. Record the pH for each solution.
Solution 1
0.1 M NaAc – 25 mL
0.1 M HAc – 5 mL
Solution 2
0.1 M NaAc – 10 mL
0.1 M HAc – 20 mL
Solution 3
0.1 M NaAc – 20 mL
0.1 M HAc – 10 mL
Solution 4
0.1 M NaAc – 5 mL
0.1 M HAc – 25 mL
54
Data/Calculations – Part 1
1.
AcpH
[Ac - ]eq 
HAc
-
[Ac ]  VAc VTotal
[HAc]eq
[HAc]  VHAc

VTotal
log
[Ac- ]eq
[HAc]eq
Sol’n 1
Sol’n 2
Sol’n 3
Sol’n 4
2. Use the data to make a graph of log term (from Henderson Hasselbalch) (x-axis) versus pH (y-axis).
What is the value of pKa for acetic acid based on your graph? Include the graph in your lab journal.
Part 2. Titration Curve
Procedure
Using a CBL pH sensor, perform a titration of acetic acid with NaOH as the titrant. After the data has
been collected, make a graph of pH versus volume of NaOH added. Determine the equivalence point of the
titration and then the corresponding pH value at ½ Veq. At the volume, the pH should equal the pKa value
for the acetic acid. Place a well-labeled titration curve in your lab book.
1. Add 50 mL of water, 2 mL of vinegar, and 1 drop of phenolphthalein to a 100 mL beaker.
2. Titrate the acid solution with 0.30M NaOH using a pipette as a microburet. Record the pH with every
0.50mL addition of NaOH.
3. Record the pH at which the solution changes color.
Data/Calculations – Part 2
1. Transfer the data to MS Excel and make a graph of the titration curve and put it into your lab journal.
2. Look at your graph keeping in mind the units for the x axis and the y axis. What was the volume of
NaOH required to reach the equivalence point (Veq) for the titration?
3. What is the volume of NaOH at ½ of Veq?
4. What is the pH at ½ Veq? How can this be used to determine the pKa of the acid?
Conclusion
Write a brief conclusion paragraph including the graphical value of pKa for acetic acid from part 1 and
the value of pKa determined from Part 2. Compare the closeness of the two values.
55
20 Determination of H3PO4 in a Cola Beverage
Introduction and Objectives
Many academic acid/base titrations involve only monoprotic acids and bases. However, experience
with polyprotic systems is also of value because species including carbonic acid, phosphoric acid and
proteins are such systems. In this experiment, a titration will be determined on the phosphoric acid
present in a common cola beverage. The results obtained here will be compared to those previous
published in literature.
Procedure
1. Determine the volume of a drop of NaOH from the pipet buret.
2. Boil about 150 mL of soda on a hot plate for about 10 minutes.
3. Make a sketch of what you expect the pH curve for the titration of H3PO4 with NaOH to look like.
Explain the shape of the curve. Include your initialled sketch in your lab journal.
3. Quantitatively add 25 mL of cooled soda to a 50 mL beaker and titrate with 0.020M NaOH using a CBL
to record pH versus NaOH volume. Record the pH after the addition of every 5 drops. You will be using
the “event with entry” mode of the CBL. For the graph, convert the number of drops to volume
measurements for each data point.
Data/Calculations
1. Transfer the titration data to MS Excel and make a graph of pH (Y-axis) vs. volume of NaOH (X-axis).
Include the graph in your journal AND explain the shape of the graph. Explain also how it is the same
and different from your sketch.
2. From the volume and molarity of the NaOH, determine the moles of NaOH required to reach the first
end point.
3. Calculate the molarity of the phosphoric acid in the soda.
4. Calculate the % diff. in the experimental molarity of the phosphoric acid in the soda and the reported
amount from a research article in the Journal of Chemical Education (J. Chem. Educ. Volume 60,
Number 5, May 1983, p420-1). Carefully read the results section for the data obtained from the
titration curve.
Conclusion
Report the molarity of phosphoric acid calculated from your experimental results and the percent
difference with the literature value in the J. Chem. Educ. paper.
56
21 Ca(IO3)2 : Solubility Product - Ksp
Introduction and Objectives
Another equilibrium constant of chemical interest is the solubility product, Ksp. This equilibrium
constant is useful in determining the extent to which a particular solid will dissolve in pure water or a
solution containing one or both of the dissociation products. In addition, if one is doing qualitative analysis
work, knowledge of this constant allows the proper concentration of precipitating ions to ensure positive
results for the analysis species.
The expression for Ksp is stated as any other equilibrium constant. For this experiment the reaction of
interest is:
Ca(IO3)2 (s)  Ca+2(aq) + 2IO3-(aq)
and the Ksp expression is:
Ksp  [Ca 2 ][IO3 ]2
Notice that the value of the solid calcium iodate is NOT included in the expression and that the value of the
iodate ion is squared.
The value of Ksp will be determined by an iodometric titration of the iodate ion. The concentration of
calcium ion is simply ½ that of the iodate. Once these two concentrations are known the calculation of Ksp
is trivial. The reaction scheme occurring during the titration is:
IO3- + 5I- + 6H+
I2 + 2S2O3-2
3I2 + 3H2O
2I- + S4O6-2
In this reaction, the iodate ion is the limiting reagent and is used to produce iodine (I2). The iodine is then
titrated with a solution of thiosulfate ion (S2O3-2) of known concentration. By knowing the mole-mole ratio
of each of the components, it is possible to calculate the original concentration of the iodate.
Procedure
1. Obtain and thoroughly clean a microburet. Fill the buret, including the tip, with the 0.050M thiosulfate
ion solution. Record the initial volume in the buret.
2. Obtain 2 mL of the 1% starch indicator in a graduated cylinder and keep it for later in the experiment.
3. Into a 250-mL Erlenmeyer flask add the following in the following order: 25 mL of DI water, 1 g KI
(ensure this dissolves before proceeding), 5 mL of Ca(IO3)2 (this volume should be pipetted), and 5 mL
of 1 M HCl. At this point the solution should be a dark brown color.
4. Begin the titration by slowly adding the thiosulfate solution and with constant swirling. Continue
adding until the solution turns a VERY pale yellow. DO NOT add so much so that the solution becomes
clear and colorless.
5. Once the solution turns pale yellow, add the starch indicator. The solution’s color should become a
dark blue. DO NOT RECORD ANY VOLUME INFORMATION AT THIS POINT. Continue the titration
DROPWISE until the solution color turns from blue to colorless. It should only take a very small
amount of titrant to produce this last color change.
6. Record the final volume of the thiosulfate titrant and its molarity and repeat the procedure.
Data/Calculations
1.
Molarity of the thiosulfate titrant solution
M
2.
Initial buret volume
mL
3.
Final Buret volume
mL
4.
Total volume of thiosulfate titrant used (in L)
L
57
5.
Moles of thiosulfate titrant used
mol
6.
Moles of iodine (I2) present in solution
mol
7.
Moles of iodate in original sample
mol
8.
Volume of iodate solution used (in liters)
L
9.
Molarity of the iodate in the original sample
M
10.
Molarity of the calcium in the original sample
M
11.
Experimental value of Ksp
Conclusion
Write a brief conclusion paragraph including the experimental value of Ksp for Ca(IO3)2.
58
22 Determination of Water Hardness
Prelab Questions
1. For the reaction: H2 + I2  2HI if the initial concentration of H2 and I2 is 0.25M what is the
concentration of HI at equilibrium. Keq = 5.2
2. What is the pH of a 3.5E-3M HNO3 solution?
3. Balance the following equation: Cu(s) + H2SO4(aq)  CuSO4(aq) + H2O(l) + SO2(g)
4. Write the expression for Keq for the reaction in #3.
5. For the reaction in #3, if 28.5 g of Cu is reacted with 35.8mL of 0.985M H2SO4, how much (in g)
CuSO4 can be produced?
O
O
HO
C
CH2
HO
C
CH2
N
O
CH2 C
CH2
CH2
OH
EDTA
N
CH2 C
OH
O
Introduction and Objectives
The composition of natural waters is primarily a result of the kinds of rocks and soil over and through
which they flow. In many sections of the US, water flows over a significant amount of rock composed of
calcite and dolomite containing calcium and magnesium ions. It is primarily these ions (but also includes
other ions such as iron) that cause water to be “hard”. This classification results from the reaction of these
ions with soap molecules leading to the formation of hard deposits or “soap scum”. Soft water is devoid of
significant amounts of these ions and thus soap scum does not form.
In order to determine the amounts of calcium and magnesium in water another titrimetric experiment
will be performed. The titrant for this experiment is a chelating agent, EDTA, ethylenediaminetetraacetic
acid. Chelate means “claw” and in essence that is what EDTA does to metallic anions present in solution. It
forms a very strong complex ion with the metal ions by making 6 bonds with them. The formula for these
complex ions is always a 1:1 ratio of EDTA to metal ion. All metal ions will complex with EDTA so it is
important to mask out other metal ions. The pH of the solution in this experiment is set to 10 so that any
other metal ions will precipitate as the hydroxide. The only complexing cations remaining are calcium and
magnesium. It has become standard practice to consider the sum of calcium and magnesium in
determining the hardness of water. Hardness is expressed in terms of the concentration of calcium
carbonate (CaCO3) with the unit of measure of parts-per-million or ppm.
An indicator is used in this experiment like in acid/base titrations. In this case, the indicator
(calmagite) is a very weak chelating agent that reacts with any calcium and magnesium present. When
these ions are present, the indicator is red in color. The EDTA first reacts with the free calcium and
magnesium in solution. However, when that has all be complexed, it is able to extract those ions from the
indicator because it is a stronger complexing agent. In its uncomplexed form the indicator color is blue.
The reaction is relatively slow so the titration should be carried out relatively slow with very careful but
vigorous swirling so that the end point is accurately observed.
In order to soften water, it is often placed through a filter system containing an ion-exchange resin
system. The resin is composed of many tiny plastic beads whose surface has been coated with either a
positive ion (ex. –NR3+) or negative ion (ex. –SO3-). There is always electrical neutrality so that they positive
ions are often neutralized with either hydroxide or chloride and the negative ions are neutralized with
hydrogen ion or sodium ion. When hard water is allowed to flow over cation-exchange resin, a reaction
such as the following occurs:
2Resin–SO3-H+ + Ca+2  2Resin–SO3-Ca+2 + 2H+
Here the hard ions are traded or “exchanged” for the soft ions (hydrogen ions). In this particular example,
only cations are exchanged hence the name cation-exchange resin. The resin having a positive ion coating
is known as anion-exchange resin.
A sample of tap water will be analyzed for its hardness and then treated with a cation-exchange resin.
A sample of the treated water will be analyzed for its hardness to observe the effectiveness of the resin.
59
Procedure
Hardness
1. Obtain and clean a 10-mL microburet. Fill the buret including the tip with 0.0050M EDTA. Record
initial volume on the microburet.
2. Place 25 mL of tap water in a 250-mL Erlenmeyer flask from the container provided by the instructor.
Add 5 mL of buffer and 3 drops of the EBT indicator. At this point, the solution color should be red. If it
is not, consult your instructor.
3. Begin the titration by slowly and carefully adding the EDTA into the flask. Carefully observe the color
change from red to blue. Once the solution turns blue, immediately stop the titration and record the
final volume on the microburet.
4. Repeat this procedure with a second 25-mL sample of tap water.
Softening
1. Place 300 mL of the tap water provided by the instructor into a 600-mL beaker. Obtain a sample of the
cation-exchange resin from your instructor.
2. Vigorously swirl the resin in the water for 10 minutes.
3. Carefully decant off 25 mL of the treated water. Add 5 mL of the buffer and 3 drops of the EBT
indicator.
4. If the solution color is blue all of the "hard" ions have been removed. (Hardness = 0 ppm).
5. If the solution color is purple or red, record the volume on the buret and add the 0.0050M EDTA
DROPWISE until the solution changes color. Record the final buret volume.
6. Repeat this procedure with a second 100-mL sample of treated water.
Data/Calculations
Hardness
1. From the volume and molarity of the EDTA, calculate the moles of EDTA required to reach the end point
of the titration.
2. Calculate the moles and grams of CaCO3 present in the water. (Remember that the EDTA is titrating the
Ca+2. However, water hardness is based on the amount of CaCO3 present. What is the mole:mole ratio
between Ca+2 and CaCO3? )
3. Calculate the parts-per-million (ppm) of CaCO3 present:
g CaCO3
x 1E6  ppm
g H2O
Softening
1. Repeat the same set of calculations for the samples of softened water.
2. Calculate the % diff between the ppm of CaCO3 in the hard versus softened water.
Conclusion
Write a brief conclusion paragraph including the average values of hardness before and after softening
the water with the ion-exchange resin.
60
23 Paper Chromatography – Metal Ions
Introduction and Objectives
It is often necessary to determine the composition of a solution. In this experiment, a formerly used
technique called paper chromatography will be used to determine the composition of a solution of metal
ions. In later experiments, other techniques will be used to determine the composition of solutions
containing both cations and anions. Remember that chromatography is a technique based on like
molecules interacting with like molecules. The components separate from one another based on their
differing affinities for the mobile phase and stationary phase used for the chromatography. In this
experiment, paper is used as the stationary phase and a solution containing HCl is the mobile phase.
Procedure
1. Obtain a piece of the chromatography paper and draw a line 2 cm from the bottom along one of the long
sides and another line about 2 cm from the top of the long side. THE LINES MUST BE DRAWN IN
PENCIL.
2. Place 6 X marks evenly spaced ensuring that the outer two are at least 2 cm from the sides.
3. Underneath each spot place one of the following designations: Cu+2, Fe+3, Ni+2, Co+2, UKN 1, and UKN 2.
4. Using the capillary tubes in each of the sample solutions spot the paper as you are instructed with each
of the solutions on its corresponding X. REMEMBER that the intent is for the spots to be as SMALL and
concentrated as possible. If one spot bleeds into another, the paper should be discarded.
5. Repeat this process at least three times and let the instructor exam the paper before proceeding.
6. When the paper is ready, the instructor will staple it into a cylinder shape.
7. Obtain one of the developing chambers having 10 mL of solution in it consisting of 9 mL of acetone and
1 mL of 6M HCl.
8. Place the paper into the developing chamber with the metal ions at the bottom. Quickly replace the
plastic wrap lid and allow the chamber to remain undisturbed until the mobile phase is at the 2 cm
mark at the top of the paper.
9. When the developing is completed, remove the paper from the chamber and QUICKLY mark all visible
spots by circling them with a pencil.
10. Some spots may still not be visible. Place the paper (still in its cylinder shape) into the ammonia
chamber to allow the spot colors to become even more pronounced. The paper should only have to be
in the ammonia chamber about 15 – 20 seconds.
11. Remove the paper and again mark any spots that did not appear originally.
12. If there are any spots that you are still uncertain of use the following table of color enhancing agents to
find the spots.
ion
reagent
Fe+3
0.2M KSCN
Ni+2
0.1M NaHDMG
spot
color
blood red
bright pink
Co+2
sat’d KSCN +
acetone
blue-green
Cu+2
0.2M K4[Fe(CN)6]
red
13. Using the color and position of the known metal ions, ID the composition of the two unknown solutions.
14. Place the original chromatograph or a copy of it in your lab book.
Data/Calculations
Include the real or a sketch of the chromatograph in your lab book and calculate the RF value for each
spot in each solution.
Conclusion
Write a brief conclusion paragraph including the identity of each component in the unknowns.
61
24 Separation of Group I Ions: Silver, Mercury & Lead
Prelab Questions
1. What is the pH of a 0.455M Benzoic acid solution?
2. If 125 g of benzoic acid (C6H5CO2H) is mixed with 255 g of sodium benzoate (C6H5CO2Na) in 800 mL
of distilled water, what is the pH?
3. What is the formula for barium nitrate nonahydrate?
4. Sketch the graph of a pH curve for the titration of a weak acid with a strong base. Indicate on the
graph how the pKa for the weak acid can be determined.
5. What pressure of O2 at 45oC in a 600 mL container would be required to completely react with
38.0 g C8H18? Reaction:
C8H18 + O2  CO2 + H2O
Introduction and Objectives
In addition to instrumental techniques, it is possible to perform a series of wet chemical experiments to
determine a material’s or solution’s composition. A relatively simple qualitative analysis scheme is
presented here and in the next few experiments for the determination of the ionic components of unknown
solutions. Three experiments will focus on cations, and one will focus on anions. The final experiment will
involve repeating the experiment in its entirety to determine the composition of a cumulative unknown
solution.
It is extremely important that the work area be kept clean and that the proper chemicals are used at the
proper time in the analysis scheme. Unclean glassware or improper chemicals could require that the entire
experiment be restarted which could mean a loss of significant lab time.
Both the written procedures and flow chart will be beneficial in helping keep all of the glassware and
information in order. Note that the numbers in bold on the flow chart correspond to the numbers in the
written procedure. It is vitally important that the procedure is carefully followed so as to avoid false
positive or negative results. The grade for this experiment will solely rest on the correctness of the ions
identified as being a part of the assigned unknown. Each unknown will have a specific letter assigned to it.
Record the letter assigned for the unknown. Without the identification of your unknown, you will
receive only 20% for the lab.
Procedure
The separation will first be performed on a known solution containing all of the ions. This will
demonstrate the positive result expected for each ion. First obtain 1 mL of each of the known solutions in a
10 mL graduated cylinder, mix thoroughly, and put 1 mL into a test tube for analysis.
Lastly, procedure will be repeated on 1 mL of an unknown solution of the ions be studied in this
experiment.
1. To the 1 mL test solution, add 3 drops of 6 M HCl. Stir the solution vigorously but carefully with a clean
stir rod and then centrifuge. Test for complete precipitation by adding in 1 drop of the HCl. If the
solution becomes cloudy, add 1 more drop of HCl and recentrifuge. Retest for complete precipitation.
After precipitation is found to be complete, the liquid layer (supernate) may be discarded FOR THIS
EXPERIMENT. NOTE: For the cumulative unknown analysis, the supernate from step 1 must be carried
into the next experiment as it has the remaining ions in it.
2. Add 3 mL of water (always DI water) to the test tube containing the precipitate from step 1. Vigorously
stir the contents in the water and place the tube in a beaker of boiling water. Continue to stir the
mixture in the boiling water for about 5 minutes. Next, let the precipitate settle (centrifuge if
necessary) and decant (pour off) the supernate into a second test tube. Add a second 3 mL of water to
the precipitate and repeat. However, this second 3 mL portion of water is to be discarded.
3. Add 1 drop of 1 M K2CrO4 to the supernate from step 2. If lead is present a yellow precipitate of PbCrO4
will form. The final confirmation of the presence of lead is produced by the addition of NaOH. Add 1020 drops of 6M NaOH. If the yellow precipitate dissolves TO ANY DEGREE to produce a clear
yellow/orange solution, the precipitate was due to lead. This step is necessary because other ions form
a yellow precipitate with chromate. However, only lead redissolves with the addition of the NaOH.
62
4. To the solid precipitate remaining from step 2, add 1 mL of 6M NH3 and then 2 mL of water.
Thoroughly mix with a clean stir rod and centrifuge. A dark (gray-black) precipitate is a mixture of
Hg2Cl2 and Hg(l) which proves the presence of Hg2+2 in the solution. Pour off the supernate into another
clean and dry test tube.
5. Add 6M HNO3 to the supernate from step 4 until the solution is acid (as tested against litmus paper). A
small excess of acid will be required. The formation of a white precipitate (that can turn purple when
exposed to sunlight) proves the presence of silver ion.
Data/Calculations
The only data necessary for this section is a detailed description of the positive result for each of the
ions and a statement of the content of the unknown.
Conclusion
Report ONLY the identity of all of the ions in your unknown and the number of your unknown
solution.
Flow Chart
63
25 Separation of Group II Ions: Iron, Aluminum & Zinc
Prelab Questions
1. What is the ΔG of a reaction if its ΔH is 234 kJ/mol and its ΔS is 160 J/molK at 25oC ? Is the reaction
spontaneous? If not, is there a temperature at which it becomes spontaneous? If so what is that
temperature (in oC)?
2. If for the reaction below, 2.0 g of Pb3(PO4)4 are produced by the reaction of 18.5 g of (NH4)3PO4 and
55.8 g of Pb(NO3)4, what is the percent yield for the Pb3(PO4)4?
(NH4)3PO4 + Pb(NO3)4  Pb3(PO4)4 + NH4NO3
3. Calculate the ΔHrxn from the ΔHfo for the reaction: 2SO3(g)  2SO2(g) + O2(g)
4. Which of the following is the strongest acid: benzoic, acetic or formic? Justify your answer.
5. Given the rate law below, suppose the concentration of A is tripled and the concentration of B is cut
in half. If the rate of the reaction was 0.270 M/s originally, what would it be under these new
conditions?
Rate = k[A]2[B]
Introduction and Objectives
See Introduction and Objectives for experiment 24.
Procedure
The separation will first be performed on a known solution containing all of the ions. This will
demonstrate the positive result expected for each ion. First obtain 1 mL of each of the known solutions in a
10 mL graduated cylinder, mix thoroughly, and use 1 mL of the mixture for analysis.
The procedure will be repeated on 1 mL of an unknown solution of the ions be studied in this
experiment. Remember to record the unknown letter.
6. Add 6M NH3 to the 1 mL test solution until the solution become basic to litmus. Then add about 1.5 mL
excess NH3. NOTE: The test solution will either be the one prepared just for this experiment or from
Step 1 in Exp. 24 as seen in the flow chart. Mix the solution thoroughly with a clean glass stir rod and
then centrifuge. Decant the (supernate) liquid layer into another test tube. NOTE: This supernate
contains any zinc ions present. Wash the iron and aluminum precipitate with 2 mL of water. Mix,
centrifuge, and decant the liquid combining it with the previous supernate solution.
7. Add about 1.5 mL of 6M NaOH to the iron and aluminum precipitate. Mix the solution thoroughly. Add
1 mL of water to the solution. Mix, centrifuge, and decant the supernate into another test tube. This
supernate will contain any aluminum ions. NOTE: At this point, there should be 3 test tubes each
containing either a solid or liquid containing one of the analysis ions.
8. With constant stirring, add 3-4 drops of 6M HCl to the test tube containing the Fe(OH)3 - the solid
remaining from Step 7 – to dissolve the Fe(OH)3 precipitate. If the precipitate doesn’t dissolve,
continue adding HCl dropwise with stirring until it does. Add 2 mL of water and 1 drop of 0.1M KSCN.
A positive result for iron will be the appearance of a blood red solution of FeSCN+2.
9. To the aluminum supernate from Step 7, add 6M HCl until it is just slightly acidic to litmus. Add 2 drops
of the aluminon reagent and then add 6M NH3 until the solution is just basic to litmus. Thoroughly mix
with a glass stir rod and let the mixture stand. A pink flocculent (floating, cloud-like) precipitate
indicates the presence of aluminum ion.
10. To the supernate from Step 6, add 6M HCl drop by drop until the solution is acidic to litmus. Add 10
drops of 0.2M K4[Fe(CN)6]. A light green precipitate confirms the presence of Zn. Centrifuge if
necessary to confirm the Zn.
64
Data/Calculations
As in the previous experiment, the only data to be recorded are the detailed observations of result
characteristics. There are no calculations.
Conclusion
Report ONLY the identity of all of the ions in your unknown and the number of your unknown
solution.
Flow Chart
For Cumulative Unknown – these ions
will come from Step 1 in Exp. 24
Fe+3 , Al+3 , Zn+2
6. NH3
Zn(NH3)4+2
Fe(OH)3 , Al(OH)3
7. NaOH
Fe(OH)3
10. K4[Fe(CN)6]
Al(OH)4-
K2Zn3[Fe(CN)6]2
(light green ppt)
other ions
9. HCl
aluminon
8. HCl
Fe+3
8. KSCN
FeSCN+
(red solution)
Al+3 complex
9. NH3
Al complex
(pink ppt)
65
26 Separation of Group III ions: Barium, Calcium & Ammonium
Prelab Questions
1. What is the [H+] in a solution having a pOH = 12.5?
2. A 50 mL solution having a pH=8 is mixed with 50 mL of 0.050M Zn+2, will Zn(OH)2 precipitate?
3. What is the electron configuration for At?
4. What is the hybridization for B in BF3? What is the molecular geometry?
5. Is BF3 polar or nonpolar? Will it H-bond with itself? With water?
Introduction and Objectives
See Introduction and Objectives to Exp. 24
Procedure
The separation will first be performed on a known solution containing all of the ions. This will
demonstrate and show the positive result expected for each ion. First obtain 1 mL of each of the known
solutions in a 10 mL graduated cylinder, mix thoroughly, and use 1 mL of the mixture for analysis. The
procedure will be repeated on 1 mL of an unknown solution of the ions be studied in this experiment.
11. Start with 1 mL of the solution to be tested. NOTE: If the cumulative analysis is being performed, start
with the supernate from Step 10 in Exp. 25. Add about 0.5 mL of 6M HC2H3O2 (HAc), 1 mL of water and
0.5 mL of 3M NH4C2H3O2. Add 3 drops of 1M K2CrO4, mix thoroughly, and centrifuge. If the solution is
not yellow add K2CrO4 dropwise but ONLY until the solution is yellow. DO NOT add an excess amount.
Decant the supernate into another test tube. The supernate will contain any calcium ion. Wash the
precipitate with 2 mL of water. Centrifuge and discard the supernate. A yellow precipitate indicates
the presence of barium ion.
12. To confirm barium, add 3 drops of 6M HCl to dissolve the yellow precipitate, and then add 1 mL of 0.1M
Na2SO4. Centrifuge the mixture. The formation of a white precipitate confirms the presence of barium.
NOTE: The supernate may be yellow, but the precipitate will be white.
13. Use the supernate from Step 12 and add 0.5 mL of 1M K2C2O4. Finally, add 6M NH3 to make the solution
basic to litmus. Let the mixture stand for at least 10 minutes if a precipitate does not form before that
time. White CaC2O4 solid indicates the presence of calcium ion. To confirm the presence of calcium,
decant the colorless to yellow supernate to waste and add 3 drops of 6M HCl and 1 mL of water to
dissolve the solid. Add one drop of 1M K2C2O4 and again make the solution basic to litmus with 6M NH3.
The CaC2O4 precipitate will reappear.
14. Use 1 mL of NH4Cl if analyzing the known set of ions or obtain a fresh 1 mL sample of the unknown
solution. Place the 1-mL sample into an evaporating dish. Moisten a piece of red litmus paper and
attach it to the underside of a watch glass large enough to cover the evaporating dish. Add 1 mL of 6M
NaOH to the evaporating dish and quickly cover it with the watch glass litmus paper side down. The
NaOH causes any ammonium ion to be converted to ammonia gas. The gas diffuses to the litmus paper
where it dissolves in the thin film of water on the paper and is converted into NH4OH turning the entire
piece of paper blue. The color change of the entire piece indicates the presence of ammonium ion.
NOTE: Care must be taken with the evaporating dish. If NaOH solution is spilled or splattered onto the
litmus paper a false positive is possible.
Data/Calculations
As in the previous experiment, the only data to be recorded are the detailed observations of results
Conclusion
Report ONLY the ions present and the unknown number.
66
Flow Chart
For Cumulative Unknown – these ions
will come from Step 10 in Exp. 25
67
27 Analysis for Common Anions
Introduction and Objectives
This experiment will involve analyzing for the presence of some common anions. The primary
difference in this and the previous three experiments is that it will not be necessary to perform individual
separation steps. Each anion has its own unique precipitating or coloring agent with which the other
anions do not interfere.
Procedure
When using the known solutions, 1 mL of each individual solution is analyzed for its specific ion. For
the unknown solution, a separate 1-mL sample of the solution will be required for each of the following
steps. Remember to record the letter of your unknown.
1. Sulfate – SO4-2 – Make 1 mL of the solution SLIGHTLY acid to litmus with 6M HCl. Add 1 mL of 0.1M
BaCl2. The white precipitate of BaSO4 confirms the presence of sulfate.
2. Chloride – Cl- - Add 1 mL of 0.1M AgNO3 to 1 mL of the solution. The formation of a white/purple
precipitate of AgCl confirms the presence of chloride.
3. Phosphate – PO4-3 – Mix 1 mL of 0.5M (NH4)2MoO4 with 1 mL of 6M HNO3. Add 1 mL of the test solution
to this mixture. A yellow precipitate of (NH4)3PO412MoO3 indicates the presence of phosphate. NOTE:
It may take several minutes for the precipitate to develop, and it may be necessary to heat the solution
to 40 oC by running the test tube under hot water.
4. Nitrate – NO3- - Make 1 mL of the test solution acidic to litmus with 3M H2SO4. Add 1 mL of freshly
prepared saturated FeSO4. Incline the test tube to 45o and SLOWLY pour about 1 mL of 18M H2SO4
down the side of the test tube. DO NOT MIX the solution. A brown ring of Fe(NO)+2 at the interface of
the two liquid layers indicates the presence of nitrate.
Data/Calculations
As in the previous experiment, the only data to be recorded are the detailed observations of result
characteristics. There are no calculations.
Conclusion
Write a brief conclusion paragraph including the identity of the ions in your unknown solution and the
letter of the unknown solution.
68
28 Cumulative Qualitative Analysis
Prelab Questions
1. Balance the following redox reaction: P4(s) + Cr2O7-2(aq)  H3PO4(aq) + Cr+3(aq)
2. The ΔHrxn for the following reaction: 2NOCl(g)  N2(g) + O2(g) + Cl2(g) is -51.2 kJ. How much heat
is generated when 45.5 g of NOCl reacts?
3. For the reaction above, if 2.5 atm of Cl2 are produced at 35oC in a 750 mL container, what pressure
of NOCl was required under the same conditions?
4. If 0.550M HX produces a pH = 4.2, what is the Ka of HX?
5. Would the reaction in #2 have a value of ΔS that is positive or negative? Justify your answer.
Introduction and Objectives
This experiment is designed to test skills for performing the entire analysis of cations. Experiments 2527 will be repeated on an unknown sample that could contain all, some, or none of the ions.
Procedure
Obtain two 1-mL samples of cation (one for ammonium ion and one for all the other ions). Remember
to record the number of your unknown. Beginning with Experiment 24, follow the flow charts and
directions for analyzing the summary unknown sample.
Data/Calculations
As in the previous experiment, the only data to be recorded are the detailed observations of result
characteristics. There are no calculations.
Conclusion
Report only the identity of all of the ions in your unknown and the number of your unknown solution.
69
29 Nuclear Chemistry
Prelab Question
1. Balance the following redox reaction for a basic solution:
I¯ (aq) + ClO¯(aq)  I3¯(aq) + Cl¯(aq)
What is the value of Ecell when [I3-]=[Cl-]=.985M, [I-]=[ClO-]=0.025M and pH=9.5.
2. What is the name for NaClO?
3. For the reaction: 2NO2  2NO + O2 the observed rate law is: rate=k[NO2]2. Which of the
mechanisms below (if any) are possible? Explain.
Mechanism I
2NO2  2NO + O2
Mechanism II
NO2  NO + O (slow)
O + NO2  O2 + NO (fast)
Mechanism III
2NO2  NO3 + NO (slow)
NO3  NO + O2 (fast)
4. What mass (in g) of NaClO are required to make 4.50L of 0.982M NaOCl? What volume (mL) of
3.25M NaOCl is required to make 4.50L of 0.982M NaOCl?
5. Which of the following has the highest boiling point? 0.1m NaCl or 0.1m BaCl2 Justify your
response.
Introduction and Objectives
The objectives of this experiment are to reexamine isotopes and become more familiar with how
radioactive isotopes decay. Several lab exercises will reinforce nuclear chemical processes.
Procedure
Complete the following experimental measurements.
Data/Calculations
1. Complete the following reactions. Name each of the species that is not an element.
a) 238U  234Th + _____
b) 234Th  234Pa + _____
c) 238U   + 234Th + 2 _____
d) 201Hg ( ___ ,) 201Au
e) 13N  13C + _____
f) _____ (,n) 234Bk
g) 238U + 12C  _____ + 6n
2. Particle Energy
Devise and test a method using paper, plastic, glass, and lead to determine which of the three
fundamental particles emitted by radioactive elements is the most energetic and the least energetic.
70
3. Isotopes of M&Mium
By knowing the abundance of an element it is possible to determine the average molar mass that
should be placed on the periodic table. For example, element X has three isotopes:
Isotope
X1
X2
X3
Atomic Mass
500
503
504
Abundance
65%
22%
13%
Average molar mass of X = (500)(0.65)+(503)(0.22)+(504)(0.13) = 501.18
Now for the new element M&Mium has 6 known isotopes. Using the following information along with
the abundance that you can determine from the “sample” of M&Mium, calculate its average molar mass.
Isotope
brown
Atomic Mass Abundance
331
Isotope
yellow
Atomic Mass
335
red
332
green
337
blue
333
orange
338
Abundance
Record the other group’s data. Do the packets represent consistent “samples” for M&Mium?
4. All Salt is NOT created equal.
Use the Geiger counter connected to the TI-CBL, determine the difference in radioactivity between table
salt and salt substitute. The voltage being measured is related to the “counter per minute” on the
Geiger counter. Note that a higher voltage corresponds to a higher “counter per minute” value and
therefore a greater radioactivity.
sample
counts per minute
Salt
Salt substitute
Which is more radioactive?
What specific isotope is responsible?
What is the abundance of that isotope?
5. Smoke Detector
Describe how a smoke detector works and what radioactive element is present in smoke detectors.
Conclusion
No specific conclusion is expected for this experiment.
71
30 Qualitative Eo Table
Introduction and Objectives
Oxidation-Reduction reactions are a key classification of chemical reactions. Every element and
chemical species has an innate ability to give or receive electrons. This tendency can be measured against a
standard reference electrode such as the standard hydrogen electrode. The voltage measured in such an
experiment allows the element or compound to be ranked with any other species similarly measured. The
tabulated ranking is called the standard reduction potential (voltage), Eo, table. Species having positive
values of Eo readily undergo reduction or accepting electrons. Species having a negative value would rather
release electrons and be oxidized.
In this experiment, a qualitative table of four species will be developed indicating their relative
tendencies to undergo reduction. The experiment is broken into two parts. In the first part, three metal
ions will be compared and ranked relative to one another. In the second part, hydrogen ion will be
examined and from its reaction results it will be ranked with the metals. The results will be compared with
numeric Eo values.
The reduction reactions of interest are:
Fe+2 + 2e
Fe(s)
+2
Cu + 2e
Cu(s)
Zn+2 + 2e
Zn(s)
2H+ + 2e
H2(g)
The final portion of the experiment will involve determining whether iodine or bromine is the better
oxidizing agent. The reactions of interest are below:
I2 (purple) + 2e
2IBr2 (brown) + 2e
2BrProcedure
Part 1
1. Obtain 6 clean and dry test tubes and put the appropriate metal and 3 mL of the appropriate solution in
each. Follow the table given below.
Test
Sol’n
Solid
Tube
#1
Fe+2
Cu
+2
#2
Fe
Zn
#3
Zn+2
Fe
+2
#4
Zn
Cu
#5
Cu+2
Fe
+2
#6
Cu
Zn
2. Examine each solution VERY carefully for a reaction characterized by a change in the color of the metal
immersed in each solution. If there is a doubt about whether or not a reaction has occurred or will
occur, the test tube sample may be slightly heated over a Bunsen burner. NOTE: One ion will react with
both metals, one ion will reaction with only one metal, and one ion will react with neither of the metals.
3. Record the observations and place the three ions in order from most reactive (highest Eo) at the top and
least reactive (lowest Eo) at the bottom.
72
Part2
1. Obtain 3 clean and dry test tubes and place 3 mL of 6M HCl into each. Add 1 piece of a different one of
the previous metals into its own test tube.
2. Again, examine each test tube for reaction. Record the observations and place the hydrogen ion in the
list ranked according to its reactivity. For example, if it reacts with all three metals, it is the most
reactive (highest Eo) and is placed at the top of the list. If it reacts with only two of the metals, it should
appear second in the list.
Part 3
1. In a clean and dry test tube add 1 mL of hexane, 1 mL of 0.1M KI, and 1 mL of a saturated aqueous
solution of Br2. Mix thoroughly and record the color of the upper (hexane) layer. The reaction of
interest is:
2I- + Br2
I2 + 2Br- (upper layer purple)
2. In a second clean and dry test tube add 1 mL of hexane, 1mL of 0.1M KBr and 1 mL of saturated alcohol
solution of I2. Mix thoroughly and record the color of the upper hexane layer. The reaction of interest
is:
2Br- + I2
Br2 + 2I- (upper layer brown)
Notice that the hexane layer of both test tubes will either be brown or purple. From this
information, determine whether iodine or bromine is the better oxidizing agent.
Data/Calculations
There are no calculations for this experiment. The data will consist of observations of the chemical
reactions.
Conclusion
Write a brief conclusion paragraph including the reactions in order of their oxidation ability and their
accepted Eo value beside each reaction.
73
31 Bleach: A Redox Titration
Introduction and Objectives
In this experiment, an indirection titration will be performed to determine the percentage of the active
ingredient NaOCl, sodium hypochlorite, in bleach. The titration is called indirect because the hypochlorite
will react with a chemical species to produce iodine that actually is titrated during the experiment. This is a
classical wet chemical technique for analysis that takes advantage of oxidation-reduction chemistry. In the
first part of the experiment, the molarity of the titration (thiosulfate – S2O3-2) will be determined and used
in the second part of the experiment to determine the percentage of the hypochlorite (OCl-) in bleach.
The reaction sequence of interest in the first part of the experiment is:
IO3- + 5I- + 6H+
I2 + 2S2O3-2
3I2 + 3H2O
2I- + S4O6-2
and for the second part the sequence is:
ClO- + 2I- + 2H+
I2 + 2S2O3-2
I2 + Cl- + H2O
2I- + S4O6-2
Procedure
Part 1
1. Obtain and thoroughly rinse a 50-mL buret. Fill the buret and its tip with the thiosulfate (S2O3-2)
solution. Record the initial volume of thiosulfate solution in the buret.
2. Using a 25-mL volumetric pipet, add 25 mL of 0.010M KIO3 to a 250-mL Erlenmeyer flask. Following
the KIO3 add, in this order: 25 mL of DI water, 2 g KI (thoroughly dissolve), and 10 mL of 2M H2SO4.
Swirl the flask and immediately begin to titrate the dark brown solution with the thiosulfate.
3. When the solution becomes very pale yellow in color add 5 mL of starch indicator and CONTINUE
titrating. It is NOT necessary to record any volumes at this point. The color of the solution should
become dark blue and require ONLY a FEW drops of the titrant to reach the clear and colorless
endpoint. Record the volume of thiosulfate in the buret at the clear and colorless endpoint.
4. Repeat the process two more times.
Part 2
1. Refill the buret and its tip with the thiosulfate (S2O3-2) solution. Record the initial volume of thiosulfate
solution in the buret.
2. Using a 1-mL volumetric pipet add 1 mL of bleach to a 250-mL Erlenmeyer flask. Following the bleach
add, in this order,: 25 mL of DI water, 2 g KI (thoroughly dissolve), and 10 mL of 2M H2SO4. Swirl the
flask and immediately begin to titrate the dark brown solution with the thiosulfate in the buret.
3. When the solution becomes very pale yellow in color add 5 mL of starch indicator and CONTINUE
titrating. It is NOT necessary to record any volumes at this point. The color of the solution should
become dark blue and require ONLY a FEW drops of the titrant to reach the clear and colorless
endpoint. Record the volume of thiosulfate in the buret at the clear and colorless endpoint.
4. Repeat the process two more times.
74
Data/Calculations
Part 1
1.
Initial buret volume
mL
2.
Final buret volume
mL
3.
Volume of S2O3-2 used
mL
4.
moles of KIO3 added
mole
5.
moles of I2 produced
(refer to chem. reaction)
mole
6.
moles of S2O3-2 required for
titration
(refer to chem. reaction)
mole
7.
Molarity of S2O3-2
M
Part2
1.
Initial buret volume
mL
2.
Final buret volume
mL
3.
moles of S2O3-2 added
(use molarity from Part 1 #7)
mL
4.
moles of I2 titrated
(refer to chem. reaction)
mole
moles of OCl- present
(refer to chem. reaction)
mole
6.
moles of NaOCl present
mole
7.
grams of NaOCl present
g
8.
grams of bleach used
(assume bleach density = 1 g/mL)
g
9.
Percent NaOCl in bleach
10.
Percent difference from bottle
5.
%
%
Conclusion
Write a brief paragraph including the average values of the concentration of thiosulfate and the percent
of NaOCl in the bleach.
75
32 Determination of F and NA
Introduction and Objectives
The purpose of this experiment is to demonstrate the principles and process of electrolysis through the
deposition of copper ions from solution to form a pure copper film. Data collected during the experiment
will be used to calculate Faraday’s Constant, F.
Faraday’s constant is significant for conversion between the amount of charge flowing within a system
and the quantity of material reacting in that electrochemical system. Electrolysis can be thought of as a
type of titration in which electrons serve as the titrant that react with the analyte at the electrode surface.
It is possible to determine the specific number of moles of analyte reacted from the moles of electrons
allowed to flow through the electrochemical cell.
The amount of charge (q) that has been passed in an electrochemical cell any time can be
determined from the current (I) in the cell and the time (t) the current has flowed in the cell as seen in
equation 1. Note: charge has the unit of coulomb.
I
q
t
(1)
The second piece of information required is the number of moles of electrons passed during the
experiment. To determine moles of electrons, the mass of the copper deposited onto the cathode will be
measured. The reduction half-reaction for the experiment is Cu+2 + 2e
Cu. From the change in
mass of the copper cathode it is possible to determine the moles of copper deposited and ultimately the
moles of electrons that passed. From equation 2, the value of Faraday’s constant will be determined and
compared to its accepted value.
F
coul
mole e 
(2)
 96,500
For this experiment the value of F can be determined also using the following equation:
F 
qT  MMCu
n  mCu
(3)
where qT is the total charge passed during the experiment, MMCu is the molar mass of copper, n is the
number of electrons from the balanced redox half reaction, and mCu is the mass gained by the copper
cathode.
Using the same information and that the charge of on one electron is 1.602 E –19C, it is also
possible to calculate the value of Avogadro's Number (NA). The value for Avogadro's number can be
calculated using the following equation:
NA 
qT  MMCu
n  mCu  qe
(4)
where the only new quantity is qe which is the charge on an electron.
Block diagram of the electrical circuit:
+
Variable
Resistor
+
+
anode
A
cathode
-
ammeter
76
Procedure
1. Gently polish the copper cathode at your workstation and rinse it with distilled water. Wipe it off
carefully so as not to leave any paper residue on the cathode.
2. Determine the mass of the copper cathode to the nearest 0.1mg using the analytical balance in the
Quantitative Analysis lab.
3. Place the cathode in the electrochemical cell at your workstation as directed by your instructor.
4. With the help of your instructor connect the cell components together and maintain the current at 150
mA (0.150A) for 45 minutes.
5. Record the current in 3-min. intervals. Use the average current for the calculations.
6. After completing the experiment remove the piece of copper from the electrochemical cell and allow it
to COMPLETELY dry.
7. Determine its final mass to the nearest 0.1 mg.
Data/Calculations
Mass of cathode – start
g
Mass of cathode – end
g
Mass change in cathode (g Cu)
g
Moles of Cu deposited
mol
Moles of electrons
mol
Time of electrolysis (sec)
sec
Current during electrolysis (I)
amp
Total Charge during electrolysis (It)
coul
Value of F
Value of NA
coul
mol e-
atoms
mole
Conclusion
Write a brief conclusion paragraph including the calculated value of F and NA.
77
33 Preparation of Aspirin
Prelab Questions
1. What is the formula for iron (III) hydrogen sulfate pentahydrate?
2. Balance the following redox reaction in an acid: Cr(s) + CrO42-(aq)  Cr(OH)3(s)
3. If 5.50 g of Cr are reacted with excess CrO4-2 in the above reaction, how many grams of Cr(OH)3 can
be produced?
4. What is the pH of a 800-mL solution 50.0g each of formic acid (HCO2H) and sodium formate
(HCO2Na)? What is the pH if 0.10M HCl is added to the solution? (Assume no change in volume.)
5. Which of the following has the higher boiling point: CH3-O-CH3 or CH3-CH2-OH? Explain.
Introduction and Objectives
Aspirin or acetylsalicylic acid is used as a painkiller (analgesic), fever depressant (antipyretic), and
swelling reducer (anti-inflammatory agent). It has been noted that since aspirin is an organic acid it can
irritate the lining of the stomach leading to ulcerated areas. Some aspirin products contain buffers that
neutralize it’s acidic effect while leaving its medicinal benefits intact.
The experiment involves a simple reaction of between the two reactants: salicylic acid and acetic
anhydride. The acid used is a catalyst. The chemical reaction is:
COOH
O
OH
+
CH3
COOH
C
O
CH3 C
H+
O
O
O
C
CH3
heat
+
CH3
C
OH
O
salicylic acid
MW = 138 g/mol
acetic anhydride
(excess)
aspirin
(MW = 180 g/mole)
acetic acid
Procedure
Aspirin Preparation
1. Weigh 4.0 g of salicylic acid and place it into a 125-mL Erlenmeyer flask.
2. IN THE FUME HOOD, add 6 mL of acetic anhydride to the flask.
3. With constant swirling, carefully add 7 drops of concentrated H3PO4.
4. Heat the mixture in a hot water bath (70-80oC) for 10 minutes. Swirl the flask often to ensure complete
reaction.
5. Remove the flask from the water bath and allow it to cool for about 5 minutes.
6. Slowly pour 20 mL of ICED DI water into the flask with CONSTANT swirling.
7. Place the flask in an ice bath and allow the crystals of aspirin to precipitate.
8. Collect the aspirin by suction filtration in a Buchner funnel.
9. Wash the solid with 20 mL of ICED ethanol.
10. Carefully remove the aspirin on the filter paper. Scrap the wet aspirin into a large beaker and allow it
to air dry.
11. When the aspirin is thoroughly dried weigh the aspirin and calculate the percent yield.
Aspirin purity
Part 1
1. Obtain a thin-film IR of your aspirin sample and compare it to an IR of pure aspirin and salicylic acid.
Part 2
2. After the aspirin is dried and weighed, about 0.100g to 5 mL of DI water in a test tube and dissolve the
aspirin.
3. Add 3 drops of 1% FeCl3. A clear and colorless solution indicates a pure aspirin sample. The more
purple the solution color the more it indicates and incomplete reaction and an impure aspirin sample.
Record the color of the solution and comment on the aspirin’s purity.
78
Data/Calculations
1.
Theoretical mass (grams of aspirin from salicylic acid)
2.
Percent yield of aspirin
g
%
Conclusions
Write a brief paragraph including the mass of the apirin, percent yield, and the purity of the aspirin.
Compare the IR spectrum your sample with that from the pure samples of salicylic acid and aspirin. Note
similarities differences in absorption bands.
IR Spectrum of Salicylic Acid
IR Spectrum of Acetylsalicylic Acid
79
34 Polymerization Reactions
Introduction and Objectives
Polymers are very large molecules produced by linking many smaller units – monomers – together.
Many important biological molecules are polymers including sugars, proteins, and DNA. The “plastic”
materials are either naturally occurring polymers or synthetic polymers.
In this experiment, two different polymers will be produced: nylon and glyptal. The formation of both
of these polymers is an example of condensation polymerization. Nylon involves the removal of HCl and
glyptal of water. The reactions are given below.
Another type of spectroscopy – Infrared (IR) – will be used to confirm the formation of each polymer.
Nylon:
O
O
C (CH2)4 C
Cl
Cl
+
H
H
N
H
O
O
(CH2)6 N
C (CH2)4 C
H
H
N
(CH2)6 N
n
H
+
nHCl
Glyptal:
O
O
H2C
CH CH2
OH OH OH
+
O
O
CH2 CH CH2 O C
C
O
C
O
C
O
O
+
O
nH2O
O
O
CH2 CH CH2 O C
C
O
O
n
Procedure
Glyptal
1. Mold a piece of Al foil that has been double folded around the bottom of a 50-mL beaker. This makes a
suitable container for the gylptal as it cools and solidifies. Ensure that the foil is NOT torn in any
location.
2. Thoroughly mix 10 g of phthalic acid and 0.5g of sodium acetate into a small beaker and transfer it into
a medium sized test tube.
3. Add 4 mL of glycerol and 1 drop of red food coloring to the test tube.
4. Clamp the test tube near the top of tube and attach it to a ring stand. Adjust the inclination to an angle
of 45o.
5. Using a Bunsen burner with a moderately cool flame, begin to heat the test tube. Sweep the flame
CONSTANTLY over the entire test tube to prevent localized heating. After gentle heating, the mixture
should melt into a red colored liquid. The solution color will slowly begin to fade.
6. Continue heating for several minutes until the solution’s color abruptly fades to a light yellow and then
for another 1–2 minutes. If the color of the solution begins to become dark, stop heating
IMMEDIATELY.
80
7. Take the assembly to a fume hood and carefully remove the clamped tube from the ring stand. Pour the
HOT polymer into the Al foil mold. The polymer may smoke and sizzle at this point.
8. Add a small object such as a penny to the polymer. This must be done QUICKLY because the polymer
will begin to harden very quickly.
9. Leave the polymer to cool and setup until the next lab.
Nylon
1. Mix 1.0 mL of 20% NaOH and 3 mL of 5% aqueous solution of hexamethylenediamine in a 50-mL
beaker.
2. Slowly add 3 mL of 5% adipoyl chloride solution in cyclohexane into the same 50-mL beaker. Add the
solution so that it layers on top of the aqueous solution. The nylon will form at the interface of these
two phases.
3. Carefully and slowly scrape the side of the beaker with a bent piece of copper wire to remove the nylon
stuck to the walls.
4. Hook the center of the film and SLOWLY lift the film out of the beaker.
5. Wrap the nylon around a glass stir rod without touching it with your fingers.
6. Slowly rotate the rod and wind out as much of the nylon as possible. Transfer the nylon into a beaker
filled with DI water. Dry it carefully with paper towels.
7. Once dry, run an IR on your nylon sample and place a copy of the spectrum in your lab book.
Data/Calculations
Compare the IR spectra you obtained to those of the known compounds. Note any similarities and
differences.
Conclusions
Write a brief paragraph discussing how the IR spectra support whether or not the two polymers were
produced.
IR Spectrum of Nylon
102
98
96
94
92
90
88
% Transmittance
100
86
84
4000
3500
3000
2500
2000
1500
Wavenumber (cm-1)
1000
500
81