Notes to Instructors Part I - Bio-Linkinformation on preparations, calculations, color photos,...

INSTRUCTORS’ NOTES

LABORATORY MANUAL FOR BIOTECHNOLOGY AND LABORATORY SCIENCE: THE BASICS

INSTRUCTORS’ NOTES, PART I

2

TABLE OF CONTENTS INTRODUCTION TO THE INSTRUCTORS’ NOTES

UNIT I SAFETY IN THE LABORATORY

Classroom Activity 1: Performing a Risk Assessment

Classroom Activity 2: Exploring Safety-Related Government Web Sites

Classroom Activity 3: Responding to Emergencies

Classroom Activity 4: Understanding the Chemicals with Which You Work

Classroom Activity 5: Personal Protection

Laboratory Exercise 1: Tracking the Spread of Chemical Contamination

Classroom Activity 6: Analyzing Safety Issues in a Laboratory Procedure

Laboratory Exercise 2: Production of Bioaerosols and Factors Affecting Aerosol Production

UNIT II DOCUMENTATION IN THE LABORATORY

Classroom Activity 7: Being an Auditor

Laboratory Exercise 3: Keeping a Laboratory Notebook

Classroom Activity 8: Writing and Following an SOP

UNIT III METROLOGY IN THE LABORATORY

Laboratory Exercise 4: Recording Measurements with the Correct Number of Significant Figures

Classroom Activity 9: Constructing a Simple Balance

Laboratory Exercise 5: Weight Measurements 1: Good Weighing Practices

Laboratory Exercise 6: Weight Measurements 2: Performance Verification

Laboratory Exercise 7: Volume Measurements 1: Proper Use of Volume Measuring Devices

Laboratory Exercise 8: Volume Measurements 2: Performance Verification of a Micropipette

Laboratory Exercise 9: Measuring pH with Accuracy and Precision

UNIT IV SPECTROPHOTOMETRY AND THE MEASUREMENT OF LIGHT

Laboratory Exercise 10: Color and the Absorbance of Light

Laboratory Exercise 11: Concentration, Absorbance, and Transmittance

Laboratory Exercise 12: Preparing a Standard Curve with Food Coloring and Using it for Quantitation

Classroom Activity 10: Beer’s Law and Calculating an Absorptivity Constant

Laboratory Exercise 13: Determination of the Absorptivity Constant for ONP

UNIT V BIOLOGICAL SOLUTIONS

Classroom Activity 11: Getting Ready to Prepare Solutions with One Solute: Calculations

Classroom Activity 12: Getting Ready to Prepare Solutions with One Solute: Ordering Chemicals

Laboratory Exercise 14: Preparing Solutions with One Solute

Laboratory Exercise 15: Preparing Solutions to the Correct Concentration

Laboratory Exercise 16: Working with Buffers



3

Laboratory Exercise 17: Preparing Breaking Buffer

Laboratory Exercise 18: Preparing TE Buffer

Laboratory Exercise 19: More Practice Making a Buffer

Laboratory Exercise 20: Making a Quality Product in a Simulated Company

UNIT VI ASSAYS

Laboratory Exercise 21: Two Qualitative Assays

Laboratory Exercise 22: UV Spectrophotometric Assay of DNA: Quantitative Application

Laboratory Exercise 23: UV Spectrophotometric Assay of DNA and Proteins: Qualitative

Applications

Laboratory Exercise 24: The Bradford Protein Assay: Learning the Assay

Laboratory Exercise 25: The Bradford Protein Assay: Exploring Assay Verification

Laboratory Exercise 26: The Beta-Galactosidase Enzyme Assay

Laboratory Exercise 27: Comparing the Specific Activity of Two Preparations of Beta- Galactosidase

Laboratory Exercise 28: Using Spectrophotometry for Quality Control: Niacin

UNIT VII BIOLOGICAL SEPARATION METHODS

Classroom Activity 13: Planning for Separating Materials Using a Centrifuge

Laboratory Exercise 29: Separation of Two Substances Based on Their Differential Affinities for Two

Phases

Laboratory Exercise 30: Separation and Identification of Dyes Using Paper Chromatography

Laboratory Exercise 31: Separating Molecules by Agarose Gel Electrophoresis

Laboratory Exercise 32: Using Agarose Gel Electrophoresis to Perform an Assay

Laboratory Exercise 33: Optimizing Agarose Gel Electrophoresis

Laboratory Exercise 34: Quantification of DNA by Agarose Gel Electrophoresis

Laboratory Exercise 35: Introduction to Ion Exchange Column Chromatography

UNIT VIII GROWING CELLS

Laboratory Exercise 36: Using a Compound Light Microscope

Laboratory Exercise 37: Aseptic Technique on an Open Lab Bench

Laboratory Exercise 38: Working with Bacteria on an Agar Substrate: Isolating Individual Colonies

Laboratory Exercise 39: Gram Staining

Laboratory Exercise 40: Preparing Phosphate-Buffered Saline

Laboratory Exercise 41: The Aerobic Spread-Plate Method of Enumerating Colony-Forming Units

Laboratory Exercise 42: Preparing a Growth Curve for E. coli

Laboratory Exercise 43: Aseptic Technique in a Biological Safety Cabinet

Laboratory Exercise 44: Making Ham’s F12 Medium from Dehydrated Powder

Laboratory Exercise 45: Examining, Photographing, and Feeding CHO Cells

Laboratory Exercise 46: Counting Cells Using a Hemacytometer

Laboratory Exercise 47: Subculturing CHO Cells

Laboratory Exercise 48: Preparing a Growth Curve for CHO Cells



4

APPENDICES TO INSTRUCTORS’ NOTES

Appendix 1

Material Safety Data Sheet for Risk Reactor

Appendix 1I

Example from Student Lab Notebook (That Exceeds Expectations) and Grading Rubric

Appendix 1II

Scheduling

ACKNOWLEDGEMENTS

Thank you to Dr. Thomas Tubon who contributed excellent ideas and comments to the sections

on mammalian cell culture. Funded in part by NSF ATE grants DUE 1003498 and 0501520.



5

INTRODUCTION TO THE INSTRUCTORS’

NOTES

A. OVERVIEW

This document contains notes for instructors whose students are using Laboratory Manual for

Biotechnology and Laboratory Science: the Basics. We have included here additional

information on preparations, calculations, color photos, safety information, a sample grading

rubric, and a sample schedule. We have also tried to provide some points to begin discussion, but

we know that every instructor will have their own methods of leading discussions and their own

knowledge to contribute. Also, class results will vary and we have found that these results in

many cases take discussions into particular directions that we have not anticipated.

These notes to instructors are meant to evolve; please feel free to send ideas to us to include.

At the moment, send comments to Lisa Seidman, [email protected]. In time, we hope

adopt a more technologically sophisticated method of sharing.

B. RATIONALE

The exercises and activities in Laboratory Manual for Biotechnology and Laboratory Science

are designed to systematically introduce students to basic biological laboratory methods with an

underlying focus on ensuring quality results. Principles of laboratory quality are woven

throughout the manual so that students develop a quality “mind-set” and learn to work

thoughtfully with awareness of the many factors in the laboratory that can cause error and

inconsistency. This approach helps prepare students to work in a company that is compliant with

the Food and Drug Administration’s Good Manufacturing/Good Laboratory Practices or with

other quality systems. However, this manual is appropriate in other settings as well -- after all, no

one wants a laboratory result that is not of good quality. Remind students that research

laboratories also produce a product—knowledge—which must be of good quality to be useful.

The laboratory manual contains simple exercises that focus on basic principles. In the past,

we tried teaching students basic laboratory methods as they were working on what we considered

to be more interesting laboratory problems. For example, students were taught about

spectrophotometry while they were exploring the properties of enzymes. We found that the

students operated the spectrophotometers in a “cookbook” fashion because they were focused on

the enzymes. We therefore found the most effective way to teach spectrophotometry is to make

the instrument the focus of the laboratory period. The time invested in learning the instrument is

time well-spent. Similarly, we found that students could learn to make solutions as part of a



6

molecular biology course, but they tended to do so as quickly as possible with little thought as to

the quality of the solutions. Since the quality of laboratory solutions is of critical importance, we

found it is necessary to teach solution-making explicitly rather than as a peripheral part of some

other topic. Thus, the activities in this manual evolved to direct students’ attention to

fundamental laboratory tasks and concepts.

Professional biotechnologists work in a wide variety of settings and so we included activities

and exercises in this manual that vary from one another in the overall skills they emphasize. Most

of the exercises are written in the step-by-step format that is common in laboratory manuals but

some of the exercises deal with instructions written in a “standard operating procedure” format

that is common in companies and production settings. Several of the exercises are more research-

oriented in that they do not include procedures but rather have the students write their own

protocols.

It would be difficult to cover all the exercises and activities in the manual in one semester.

Our program is organized so that nearly all the exercises in Units I, II, III, IV, V, and VI are

completed in a one semester basic laboratory methods course that meets seven hours per week

(one hour lecture, two 3 hour labs) and a companion safety course that meets four hours/week.

Every student must complete the basic laboratory methods course before moving into more

advanced courses. We complete the other exercises in this manual at the beginning of courses

that are more specialized, including Protein Bioseparations, Molecular Biology, and Cell Culture.

C. STRATEGIES WE HAVE FOUND USEFUL

i. Documentation

Documentation is at the heart of quality practices and so documentation is a guiding theme

throughout these laboratory activities. We have found that it works to teach documentation by

having the students keep a laboratory notebook. Lab notebooks must be kept according to all the

standard principles of documentation (e.g. they must record events with verifiable dates; records

must be attributable to a particular individual). If students can keep a good laboratory notebook,

they should be able to easily learn other forms of documentation.

The students rely heavily on their laboratory notebooks during discussions. In these

discussions, every student or group reports back their data and talks about it. The instructor also

collects, reads, and grades the lab notebook entries for every student for almost every laboratory

exercise. We therefore have students purchase the type of laboratory notebook that makes a copy

of every page.

We used to require students to write one or two formal lab reports in the first basic lab

methods course. We now teach this skill in other courses because maintaining a complete and

usable lab notebook is challenge enough for beginning students. We find that if students learn to

keep a good laboratory notebook, they are able to write a materials and methods section for a



7

report later. They also learn to organize their results into interpretable tables. We require students

to write the answers to discussion questions using a word processor, and this provides an entry

into writing a discussion section of a formal report.

ii. Materials Students Hand in; Grading

Grading and assessment will, no doubt, vary between institutions and instructors. It works for us

to routinely collect and grade copies of the students’ laboratory notebook and their answers to the

Laboratory/Meeting Discussion questions. It is possible to collect the copies of the laboratory

notebook pages at the end of the laboratory period so that students can not change them at home.

A few of the exercises and activities have additional materials to hand in, such as SOPs. We

have the students write the answers to the discussion questions on a word processor to

distinguish between the type of documentation kept in a lab notebook and the type of synthesis

that typically occurs in a formal report.

iii. Face-to-Face Discussions

We have found that a class discussion at the end of nearly every laboratory exercise is one of the

most valuable teaching tools in our basic lab methods course. The students all record their data

on the blackboard. Very often students are unaware that they have made a mistake until they see

the results of others. Importantly, as students review everyone’s results, they come to understand

the causes and consequences of variability. During discussions, the instructor will usually come

back to key questions:

• What is the correct way to perform this task; what factors must be controlled?

• How do you know that you did the task properly?

• What documentation is associated with this task?

• How can variability be reduced?

iv. Other Associated Resources

A number of on-line resources can be found on the Bio-Link website (Bio-Link.org) in the

Instructional Materials Clearinghouse including: laboratory notebook guidelines, historical

laboratory notebook entries, legal considerations, activities relating to regulatory affairs, etc.

Extensive background for all the exercises and activities in this laboratory manual can be found

in Basic Laboratory Methods for Biotechnology by Lisa A. Seidman and Cynthia J. Moore, Ed.

2, ISBN: 0-321-57014-6, Publisher: Benjamin Cummings, Copyright: 2009. For students who

have difficulty with calculations, Basic Laboratory Calculations for Biotechnology (Seidman,

Benjamin Cummings, ISBN 0-13-223810-1, Copyright: 2008) provides a number of solved

problems and discussion of typical laboratory calculations.



8

UNIT I

SAFETY IN THE LABORATORY

UNIT I SAFETY IN THE LABORATORY

Classroom Activity 1: Performing a Risk Assessment

Classroom Activity 2: Exploring Safety-Related Government Web Sites

Classroom Activity 3: Responding to Emergencies

Classroom Activity 4: Understanding the Chemicals with Which You Work

Classroom Activity 5: Personal Protection

Laboratory Exercise 1: Tracking the Spread of Chemical Contamination

Classroom Activity 6: Analyzing Safety Issues in a Laboratory Procedure

Laboratory Exercise 2: Production of Bioaerosols and Factors Affecting Aerosol

Production



9

CLASSROOM ACTIVITIES 1 AND 2: PERFORMING A RISK

ASSESSMENT AND EXPLORING SAFETY-RELATED GOVERNMENT

WEB SITES

Comments: These two classroom activities are straightforward and students usually have no

difficulties. It is interesting to us as teachers to hear their stories of accidents. Sometimes

students have worked in laboratory environments and have valuable stories to share.

This is a good place to introduce the idea of looking for root causes of problems. Analyzing

problems to find their root cause is an important part of a CAPA, Corrective and Preventive

Action, program. Companies that are compliant with the Food and Drug Administration’s

regulations governing pharmaceutical manufacturing must have a CAPA program in place to

respond to problems in production. Although the context in this classroom activity is problems

that relate to safety, the concept of a problem having one or more “root causes” is still applicable.

CLASSROOM ACTIVITY 4: UNDERSTANDING THE CHEMICALS

WITH WHICH YOU WORK

Comments: Students may ask why sodium azide is so toxic. Azide is an inhibitor of cytochrome

oxidase, an important enzyme involved in respiration. This is why it inhibits the growth of

bacteria and is useful as a preservative. Therefore, sodium azide severely affects organs that are

very dependent on respiration, such as the heart and the brain.

MSDSs have the difficulty that the language used to describe chemicals is often similar for

chemicals that are hazardous and those that are not. Thus, there is much similarity in the two

MSDSs, even though one is for a hazardous chemical, sodium azide, and the other is for table

salt. Students therefore need to spend time reading the two MSDSs to find the important

differences between them.

CLASSROOM ACTIVITIES

2. Compare the NFPA ratings for health, flammability, and reactivity for the two chemicals.

Based on these ratings alone, which chemical is most hazardous? Why? The background section

already stated that sodium azide is much more hazardous than sodium chloride. This question

asks the student to look at the NFPA ratings. These ratings can be found in Section 3, Hazard

Identification. Sodium azide is identified as “highly toxic” with a health rating of 4. Sodium

chloride has a health rating of 0. This is significant because a health rating of 4 is reserved for

chemicals with significant potential to cause harm. Sodium azide has a rating of 2 for reactivity

as compared to 0 for sodium chloride. The two chemicals have equal ratings for flammability.



10

3. Look for the words “emergency” or “emergency overview” in each MSDS. 3.1. Compare the

two chemicals with respect to emergency measures described. 3.2. Describe those emergency

measures. Note the similarity of the first aid measures described in the two MSDSs. The exact

same wording is used except that, in the case of sodium azide, “Call a physician immediately” Is

used instead of “Call a physician.”

4. Find the section that describes toxicity and compare the toxicity of the two chemicals.

4.1. What is the LD50 for sodium chloride in rats? 3000 mg/kg

4.2. What is the LD50 for sodium azide in rats? 27 mg/kg

4.3. If an experimenter tested 300 rats and gave them 27 mg/kg of sodium azide, how many

rats would be expected to die in the experiment? 150 rats, since the LD50 is the amount of

the chemical that would kill 50% of the rats, on average.

4.4. What are the words used in the “Hazard Identification” section to describe the toxicity of

sodium azide in the United States? Highly toxic (USA), Very Toxic (European Union)

4.5. Is sodium azide toxic by inhalation? Yes

4.5.1. At what level is inhalation toxicity reported? In a mouse, sodium azide is toxic at

32.4 mg/m3.

4.5.2. How could you avoid exposure by inhalation? Wear a mask or respirator to weigh

out the chemical. Use a chemical fume hood. Alternatively, if possible, only work with the

chemical in solution form.

4.6. What are the routes of entry for sodium azide into the body? Oral (swallowing);

absorption through skin, eyes; inhalation.

4.6.1. Which route(s) do you think is the biggest danger to you as a laboratory technician?

Absorption through skin. Students may also say inhalation.

4.6.2. How could you protect yourself ? For absorption hazard, wear PPE, including

gloves, when using solutions containing sodium azide. For inhalation, see 4.5.2 above.

5. Find the section that deals with disposal. Compare the language used for NaN3 and NaCl.

5.1. What do you think could happen if sodium azide solution was poured down the drain of

the laboratory sink? Since sodium azide is reactive as well as toxic, contact with metals

(pipes) should be avoided (Section 10). Toxicity is also a major concern so a licensed

professional should be consulted.

5.2. How does the MSDS direct you to dispose of sodium azide? The disposal

recommendations for sodium azide specify a licensed professional waste disposal service but

sodium chloride can be disposed of “by methods consistent with local, state and federal

regulations.”

6. Your supervisor asks you to make a 10% stock solution of sodium azide in water. When you



11

get the stock bottle, you notice the skull-and-crossbones on the label:

6.1. What concerns do you have about making this solution? Inhalation, when weighing out

the chemical, would be the first concern.

6.2. After reading the MSDS for sodium azide, what could you do to protect yourself? Wear

full PPE and use a chemical fume hood and respirator to weigh the chemical.

6.3. What should you do to protect others in the laboratory from exposure to sodium azide?

Other than measures described in 6.2, accurately label all solutions and reagents containing

the chemical.

6.4. What should you do to protect the environment from sodium azide? Employ a licensed

professional to dispose of waste, as described in Section 13 of the MSDS.

6.5. How would you label the solution that you make? Skull and crossbones, “CAUTION”,

“TOXIC,” etc. Use a red laboratory marker.

7. You are performing an experiment of a certain type for the first time. The procedure you are

following calls for incubating your sample with Blocking Buffer. You find the buffer in the

refrigerator, and there is a label on the buffer bottle:

Blocking Buffer

PBS with 3% BSA and 0.02% azide

Based on the presence of the sodium azide in the solution, what precautions do you think you

should take when performing this experiment? Since sodium azide is already in solution at a low

concentration, the possible hazardous route of entry would be through skin absorption. In this

case, regular PPE and gloves would be adequate protection.

CLASSROOM ACTIVITY 5: PERSONAL PROTECTION

1. As an individual, provide a rationale for the following safety rules; that is, explain what might

happen if the rule is not followed. Possible answers include the following:

1.1. Do not wear necklaces or neckties in lab.

• They might be contaminated by microbes or might contaminate a culture by dangling in it.

• They might get caught in moving equipment, such as a centrifuge, and are a choking

hazard.

• They might dangle in a hazardous substance, such as an acid or base.

1.2. Do not wear loose clothing in lab.

The same issues as apply to 1.1 apply but with less possibility of choking



12

1.3. Avoid wearing jewelry in lab.

A hazardous substance, such as an acid or microbe, might be trapped underneath holding the

hazard next to the skin

1.4. Do not launder your lab coat at home (once you are no longer a student).

Potentially contaminated items from the laboratory should never be taken home where they

might affect not only you but also family and friends.

1.5. Do not wear open-toed shoes in lab.

• Hazardous substances might spill on ones feet.

• Heavy objects might drop on ones feet.

1.6. Keep fingernails short.

• Long fingernails can tear through gloves.

• The nails provide a surface that might become contaminated.

1.7. Do not leave Bunsen burners unattended.

The possibility of burners causing burns and fires is obvious. Students should realize that

another person may not realize that a burner is lit and might accidentally reach across it. Also,

invisible fumes from volatile chemicals used elsewhere in the laboratory can catch fire.

1.8. Use water baths, not Bunsen burners, when organic solvents are in use.

Organic chemicals are flammable. The obvious hazard is the possibility of fire.

1.9. Wear face shields when using UV light.

The UV light in these laboratory sources is dangerous to the skin and eyes. Face shields

protect the skin from severe “sunburn” and also protect the eyes from light that might be

reflected underneath glasses. The light sources are very strong and can cause serious burns

quickly.

1.10. Always dispose of sharps items, like broken glass, in the appropriate container.

This rule protects not only laboratory workers, but also custodial staff who might

inadvertently be injured by sharp objects.

2. Determine what type(s) of gloves are available in your laboratory and of what material they

are made. Consult the manufacturer’s information regarding these gloves. Will these gloves

protect you from all chemicals? What are the limitations of these gloves? The answer depends

on the type of gloves available. In general, all glove materials are vulnerable to certain classes

of chemical. Students will explore the compatibility of glove materials with chemicals in

Classroom Activity 6. Also, inexpensive gloves may have small holes.



13

LABORATORY EXERCISE 1: TRACKING THE SPREAD OF CHEMICAL

CONTAMINATION

Comments. Fluorescent compounds are often used to practice good technique because they

demonstrate that contamination can occur that is not immediately visible and of which the

worker is unaware. However, using fluorescent compounds will not disclose low levels of

contamination and might therefore lead students to think there technique is perfect when it is not.

Students should know that when working with pathogens, for example, a very low level of

exposure might be hazardous.

LAB MEETING/DISCUSSION QUESTIONS

1. What precautions do the MSDSs recommend for the materials used in this lab? Students

should be cautious with these chemicals, as with all laboratory chemicals.

• Fluorescein. Irritant, mutagenic in yeast and bacteria. NFPA ratings, 2 for health and 1 for

flammability.

• Rhodamine. Harmful if swallowed, skin irritant, can cause severe irritation to the eyes. NFPA

rating 2 for health and 0 for flammability

• Tide detergent is rated 1 for health and 1 for flammability. It can cause transient irritation to

the eyes and transient irritant to the respiratory system if inhaled over an extended period of

time.

• We do not have MSDSs for DayGlo. The MSDS for Risk Reactor is provided in Appendix 1

of this document. It states that Risk Reactor is rated 1 for health, 1 for flammability, and 0 for

reactivity. More information about Risk Reactor products is available at

www.riskreactor.com.

2. What areas exhibited contamination? Do you think that you could minimize the likelihood of

having these areas of contamination in the future? What precautions would you need to take? It

is common for one or more students to observe contamination on their gloves. This is an

excellent opportunity to point out that gloves that are contaminated have the potential of

spreading contamination around the laboratory. If students use gloves they must change them

frequently, whether or not they are aware that the gloves are contaminated.

4. When working with hazardous chemicals in the context of a procedure that you have not

previously performed, it can be helpful to perform a “dry-run” of the process. This exercise was

essentially a dry-run. Imagine that you must make Dilutions I–IV with chemicals that are

hazardous by inhalation and ingestion and that are not fluorescent or visible. What changes to the

procedure, as written, would you make on the basis of your dry-run results? The most obvious

change is the use of a chemical fume hood to protect oneself from hazards due to inhalation.

Students may come up with other ideas based on their experiences in the exercise.



14

CLASSROOM ACTIVITY 6: ANALYZING SAFETY ISSUES IN A

LABORATORY PROCEDURE

Comments. The purpose of this activity is to introduce students to the concept of developing a

safety “mindset” and the importance of thinking through the safety issues involved in a

procedure before launching into that procedure in the laboratory. Thinking through even this

relatively short procedure requires investigating MSDSs, issues of chemical compatibility,

equipment usage, and disposal.

The particular procedure in this activity is not commonly used anymore – probably at least in

part to the safety issues involved. We use it because it brings up a number of issues that require

thought and can be used to begin useful classroom discussion.

The “correct” answers in some cases will vary depending on local regulations and on the

institution’s chemical hygiene plan. Note also that some of the difficulties in this procedure are

alleviated by having institutional procedures for handling hazardous chemicals.

ANALYSIS

1. Identification of Hazardous Chemicals

1.1. List the chemicals in this procedure. Phenol, chloroform, DNA, water, sodium acetate,

ethyl alcohol.

1.2. Consult the MSDS for each chemical to determine special precautions to be taken when

handling this material. You can find the MSDS for each chemical online at the Sigma-Aldrich

web site.

• Phenol is one of the most hazardous chemicals in the biologist’s stock room. Students

should note that even though all MSDSs contain what seem to be dire warnings, the

MSDS for this compound looks different. Notably, the NFPA rating for health is 4. The

phenol MSDS has extensive toxicological information. In addition to its health hazards,

phenol is rated 2 for flammability. Phenol is extremely hazardous by all methods of

exposure. It causes severe burns to skin (possibly without the sensation of burning). It

damages the lungs and in extreme cases can be fatal by inhalation. It can cause blindness.

It is potentially lethal if ingested. Note also that adding chloroform to phenol enhances the

ability of phenol to be absorbed by the skin.

• The NFPA rating for health for chloroform is 2. It is an irritant and is carcinogenic. It is

not flammable or reactive. The neurological effects of inhaling chloroform are familiar

and most students are aware of this hazard. Inhalation can cause dizziness, headache, and

drowsiness. Exposure to high concentrations can cause serious neurological effects

including delirium and unconsciousness. Inhalation of 25,000 ppm chloroform for five



15

minutes can cause severe burning to the respiratory tract and can be fatal. Observe,

however, that the amount of chloroform used in this procedure is relatively low.

• Aqueous DNA solutions should be handled with respect, as should all laboratory

chemicals, but are not considered to be hazardous. Students may not be able to find an

MSDS for this chemical.

• Sodium acetate has an NFPA rating of 2 for health and 1 for flammability. It is an irritant

for the skin and eyes and is potentially hazardous if ingested. Note, however, that sodium

acetate is also used in foods as a preservative or for flavor. This may be confusing to

students because it has the same NFPA rating as chloroform, which is more of a concern.

• Ethyl alcohol (ethanol) is a familiar chemical to most students. The particular grade

specified in the procedure is not denatured and therefore does not have isopropyl or methyl

alcohol added to it. (Denatured alcohol is more toxic by ingestion than pure alcohol.) The

NFPA ratings for ethyl alcohol are 2 for health and 3 for flammability, so fire hazard is

significant. Ethyl alcohol is a health hazard by all means of exposure. The quantity in this

procedure, however, is low.

2. Selection of Protective Equipment and Attire

2.1. At what steps should a fume hood be used? The steps involving phenol should be

performed in a fume hood.

Based on the MSDS and the fire hazard, one might consider using ethyl alcohol in a fume

hood, but common experience makes this seem unnecessary. This issue might be worth

discussing with students.

2.2. What type of protective wear should be worn? Laboratory coats, gloves (see question

2.3) and chemical goggles should be worn.

2.3. Research the appropriate type of gloves to use. Assume that the school gets its gloves

through Lab Safety Supply (http://www.labsafety.com/refinfo/ezfacts/ezf166.htm), and that

there are Microflex disposable latex and disposable nitrile gloves in the lab. If these gloves are

not appropriate, assume that nondisposable gloves from Ansell are also in the lab. Links to

these brands are available on the Lab Safety Supply site above. Disposable latex and nitrile

gloves are not recommended for use with chloroform and phenol. Ansell’s laminate film

glove will work with both chemicals. Note, however, that many laboratory workers would use

disposable latex or nitrile gloves for this procedure because the amounts of each chemical are

small and the time of exposure is minimal. A document from UCLA’s environmental health

and safety office (http://ehs.ucla.edu/PhenolFactSheet.pdf) notes that the breakthrough time

for nitrile gloves exposed to chloroform is 3 minutes. They suggest the use of neoprene gloves

for larger amounts of phenol-chloroform or wearing two pairs of nitrile gloves and changing



16

them immediately if they contact the chemicals. Each instructor will need to decide how best

to handle this question.

3. Choose Equipment

3.1. List the equipment required for this procedure. Clinical centrifuge, high speed

centrifuge, centrifuge tubes, freezer, 5mL serological pipets and bulbs. The serological pipets

should be glass, Microfuge tubes can be used to store the DNA pellet.

.

3.2. Choose tubes that are compatible with the chemicals in use:

• Assume that the centrifuge tubes in the lab are made from polycarbonate or

polypropylene.

• Assume that the centrifuge tubes in the lab are from Nalgene.

Will these tubes be acceptable? Find out by checking the chemical compatibility tables in the

Nalgene reference guide

(http://www.nalgenelabware.com/pdf/NalgenePlasticsTechGuide0209.pdf) at the Nalgene

web site.

Ethanol and sodium acetate are not a problem with these tubes. Chloroform and phenol,

however, are not compatible with polycarbonate and polypropylene. One possibility is to use

Corex centrifuge tubes, DuPont catalog # 00156, for the steps with phenol and chloroform

that are performed in a clinical centrifuge. Possibly people would use plastic tubes for this

procedure even though they are not recommended, because the procedure is relatively quick.

However, it would be advisable to test the tubes first by placing phenol and chloroform in

them and placing the tubes in a beaker in the hood. This test can be performed for a relatively

short time (e.g., half an hour) because individual tubes will not be exposed to the chemicals

for longer than this.

4. Research Hazardous Waste Disposal Methods

4.1. Consider how waste should be handled during the experiment. Students should be aware

that methods of handling waste are governed by local, state, and federal regulations. Most

institutions have developed systems for dealing with hazardous waste that comply with the

appropriate regulations. Therefore, laboratory workers often rely on a preexisting system of

managing waste. In a classroom setting, the instructor and laboratory manager will tell

students where to put chemical wastes. Students should be aware that they should never mix

chemical wastes without checking with the instructor first and they should never dispose of

chemicals in the waste basket or sink without permission.

For specific information about wastes, we suggest that you obtain one or more copies of the

Flinn catalog that students can share. The Flinn catalog has extensive information about



17

handling hazardous chemicals.

Some notes based on the Flinn catalog:

This procedure will generate small volumes of ethyl alcohol with dissolved sodium acetate.

Assuming that the laboratory’s drains are connected to a sanitary sewer system with a water

treatment plant, then ethyl alcohol can go down the sink drain. Wastes that can be disposed of

in the drain should be first dissolved in water and then rinsed down the sink with a tenfold

excess of water. Solid sodium acetate can be put in normal waste for disposal in a landfill and

so, presumably, is of low enough hazard to be washed down the drain with the alcohol.

Small quantities of chloroform can be allowed to evaporate in a chemical fume hood.

However, in this case the chloroform is mixed with phenol. Phenol is considered to be a

hazardous waste and is best bottled and disposed of by a licensed hazardous waste handler. In

this case, a glass bottle should be labeled for storing the phenol/chloroform waste. The bottle

should be stored in the chemical fume hood until it is removed by the professional waste

handlers. The waste would be transferred into the bottle inside the fume hood.

LABORATORY EXERCISE 2: PRODUCTION OF BIOAEROSOLS AND

FACTORS AFFECTING AEROSOL PRODUCTION

Comments. Many of the same issues pertain to this laboratory exercise as do to Laboratory

Exercise 1.

Two particular ideas to emphasize in this exercise are the ease with which aerosols are

produced and the idea that, in some situations, exposure to very low levels of pathogens can be

harmful. Laboratory workers have been infected by pathogens due to exposure through the skin

(e.g., by needlesticks). This type of exposure is certainly a concern, but bioaerosols are a serious

danger because they are invisible and people often are unaware of their presence.


2. What do the results of this laboratory suggest about the need for proper cleanup following

handling of biohazardous liquids? Surfaces should be routinely disinfected because invisible

airborne particles can land on them. It is also good practice to assume that instruments, such as

centrifuges and pipettes, also become contaminated and should be routinely disinfected.

3. In what ways can you envision personnel (both laboratory and nonlaboratory) contracting



18

laboratory-acquired infections? As noted above, exposure via sharp items is one of the major

causes of laboratory acquired infection. However, exposure by inhalation of aerosols and by

contact with contaminated surfaces is also possible. Ingestion becomes a concern when workers

are careless about eating and drinking in the laboratory, fail to wash their hands after finishing

work, or use other poor practices.

4. Can you suggest ways to minimize bioaerosol production in any of the laboratory procedures

investigated? For each task that you performed, are there alternative procedures that would

reduce bioaerosol production? Students will come up with various tactics to reduce aerosol

formation in routine laboratory work. You might want to point out that it is impossible to

completely eliminate aerosols using standard methods, such as we use in a teaching laboratory.

Therefore, scientists who work with pathogens use a number of sophisticated practices and have

special equipment to protect themselves and the environment. For example, centrifuges that are

used with pathogens often have special tubes, bottles, and rotors to safely contain pathogenic

samples. This is because the act of centrifugation is very likely to produce aerosols. There are

special hoods that are used to work with pathogens. In extreme cases, people wear suits that

cover their bodies entirely and have special breathing apparatuses. There are also facilities with

specialized engineering features devoted exclusively to handling the most dangerous pathogens.



19

UNIT II

DOCUMENTATION IN THE LABORATORY

Classroom Activity 7: Being an Auditor

Laboratory Exercise 3: Keeping a Laboratory Notebook

Classroom Activity 8: Writing and Following an SOP



20

CLASSROOM ACTIVITY 7: BEING AN AUDITOR

PART A: Suggestions for discussion of the scenes in Tour of the Facility (refer to

Table 2.2 on pages 50-51 in the manual):

Scene 1: The analyst is using pencil in a laboratory notebook. This violates the principle that

records should not be capable of being altered, either accidentally or intentionally.

Scene 2: The technician is collecting data on a sheet of paper that is not in a bound laboratory

notebook. This violates the principle that records must record events with clear and verifiable

dates because the pages can be put into a binder in any order. It also violates the rule that records

should not be capable of being altered, either accidentally or intentionally because pages can

simply be thrown away from a three ring binder.

Scene 3: The scientist has crossed out an entry in such a way that it cannot be read. This violates

the rule that records should not be capable of being altered, either accidentally or intentionally.

Scene 4: In this case, the analyst has left a lot of white space at the top of the page. Entries

should be continuous to record events with clear and verifiable dates. If there is white space in

the notebook, there is room for entries to be made at a later time. Students have often been taught

to make their lab notebooks “neat” and they therefore like to leave space to fill in information at

a later date -- this is a poor practice.

LABORATORY EXERCISE 3: KEEPING A LABORATORY NOTEBOOK

Comments. This activity helps students become accustomed to recording details such as dates,

equipment used, ID numbers on instruments, chemical names, catalog numbers, lot numbers,

methods, and results. Although lab notebooks may or may not be used in laboratories where

samples are tested (e.g. a quality control or forensics laboratory), the skills learned by keeping a

good laboratory notebook are readily transferable to other forms of documentation.

As part of this first lab exercise, students make “slime” because it is fun and is simple enough

not to distract students from the tasks of documentation. Other laboratory tasks will work as well.

Witnessing another person’s laboratory notebook is required in many professional settings,

for example, wherever work might lead to a patent. It is therefore good experience for students to

witness their classmates’ notebooks. It is also a good way for students to learn from others. We

put together the checklist on page 65 of the manual to facilitate this process of witnessing. In our

classes, we do not witness notebooks routinely because of time constraints, but it might be a

good idea to do so.



21

CLASSROOM ACTIVITY 8: WRITING AND FOLLOWING AN SOP

Comments. Standard operating procedures, SOPs, are critical parts of documentation in almost

any biotechnology company. SOPs are more common in industry than in academic research

laboratories but SOPs or analogous documents can be used to direct researchers in routine

operations.

Alternatives:

This laboratory exercise requires students to write an SOP that someone else can follow. The key

feature of the exercise is that students produce a product based solely on what their partner

writes down in the SOP. If the SOP is poorly written, the product will not be correct. We

sometimes have the students write SOPs for building toothpick structures that are held together

with Elmer’s Glue. (In this case, the instructor provides original models for the students to copy.)

Toothpicks are inexpensive and many students enjoy building structures. Note, however, that

building toothpick models requires some patience, some manual dexterity, and a bit of finesse.

Some students find this frustrating, but we remind students that many laboratory tasks also

require some manual dexterity and persistence (think of cutting thin sections for microscopy,

pouring polyacrylamide gels, or loading samples into an agarose gel.) We assure students that

they are graded on their documentation -- not their skill in working with toothpicks. If you use

toothpicks, tell the students that the glue must be fairly tacky (dry) before the toothpicks will

stand up. The option that is provided in this laboratory manual is to use marshmallows to hold

the models together thus avoiding the aggravations of glue. Marshmallow models tend to be

simple. There are many other options for this exercise, such using tinker toys and commercially

available model kits. Some teachers have students write SOPs for making coffee, peanut butter

sandwiches, popcorn, and other familiar jobs.

Students should build the model at the same time as they write an SOP for that model because

a person should never write an SOP to do a task she or he has not done. This means that students

construct two models: They construct the one for which they write the SOP and they construct

one from their partner’s SOP. This exercise as written usually requires a three hour laboratory

period.



22

UNIT III

METROLOGY IN THE LABORATORY

Laboratory Exercise 4: Recording Measurements with the Correct Number of

Significant Figures

Classroom Activity 9: Constructing a Simple Balance

Laboratory Exercise 5: Weight Measurements 1: Good Weighing Practices

Laboratory Exercise 6: Weight Measurements 2: Performance Verification

Laboratory Exercise 7: Volume Measurements 1: Proper Use of Volume Measuring

Devices Laboratory Exercise 8: Volume Measurements 2: Performance Verification

of a Micropipette

Laboratory Exercise 9: Measuring pH with Accuracy and Precision

GENERAL DISCUSSION POINTS FOR THIS UNIT

1. How much variability is there in class results?

There is always some variability in measurement and test results; if there is no variability the

system is either malfunctioning or insensitive. Students should be alert to an instrument that

keeps providing the same value over and over again – it may be malfunctioning. However,

variability that is due to mistakes should obviously be avoided. It is not unusual for students to

observe that their own data varies substantially from the values of the rest of the class. This is a

good jumping off point for a discussion of differences in technique, possible errors, and a

discussion of how much variability is acceptable.

2. Where do variability and loss of accuracy come from?

Categories to discuss:

• The equipment (e.g., maintenance, calibration, inherent design features)

• The analyst’s technique (e.g., maintaining constant temperature, reading the display, using a

pipette)

• The sample (e.g., chemical reactions that occur while the sample is being analyzed)

• Random variability (subtle fluctuations that cannot be predicted or avoided)



23

3. Why is variability a problem?

Variability is always a problem for product quality. An obvious example is a drug product that

varies in the concentration of active product. Variability is a problem in a research setting also.

The goal of a scientist is to control all the potential variables in an experiment except those being

studied. If the components of solutions vary, if measurements are being made in differing ways,

if samples are being treated differently, then the results of an experiment are likely to be

meaningless, or at least hard to interpret.

4. How can variability be reduced?

Variability is reduced by the use of standard operating procedures, documentation (so one knows

what parameters to use each time), proper use of instruments, maintenance and calibration of

instruments, control of environmental factors, training and education, etc.

5. How can you achieve the most accurate measurements?

Achieving accuracy will require the same understanding and use of good technique as controlling

variability. You might point out to students, however, that a device that is improperly calibrated

might provide good precision (low variability) but it will not provide accurate results. Similarly,

using a device in a consistent but incorrect way will provide good precision but inaccurate

results. In each section (e.g., volume measurements, weight measurements) point out how

understanding that particular measurement is necessary to avoid error. For example, when

weighing in the milligram range or measuring the pH of buffers, controlling temperature is

important. If an analyst does not understand temperature effects, his or her measurements may be

erroneous.

6. How can you determine and document the accuracy of your measurements?

Some ideas to discuss:

• Determining accuracy requires knowing the “right” answer. It isn’t possible to know, without

any uncertainty, the right answer for a sample. But it is possible to check a measuring system

by using standards, such as standard weights for a balance and standard pH buffers to check

pH meters. We frequently have students obtain measurements for a standard and then

calculate the percent error of the measurement as a way to evaluate and document the

accuracy of a measuring system.

• In many cases in the classroom, students are given “unknowns” where the instructor “knows”

the correct answer. The values provided by the instructor in an answer key can be used by the

students to assess the accuracy of their measuring system and their technique. Most analysts

would then assume that their values for real samples will achieve a similar level of accuracy.

• It is often worth discussing the reliability of the instructor’s answer key. The instructor may



24

(or may not) be using recently calibrated equipment. The instructor is more knowledgeable

about the proper operation of devices and is therefore more likely to provide accurate results

– but is it possible for instructors to make errors? The instructor may be using materials

provided by a manufacturer who, in turn, provides some assurance of accuracy. Can these

assurances be trusted?

7. How close is “close enough?”

This is a challenging question because the answer is: “it depends…” In a work setting, the limits

of accuracy and precision are determined based on experience with the measurement or assay

and knowledge of the features required in a final product. For example, weighing microgram

amounts of a drug in a pharmaceutical quality control laboratory requires better accuracy (and

sensitivity) than weighing apples at the grocery store. There are commonly used statistical

quality control methods that help to set limits on accuracy and precision (e.g., see pages 287-289

in Basic Laboratory Methods for Biotechnology, Seidman and Moore).

8. The U.S. Code of Federal Regulations outlines the following requirement for

pharmaceutical companies: “The calibration of instruments, apparatus, gauges, and recording

devices at suitable intervals in accordance with an established written program containing specific

directions, schedules, limits for accuracy and precision, and provisions for remedial action in the event

accuracy and/or precision limits are not met (21 CFR 211.160, rev. April 2000).” Suggest a

program for ensuring that the measuring devices in your classroom are compliant with

this regulation.

Strategies include but are not limited to:

• SOPs that specify how and when items should be calibrated and maintained

• Established schedules for maintenance and calibration

• Documentation of maintenance and calibration for each measuring item

• A method to ensure that any measuring item used has been calibrated and maintained within

the specified timeframe

• Methods to verify the performance of each measuring item at regular intervals (e.g., checking

the weight of a mass standard every morning and noting the result in a logbook)

• SOPS to tell analysts how to recognize a questionable measurement

• SOPs that specify how to investigate, correct, and prevent further measurement errors



25

LABORATORY EXERCISE 4: RECORDING MEASUREMENTS WITH

THE CORRECT NUMBER OF SIGNIFICANT FIGURES

Comments: It is important that students learn to read measuring devices with the correct number

of significant figures. If the value from a calculation is not rounded according to significant

figure conventions, the problem can be corrected later – but only if the initial readings are

recorded correctly.


1. This question relates to the temperature data in Figure 3.6 in the manual. The thermometer

data can be used to illustrate several important points.

• First, the students should all read the thermometers with the same number of significant

figures; all the figures that are certain plus one that is estimated.

• Second, technique matters. The students’ values should only vary in the last decimal place.

The student data in Figure 3.6 vary more than this, possibly because some of the students

partially removed the thermometers from the beakers in order to see them better. The

thermometers should not be lifted from the water nor should they be touched to avoid

introducing error. It can be illustrative to have the students read the values a second time,

after discussing the importance of technique, to see if the variability in the readings is

reduced.

• A third point is that alcohol thermometers for student use are manufactured, at best, to be

accurate only + 1° C. This is consistent with the student results in Figure 3.6. Students may

consider whether they can achieve a better value for the water temperature by averaging the

results of all the thermometers used.

2. This question relates to the length data in Figure 3.6. There tends to be less variability in the

measurements made with rulers, probably because subtleties in technique are not as important as

with thermometers and students are more familiar with rulers.

3. This question relates to the volume data in Figure 3.6. All the students should have the same

number of significant figures for each measurement. The values should be read from the bottom

of the meniscus. Ideally the meniscus should be at eye level, but you may want to discuss

whether there are times not to bring the cylinder close to the eyes to avoid exposure to liquids

and for convenience.

4. Put the values for your class on the blackboard. The students can compare their data to that of

the class shown in Figure 3.6 and can discuss the same issues based on their own data. Most



26

importantly, be sure that everyone reads every measurement with the same number of significant

figures. See also the general comments on pages 22-24.

5. Explain in your own words how significant figures relate to the uncertainty of a measurement.

The International Standards Organization (ISO) points out that the word “uncertainty” means

doubt, and thus in its broadest sense “uncertainty of measurement” means doubt about the result

of a measurement. One of the causes of uncertainty in any measurement is the fact that no

measuring device can be read to an infinite number of places. This means that we cannot be sure

about the value of the measurement past the places read by the instrument. Manufacturers put

numbers or displays onto their instruments that we read. The manufacturers should create these

displays so that their readings are accurate with little or no doubt except for the last decimal

place that is estimated.

CLASSROOM ACTIVITY 9: CONSTRUCTING A SIMPLE BALANCE

Comments: Points to raise:

• This activity shows how “balances” got their name

• Observe that weighing compares the force of gravity on the sample with the force of gravity

on standards

• A more sensitive balance allows weight measurements that are more exact and have more

significant figures. It is the fineness of the knife edge, the edge on which the beam is

balanced, that determines the sensitivity of the balance. Figure 1 below illustrates a balance

of higher sensitivity. Figures 2A and 2B show that the more sensitive balance is able to

respond to the addition of ½ an M&M. Figures 3A and 3B illustrate a less sensitive balance

that barely responds to the addition of a whole M&M.

In early laboratory balances, the beam was balanced on a very thin edge (knife edge). This

was one of the problems in designing these early laboratory balances – a fine knife edge

means the balance is fragile and extremely sensitive to motion. For this reason, balances with

a beam have a way to arrest the beam when materials are added to or removed from the

weighing pans.



27


2. What does it mean to “weigh” something? The devices the students construct illustrate the

basic principle of weighing, that is, comparing the effect of gravity on a sample and a standard.

In this activity, the standards are M&Ms so the weight of the sample is expressed in units of

M&Ms. In the laboratory, we express weight in units of grams. (In the laboratory, the standards

are objects of known mass. It therefore seems that in the laboratory we should be measuring the

mass of the sample when it is compared to a mass standard. In reality, there is a slight error in

this comparison because the standards are usually denser than the samples. The standards

therefore experience less buoyancy by the air than the samples. In practice, this slight error due

to buoyancy differences is generally ignored by biologists, but it means that in practice we do not

find the true mass of our samples, rather, we find their weight.)

3. How can you maximize the accuracy of your balance? The students should see that where they



28

place the weighing pans on the lever and where they place the items in the weighing pan will

affect the accuracy of their measurements. (The same is true in a modern balance where the

sample should be placed in the middle of weighing pan.) The students will likely also talk about

the accuracy with which M&Ms are manufactured – a factor over which they have no control.

The students are likely to try cutting the M&Ms to obtain fractions of an M&M unit and this is a

task that can be controlled to some extent.

4. How can you maximize its precision? A discussion of technique is appropriate; for example,

some students will hold the lever steady while adding objects to the weigh pan. Others will not.

Some students are more careful about placement of samples and standards in the weigh boats and

placement of the ruler on the center block.

5. What does the term “sensitivity” mean with regard to weighing? Which group constructed a

balance with the most sensitivity? What are the disadvantages to a balance that is very sensitive?

The sensitivity of these balances is largely determined by the fineness of the center block edge.

The more sensitive devices are more difficult to control and use and are more sensitive to

movement and jostling.

6. Older style mechanical analytical balances almost always have an arrested position where the

beam is locked so it will not move. Standards and samples are added to the beam while the

balance is in the arrested mode. Based on your observations, why is this true? This is evident

after the students try to place their items on the weighing pan.

8. How many significant figures do your measurements have? We consider the number of

significant figures to be 2, for example, 3.2 M&M units, assuming that the students cut the

M&Ms to subdivide them. Do you have a different answer for this question? If so, please let us

know.

9. What factors would a user of your balance need to consider in order to obtain the best

possible measurements? Students usually talk about centering the ruler and pans, holding the

lever steady while adding items to the balance, cutting the M&Ms carefully, etc.



29

LABORATORY EXERCISE 5: WEIGHT MEASUREMENTS 1: GOOD

WEIGHING PRACTICES

Comments. Part A. This is a good opportunity to introduce students to the instruction manuals

that come with equipment. (They will probably figure out quickly why most companies write

their own SOPs for equipment use, even when a manual from the manufacturer exists.) Students

can take turns learning how to use a particular balance, based on the instruction manual, and then

teaching other students how to operate that piece of equipment.


1. What do the terms “accuracy” and “precision” mean in the context of weight measurements?

See the general comments about metrology on pages 22-24.

2. What was the effect of temperature? Students often predict that handling an object will cause

its weight to increase due to oil from their fingers. However, the measured weight actually goes

down due to air currents. (Mettler’s technical literature discusses this phenomenon.) It is

common to see variability in the results of the effect of temperature on weight; some students see

the weight go up with handling, others report that it goes down. Perhaps some of this variability

may be due to student expectations of what should happen. Students generally find, however, that

handling samples before weighing will affect the results.

3. What factors in addition to temperature will affect the accuracy of your weight

measurements? Modern analytical balances are precise instruments that are capable of accurately

measuring the weight of an object to at least the nearest 0.1 mg. Major causes of inaccuracy in

weighing are poor technique (e.g., failing to close balance draft shield doors, failing to place the

sample in the middle of the weighing pan), vibration and jostling, temperature effects, failure to

level the balance, and lack of balance maintenance and calibration. Students should also note that

samples can change weight due to loss or absorption of moisture. Static charge is another major

cause of difficulty for some samples. See also General Comment #2 on page 22.

4. What did you find out about the accuracy of your weight measurements? See the general

comments on pages 22-24. Students generally realize quickly that if one individual or group, or

one item of data is substantially different than all the others, it should be investigated.

5. How does the design of the balance(s) you used for the current laboratory exercise compare

to your M&M balance? Students realize that electronic balances do not look like double pan

balances. Point out that modern balances place the standard on the weighing pan before placing

the sample on the same pan, rather than have them both on different pans at the same time.

Modern electronic laboratory balances seldom have a beam. However, all these balances

compare the effect of gravity on standards and samples.



30

LABORATORY EXERCISE 6: WEIGHT MEASUREMENTS 2:

PERFORMANCE VERIFICATION

Comments: This exercise verifies the accuracy and performance of a balance. It also introduces

the ideal of linearity. Linearity refers to the instrument’s ability to deliver identical sensitivity

throughout its weighing range. If the response of a balance is not linear, then its results will not

be accurate throughout its range. A simple way to test the linearity of a balance at its midpoint is

to weigh two stable objects separately, each of approximately one half the weighing capacity of

the balance. The sum of the two readings should equal the reading obtained when both objects

are weighed together. The procedure in the manual is a little more complicated because it also

checks linearity at 25% and 75% of capacity. However, the simplified method is sufficient to

teach the principle of linearity and may be convenient if four suitable objects, each roughly ¼ of

the capacity of the balance, cannot be obtained.

Be sure that students understand the difference between calibrating their balance and doing a

performance verification of their balance. In the first case, the response of the balance is adjusted

so that the balance’s readings are in conformance with international standards. Some balances

cannot be calibrated by the user. This is the case with older style mechanical analytical balances.

Laboratory balances that cannot readily be calibrated by the user are typically calibrated

annually, or at another time interval, by a qualified technician.


1and 2. Discuss the results. See the general comments on pages 22-24.

Here are typical (and acceptable) class results for analytical balances. (These balances are

professionally maintained and calibrated once a year, but do experience heavy classroom use.)

3. Explain linearity and its importance in your own words. Answers vary. Students should

understand that nonlinearity affects the accuracy of a balance.



31

4. Imagine that an auditor or inspector comes to your laboratory and asks, “How do you know

that your weight measurements are correct”? How could you answer? Be sure that students

understand traceability in this context. Traceability refers to the “genealogy” of their standards,

going all the way back to the kilogram standard in France.

5. Complying with the CFR. See the general comments for metrology on page 24.

LABORATORY EXERCISE 7: VOLUME MEASUREMENTS 1: PROPER

USE OF VOLUME MEASURING DEVICES


1 and 2. What factors affect the accuracy of volumes dispensed by a micropipette? See the

general comments on pages 22-24.

LABORATORY EXERCISE 8: VOLUME MEASUREMENTS 2:

PERFORMANCE VERIFICATION OF A MICROPIPETTE

Comments. It is good practice to verify the performance of equipment, particularly

micropipettes. In a regulated company, this is done on a regular basis. This exercise also

provides a good illustration of how standard deviation is applied in a quality environment.


1 and 2. Discuss the class’s results. Note that pipetting technique is very important. Therefore,

when a micropipette does not meet its specifications, it is difficult to know whether the device is

out of spec or the operator is not using it properly. We find it useful to observe each student’s

pipetting technique one at a time before beginning the performance verification lab.

3. See the general comments on pages 22-24.



32

LABORATORY EXERCISE 9: MEASURING PH WITH ACCURACY AND

PRECISION


1. Discuss the results for Part B.

See the general comments on pages 22-24.

2. Discuss stabilization of pH measurements.

pH readings require time to stabilize and students should learn not to take a reading until it is

stable. Some modern meters have an “autoread” feature that will automatically lock onto a

reading when the values have stabilized. We find that meters with the autoread feature provide

more consistent readings than meters without this feature.

3. Discuss the pH of water.

This activity is meant to show students how pH meters respond to a “difficult” sample whose pH

is inherently difficult to measure, in this case, pure water. The pH reading for pure water may

never stabilize. Tap water contains various impurities that should provide a more stable pH

reading.

Note that pure water reacts with carbon dioxide in the air to form carbonic acid, causing its

pH to be less than 7. The pH is actually changing over time.

4. Discuss the effects of temperature on pH measurements.

Students generally obtain good results with this part; that is, their graph of temperature versus pH

for Tris forms a straight line with most of the points on the line. The slope of their lines should

be similar to the published value, which is -0.28 pH units/°C.

5. Discuss the effect of dilution on the pH of buffers.

Dilution has some effect on the pH of both Tris and phosphate buffers. This is important to know

because it is common to maintain concentrated stock solutions that are diluted before use.

Usually the change in pH with dilution is small enough to be inconsequential, but there might be

situations where pH must be tightly controlled.

We find that students usually have a lot of variability in their results for this section; we are

not sure why. Therefore our laboratory manager’s data are shown here.

:



33

1M Tris pH 9.0 made with Tris base and brought to pH 9.0 with 3M HCl

0.1M pH 8.91

0.01 M pH 8.91

0.001 M pH 8.98

0.0001 M pH drifted downwards and was not stable

1M Tris made by combining Tris base and Tris-HCl

Initial pH 9.14

0.1M pH 8.99

0.01 M pH 9.02

0.001 M pH 8.95

0.0001 M pH drifted downwards and was not stable

1M Phosphate Buffer

Initial pH 6.75

0.1M pH 6.93

0.01 M pH 7.15

0.001 M pH 7.26

0.0001 M pH 7.35 drifted downwards and was not stable

6. How will knowing about the factors that affect pH measurements affect your methods of

measuring pH in the future? See the general comments on pages 22-24.

7. Discuss Part D.

Provide the students with a two or three bottles of the same solution from which they can take a

sample. The sample should be properly labeled. It is possible to use a “difficult” sample, such as

concentrated sodium chloride, pure water, or Tris buffer that is recently removed from the

refrigerator. However, it may be challenging enough to get a consistent reading from a room

temperature buffer. This is especially true if your lab is equipped with multiple pH meters and

electrodes. The most important lesson in this activity is that students do not take a single quick

pH reading and assume that it is correct. Measuring pH requires careful technique. It is difficult

to obtain consistent pH readings across instruments and operators. Students might want to strive

to have all their class readings + 0.1 pH unit, a value that is sometimes used in industry to set the

limits for pH.



34

8. Compare and contrast pH measurements and weight measurements:

It is more difficult to get accurate and consistent results from pH meters than from balances or

volume measuring devices. Temperature effects may be significant and many samples are

inherently difficult to pH. It is important that students be aware of the factors that affect

measured pH values.

9. Discuss the class results shown in Figure 3.30.

These are fairly typical results for a beginning class. One point that can be made about these data

is that when presenting data to others it is important to clearly label and explain graphs and

tables. It is possible also to discuss the variability between groups. For example, the pH readings

for the 1M Tris range from 7.85-8.23 when students were studying the effect of dilution. The

class was all working from the same stock bottle. Students can consider whether this is

acceptable variability – we would not consider this to be acceptable. We have observed that

when students are told to focus on getting consistent and accurate readings for a sample (as they

are told to do in Part D of this laboratory exercise) the variability between groups and individuals

is low. But, when students are directed to pay attention to something else (in this case, the effect

of dilution) the variability between individuals and groups is higher. Students thus should be

directed to always pay attention to such factors as temperature, stabilization time, proper

calibration, etc. in order to obtain meaningful readings.

10. Complying with the CFR. See the general comments for metrology on page 24.



35

.

UNIT IV

SPECTROPHOTOMETRY AND THE

MEASUREMENT OF LIGHT

Laboratory Exercise 10: Color and the Absorbance of Light

Laboratory Exercise 11: Concentration, Absorbance, and Transmittance

Laboratory Exercise 12: Preparing a Standard Curve with Food Coloring and Using

it for Quantitation

Classroom Activity 10: Beer’s Law and Calculating an Absorptivity Constant

Laboratory Exercise 13: Determination of the Absorptivity Constant for ONP

Unit IV is placed after metrology because it continues to explore principles of measurements.

However, students begin working with solutions in this unit and solutions are not formally

introduced until Unit V. In Unit IV students are not expected to be able to perform most solution-

related calculations but they should understand the general concept of concentration. You may

want to complete Unit V before Unit IV. We have found that there are advantages and

disadvantages to either order of units.



36

LABORATORY EXERCISE 10: COLOR AND THE ABSORBANCE OF

LIGHT

Comments. This is a relatively simple exercise that explores the relationship between color and

the absorbance of light of different wavelengths. This exercise also introduces students to

spectrophotometer operation. The exercise will need to be modified based on the models of

spectrophotometers available.

Part A works in a “Spec 20” style instrument because the light remains on when the cuvette

chamber door is open, as long as a cuvette is present. It is a nice illustration of the relationship

between color and wavelength, if you have the equipment to do it. It will not work in many

models of spectrophotometer.

We often use red food coloring as a readily available, cheap, and safe sample. Other instructors

use samples like Kool Aid that are also safe and inexpensive.

Red food coloring from McCormick is a combination of FD&C dyes 3 and 40. We do not

know the exact proportions of the dyes and have not been able to find out. The MW for FD&C 3

is 879.92. FD&C 40 has a MW of 496.42. Food colorings are regulated by the FDA and

guidelines are in the Code of Federal Regulations. An excerpt from the regulations for food

coloring is:

PART 74--LISTING OF COLOR ADDITIVES SUBJECT TO CERTIFICATION--Table of

Contents Subpart A--Foods

Sec. 74.340 FD&C Red No. 40.

(a) Identity. (1) The color additive FD&C Red No. 40 is principally the disodium salt of 6-

hydroxy-5-[(2-methoxy-5-methyl-4- sulfophenyl)azo]-2-naphthalenesulfonic acid.

(2) Color additive mixtures for food use (including dietary supplements) made with FD&C Red

No. 40 may contain only those diluents that are suitable and that are listed in part 73 of this

chapter as safe for use in color additive mixtures for coloring foods...

(b) Specifications. FD&C Red No. 40 shall conform to the following specifications and shall

be free from impurities other than those named to the extent that such other impurities may be

avoided by good manufacturing practice...



37

Here are typical results for green and yellow food coloring absorbance spectra:

Spectra of Yellow and Green Food Coloring. Green food coloring appears to be a combination of yellow and blue dyes, based on the presence of two peaks. The unlabeled peaks to the left are in the UV range of the spectrophotometer. The furthest left peak, at about 200 nm, is present with every sample and we think is due to the cuvette and/or to water. The peaks at 250 nm are likely due to the food coloring.


Key points to discuss with this exercise include:

• The peaks of the food coloring dyes are those predicted by Table 4.2 in the lab manual.

• Dyes that are a combination of more than one compound can have more than one peak.

• The absorbance and transmittance spectra are inverse; when absorbance peaks, transmittance

is at its minimum.



38

LABORATORY EXERCISE 11: CONCENTRATION, ABSORBANCE,

AND TRANSMITTANCE

Comments. The laboratory part of this exercise is simple for the students to perform because

the instructor provides all the red food coloring solutions. The challenge for the students is to

explore the relationship between absorbance and transmittance and the equation for the line

when concentration is plotted versus absorbance. Their analysis begins in this laboratory exercise

and continues in Classroom Activity 10.

For 15 years, we have used McCormick red food coloring and have assumed that it is present

at a concentration of 25,000 ppm when purchased (based on anecdotal information that may or

may not have been correct). Using this assumption, solutions with concentrations from 1-30 ppm

have absorbances in the linear range of most spectrophotometers. If you would like to be certain

of the initial concentration of dye, we suggest using allura red (FD&C red #40), one of the

components of red food coloring. You can begin with allura red powder and dissolve it to a

known stock concentration. Our experience suggests that concentrations of allura red in the range

of 1-20 ppm should provide a linear graph.

When students take the negative log of the transmittance, be sure they do so with

transmittance, not with percent transmittance.


1. Based on your graph of absorbance versus concentration, what do you predict would be the

absorbance of a solution with 23 ppm red food coloring at 517 nm? The answer to this question

depends on the spectrophotometer used. With a high resolution spectrophotometer the graph is

likely to be linear at 23 ppm and it is easy to read off the expected absorbance. However, some

spectrophotometers may not provide a linear response at this concentration and students should

be cautioned not to make predictions above the linear range.

2. and 3. Students can read these values from their graphs.

4. Based on your graph, would 20 ppm red food coloring have twice the absorbance as 10 ppm

red food coloring? If the graph is linear at concentrations as high as 20 ppm, then the absorbance

at 20 ppm should be twice that at 10 ppm because absorbance has a linear relationship with

concentration.

5. Based on your graph, would 10 ppm red food coloring have half the transmittance as 20 ppm

red food coloring? No, because transmittance does not have a linear relationship with

concentration.



39

LABORATORY EXERCISE 12: PREPARING A STANDARD CURVE

WITH FOOD COLORING AND USING IT FOR QUANTITATION

Comments. This exercise involves making a standard curve using food coloring. As part of this

activity, students check what happens to the graph if standards are used that are too concentrated

to be in the linear range. Depending on your spectrophotometer, concentrations above about 20

ppm may begin to fall in the nonlinear range. A high-end spectrophotometer may allow you to go

as high as 50 ppm before the graph becomes visibly nonlinear. An unknown sample that is too

concentrated should be provided. Students should dilute the sample and compensate for the

dilution in their calculations. There is extensive information and many practice problems relating

to dilutions in the textbook Basic Laboratory Calculations for Biotechnology, as referenced in

the Introduction to this document.

Notes Regarding Dilutions and Calculations:

• Calculations are shown in the table below and are based on the assumption that the original

stock of food coloring from the store is 25,000 ppm.

• Prepare two or three unknowns for each group. Values between 2 and 20 ppm should be

usable without dilution. Another unknown that is above the linear range, for example, 75

ppm, is recommended.

• If the students pipette accurately, their standards should form a “perfect” line with all the

points right on the line, or very close to it.



40

Table 4.3 (page 145 in the laboratory manual) completed

PREPARING STANDARDS FROM A 25,000 PPM STOCK

Dilution needed Solution to use mL of this solution

needed

mL water

needed

Prepare 10 ml of 1000 ppm stock

from the original stock

25,000 ppm stock 0.4 9.6

Prepare 40 mL of 100 ppm stock.

1000 ppm 4.0 36

Prepare 40 mL of 10 ppm stock.

100 ppm 4 36

Prepare 10 mL of a 200 ppm

standard

1000 ppm 2 8


standard

100 ppm 10 0


standard

100 ppm 7.5 2.5


standard

100 ppm 5.0 5.0


standard

10 ppm 10 0


standard

10 ppm 8.0 2.0


standard

10 ppm 6.0 4.0


standard

10 ppm 4.0 6.0


standard

10 ppm 2.0 8.0


1. What happens if the concentration of a standard is so high that its absorbance cannot be read

accurately? What should you do if this is the case? Explain the term “linear range” in your own

words. Samples that are not in the linear range of the test need to be diluted and measured again.

The dilution then needs to be considered in calculating the final answer. (Later on, when students

are performing assays [such as the protein assay in Laboratory Exercises 24 and 25] if the result

for a sample is out of the linear range of the assay, the entire assay should be run again,

beginning with the preparation of the standards.)

There are two reasons why samples may exceed the linear range of the assay. When working



41

with these simple food coloring samples, the linear range is determined by the range of

absorbance the spectrophotometer can accurately read. More expensive instruments can read

absorbances above 2 and even close to 3, while less expensive instruments plateau below 2.

However, assays that involve chemical reactions are also limited by the relative amounts of

reactants and enzymes and so may have a narrower linear range than the spectrophotometer. This

is the case with the protein and enzyme assays that the students perform later in Laboratory

Exercises 24-27.

3. Discuss your results for your unknowns. Students often get so caught up in preparing

standards and checking if their points lie on a straight line that they forget that the purpose of a

standard curve is to find the concentration of analyte in samples. Students need to check to see if

the results they obtain for their unknowns are correct according to the teacher’s answer key.

CLASSROOM ACTIVITY 10: BEER’S LAW AND CALCULATING AN

ABSORPTIVITY CONSTANT

This exercise introduces Beer’s Law through analyzing data the students collect. Our student

intern’s data for McCormick brand red food coloring is shown below. These data are based on

the assumption that the initial concentration of the dye is 25,000 pppm (which we have been

unable to confirm). For these data:

Question 13.1. (page 150 in the laboratory manual) First, determine the slope of your line on

your graph of concentration versus absorbance.

The equation for the slope of a line is:

Y2 – Y1

X2 – X1

Be sure to include the units.

Absorbance has no units, so your units for the slope will be 1/ppm.

The slope of the line for this intern’s data ≈ 0.0619/ppm.

14. Next, determine the absorptivity constant for red food coloring at the wavelength you

used based on the slope of your standard curve. Assume the path length through your

cuvette was 1 cm. (This is standard for normal cuvettes.)

From Beer’s Law: m = (αb)

Rearranging this equation:



42

α = m/b

Now you can substitute into this equation your slope and 1 cm for the path length.

Solve for the absorptivity constant.

Absorptivity constant for red food coloring = slope = α = 0.0619

path length ppm-cm

SSttuuddeenntt IInntteerrnn SSttaannddaarrdd CCuurrvvee ffoorr MMccCCoorrmmiicckk BBrraanndd RReedd FFoooodd CCoolloorriinngg BBaasseedd oonn tthhee

AAssssuummppttiioonn tthhaatt tthhee IInniittiiaall CCoonncceennttrraattiioonn iiss 2255,,000000 ppppmm..



43

LABORATORY EXERCISE 13: DETERMINATION OF THE

ABSORPTIVITY CONSTANT FOR ONP


CALCULATIONS FOR PREPARING STANDARDS FROM A 10 mM STOCK

Concentratio

n needed (C2)

Volume

to

make

(V2)

Concentration

of Stock

solution to use

(C1)

mL of stock solution

needed (V1)

mL Z buffer

needed

2.0 mM

5 mL 10 mM stock Example calculation:

C1V1 = C2V2

10 mM (?) = 2.0 mM (5

mL)

? = 1.0 mL of stock

Combine 4.0 mL of

Z buffer with 1 mL of

the 10 mM stock of

ONP

1.0 mM 5 mL 10 mM C1V1 = C2V2

10 mM (?) = 1.0 mM (5

mL)

0.5 mL = 500 µL

4.5 mL = 4500 µL

0.80 mM 5 mL 10 mM 0.4 mL = 400 µL 4.6 mL = 4600 µL

0.50 mM 5 mL 10 mM 0.25 mL = 250 µL 4.75 mL = 4750 µL

0.20 mM 5 mL 10 mM 0.1 mL = 100 µL 4.9 mL = 4900 µL

0.10 mM 5 mL 10 mM 0.05 mL = 50 µL 4.95 mL = 4950 µL

0.05 mM 5 mL 10 mM 0.025 mL = 25 µL 4.975 mL = 4975 µL

0.01 mM 5 mL 10 mM 0.005 mL = 5 µL 4.995 mL = 4995 µL


2. What can you conclude about the effect of pH on the intrinsic ability of ONP to absorb light at

420 nm? The pH effect is dramatic and should be easy to see, even if students do not get an

absorptivity constant similar to the published value.

3. The published absorptivity constant for ONP at pH 9.0 is 4500 L/mole-1cm at 420 nm. How

do your values compare to the published value? How does the average class value compare?

We have had varying results when performing this exercise with groups of students. The student

intern data shown in the Excel graph below provides an absorptivity constant of

≈ 4280 L/mol-cm, which corresponds reasonable well to the literature value. (Absorptivity

constants will vary if spectrophotometers are not correctly calibrated.) Similarly, the teachers’

data shown below gives an absorptivity constant of 4800 L/mol-cm at pHs above 9. The teachers’

graph was based on data from a group of teachers in a teacher enhancement workshop. In



44

contrast, the student data in the table below gives absorptivity constants that are substantially

lower. It may be important to use freshly purchased ONP. Be sure to prepare all your standards

and stock solutions in Z buffer, otherwise, we find the absorptivity constant is lowered. Also, be

certain students are performing their calculations correctly.

In any event, the effect of pH is significant and students consistently observe it. With ONP,

the pH effect is obvious because the sample is more brightly colored at higher pHs. Students

should be aware that there may be a pH effect on absorbance that is invisible when working in

the UV range, yet could have a significant impact on results.

y = 4280.2x + 0.0115

0

0.5

1

1.5

2

2.5

0 0.0001 0.0002 0.0003 0.0004 0.0005

Concentration ONP (moles/L)

Ab

so

rban

ce

Student Intern Data for Absorptivity Constant for ONP. These data were obtained after the sodium carbonate was added. The pH of the standards was 9.5, as estimated with pH paper. The equation for the line is shown at the top. A path length of 1 cm was used.

pH effect on o-nitrophenol absorbance at 420 nm

y = 0.0048x - 0.0049

R2 = 0.9999

y = 0.0038x + 0.0007

R2 = 0.9999

y = 0.0013x + 0.0022

R2 = 0.9991

y = 0.0004x + 0.011

R2 = 0.9532

-0.2

0

0.2

0.4

0.6

0.8

1

1.2

0 50 100 150 200 250

micromolar concentration

ab

so

rba

nc

e

10.2

9.5

8.5

7.5

6.5

5

Linear (10.2)

Linear (7.5)

Linear (6.5)

Linear (5)

Teachers’ Data for ONP Absorptivity Constant.



45

Laboratory Exercise 13

Sample of Class Data

pH

Absorptivity Constant

(L/mol-cm)

7.17 496

10.62 1058

7.0 561

10.5 1262

6.91 457

6.6.1 453

9.9 922

6.9 507

10.0 1128



46

UNIT V

BIOLOGICAL SOLUTIONS

Classroom Activity 11: Getting Ready to Prepare Solutions with One Solute:

Calculations

Classroom Activity 12: Getting Ready to Prepare Solutions with One Solute:

Ordering Chemicals

Laboratory Exercise 14: Preparing Solutions with One Solute

Laboratory Exercise 15: Preparing Solutions to the Correct Concentration

Laboratory Exercise 16: Working with Buffers

Laboratory Exercise 17: Preparing Breaking Buffer

Laboratory Exercise 18: Preparing TE Buffer

Laboratory Exercise 19: More Practice Making a Buffer

Laboratory Exercise 20: Making a Quality Product in a Simulated Company

GENERAL DISCUSSION POINTS FOR THIS UNIT

1. How can you determine and document the quality of your solutions?

It is very helpful when teaching solution preparation to have a method to determine if solutions

have been made correctly. We routinely use conductivity as a simple check for student solutions.

Solutions that have the same concentrations of solutes are expected to have the same

conductivity, if the solutions’ conductivity is measured under the same conditions (e.g., using the

same meter at same temperature). If a student makes a 10X error in a calculation, the

conductivity of their solution will likely be about 10X different than everyone else’s. It is not

unusual for students to observe that their own data varies substantially from the values of the rest

of the class. This is a good jumping off point for a discussion of differences in technique,

possible errors, and a discussion of how much variability is acceptable. Inexpensive conductivity

meters will work sufficiently well and we have found them for less than $100. However, these



47

inexpensive meters sometimes cannot be calibrated to external standards. If the meters are only

used to compare solutions with one another, the lack of calibration is not a serious problem.

Standard deviation is commonly used to evaluate variability. The acceptable range for

standard deviation is usually determined through experimentation. Students can learn to calculate

standard deviation as a measure of variability although it is a more complex problem to

determine how much variability is acceptable.

2. Why must solutions be made properly and consistently? Biotechnologists rely on

laboratory solutions for almost every task they perform, from conducting molecular biology

experiments to manufacturing biopharmaceuticals. It is essential that these solutions be made

correctly to ensure the quality of the final products. For example, cells will not survive and

reproduce if the concentrations of nutrients and salts are incorrect. Molecular biology

experiments will not work if the wrong conditions are provided for enzymes to function. Many

enzymes and reactions are sensitive to the level of salt ions in the solution. PCR, restriction

digests, transformation, transfection, labeling reactions, and a host of other molecular biology

techniques are sensitive to salt ions and will not work or will work sub-optimally if salt ions are

not controlled. Moreover, inconsistency in solution-making can lead to inconsistent results. It is

extremely frustrating to have an experiment work one day and not the next, or to try to reproduce

someone else’s work when they have made an error in reporting the materials and methods they

used.

3. What should the student do if he or she accidentally adds too much water

when bringing a solution to volume? Begin again; otherwise, the concentration of solutes

will be incorrect. It may be possible to “rescue” the solution by recalculating the amounts of

solutes required and adding those amounts in a situation where the solutes are valuable.

The conductivity will go down if the volume of water is too high.

4. What should the student do if he or she accidentally overshoots the pH of a

solution when adjusting its pH? Begin again. Even though it is possible to return the pH to

the proper level by adding more acid or base, this will change the concentration of ions in the

solution and will introduce inconsistency into the solution preparation.

5. What is required to properly prepare solutions?

Some thoughts to consider:

• Understand and be able to reliably perform the calculations required to determine how much

of each solute is required

• Be able to select the proper balance and use it correctly

• Be able to measure volumes with accuracy and precision



48

• Be able to measure pH with accuracy and precision

• Be able to check the solution’s quality. We recommend the measurement of conductivity as a

quick and simple quality control method.

6. How close is “close enough?”

This is a challenging question because the answer is: “it depends…” In a work setting, the

features of a solution are determined based on experience with the system.

7. The U.S. Code of Federal Regulations (CFR) outlines the following requirement

for companies that perform laboratory testing of potential drug compounds: All

reagents and solutions in the laboratory areas shall be labeled to indicate identity, titer or

concentration, storage requirements, and expiration date. Deteriorated or outdated reagents

and solutions shall not be used. (21 CFR 58.83, rev. 2001)

How might you meet these requirements in a workplace?

Strategies include but are not limited to:

• SOPs that specify how solutions should be labeled and stored

• SOPs that are based on experimental data that specify how to set expiration dates

• SOPs that specify how to check reagents for expiration date and tell analysts what to do with

expired materials

A general comment: It is surprising how much time and practice is required to learn to make

biological solutions consistently and accurately. We find in our program that the nine activities

and exercises in this unit are a good start but are not sufficient to ensure mastery. We continue to

review solution preparation in later courses. This leads us to wonder how many research studies

cannot be reproduced because of problems in the way the solutions are prepared and their

components reported. We’ll never know -- but we do wonder….



49

CLASSROOM ACTIVITY 11: GETTING READY TO PREPARE

SOLUTIONS WITH ONE SOLUTE: CALCULATIONS


CALCULATIONS OF AMOUNT OF SOLUTE REQUIRED

NAME OF

SOLUTION

RECIPE AS

GIVEN IN

TECHNICAL

MANUAL

VOLUME NEEDED AMOUNT SOLUTE

NEEDED (SHOW

CALCULATIONS)

AMOUNT

NEEDED FOR

WHOLE

CLASS

Saline

solution

0.91% NaCl

(w/v)

50 mL 0.4550 g

0.4550 g X

number of

groups or

students

Calcium

chloride

1 M 50 mL Dihydrate FW = 147.01

1 mole/L X 147.01 g/1

mole X 0.050 L = 7.3505 g

Monohydrate FW = 111.0

g

1 mole/L X 111.0 g/mole X

0.050 L = 5.5505 g

7.3505 g X

number of

groups or

students or

5.5505 g X

number of

groups or

students or

Magnesium

chloride

15 mM MgCl2 20 mL Anhydrous FW = 95.2

0.015 mole/L X 95.2

g/mole X 0.020 L =

0.0286 g

0.0286 g X

number of

groups or

students

BSA

standard

1 mg/mL 15 mL

(Do the calculation

assuming you will make

15 mL. However, if you

weigh out a bit too

much or too little BSA,

you can adjust the

volume.)

For 15 mL need

15 mg = 0.0150 g

15 mg X

number of

groups or

students



50

CLASSROOM ACTIVITY 12: GETTING READY TO PREPARE

SOLUTIONS WITH ONE SOLUTE: ORDERING CHEMICALS

Table 5.5 (page 170 in the laboratory manual) completed CHEMICAL CATALOG INFORMATION

Column 1

Column 2

Column 3

Column 4

Column 5

Column 6

Column 7

Column 8

Column 9

Com-

pound

Type Liquid or

solid?

Anhydrous

or hydrated?

Amount

needed for

whole class

Manu-

fac-

turer’s

catalog

number

Cost Grade or

descrip-

tion

Considerations

Sodium

chloride

Salt

Solid crystals No Assuming 20

students or

pairs

About 10 g

S3014-500 g

Crystalline This grade is

intended for

molecular biology

and so is an

obvious choice. It

has been tested

for enzymes that

would degrade

DNA, RNA, or

proteins.

Calcium

chloride

Salt

Solid Yes

Calcium

chloride

anhydrous,

calcium

chloride

dihydrate,

calcium

chloride

hexahydrate

Assuming 20

students or

pairs

147 g (dihydrate)

C1

016-100

g

223506-500 g

$33.00 An-

hydrous

Dihydrate

Magne-

sium as

magne-

sium

chloride

Salt

Solid Yes, can be purchased as anhydrous or hexahydrate

Assuming 20

students or

pairs

Less than 1 g

(anhydrous)

M2670-100G

$26.50

$60.30

Sigma Ultra

The hexahydrate

is commonly

used. Any of the

grades of

magnesium

chloride

hexahydrate will

probably be

adequate.



51

BSA Protein (not enzyme)

Solid, but can be purchased in solution

NA Assuming 20 students or pairs, less than 1 g

Sigma A2153-10 g or A7906-10 g A4503-10 g

29.90 Lyophilized powder

BSA is isolated from a natural source, bovine serum. It, therefore, must be purified before use. All three of the grades shown on this table appear to be suitable for use as a protein standard.

LABORATORY EXERCISE 14: PREPARING SOLUTIONS WITH ONE

SOLUTE


1. Suppose you are bringing your solution to volume and you accidentally add too much water.

a. What should you do?

b. Would you expect the conductivity of the solution to go up or to go down? a. Begin again;

otherwise, the concentration of solutes will be incorrect. It may be possible to “rescue” the

solution by recalculating the amounts of solutes required and adding those amounts in a situation

where the solutes are valuable. b. The conductivity will go down if the volume of water is too

high.

2. How should you decide which type of balance to use to weigh out the solutes for each of

these solutions? Encourage students to think about the best balance for a given situation based on

its capacity (and possibly other features, like automatic printouts). For example, a top loading

balance with a capacity of 1 kg will weigh a sample of a few grams and is convenient to use, but

an analytical balance will usually be a better choice for such a small sample.

4. Work flow is the organization and order of one’s work in the laboratory or other workplace. How

could you organize your work to be most efficient and accurate when making a series of

solutions? Skilled laboratory professionals are organized and efficient. Students can be

encouraged to think of ways to be more efficient, for example, by labeling all their containers

before beginning their work, assembling the materials required, scheduling the use of shared

equipment, etc. All of these strategies rely on preplanning.

5. Discuss class results. See the general comments on pages 46-48.

6. The U.S. Code of Federal Regulations (CFR). See the general comments on page 48.



52

LABORATORY EXERCISE 15: PREPARING SOLUTIONS TO THE

CORRECT CONCENTRATION

Comments. The concept of this exercise is simple but it is effective in uncovering student math

and solution-making misconceptions. If students make their solutions correctly, then all groups

will have similar conductivities for each solution and each groups’ data will be linear when

plotted with concentration on the X axis and conductivity on the Y axis. Note, however, that this

exercise requires a conductivity probe (or more than one probe) with a wide range of response.

There are a couple of logistic issues to resolve with this exercise. First, it is glassware

intensive. However, almost any combination of beakers and flasks can be used. Second, it

requires that each group has access to an analytical balance for sufficient time to weigh out 17

samples (If the exercise is performed as written). Another bottleneck is the use of a conductivity

meter to check all the solutions. Using fewer samples may be necessary to alleviate resource

problems.


SOLUTIONS TO PREPARE

Solution Concentration Final Volume

Desired

Solute Desired

(g)*

A 1.000 M 100 mL 5.844

B 1.0 g/L 100 mL 0.1

C 2.0 g/100 mL 100 mL 2.0

D 3.0 g/100 mL 100 mL 3.0

E 5.0 g/100 mL 100 mL 5.0

F 0.10 g/mL 100 mL 10.0

G 0.070 g/mL 100 mL 7.0

H 0.010 g/mL 100 mL 1.0

I 0.0010 g/mL 100 mL 0.10

J 0.00010 g/mL 100 mL 0.010

K 0.000010 g/mL 100 mL 0.0010

L 1.0 mg/mL 100 mL 0.10

M 1.0 % 100 mL 1.0

N 0.10% 100 mL 0.10

O 1.000 M 200 mL 11.688

P 1.0 g/100 mL 200 mL 2.0

Q 1.0 % 200 mL 2.0

*Students may not be able to weigh out exactly the desired amount of solute and should get in the habit of

recording what is actually weighed out as well as the desired amount.



53

Laboratory Exercise 15, Sample of Student Data

Solution Concentration Volume Conductivity

(mSi/cm)

A 1M 100 mL 64.3

B 1 g/L 100 mL 1.612

C† 1 g/mL 100 mL sample was lost

D† 0.7 g/mL 100 mL sample was lost

E† 0.5 g/mL 100 mL sample was lost

F 0.1 g/mL 100 mL 98.7

G 0.07 g/mL 100 mL 74.8

H 0.01 g/mL 100 mL 13.94

I 0.001 g/mL 100 mL 1.540

J 0.0001 g/mL 100 mL 0.1456

K 0.00001 g/mL 100 mL 45.4

L 1 mg/mL 100 mL 0.0194

M 1 % 100 mL 13.49

N 0.1% 100 mL 1.683

O 1 M 200 mL 63.2

P 1 g/100 mL 200 mL 13.69

Q 1 % 200 mL 13.44

These values are from a year where we used slightly different concentrations for some of the samples

Laboratory Exercise 15

Sample of Student Data Organized by Concentration

Solution Concentration Conductivity

(mSi/cm)

A 1M 64.3

O 1 M 63.2

K 0.001 g/ 100 mL 0.0454

J 0.01 g/100 mL 0.1456

B 0.1 g/100 mL 1.612

I 0.1 g/100 mL 1.540

L 0.1 g/100 mL 0.0194

N 0.1 g/100 mL 1.683

H 1 g/100 mL 13.94

M 1 g/100 mL 13.49

P 1 g/100 mL 13.69

Q 1 g/100 mL 13.44

G 7 g/ 100 Ml 74.8

F 10 g/100 mL 98.7



54

Laboratory Exercise 15, Student Data; Slopes for Comparison

Group Slope of all data,

(µSi/cm/mg/mL)

Slope of lower

concentration data,

(µSi/cm/mg/mL)

Slope of higher

concentration data,

(µSi/cm/mg/mL)

L. 1,114 1,564 946

T. & M. 1,086 1,566 858

M. & D. 1,210 1,500 1,137

J. & R. 1,087 1,514 917


1, 2, and 3. Discuss Data. See the general comments on pages 46-48.

4 and 5. Quality Control. It is important to have a means to check whether a system is

functioning properly. In a production setting there is likely to be a quality control department in

charge of this task. In a research setting, the individual researcher needs to put systems in place

to catch errors and control the experimental variables.

LABORATORY EXERCISE 16: WORKING WITH BUFFERS

A. PREPARE 3 M HCl FROM 6 M STOCK

A.1. Calculate how much 6 M HCl is needed to make 100 mL of 3 M HCl:

50 mL

B. PREPARE 2 N NaOH FROM PELLETS

B.1. Calculate how much NaOH is required to make 100 mL of 2 N NaOH:

8.0 g

C. PREPARE 100 ML OF A STOCK SOLUTION OF 1 M TRIS, pH 8.0

C.1. Calculate how much Tris base is required to make 100 mL of 1 M solution:

12.114 g


1, 2, 3, and 4. Discuss Data. See the general comments on pages 46-48.



55

LABORATORY EXERCISE 17: PREPARING BREAKING

BUFFER

Comments. 5g of glycerol is not the same as 5 mL, so there is ambiguity in the recipe as

written. Should it be 5% w/v or 5% v/v? We do not know the answer to this and either way

works for extracting proteins. An SOP would clarify this issue. We do not discuss this in the lab

manual so as to avoid confusion; however, you might want to point it out to your students.

Glycerol is hard to pipette. A 10 mL graduated cylinder might work better than a pipette.

BREAKING BUFFER∗∗∗∗

0.2 M Tris, pH 7.6

0.2 M NaCl

0.01 M MgAc

5% glycerol

SOLUTE CALCULATIONS FOR MAKING 100 mL

0.2 M Tris

Proportion Strategy:

1 M = 121.1 g = ? ? = 12.11 g

1000 mL 100 mL

12.11 g = ? ? = 2.422 g

1 M 0.2 M

0.2 M NaCl

Formula Strategy:

58.44 g/mole X 0.2 mole/L X 0.1 L = 1.1688 g

0.01 M MgAc

Formula Strategy:

214.4 g/mole X 0.01 mole/L X 0.1 L = 0.2144 g

5% glycerol

5% = 5 mL/100 mL so need 5 mL, assuming the recipe is v/v



∗ This buffer normally contains 5mM DTT, dithiothreitol, to help stabilize protein structure. DTT is

somewhat expensive and toxic and therefore is not included in this recipe.



56

LABORATORY EXERCISE 18: PREPARING TE BUFFER

Comments. This solution is a staple in molecular biology but it is a difficult solution to prepare

because EDTA is hard to get into solution. Allow ample time for EDTA to dissolve (at least half

an hour but more likely an hour).

The recipe given for TE buffer, pH 8.0 is found in Molecular Cloning: A Laboratory

Manual”, Second edition, by Sambrook, Fritsch, and Maniatis, Vol. III, Appendix B20, Cold

Spring Harbor Press, NY, 1989. Students need to be able to interpret this brief recipe.

Step A.1. Prepare 100 mL of 0.5 M EDTA stock, pH 8.0. (Ethylenediaminetetraacetic Acid)

This requires 18.62 g of EDTA (disodium salt FW 372.24). It may help to use a little heat to help

the EDTA dissolve, however, the final pH should be read at room temperature. It takes about

2.2 g of sodium hydroxide pellets to bring the solution to pH 8. EDTA will not dissolve until the

pH is about 8. Stir continuously while dissolving.

Step A.2. Prepare 100 mL of a stock solution of 1 M Tris Buffer that is pH 8.0. This requires

12.11 g of Tris base (FW 121.1). Dissolve the Tris base in about 70 mL of purified water. Adjust

the pH by slowly adding HCl (3 or 6 M HCl works, but have students be cautious with this acid)

while monitoring the pH. Stir continuously with a stir plate, if possible. When the solution is at

pH 8.0, add water to a total volume of 100 mL. Work at room temperature.

To make 50 mL of TE buffer:

For Tris:

C1V1 = C2V2

1M (?) = 0.010 M(50 mL) ? = 0.5 mL = 500 µL

For EDTA:

C1V1 = C2V2

0.5 M(?) = 0.001 M (50 mL) ? = 0.1 mL = 100 µL

Combine 0.5 mL of the 1 M Tris stock solution and 0.1 mL of the 0.5 M EDTA stock, mix, and

bring to a final volume of 50 mL. Check the pH but do not readjust it.



3. TE is used routinely in procedures that involve dissolving and manipulating DNA. What might

be some consequences of using improperly prepared solutions of TE during experiments that

involve DNA? Students might speculate that the DNA will be degraded, that enzymatic reactions

will proceed slowly, or not at all, or that interactions of DNA with other molecules will not occur

properly.



57

LABORATORY EXERCISE 19: MORE PRACTICE MAKING A BUFFER

Comments. This activity can be used to assess the students’ documentation and calculation

skills. We emphasize the importance of documentation when students perform this activity. We

store their completed solutions and collect their copies of the laboratory from their laboratory

notebooks. Later, for part of a laboratory exam, we give them back their laboratory notebook

pages for this activity and ask them to reproduce exactly the same solution. We check the

conductivity of the new solution against that of the original one and assign points based on their

documentation and the similarity of their conductivities. The reagents that we typically provide

for this activity include: Tris buffer, sodium chloride, magnesium chloride, sucrose, and glycerol.

You may want to make some suggestions regarding concentrations ranges to avoid wasting

solutes. For example, reasonable ranges for Tris concentration might be 0.1 M to 2 M.

LABORATORY EXERCISE 20: MAKING A QUALITY PRODUCT IN A

SIMULATED COMPANY

Comments. This lab is simple in concept but complex in execution. There are a many factors to

consider and it takes a long time to develop the documentation and processes even to make a

single, simple buffer that is compliant with GMP requirements. However, this activity is very

effective at getting students to really think about what is involved in making a product. We have

used as many as three lab periods (each 3 hours) on this one activity and it has been valuable.

Make sure the students actually make the buffer, more than once, themselves. Three times is a

reasonable requirement. They should make it enough times to experience variability and to think

about why variability is there and how to reduce it. Have students set specifications, as a range,

for critical parameters including pH and conductivity. Specifications are based on the purpose of

the product and on the range of variability that can normally be expected for that item. It is

common for students to set the specifications for their product narrowly to begin with and then

discover that they cannot consistently meet their specifications.

We had a situation one year where one group had varying pHs and conductivities each time

they made the solution. By watching another group, and examining their own documentation,

they realized that one person in the group was not bringing the solution to the proper final

volume.

Some groups have done an excellent job with documentation, creating SOPs, labels, and

equipment logbooks. It takes two or three hour laboratory periods to make significant progress

on this activity; we have never “completed” it.



58

End Part I

Continued in Part II

Date post:	10-Oct-2020
Category:	Documents
Upload:	others
View:	0 times
Download:	0 times

Notes to Instructors Part I - Bio-Linkinformation on preparations, calculations, color photos,...

Documents