Bergen Community College Physical Sciences Department
General Chemistry II CHM 241
LABORATORY MANUAL 2019
Dr. Ara Kahyaoglu Prof. Jean Acken Associate Professor Assistant Professor
BERGEN COMMUNITY COLLEGE (BCC) is committed to providing quality material that promotes the best in inquiry-based
science education. However, conditions of actual use may vary, and the safety procedures and practices described in this resource
are intended to serve only as a guide. Additional precautionary measures may be required. BCC and the authors do not warrant or
represent that the procedures and practices in this resource meet any safety code or standard of federal, state, or local regulations.
BCC and the authors disclaim any liability for personal inquiry or demand to property arising out of or relating to the use of this
resource, to include any of the recommendations, instructions, or materials contained therein.
Bergen Community College
400 Paramus Road, Paramus, NJ 0765
Bergen Community College 2 General Chemistry II Laboratory
To the instructor,
This manual is available as a free download from the Bergen Community College website.
The Science Department’s aim has been to provide low cost, safe, and interesting, yet
Students are to complete the pre-laboratory exercises prior to the laboratory session. The
departmental format for the laboratory report can be found in the course syllabus.
Much effort has been made by the chemistry faculty to review this manual and make it as
error-free and accurate as possible. However, some errors will have escaped our notice.
Your help in forwarding to us any errors, inaccuracies, and/or suggestions will be greatly
appreciated. We will definitely welcome your comments and suggestions. We will make the
changes and improvements as soon as we can to make the updated manual ready for
succeeding semesters. We can be contacted at [email protected] and
Good luck and have a good semester. We look forward to hearing from you.
To the students,
This manual is available to BCC students for free download from the Bergen Community
College website. The Science Department’s aim has been to provide low cost, safe, and
interesting, yet relevant experiments that illustrate the concepts presented in the lecture
Safety is everyone’s number one priority. Do not hesitate to ask your instructor if you do not
understand the procedure.
Keep in mind that your instructor expects you to be prepared for every laboratory session.
We appreciate the efforts of professors PJ Ricatto, Linda Box, Gary Porter, Brent Chapman,
Frank Ramdayal, Farah Rezae and Riwa Dandan through their suggestions and corrections.
Dr. Ara Kahyaoglu, author
Prof. Jean Acken, contributor and editor
Bergen Community College 3 General Chemistry II Laboratory
TABLE OF CONTENTS
Course Schedule 4
Laboratory Safety 5
Integrity of Data Guidelines 8
Experiment 1. Heat of Fusion 9
Experiment 2. Intermolecular Forces 14
Experiment 3. Spectroscopy 23
Experiment 4. Percent Copper in Brass 31
Experiment 5. Freezing Point Depression 39
Experiment 6. Chemical Kinetics 50
Experiment 7. Le Châtelier's Principle 63
Experiment 8. Coordination Number 75
Experiment 9. Identification of a Weak Acid 83
Experiment 10. Solubility Product 94
Experiment 11. Qualitative Analysis of Cations 101
Experiment 12. Titration of Hydrogen Peroxide 109
Experiment 13. Electrochemistry 119
Experiment 14. Bicarbonate-carbonate mixture 128
Appendix A. Common Laboratory Equipment 135
Appendix B. Volumetric Glassware 136
Appendix C. Graphing 137
Appendix D. Titration 139
Appendix E. Filtration 140
Appendix F. Periodic Table 141
Bergen Community College 4 General Chemistry II Laboratory
Fifteen Week Semester Twelve Week Semester
1. Lab Safety and Exp. 1 1. Lab Safety and Exp. 1
2. Exp. 2 2. Exp. 3 and Exp. 2, Parts A&B (or Part C)
3. Exp. 3 3. Exp. 4 and Exp. 2, Part C (or Parts A&B)
4. Exp. 4 4. Exp. 5
5. Exp. 5 5. Exp. 6
6. Exp. 6 6. Exp. 7 and Exp. 8 part A
7. Exp. 7 7. Exam 1 and Exp. 8 part B
8. Exam 1 and Exp. 8 part A 8. Exp. 9 and Exp. 8 part C
9. Exp. 8 parts B and C 9. Exp. 10*
10. Exp. 9 10. Exp. 11
11. Exp. 10* 11. Exp. 12
12. Exp. 11 12. Exam 2 and Exp. 13
13. Exp. 12
14. Exp. 13 or Exam 2
15. Exam 2 or Exp. 13
*The NaOH solution standardized in Experiment 9 is used again in this experiment.
Bergen Community College 5 General Chemistry II Laboratory
BERGEN COMMUNITY COLLEGE
SAFETY REGULATIONS FOR THE CHEMISTRY LABORATORY
1. Read these safety regulations carefully and be sure you understand them. Before each
laboratory session, your instructor will discuss any safety hazards that may be associated
with that day’s experiment. Therefore, it is imperative that you come to lab on time.
2. Due to safety concerns students who arrive after the pre-lab presentation may not be allowed to perform that particular lab experiment.
3. It is strongly suggested that you obtain a hall locker from the Security Office. Only your lab manual, notebook, and calculator are allowed on the lab bench.
4. Report all accidents, no matter how minor, to your instructor at once. No one in the lab is permitted to give out bandages or medication. You must see the College Nurse.
5. Safety glasses or goggles are required and must be worn by everyone in the lab when experiments are being conducted. Contact lenses are not recommended in the chemistry
lab. Safety glasses are provided by the college, but students may purchase their own.
6. Do not perform any unauthorized experiment.
7. Do not taste anything in the laboratory. Never eat, drink or smoke in any of the labs.
8. You must tie back long hair. Do not wear open-toed shoes, shorts, fuzzy sweaters, loose sleeve shirt or any dangling jewelry. You must cover bare midriffs. You are advised to
wear a lab coat or old clothing to the lab.
9. Do not fill pipettes by mouth. Rubber bulbs or pipette pumps are provided. The instructor will demonstrate how these are to be used.
10. Exercise care when noting the odor of fumes. Use ‘wafting’ if you are directed to note an odor.
11. Do not force glass tubing or a thermometer into rubber stoppers. Lubricate with water and introduce it gradually and gently into the stopper, or insert through a cork borer.
Protect your hands with toweling when inserting without a cork borer.
12. Never point a test tube containing a reaction mixture (especially when heating) toward yourself or another person.
13. No ‘fooling around’ in the laboratory. A less than serious approach to lab work may result in an accident.
14. Before connecting or disconnecting electrical equipment, make sure that the switches are in the off position.
Bergen Community College 6 General Chemistry II Laboratory
15. Never work in the laboratory alone.
16. Make sure all apparatus is properly supported on the workbench.
17. Read the label on every bottle twice before using it in the laboratory. Many chemical names are very similar but are very different chemically.
18. Replace caps and stoppers on bottles immediately. Return spatulas to their correct place immediately after use. Do not mix them up.
19. Do not remove or relocate any chemical that has been placed in the hood. Sample it in the hood.
20. Never light a Bunsen burner with a cigarette lighter. Use the strikers that are provided.
21. Students are responsible for keeping their work area neat and orderly. All spills are to be cleaned up immediately using the spill kits located on the instructor’s desk. Solid
chemical waste should be and placed in the appropriately labeled container. Liquid
chemical waste should be poured into the appropriately labeled container. All waste
material should be left in the hood for subsequent disposal. If there is doubt about proper
disposal, ask the instructor.
22. Wash all glassware immediately after use. Place clean glassware on drying rack or in the designated bin on the counter.
23. Dispose of broken glassware in the labeled broken glassware boxes.
24. Wash your hands before leaving the laboratory.
25. You must notify your instructor of any chemical to which you are allergic.
26. If you are pregnant or planning to become pregnant this semester, you must notify your physician that you are enrolled in a chemistry lab course. You and your physician must
decide whether or not it is appropriate for you to remain in the course.
Note the location of the following safety equipment so that you can get to it quickly
in an emergency.
SAFETY EQUIPMENT LOCATION
Bergen Community College 7 General Chemistry II Laboratory
SAFETY IN THE LABORATORY
1. Safety glasses must be worn by everyone working in the lab. T F
2. Only major accidents in the lab need to be reported T F
3. Material Safety Data Sheets are provided in the lab T F
4. Eating and drinking are permitted in the lab T F
5. It is OK to taste a chemical as long as it smells good T F
6. Only authorized experiments are to be performed T F
7. You should wear shoes at all time in the lab T F
8. In order to save time, it is permissible to weigh hot objects T F
9. Broken glassware should be disposed of in the appropriate box T F
10. Working alone in the lab is an acceptable practice T F
A typical Chemistry Laboratory safety YouTube video link is given below: (hold Ctrl Key
and hover the mouse over the link) https://www.youtube.com/watch?v=UKovNdse5MU
Please complete sign this attached form. Remove it from the safety regulations and hand it to
your Laboratory Instructor.
I, the undersigned, have read the Divisional Safety Regulations for the Chemistry
Laboratories. I understand them and will abide by them.
Print your name: ________________________________________________________
Course Name and Number:_______________________________________________
Bergen Community College 8 General Chemistry II Laboratory
INTEGRITY OF DATA GUIDELINES
One purpose of a laboratory course is to reinforce the concepts covered in the lecture
course. A second, equally important purpose, is to experience working in a chemistry lab, and
to learn about practices and procedures that are employed in such an environment. In addition
to specific laboratory procedures that will be covered in the array of experiments, there are
two universal practices in all laboratory settings- Laboratory Safety, which was discussed in
the previous pages, and Integrity of Data Guidelines.
These guidelines are used in all laboratory settings- from the traditional research
laboratory to hospitals and the physician’s office. The purpose of the guidelines is to ensure
that data is recorded in such a way that its veracity, or authenticity, cannot be questioned.
Taken as a whole, these practices protect the integrity of the data by preventing it from being
changed or recorded in error. Students are expected to follow these integrity of data guidelines
when collecting and recording data. The guidelines are as follows:
1. Data sheets must include the date and the student’s name.
2. Data is recorded in blue or black non-erasable ink; no white-out is permitted.
3. If a mistake is made while entering data, a single line is used to cross it out
and the correct entry is made nearby. (The original entry must be legible.)
4. No transcription is permitted. (Data is recorded directly into the data sheets.)
5. Data is recorded at the time it is observed.
In most laboratories today, notebooks are electronic rather than paper. Although this
renders a different set of guidelines, their purpose is the same- to ensure the authenticity of
data. Laboratory notebook software does not permit a change to be made once data has been
entered. In instances where a change is required, there is a record of the original entry. When
a measurement is recorded on a scrap of paper, that original data is scanned and becomes a
part of the notebook. These and other practices concerning electronic lab notebooks, along
with the guidelines described above regarding paper notebooks, work together to protect the
integrity of experimental data.
Bergen Community College 9 General Chemistry II Laboratory
Heat of Fusion
OBJECTIVE: To determine the heat of fusion of water.
When the solid phase of a molecular substance is converted to the liquid phase, energy,
in the form of heat, must be added in order to break the attractions between the molecules.
These intermolecular forces in a solid hold the molecules locked into position. Although the
molecules vibrate in place, they do not move relative to each other, i.e. they have no
translational movement. In contrast, the molecules in the liquid phase, although close to one
another, do have translational movement. They are constantly making and breaking
intermolecular attractions as they move about in random translational motion.
As heat is added to a molecular substance in the solid phase, the kinetic energy of the
molecules increases, resulting in greater vibrational motion, and evidenced by an increase in
temperature. This process continues until the melting point is reached, when molecules begin
to have sufficient energy to break the attractive forces holding them in position, and the
substance begins to melt. At this point, added energy results in breaking attractive forces rather
than in increased movement, and the temperature remains constant throughout the melting
process. When the entire sample has become a liquid, added heat increases the kinetic energy
and the temperature increases once again.
A similar transition occurs in converting a substance from the liquid to the gas phase.
As heat is added once the boiling point has been reached, this energy is used to break
intermolecular forces between molecules in the liquid phase. Again, during the process of
vaporization, the temperature remains constant.
These relationships can be summarized in a heating curve, as illustrated in the figure
on the following page.
The amount of heat required to convert a substance from the solid to the liquid phase
is quantified as the heat (or enthalpy) of fusion, ∆Hfus. It is a physical property, and can be
reported as heat per gram of substance or per mole of substance. The latter is often referred to
as the molar heat of fusion.
Experiment 1 Heat of Fusion
Bergen Community College 10 General Chemistry II Laboratory
In this experiment, the heat of fusion of water, in joules/gram, will be determined using
a coffee cup calorimeter where a sample of ice has melted in tap water. The amount of heat
given up by the tap water in the calorimeter as it cools (qwater) will be absorbed as heat by a
sample of ice as it melts (qfusion) and as this melted ice warms (qmelted ice). Assuming no loss of
heat to the surroundings, the sum of these must equal zero.
qwater + qfusion + qmelted ice = 0
qfusion = – qwater – qmelted ice (eq. 1)
Values for both qwater and qmelted ice are obtained from the following equations, where m
represents mass, c represents the specific heat of water (4.18 J/g ºC) and ∆T represents the
change in temperature.
q = m c ∆T (eq. 2)
∆T = Tfinal – Tinitial (eq. 3)
Once the heat of fusion is determined, the experimental error can be found as follows.
𝑃𝑒𝑟𝑐𝑒𝑛𝑡 𝐸𝑟𝑟𝑜𝑟 = |𝑇ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝑉𝑎𝑙𝑢𝑒 − 𝐸𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙 𝑉𝑎𝑙𝑢𝑒|
𝑇ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝑉𝑎𝑙𝑢𝑒 × 100
Three determinations will be made using different sample sizes. A graph of the heat,
in joules, absorbed in melting the ice (qfusion) as a function of the mass of melted ice, in grams,
will be constructed. The slope of the line represents the heat of fusion of water.
at u re
Figure 1: Heating Curve
Experiment 1 Heat of Fusion
Bergen Community College 11 General Chemistry II Laboratory
REAGENTS: Ice EQUIPMENT: 150-mL beaker Tap water 100-mL graduated cylinder
coffee cup calorimeter 400-mL beaker
thermometer, or thermocouple
1. Measure and record the mass of a 150-mL beaker. Set it aside ready to use in step 5.
2. Tare a coffee cup calorimeter. Add 100 mL tap water using a graduated cylinder.
Measure and record the mass. Place the calorimeter in a 400-mL beaker for stability.
3. Measure and record the initial temperature of the tap water in the calorimeter.
4. Add sufficient ice to fill the volume of water, and gently stir with the thermometer.
5. When the temperature reaches between 0ºC and 5ºC, record the final temperature.
Immediately pour the water into the 150-mL beaker, leaving the unmelted ice behind.
6. Measure and record the mass of the beaker and contents.
7. Repeat steps 1 – 6 using 70 mL tap water, and again using 40 mL.
Disposal: Water may be disposed of down the drain.
A. Perform the following calculations for each of the three determinations.
1. Determine the mass of the contents of the beaker. This is the mass of the original tap water
in the calorimeter, plus that of the melted ice.
2. Determine the mass of the melted ice by subtracting the mass of the tap water from the
mass of the beaker contents.
3. Determine the temperature change for the tap water, ∆Twater, using eq. 3.
4. Determine the temperature change for the melted ice, ∆Tmelted ice, using eq. 3. The initial
temperature for the ice is assumed to be 0 ºC.
5. Determine qwater using the mass and temperature change of the tap water and eq. 2.
6. Determine qmelted ice using the mass and temperature change of the melted ice and eq. 2.
7. Determine qfusion using eq. 1.
B. Prepare a graph in Excel* of qfusion, in joules, as a function of the mass of melted ice, in
grams. Determine the heat of fusion for water from the graph.
*See Appendix C for directions on graphing.
Bergen Community College 12 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 1: Heat of Fusion
Data: Determination 1 Determination 2 Determination 3
Initial mass of beaker ____________ ____________ ____________
Mass of tap water ____________ ____________ ____________
Initial temperature ____________ ____________ ____________
Final temperature ____________ ____________ ____________
Final mass of beaker ____________ ____________ ____________
Mass of beaker contents ____________ ____________ ____________
Mass of melted ice ____________ ____________ ____________
ΔT of tap water ____________ ____________ ____________
ΔT of melted ice ____________ ____________ ____________
Heat for water, qwater ____________ ____________ ____________
Heat for melted ice, qmelted ice ____________ ____________ ____________
Heat for fusion of ice, qfusion ____________ ____________ ____________
Heat of fusion of ice _______________________
Bergen Community College 13 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 1: Heat of Fusion
1. The heat of fusion of water is 333 J/g. Determine the percent error using equation 4.
2. Determine the amount of heat required to raise the temperature of a 22.5-gram
sample of copper from 125 ºC to its melting point of 1084 ºC, and then melt the
copper. (The specific heat of copper is 24.4 J/mol ºC and its heat of fusion is
3. If the ice had begun at a temperature lower than 0 ºC, would the calculated value of
the heat of fusion have been higher, lower, or unchanged? Briefly explain.
Bergen Community College 14 General Chemistry II Laboratory
OBJECTIVE: To relate intermolecular forces of molecules to physical properties.
The attractive forces between molecules and their neighbors are called intermolecular
forces. These forces are much weaker than the intramolecular forces within a substance- the
covalent (or ionic) bonds. Intermolecular forces are the attractions that need to be overcome
for a molecular solid to melt, and for a liquid to vaporize. Therefore, the strength of these
attractive forces influence a substance’s physical properties. The stronger the intermolecular
forces, the higher melting point, boiling point, heat of vaporization, and other properties. The
following table summarizes the boiling points of some molecular compounds.
Compound Formula Polarity Molecular Structure Boiling Point
Methane CH4 Nonpolar
Propane C3H8 Nonpolar
Butane C4H10 Nonpolar
Hexane C6H14 Nonpolar
Acetone C3H6O Polar
Ethanol C2H6O Polar
Water H2O Polar
Experiment 2 Intermolecular Forces
Bergen Community College 15 General Chemistry II Laboratory
There are three general types of intermolecular forces. All substances exhibit London
Dispersion Forces (LDF), and they are generally the weakest of the three types. These London
forces are due to the attractions between small, temporary dipoles that arise from the constant,
random movement of the electrons in a substance. As molar mass increases, the size of the
electron cloud increases as well. It becomes more easily distorted, and produces temporary
dipoles of greater magnitude. This causes the attractions to be stronger, requiring more energy
for both fusion and vaporization. For halogens, this results in increasing melting and boiling
points, shown by the fact that at room temperature F2 and Cl2 are gaseous, Br2 is liquid and I2
is solid. The extent to which the electron cloud can be distorted is called polarizability.
Dipole-dipole forces exist between molecules that are polar. Since the dipoles are
permanent, these attractions are generally stronger than London Dispersion Forces. This
means that a polar molecule with similar molar mass as a nonpolar molecule will have higher
melting points and boiling points. Not all molecules containing polar bonds are polar. The
polar bonds must be unevenly dispersed in the molecule in order to produce a polar molecule.
CO2 and CBr4, for example, have polar bonds but are not polar molecules.
The third type of intermolecular force is hydrogen bonding, a specific type of dipole-
dipole attraction that is stronger than other dipole-dipole attractions. Hydrogen bonds form
when a hydrogen atom is covalently bonded to a very electronegative atom. This causes its
electron to be drawn away from its nucleus. The positive hydrogen is then attracted to the very
electronegative atom in a neighboring molecule. In order to observe hydrogen bonding, the
hydrogen atom must be covalently bonded to fluorine, oxygen or nitrogen. A hydrogen atom
bonded to a carbon atom cannot create a hydrogen bond. It’s
important to note that, despite its name, a hydrogen bond is
an intermolecular force, not a bond. The figure to the right
illustrates H-bonding between water molecules. H-bonding
is important in biochemistry; the structure of a biopolymer is
largely determined by the formation of hydrogen bonds.
The relative strengths of the three types of intermolecular forces, and thus boiling
points, are generally as follows:
London Dispersion Forces < Dipole-Dipole Forces < H-Bonding
However, this is not always true. Since molar mass is also a factor, a large non-polar molecule
can have a higher boiling point than a compound that interacts with dipole-diploe forces, or
even a substance with H-bonding. For example, octane, a component of gasoline, has a boiling
point of 125oC- much higher than acetone (dipole-dipole) and H2O (H-bonding). This is due
to the polarizability of the large electron cloud.
To make comparisons of the intermolecular forces of a substance, evaporation rate can
be used instead of boiling point. Evaporation rate is the ratio of the change in temperature to
the change in time as a substance evaporates. A faster rate of evaporation translates to a lower
boiling point and, in turn, weaker intermolecular forces.
Experiment 2 Intermolecular Forces
Bergen Community College 16 General Chemistry II Laboratory
Boiling point is not the only physical property affected by the type of intermolecular
forces a substance has. Solubility is also dependent upon the polarity of a molecule. The term
“like dissolves like” suggests that polar solutes dissolve in polar solvents and nonpolar solutes
dissolve in nonpolar solvents. Therefore, polarity, and the associated intermolecular forces,
determine a substance’s solubility in water and in other solvents.
The solubility of a solid in a liquid is readily observed. When liquids mix forming a
homogeneous solution, they are said to be miscible; if they do not mix, they are immiscible.
If two liquids are miscible, there is no observable interface between the two. If the two liquids
are immiscible, two distinct layers are seen.
In this experiment, both miscibility and evaporation rates of acetone, ethanol, hexane,
and water will be determined. Salt solubility in an ethanol-water mixture will also be observed.
REAGENTS: acetone EQUIPMENT: Thermometers ethanol filter papers
hexane rubber band
distilled water tape
sodium chloride stop watch
5 test tubes containing a 50% by volume: small test tubes
water and acetone wood block
water and hexane
hexane and acetone
hexane and ethanol
ethanol and acetone
– Do not pour any materials into the sink!
– Wash hands and laboratory bench after the experiment.
– Acetone: Extremely flammable liquid and vapor. Vapor may cause flash fire. Causes eye
irritation. Breathing vapors may cause drowsiness and dizziness. Causes respiratory tract
irritation. Aspiration hazard if swallowed. Can enter lungs and cause damage. Prolonged or
repeated contact may dry the skin and cause irritation.
Hexane: Extremely flammable liquid and vapor. Vapor may cause
flash fire. Breathing vapors may cause drowsiness and dizziness.
Causes eye, skin, and respiratory tract irritation. May be harmful if
absorbed through the skin. Aspiration hazard if swallowed and enters
lungs causing damage. Possible risk of impaired fertility. Long-term
exposure may cause damage to the nervous system of the extremities.
Bergen Community College 17 General Chemistry II Laboratory
Date: ____________________ Name: ________________________
Experiment 2: Intermolecular Forces
1. Which of the substances used in this experiment must be handled in the fume hood?
2. Identify the strongest type of intermolecular forces in acetone, ethanol, water and hexane.
(Structures listed on page 15.)
3. Predict the relative strength of the intermolecular forces in the four liquids above.
______________ < _______________ < …
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