Spectrophotometry of co2+ post lab answers

Bergen Community College Physical Sciences Department

General Chemistry II CHM 241

LABORATORY MANUAL 2019

Second Edition

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

PREFACE

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

relevant experiments.

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 akahyaoglu@bergen.edu and

jacken@bergen.edu.

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

course.

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

mailto:akahyaoglu@bergen.edu
mailto:jacken@bergen.edu
Bergen Community College 3 General Chemistry II Laboratory

TABLE OF CONTENTS

Page

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

COURSE SCHEDULE

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.

Safety Regulations

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

FIRE EXTINGUISHER

SAFETY SHOWER

EYEWASH

EMERGENCY PHONE

FIRE ALARM

NEAREST EXIT

Safety Regulations

Bergen Community College 7 General Chemistry II Laboratory

SAFETY IN THE LABORATORY

True False

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: ________________________________________________________

Signature: ____________________________________________________________

Date: ____________________________________________________________

Course Name and Number:_______________________________________________

https://www.youtube.com/watch?v=UKovNdse5MU
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

Experiment 1

Heat of Fusion

OBJECTIVE: To determine the heat of fusion of water.

BACKGROUND:

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

Rearranging,

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.

T em

p er

at u re

Heat added

melting

vaporization

Figure 1: Heating Curve

(eq. 4)

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

PROCEDURE:

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.

CALCULATIONS:

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 ____________ ____________ ____________

Results:

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

POST-LAB QUESTIONS:

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

13.26 kJ/mol.)

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

Experiment 2

Intermolecular Forces

OBJECTIVE: To relate intermolecular forces of molecules to physical properties.

BACKGROUND:

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

- 161oC

Propane C3H8 Nonpolar

- 42oC

Butane C4H10 Nonpolar

10oC

Hexane C6H14 Nonpolar

70oC

Acetone C3H6O Polar

56oC

Ethanol C2H6O Polar

77oC

Water H2O Polar

100oC

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

SAFETY ALERT:

- 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

PRE-LABORATORY QUESTIONS

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.

______________ < _______________ < _______________ < _______________

4. Water (MW = 18 g/mol) has higher boiling point than ethanol (MW = 46 g/ mol) and acetone

(MW = 58 g/mol). Why does water, such a small molecule, have such a high boiling point?

5. What is the meaning of the term "like dissolves like"?

Experiment 2 Intermolecular Forces

Bergen Community College 18 General Chemistry II Laboratory

PROCEDURE:

Part A: Evaporation Rate of Acetone and Hexane

MUST BE PERFORMED IN A FUME HOOD!

1. Obtain a thermometer that reads to the tenth of a Celsius degree.

2. Add a few mL acetone into a small test tube.

3. Wrap a piece of filter paper around the end of a thermometer; secure it with a rubber band.

4. Place the thermometer into the test tube. Wait approximately 10-20 seconds, until the

temperature is stabilized, and record this as the temperature at time = 0 seconds.

5. Remove the thermometer from the test tube, touching the tip of the filter paper to the tube

to leave behind any excess liquid. Tape it to a woodblock in the fume hood, so that the

thermometer lies horizontally, and the filter paper does not touch the workbench.

6. Record the temperature every 30 seconds for 5 minutes.

7. Repeat steps 1 – 6 using hexane.

Disposal: Return any remaining acetone and/or hexane to the appropriate collection

container. Leave the filter paper in the hood to dry, then dispose of in trash bin.

Figure 1: Thermometer, filter paper and rubber band in test tube with

liquid (left), and taped to wood block (right).

Experiment 2 Intermolecular Forces

Bergen Community College 19 General Chemistry II Laboratory

Part B: Evaporation Rate of Water and Ethanol

Working at the lab bench, not in the fume hood, follow steps 1 – 6 as described in part A

using ethanol, and again using water.

Disposal: Return any remaining ethanol to the appropriate collection container. Leave the

filter papers on the lab bench to dry, then dispose of in trash bin.

Part C: Solubility

1. Obtain a set of five prepared test tubes containing mixtures of the following:

water and acetone

water and hexane

hexane and acetone

hexane and ethanol

ethanol and acetone

2. Prepare a sixth test tube by adding equal amounts (about 1 mL) of ethanol and water.

3. Inspect the test tubes and record the solubility of each mixture: S = soluble, I = insoluble.

4. Add a pea-sized quantity of sodium chloride to the ethanol/water mixture. Gently shake

the tube. Make and record observations.

Disposal: Place the ethanol-water-salt mixture in the designated disposal container.

CALCULATIONS:

Parts A and B:

1. Subtract each temperature from the initial temperature in order to find ∆T.

2. Prepare (by hand) a graph* of ∆T as a function of time for each of the liquids on one piece

of graph paper.

3. Determine the relative evaporation rates for the four liquids by comparing the curves. A

faster evaporation rate generates a steeper curve.

*See Appendix C for directions on graphing.

Bergen Community College 20 General Chemistry II Laboratory

Date: ____________________ Name: ________________________

Experiment 2: Intermolecular Forces

Data: Parts A and B: Evaporation Rates

Results: Parts A and B: Evaporation Rates

Experimentally determined relative strengths of intermolecular forces:

_______________ < ________________ < ________________ < ________________

time,

seconds

Acetone Hexane Ethanol Water

T, °C ∆T, °C T, °C ∆T, °C T, °C ∆T, °C T, °C ∆T, °C

0.

30.

60.

90.

120.

150.

180.

210.

240.

270.

300.

Bergen Community College 21 General Chemistry II Laboratory

Date: ____________________ Name: ________________________

Experiment 2: Intermolecular Forces

Data: Part C: Solubility

Acetone Ethanol Hexane

Water

Hexane N/A

Ethanol N/A N/A

Observation: Solubility of NaCl in Ethanol/Water Mixture

Results: Part C: Solubility

Experimentally determined polarity of liquids: Given that water is polar, list the other liquids

as polar or nonpolar.

Polar: _____________________________________________________

Nonpolar: __________________________________________________

Bergen Community College 22 General Chemistry II Laboratory

Date: ____________________ Name: ________________________

Experiment 2: Intermolecular Forces

POST-LAB QUESTIONS:

1. Explain the experimentally determined relative strengths of intermolecular forces. Did the

results match what was predicted in the Pre-Lab questions? Discuss any discrepancies.

2. Explain the polarity of the four liquids studied. Did the experimental miscibility results

match the polarities listed in the table on page 15? Discuss any discrepancies.

3. Explain why hexane and water do not mix.

4. Two layers are observed when sodium chloride is added to the ethanol-water mixture, a

phenomenon known as "salting out". Explain why the water and ethanol separate.

Bergen Community College 23 General Chemistry II Laboratory

Experiment 3

Spectroscopy

OBJECTIVE: To determine the molar concentration of Co+2 using spectroscopy.

BACKGROUND:

The field of Spectroscopy deals with the interaction between electromagnetic radiation

and matter. In the ultraviolet and visible portions of the spectrum, light is absorbed by a

substance when electrons move from an orbital of lower energy to one of higher energy. The

wavelength of the absorbed light corresponds to the difference in energy of the two orbitals.

When transition metal ions are dissolved in water, the energy differences of the d orbitals

correlate to wavelengths in the visible portion of the electromagnetic spectrum. Since only

certain visible wavelengths are absorbed, these solutions are colored. The observed color

corresponds to the wavelengths that are transmitted, passing through the sample without being

absorbed. Each aqueous transition metal ion has its own characteristic color. An Absorption

Spectrum shows the amount of absorbed light across a wide range of wavelengths. Because

the amount of light absorbed depends on concentration, visible spectroscopy is often used to

determine the concentration of aqueous transition metal ions.

A spectrophotometer is used to measure the amount of light that passes through a

sample. It consists of a light source, a diffraction grating to separate the light into the individual

wavelengths, a lens to focus the light, and a detector that measures the amount of light that has

passed through the sample. The light passing through the sample is quantified as the

transmittance, T, and is the ratio of the intensity of the light emerging from the sample, I, to

the intensity of the light entering the sample, I0. Oftentimes, the percent transmittance, %T, is

used. These relationships are shown below.

I

T = %T = T x 100

I0 I I0

The amount of light absorbed (removed) by a sample is quantified as Absorbance, A, and is

related to the Transmittance by the following relationship.

A = - log (%T/100) As absorbance is a logarithmic relationship, it has no unit.

When making absorbance measurements, the spectrophotometer must first be calibrated using

a solution that contains all of the components of the solution to be analyzed, other than the

absorbing species. Such a solution is referred to as a blank.

Experiment 3 Spectroscopy

Bergen Community College 24 General Chemistry II Laboratory

An absorbance spectrum is a

graph of absorbance as a function of

wavelength. The wavelength having the

maximum absorbance, λmax, is called the

optimal wavelength, or the analytical

wavelength. When making absorbance

measurements for the determination of

concentration, this wavelength is always

used. Because the absorbance value is

large, this increases the precision of the

measurements. Because this represents

the top of the curve rather than the

shoulder, it increases the accuracy.

The absorbance of a solution depends on the distance the light travels in the sample

and the molar concentration of the absorbing species, as well its identity. This relationship,

known as the Beer-Lambert Law, or more commonly, Beer’s Law, is A =  × l × c, where  is the molar absorptivity, a characteristic specific to the absorbing species, l is path length, and

c is the molar concentration. Since throughout the course of an experiment, both the path

length and the molar absorptivity remain constant, the absorbance depends solely on the

concentration of the absorbing species. Although possible to determine concentration by

comparison of an unknown’s absorbance to that of just one solution, it is much more common

to make several absorbance measurements in order to increase accuracy.

A Beer’s Law plot is a graph of absorbance

as a function of concentration. Absorbance values

for several standard solutions, those with known

concentration, are obtained at λmax and graphed.

The concentration of an unknown solution is found

using its absorbance and the equation of the line. If

the graph has been prepared by hand, the

absorbance of the unknown is located on the y-axis,

and the corresponding value on the x-axis is found.

It’s important to note that the Beer-Lambert relationship holds only for solutions that

are sufficiently dilute. If the unknown solution’s absorbance is greater than those used in the

plot, a dilution is required in order to determine concentration. Extrapolations are not possible.

In this experiment, the concentration of a cobalt(II) nitrate solution will be determined.

First, a solution will be prepared from solid cobalt(II) nitrate. Next, absorbance values of this

solution at various wavelengths will be measured, from which the analytical wavelength will

be determined. A series of standard solutions will then be prepared by dilution. Absorbance

values for these solutions will be measured at the analytical wavelength, from which a Beer’s

Law Plot can be constructed. Lastly, the absorbance of an unknown solution will be measured,

and its concentration determined from the Beer’s Law Plot.

A b

so rb

an ce

Wavelength

Absorbance Spectrum λ max

y = 6.4848x R² = 0.9997A

b so

rb an

ce

Concentration

Beer's Law Plot

Bergen Community College 25 General Chemistry II Laboratory

Date: ____________________ Name: ________________________

Experiment 3: Spectroscopy

PRE-LABORATORY QUESTIONS

1. Why are some solutions colored, and others are colorless?

2. How does a solution interact with light?

3. What is meant by the term analytical, or optimal, wavelength?

4. What is plotted on the x and y axes of an absorbance spectrum?

What is plotted on the x and y axes of a Beer’s Law plot?

5. What is the absorbance (A) of a solution with percent transmittance (%T) of 25.1?

Experiment 3 Spectroscopy

Bergen Community College 26 General Chemistry II Laboratory

REAGENTS: EQUIPMENT:

Distilled water Beaker, 50-mL

0.100 M Co(NO3)2 Volumetric flask, 50-mL

(made by dissolving 29.10 g Beral pipettes

Co(NO3)2∙6H2O in water to Volumetric pipette, 25-mL yield 1.00 L solution.) Spectrophotometer

Cuvettes

Lint-free Wipes

SAFETY ALERT:

Cobalt(II) nitrate: May cause an allergic skin reaction, allergy, asthma symptoms or breathing

difficulties if inhaled. Suspected of causing genetic defects. May cause cancer by inhalation.

PROCEDURE:

Part A: Determination of the Optimal Wavelength

1. Turn the spectrophotometer on. Allow it to warm up for 10-15 minutes before use.

2. Set the spectrophotometer wavelength to 430 nm.

3. Obtain 40-50 mL of the 0.100 M Co+2 stock solution.

4. Prepare a cuvette by adding the Co+2 solution until ¾ full. Clean with a lint-free wipe.

5. Prepare a blank by adding d-H2O to a cuvette until ¾ full. Clean with a lint-free wipe.

6. Place the blank in the spectrophotometer and zero the instrument.

7. Replace the blank with the Co+2 cuvette. Read and record the absorbance.

8. Set the wavelength to 450 nm. Zero the instrument with the blank, replace with the Co+2,

read and record the absorbance.