Introduction to Chemistry 151L
This course is the first semester of general chemistry laboratory. It is taken by students enrolled in either Fundamental Chemistry (CHM 130) or University Chemistry (CHM 151). The lab will provide you with instruction in common laboratory manipulations, data collection and interpretation, record keeping, and stoichiometric (chemical) calculations. You will become familiar with some of the chemist's basic laboratory equipment and will learn why and when this equipment is used. This lab should help you connect theory with the real world. You will be concentrating on the correct use of various kinds of laboratory equipment and procedures, but the lab will reinforce chemical principles covered in your lecture.
You will analyze chemical unknowns for several of the seven experiments done for this course. If you use the equipment properly, record observations accurately, and calculate correctly, you will obtain an "acceptable" answer (close to the true value) and receive a PASS. On the other hand, if an "acceptable" answer is not reported, you will receive a PARTIAL CREDIT or a REPEAT. If you make errors while doing the experiment, do the calculations wrong, or record the wrong unknown number; you most likely will not get a "PASS". Calculation checks are available to verify that computations are done correctly. You will need to make sure that data and the correct unknown number are recorded. But you will only know if the experiment was done correctly when the value for your unknown is graded. An "acceptable" answer for an unknown is one that correctly identifies it or comes within a range of possible answers obtained by students who carefully analyzed a sample of your unknown using the same procedures and equipment that you will use to determine a value.
For example, if you measured the density of an alcohol-water mixture and calculated the percent alcohol in the mixture then the "pass" range would be determined as follows: Two students did a total of 36 trials using 6 different pipets and 2 different balances. For a given unknown their answers had a standard deviation of ±0.5% alcohol. Statistically, there is a 99 percent probability that other students doing the work carefully would obtain answers within ±1.4% alcohol of their average result. However, you would be given a PASS if you were within ±3% alcohol of the "right" answer (the average answer obtained by the Chemistry Department student assistants). You would be given PARTIAL CREDIT (approximately one-half the number of points given for a PASS) if you were within ±6% of the "right" answer. Each experiment with numerical answers has a similar range of acceptable answers that are more than broad enough to account for expected experimental variations. For best results, relax and work carefully.
There are videos available on the use of the balance, pipet, buret, and other equipment. Make an effort to see them before you use them. Computer software is available online (address in syllabus) to prepare you for an experiment, check your calculations, and test your understanding of an experiment. If you have to repeat an experiment, get help from your teaching assistant or instructor and have them check your technique. You may repeat an experiment any number of times to obtain the PASS grade, but be sure to use a new unknown and a new report sheet each time you repeat an experiment. Staple the old report sheet to the back of the current one.
The syllabus will explain the grading scale. Save all graded unknown report sheets in case there are problems and record your results and unknown number in the space provided at the end of each experiment before the report sheet. Do this before turning in your results.
Learning Objectives
I. Technique - Laboratory Procedures and Equipment
A. The proper use of equipment for the measurement of various physical properties is very important. Considering the ease and speed of use and the desired precision and accuracy, you should be able to choose the right piece of equipment and follow proper technique to use it. You must be familiar with the following:
1. Volume measurement: volumetric pipet, graduated cylinder, beaker, Erlenmeyer flask, Mohr pipet, buret, volumetric flask, and other measuring equipment. See the section on significant figures in the lab manual for the precision of each of these. This includes interpolation and estimating between the marks on a scale; using appropriate number of significant figures in reporting measured quantities (Experiments 1, 2, 3, 4, 6, 7)
2. Mass measurement: proper weighing technique and use of the balance (Experiments 1, 2, 3, 4, 6, 7).
3. Temperature measurement: use of thermometer (Experiments 1, 3, 7).
4. Measuring the speed of a chemical reaction (Experiment 3).
5. Heating procedures: use of Bunsen burner & hot plate (Experiments 2, 3, 4, 7).
6. Qualitative analysis: using flame tests, observing and describing the formation and dissolving of precipitates and gases (Experiments 2, 3, 5).
7. Quantitative techniques: titration, pipeting, and weighing (Experiments 1, 2, 3, 4, 6, 7).
8. Proper cleaning of glassware, mixing of solutions, and use of reagents without causing cross-contamination (Experiments 1, 2, 5, 6).
9. Laboratory safety: safe use and disposal of chemical reagents and safe use of lab equipment; awareness of lab hazards and safety equipment procedures.
10. Collection of experimental data.
B. Recognition of other equipment found in your drawer and used in the lab.
II. Calculations
A. The basic calculations used in this lab are important for your future work in chemistry and other sciences. The following should be familiar at the end of this course:
1. Percent by mass (Experiments 2, 4, 6).
2. Density (Experiments 1, 2, 7).
3. Mass by difference (Experiments 1, 2, 3, 4, 6, 7).
4. Using a chemical formula to calculate molar mass (Experiments 2, 4, 6, 7).
5. Conversion from grams to moles or moles to grams using the molar mass (Experiments 2, 4, 6, 7).
6. Use of concentration units such as molarity (Experiments 2, 3, 6).
7. Use of ideal gas law: PV = nRT and calculation of molar mass (Experiment 7).
8. Calculation of the simplest formula, molar mass (Experiments 3, 4, 6, 7).
9. Simple dilution calculations using M1V1=M2V2 (Experiments 3 and 6).
10. Use of chemical equations in calculations (Experiments 1, 3, 4, 6, 7).
11. Calculation of the median and the average from a set of data taken in the lab (Experiments 1, 2, 4, 6, 7).
B. Data recording: proper format and rules.
C. Significant figures should be understood and the appropriate number of digits should be used for recording data, doing calculations, and predicting precision; in addition, mean or average, median, range, and error analysis will be introduced.
D. The use of unit cancellation or dimensional analysis in solving problems, which has been introduced in the lecture course, should be used in solving problems and checking work in this lab course (see appendix in this manual).
E. Graphing data and using linear regression and other functions to predict values.
III. Concepts
A. You will learn to understand the following concepts:
1. Atomic emission (Experiment 2).
2. Concentration (Experiments 2, 3, 5, 6, 7).
3. Density (Experiments 1, 2, 7).
4. Hydrates (Experiment 4).
5. Stoichiometry (Experiments 1, 4, 6, 7).
6. Qualitative analysis (Experiments 2, 5).
7. Acid-base reactions as used in titrimetry (Experiment 6).
8. Use of indicator (Experiment 6).
9. Ideal gas law (Experiment 7).
10. Significant figures.
11. Treatment of Data.
Safety in the Chemistry Laboratory
Laboratory safety involves the prevention of and response to laboratory emergencies. Good prevention is far better than someone getting hurt. This begins with always being aware of chemical and laboratory hazards. Hazard codes, chemical labels, and material safety data sheets are the first sources of information that help us prepare to work safely in a laboratory. This information can be used to do risk assessment on an experiment you are about to do. Certain rules need to be followed to keep you safe and you must know what to do in case of an emergency. Chemical waste management is another important aspect of a safe laboratory and a key regulatory compliance issue.
Chemical Labels and Hazard Codes
The first source of information is the label on a chemical bottle. Read the label carefully before using a chemical. A commercial chemical bottle will have extensive information on the label such as the chemical name and formula, physical properties, purity, molar mass, hazards, safety precautions, suggested protective equipment, and other information. A hazard code may also be included on the label.
The chemistry department has adopted the “Baker” hazard code classification system to inform users of potentially hazardous chemicals. This system is designed to provide information to people who handle chemicals in laboratories. Hazards are classified according to four types: health (toxic), flammability (fire), reactivity (explosive or reactive), and contact (corrosive). The intensity of the hazard is indicated by using a number from "0" (no hazard) to "4" (extreme hazard). This information is conveyed using either a four-colored label found on "J.T.Baker" chemical products or as a series of four digits. The label on chemical bottles may look like this:
The four-digit hazard code used in the lab manual would look like this:
1321
For example, the code listed above for acetone indicates a slight health hazard (1), a high flammability hazard (3), a moderate reactivity hazard (2), and a slight contact hazard (1). Hazard codes will be listed after the chemical inside parentheses: (1321)
The National Fire Protection Association (NFPA) is a hazard code system that was adopted in 1975 to communicate hazards to emergency responders. This system uses a label that you may be familiar with since it appears on entrances to stores containing hazardous chemicals and on chemical containers. The NFPA may differ from the “Baker” code since it provides information to firefighters while the “Baker” code provides hazard information in a laboratory situation. The codes are very similar except the white section in the NFPA code refers to special or specific hazards of importance to firefighters such as “ox” for oxidizing agent.
The Material Safety Data Sheet (MSDS) should be read to obtain additional safety information before the chemical is used. These sheets are available in Room 212 for all chemicals used in the chemistry department; they must not be removed from that room. The internet is a great resource for MSDS and general safety information.
How to Protect Yourself
1. Eye Protection MUST BE WORN IN THE LABORATORY AT ALL TIMES unless otherwise notified by the instructor or TA. Avoid rubbing your eyes in lab unless you wash your hands first. Use extra caution when using corrosive chemicals. Indirectly vented or nonvented goggles are the required eye protection for this lab course. Safety glasses or directly vented goggles are not acceptable. Do not modify or remove the vents on goggles.
2. Skin protection should be employed where appropriate; you may be required to wear long pants. Avoid wearing shorts. The use of a lab coat or plastic apron is recommended. Closed toed shoes must be worn at all times in the laboratory for protection against broken glass and spilled chemicals. Open-toed shoes or sandals are not appropriate footwear in lab areas. Disposable gloves are available for the handling of hazardous chemicals. After completing lab work for the day, wipe down your work area with a wet sponge to clean up any spilled chemicals. Rinse and wring out the sponge and then wash your hands.
3. Protection from fumes or fine powders: Never allow hazardous chemical fumes or dust to escape into the open room; use fume hoods when necessary or specified. Be sure to use the fume hoods correctly, following the instructions provided by your TA or instructor. Avoid putting your head inside the fume hood and close the sash or fume hood window when it is not in use.
4. Protection from internal poisoning: Never "pipet by mouth", eat, drink, or smoke in the laboratory. These activities are prohibited. Wash your hands after you have completed lab work.
5 Protection from hot surfaces: Use the appropriate types of tongs to handle hot objects. Test tube holders are too weak for carrying flasks.
6. Protection from fire and explosion: Never allow flammable vapors to escape into the open room (see No. 3). Ether is especially dangerous in this respect. Never use an open flame while flammable liquids are being used in the room. Hot plates should be used with care, as they are an ignition source. Flammable volatile liquids should be used in fume hoods and stored in solvent cabinets when possible. Long hair should be tied back to keep it away from open flames.
7. Protection from cuts: When manipulating glassware or ceramic ware, protect your hands with a cloth towel. Clean up broken glass immediately. Do not pick up broken glass with bare hands. Use a broom and dustpan to dispose of glass in the "Broken Glass Container". Do not clean up broken mercury thermometers without help from your TA since mercury requires special disposal procedures.
8. Protection from the unexpected: Always read all labels noting the chemical name, formula, concentration, and warnings (including hazard codes) carefully and double check to make sure you have the correct chemical and concentration. Follow directions in the experimental procedure exactly. Remove obstacles by keeping lockers closed, lab stools out of aisles, and backpacks and coats stored on coat rack. For unassigned lab work, you must have the approval of the instructor. Carefully follow hazardous waste disposal instructions given later.
9. Safety Violations: Any student who does not follow the above guidelines will be given one warning and will then be removed from the lab for the day for any subsequent violations. There may also be grade deductions or permanent removal from the lab for serious violations.
What to Do In Case Of Accident
1. During your first lab period, locate the position of the fire extinguishers, eyewashes, safety shower, first aid kit, phone, fire alarm pull stations, exits, hallway showers, and any other safety equipment.
2. In all cases of accident or injury, notify the TA and the instructor.
3. For any serious fire or injury: Call the POLICE DEPARTMENT (33000) from any campus phone. Security is in the best position to summon fire or ambulance service. Call the Flagstaff Fire Department (8-774-1414) or dial 8-911 if Security cannot be reached. Use the FIRE ALARM PULL STAIONS (red box by every stairwell entrance) to clear the building of personnel. THE LOCAL FIRE ALARM IN THE LAB BUILDING WILL SUMMON HELP BUT STILL ALWAYS CONTACT CAMPUS SECURITY FROM A SAFE LOCATION.
4. In case of a small fire: Immediately get help from your TA or instructor. Fire extinguishers are rated for ABC type fires in chemistry where A is combustible (paper, etc.), B is flammable liquids, C is electrical, and D is combustible metals. Use dry sand for D type fires or a special extinguisher rated for these fires. To use an extinguisher remember “PASS”: Pull the pin, Aim the hose, Squeeze the handle, and Sweep the base of the flames. If a person's clothing is on fire, they should immediately stop-drop-roll, use the safety shower if it is close, or smother the fire with a lab coat or fire blanket. Cover beaker fires with a watch glass or larger beaker to remove oxygen and put out the fire. Cool minor burns in cold water immediately.
5. In case of chemical contact: If the area of contact is small, flush it well under the nearest water tap for 15 minutes. Eyes must be flushed immediately using the eyewash at one of the sinks or the eyewash by the safety shower keeping the contaminated eye(s) open. In case of large areas of contact, start rinsing the person using the safety shower and remove contaminated clothing. After decontamination, the person will be taken to another shower facility if available in the building. Immediately inform the instructor or TA in any case.
6. In case of mercury spillage: To dispose of this hazardous material properly, notify your TA and he or she will collect the mercury using a special spill kit.
7. Chemical spill: In case of a chemical spill notify your TA and ask for help. If the substance spilled is flammable, turn off all burners, hot plates, or electrical devices. Notify the instructor, the laboratory manager, or the stockroom manager. Clean-up materials are available in the lab or stockroom.
Hazardous Waste Disposal
The Resource Conservation and Recovery Act (RCRA) mandates the proper disposal of hazardous waste. Disposal of many waste chemicals by putting them down the sink is now illegal. Regardless of regulations, the proper management of hazardous waste is of particular importance to the people of Arizona where the contamination of groundwater by hazardous waste could have grave consequences. Please carefully follow the instructions below to protect our groundwater and keep your lab safe. Hazardous waste is determined by four properties:
TOXIC: A poisonous substance, potentially harmful to human health, can cause cancer or birth defects, or can contaminate, harm or kill wildlife.
FLAMMABLE: substances, which can explode, ignite, or emit toxic gases or fumes if exposed to a source of ignition.
REACTIVE: An unstable substance which can react spontaneously if exposed to heat, shock, air, or water. Reactions may include fires or explosions. The research director or instructor for the lab must neutralize any reactive substance before it can be accepted for disposal.
CORROSIVE: A substance that could corrode storage containers or damage human tissue upon contact. (For example, acids and bases, pH <4 or >10)
Chemical waste that does not fit into the above categories may be flushed down the drain with large amounts of water. The instructor or TA must be consulted if there is uncertainty with regard to the collection of a chemical waste.
All waste bottles are labeled and color-coded with tape. The label will include an experiment number and a hazardous waste description that will help you decide which bottles to put your waste into. Find the correct waste bottle for your experiment number and for the type of chemical waste you have; make sure the description of the composition fits the waste you are adding to the bottle. Using the wrong waste bottle could create a safety hazard and will be treated as a safety violation. The following table should help. Some nonhazardous chemical waste from experiments you do may be put down the drain. Avoid using cup sinks or water troughs to dispose of chemicals, instead use large sinks in the lab.
|
Lab-Exp. |
Colors |
Description of Waste |
Comment |
|
5 |
Blue |
Aqueous Cations, Anions: .1% NH3, .5% Ag, .1% Cd, .1% Pb, .1%Ba .1% Al, 1% NO3-, 98% Water |
Keep in hood |
|
2,3,6,7 |
White |
Corrosive Liquids |
Keep in hood |
Acidic or basic used chemicals (pH <4 or >10) will be disposed of by neutralization in a fume hood. Waste bottles are also color-coded using the following scheme:
Blue - health hazard, poison
Red - flammable hazard, organic liquid
Yellow - reactivity hazard (strong oxidizers, etc.)
White - contact hazard, corrosive
Orange - low hazard materials with hazard codes of 2 or less
Handling Reagents
The liquids, solids, and solutions used in a laboratory are called reagents. You must become well acquainted with these reagents, their containers, and their proper use. The reagents are kept on a separate bench away from your work area. Some reagents, such as concentrated NH3 must be kept in the fume hood because it generates toxic fumes. The reagents are grouped according to experiment, starting with Experiment 1 and ending with Experiment 7. When you need a reagent please follow these rules:
1. Be sure to use the correct reagent, especially noting the concentration. Find the reagent, check the concentration, and then carefully read the label again to be sure you have the right one. Note the hazard code and take necessary precautions.
2. Use the reagent at the reagent bench. Do not take the reagent container to your work area.
3. Please conserve and take only what you need.
4. Do not contaminate the reagents. Always use a clean spatula for solids and clean glassware for liquids. Never put a pipet, etc. into a liquid reagent, instead pour what is needed into a clean, dry container and take it to your work area.
5. NEVER return unused reagents, liquid or solid, to the reagent bottles. Discard or share any excess.
6. Put the lid back on the reagent container snugly and put it back in the correct location.
7. Clean up any reagent you spill with a wet sponge, rinse out the sponge at the sink, and then wash your hands.
8. Use great care with corrosive chemicals (strongly acidic or basic solutions). Always wear safety goggles! Rinse your hands with tap water after using corrosive chemicals, especially if you feel a burning or slimy sensation on your skin. Wear the gloves provided in the laboratory if called for. All strong acids and bases will be disposed of in the hood as noted later in experimental procedures.
9. When washing glassware, often all that is needed is to rinse well with hot tap water 4 or 5 times. If the glassware is really dirty use detergent and hot tap water. You should always rinse your test tubes, pipets, buret, volumetric flask, and other glassware out with hot tap water several times and then a small amount of pure water (RO) before storage. You do not need to dry the inside of glassware.
10. Avoid using water troughs or cup sinks to dispose of nonhazardous chemicals, instead use large sinks available in the lab. Be sure to follow the instructions in the experiments with regard to the disposal of chemicals.
Data Recording
Recording Experimental Measurements
Normally a laboratory notebook is used to collect data in a lab. It would serve several purposes; the most important of which is to be the permanent, understandable record of data collected and observations made during an experiment. In this course a report sheet is used instead of a laboratory notebook to provide more structure in data collection. The same rules used in a lab notebook will be used in this course. You should be able to look at your report sheet a year from now and be able repeat the experiment or calculations. Calculations should also be shown on the report sheet or an attached paper. Dimensional analysis or unit cancellation should be used to do all calculations in this course. If the same calculation is repeated several times, the calculation can be shown once. If a spreadsheet or graph is used to do the calculations, a copy of it must be stapled to the report sheet.
Before you start an experiment key aspects of the laboratory procedure should be outlined on the report sheet or an attached paper. Any procedures not in the lab manual including changes to procedures listed in the manual must be noted. Key data must also be recorded in the laboratory manual in case the report sheet is lost. Certain rules need to be followed when keeping a report sheet:
1. Record all data and observations directly on the report sheet in pen so that the data cannot be erased or changed. This is by far the most important rule in recording data. Do not transcribe data from other pieces of paper, i.e., DO NOT record data on scraps of paper and then recopy the data onto the report sheet. Write down exactly what you are doing and your observations as you are doing the experiment. Errors in your procedure can be caught this way. Points may be taken off for writing data in places other than the report sheet. If you run out of room on a report sheet you may staple a piece of paper to the report sheet to record more data.
2. Clearly identify all data, graphs, axes and use correct units.
3. A ball point pen is normally used for all entries in a lab notebook but for this course you may use pencil to record calculated values if you wish on your report sheet but a pen must be used for all measured data (mainly mass and volume data) and observations. Do not white out, erase, or obliterate any entry; simply cross out mistakes with a single line (the mistake should still be readable) and give a short note to explain the nature of the mistake, e.g., "misread." Sometimes you will find later that the entry was not a mistake after all and will want to retrieve the data. So never obliterate or destroy data no matter how bad it looks!
4. Before an experiment is started, the entire experimental procedure must be read. As you read it, note the objectives and key points of the experimental procedure on your report sheet or a separate piece of paper. This will prepare you for the experiment before you come to lab.
5. Another important facet of scientific experiments involves the propagation of accuracy (or inaccuracy) of measurements through the calculations to the results. Use the correct number of significant figures, as outlined below, during the collection of data and calculations.
Recording Experimental Data Using Correct Significant Figures
It is important to take data and report answers such that both the one doing the experiment and the reader of the reported results know how precise the final answer is. The simplest, way of expressing this precision is by using the concept of significant figures where a significant figure is any digit that contributes to the accuracy of an experimentally measured number or to a number calculated from experimentally measured numbers. Please refer to the chemistry textbook for a discussion of the use of significant figures.
This laboratory course emphasizes learning the precision of the glassware and the balances. A larger number of significant figures can and should be carried when you are using a volumetric pipet or flask, pipettor, or buret than when you are using a beaker, erlenmeyer flask, or graduated cylinder. Different equipment in the laboratory is used to achieve different levels of precision. For example, the volume of a liquid sample can be measured using different types of glassware, some of which are more precise than others. This is shown in the following table.
|
|
Precision of Volume Measurement |
|
Equipment |
for Various Equipment |
|
250 mL Beaker |
±10 mL |
|
250 mL graduated cylinder |
±1 mL |
|
100 mL volumetric flask (class A) |
±0.08 mL |
|
25 mL graduated cylinder |
±0.2 mL |
|
10 mL measuring pipet (Mohr) |
±0.05 mL |
|
25 mL buret |
±0.02 mL |
|
5 mL, 10 mL volumetric pipet |
±0.01 mL |
When an approximate volume is needed, a beaker or graduated cylinder will be used, but when an accurate volume is needed, a pipet or buret will be specified for use. Recognizing when to make an accurate measurement and when to be satisfied with an approximate measurement can save much time. Frequently, the written directions will give clues to the needed precision by using the words "approximately" or "about" when the precision is not important and "exactly" or "precisely" when the precision is important. Another clue would be the number of significant figures used to write a number. For example, a volume of 10.00 mL would require the use of a 10 mL volumetric pipet or pipettor to measure the desired volume. On the other hand, a volume expressed as “about 10 mL” would require only a small beaker or graduated cylinder to measure the volume. It is also important to note that glassware used for accurate measurements is calibrated at a specific temperature, which is noted on the glassware.
When a measurement is made, the question arises: "How many digits or figures should be recorded?" The answer is straightforward: For a measured number record all digits, which are known with certainty, and the last digit, which is estimated. Many of the measurements in this course involve the estimation to the nearest one-fifth or one-tenth of a scale marking. For example, in Experiment 1 a 25mL graduated cylinder, which has scale markings every 0.5 mL, should be read to the nearest 0.1 mL, estimation to the nearest one-fifth of a division. The graduated cylinder does not need to be used to this accuracy at all times; for example, in Experiment 5 the volumes are approximate: the nearest 0.5 mL is more than adequate.
NOTE: Whenever estimation between markings is being done and the reading is "on the mark," the last digit should be included to convey the idea of accuracy to the reader. For example, with a buret, which has markings every 0.1 mL, a reading on the mark of 11.3 mL would be recorded as 11.30 mL; otherwise, the reader will not know that the buret was really read to the nearest 0.01 mL.
Reporting Answers in Addition and Subtraction
When experimental data have been recorded correctly, the uncertain or estimated digit is the last digit. The calculated sum or difference of experimental measurements must be carried out only to the place where the first digit of uncertainty enters the calculation. Example: Add 14.75, 1.475, and .001475 (all of which are experimental numbers). The digits of uncertainty are underlined.
14.75
1.475
0.001475
--------------
16.226475
Since the answer may include only the first digit of uncertainty, it should be rounded off to that digit and reported as 16.23. It helps to line the numbers up by the decimal point.
Reporting Answers in Multiplication and/or Division
1. All measurements should be recorded to the appropriate number of digits as discussed in the section on recording experimental data.
2. The position of the decimal point is ignored in counting the number of significant figures.
3. All digits except zero are always significant.
4. Zeros may or may not be significant. Any zero to the left of the first non-zero digit is never significant (0.0256 has 3 significant figures because neither zero is significant).
a. Any zero to the right of the first non-zero digit is always significant if there is a decimal point (2.5070 has 5 significant figures since both zeros are significant).
b. If there is no decimal point, zeros to the right of non-zero digits are significant unless it is stated otherwise (the number 25000 has 5 significant figures unless some other precision is stated, such as 25000 ± 100). Numbers with "trailing" zeros (zeros to the right of all other digits) should be written in standard exponential form to remove questions (2.50 x 104 has 3 significant figures; 2.5000 x 104 has 5 significant figures).
4. In multiplication and/or division, the answer should be reported to the same number of significant figures as the value in the computation with the least number of significant figures
Example: Find the answer to the following multiplication/division problem to the correct number of significant figures.
0.085 has 2 significant figures; 0.08206 has 4; 366 has 3; and 0.782 has 3. A calculator shows the answer to be 22.989865, so the answer should be reported as 23.
Interpretation of Data
Significant figures are excellent to express the precision of raw data but not always so good to express the precision of calculated values. As a general rule in this laboratory course you should always use at least four significant figures for calculated values to avoid rounding errors. Once the final answer is calculated, it can be expressed using correct significant figures. In order to interpret how good your results are, certain terms need to be understood. You will need to understand the following definitions.
1. Accuracy: The term "accuracy" describes the nearness of a measurement to its accepted or true value. In CHM 151L, the accuracy of your work becomes known when your unknown is graded. A PASS grade indicates good accuracy, a PARTIAL CREDIT grade indicates marginal accuracy, and a REPEAT grade indicates that your results had poor accuracy.
2. Precision: The term "precision" describes the "reproducibility" of results. It can be defined as the agreement between the numerical values of two or more measurements that have been made in an identical fashion. Good precision does not necessarily mean that a result is accurate.
3. Range: The "range" is one of several ways of describing the precision of a series of measurements. The range is simply the difference between the lowest (or lower) and the highest (or higher) of the values reported. As the range becomes smaller, the precision becomes better.
Example: Find the range of 10.06, 10.38, 10.08, and 10.12.
Range = 10.38 – 10.06 = 0.32
4. Mean: The "mean" or "average" is the numerical value obtained by dividing the sum of a set of repeated measurements by the # of individual results in the set.
Example: Find the mean of 10.06, 10.38, 10.08, 10.12
(Note that the value 10.38, which is far greater than the other values, has a large influence on the mean, which is larger than three out of the 4 individual values.)
5. Median: The "median" of a set is that value about which all others are equally distributed, half being numerically greater and half being numerically smaller. If the set has an odd number of measurements, selection of the median may be made directly. (Example: the median of 7.9, 8.6, 7.7, 8.0 and 7.8 is 7.9, the "middle" of the five). For an even number, the average of the central pair is taken as the median (Example: the median of 10.06, 10.38, 10.08, and 10.12 is 10.10--the average of the middle pair of 10.08 and 10.12). Notice in the example that the median is not influenced much by the value 10.38, which differs greatly from the other three values. For this reason, the median is usually better to use in reporting results than the mean for small data sets.
6. Error: The absolute error of an experimental value is the difference between it and the true value. For example if the experimental value is 30.9 and the true value is 26.5, the error would be 30.9–26.5 or 4.4.
7. Relative percent error would be the error divided by the true value times 100: (4.4/26.5)x100%=16.6% or 17%.
Experiment One: Measurement of Mass,
Volume, Density, and Fermentation
In this experiment, you will master various techniques for measuring mass and volume and apply these techniques to the study of density and fermentation. In part A, you will become familiar with the use of electronic balances to measure mass.
In Part B you will use a graduated cylinder and the balance to determine the density of various metal samples. If you determine the correct density for the metal samples then you may assume you are using the graduated cylinder and balance correctly. You will also be measuring the density of an unknown metal sample.
Density is a characteristic physical property of matter that can be used to help identify a substance. The density of any given substance or object refers to its mass per unit volume. It is mathematically expressed as
and commonly has units of grams per cubic centimeter (g/cm3) or its equivalent, grams per milliliter (g/mL). For example, aluminum has a density of 2.70 g/cm3 and nickel has a density of 8.90 g/cm3, so measuring the density of a piece of metal would be one way to tell which metal it could be. Other properties, such as appearance, melting point, and chemical reactivity, could then be used to confirm the identification.
In part C of this experiment you will study the design, use, and precision of glassware used to measure volume. This will be done by measuring the mass of water delivered by various types of glassware, bottle top dispensers, and pipettors. The density of water (known at various temperatures) can be used to calculate the volume of water delivered to test your technique and analyze the precision of glassware.
Another part of this experiment involves the fermentation of apple juice to cider. This part of the experiment will begin in the first week of the experiment and be completed the second week. Different fermentation conditions may be used to investigate the nature of this reaction as time allows. More calculations using this data will be done in a later experiment.
Fermentation has been utilized for many years in the preparation of beverages. Materials from Egyptian tombs demonstrate the procedures used in making beer and leavened bread. The history of fermentation, whereby sugar is converted to ethanol by the action of yeast, is also a history of chemistry. Van Helmont coined the word “gas” in 1610 to describe the bubbles produced in fermentation. Leeuwenhoek observed and described the cells of yeast with his newly invented microscope in 1680. Joseph Black discovered carbon dioxide and showed it to be a product of fermentation in 1754. Lavoisier in 1789 showed that fermenting sugar produces ethanol and carbon dioxide, and he made quantitative measurements of the amounts consumed and produced [1].
In 1815, after the mole concept was established, Gay-Lussac showed that one mole of glucose produces exactly two moles of ethanol and two moles of carbon dioxide. But the process of fermentation puzzled some great chemists. Kutzing wrote in 1837, “It is obvious that chemists must now strike yeast off the role of chemical compounds, since it is not a compound but an organized body, an organism.” And then there were chemists such as Berzelius, who believed that yeast had a catalytic action [1].
It remained for Pasteur to show that fermentation was a physiologic action associated with the life processes of yeast. In his classic paper of 1857, he described fermentation as the action of a living organism. However, other chemists disputed his findings because the conversion of glucose to ethanol and carbon dioxide is a balanced equation. So a search was begun to find the substance in yeast that might cause the reaction. It took 40 years before a clever experiment by Eduard Buchner ended the search. He made a cell-free extract of yeast that still caused the conversion of sugar to alcohol. This cell-free extract contained the catalysts, which we now call enzymes that are necessary for fermentation. This discovery earned Buchner the 1907 Noble Prize. In 1905 Harden discovered that inorganic phosphate, added to the enzymes, increased the rate of fermentation and was also consumed itself. This result led to the isolation of many other types of sugars. Clearly, the history of biochemistry is closely associated with the study of alcoholic fermentation [1].
Ancient peoples discovered many of the essential reactions of alcoholic fermentation completely by accident. How did they discover that the starch of wheat or barley could be converted to sugar by the enzymes in malt? When grain germinates, enzymes are produced that turn the starch into sugar. The process of malting involves letting the grain start to germinate and then heating and drying the sprouts to stop the process before the enzymes are used up. The color of the malt depends on the temperature of the drying. The darkest is used for stout and porter, the lighter, for brown, amber, and pale ale. At some point hops were added to beer when it was discovered that the resulting beverage did not spoil so rapidly [1].
The process of fermentation is used to make other products. Fermentation products have both industrial and medicinal applications. Penicillin, other antibiotics, and riboflavin vitamins are all created by fermentation. Citric acid created by fermentation is used in metal cleaners and in food preservation. This acid naturally occurs in citrus fruits, but getting it directly from the fruits is uneconomical. By fermenting molasses, the same acid is obtained cheaply and efficiently.
Different fermentation processes are used depending on the product. Some forms of fermentation are naturally occurring. That is, the yeast, mold, or bacteria is already in the reactant. In other situations the ferment, that is the yeast, mold, or bacteria, must be carefully added under specific conditions. But what happens is always the same. The sugar molecules in the substance are converted to something else, such as alcohol, or lactic acid.
In this experiment alcoholic fermentation will be explored through a traditional production of cider. This fairly simple method calls for picking apples, aging them for a week and then crushing them using a cider press to collect the juice. The freshly pressed juice is then fermented by immediately storing the juice in a wooden barrel and allowing wild yeast to ferment the juice. The fermentation starts in 1-2 days and continues for several weeks, during which time the barrel is topped off with more juice. Once fermentation is over, the barrel is sealed and matured for 5-6 months [2].
This traditional cider is often served completely flat and may be cloudy. It may also be served as a naturally conditioned cask cider, analogous to real ale. In France, cider is produced by the Charmat process (often used to produce sparkling wines) and is highly carbonated and more like an apple wine than traditional English cider [2].
The cider made in this experiment will be prepared using bread yeast or, if available, a dry wine yeast.
The Experiment
Risk Assessment – low hazard: no hazardous chemicals in use.
Measurement of Mass
There are many different types of balances or scales available to measure the mass of an object. The selection of the balance depends on the mass of the object or sample and the precision needed for the measurement. In this course the measurement of mass will be done using top loading electronic balances. Our balances accurately measure mass to ±0.001 g or ±1 mg, so always record masses determined on these balances to this precision (three digits to the right of the decimal place in grams) even if the last digit is zero. All mass measurements will be done in grams. Please note that balances will change in the last digit ±0.002 g; this is to be expected.
There are many different types of electronic top balances but they all use same two simple procedures. To simply weigh an object tare the balance to zero and then place the object on the balance to measure its mass. Weighing-by-difference is used to measure the mass of a sample being transferred from one container to another and will be used in experiment 6. A few rules need to be followed when using a balance.
Figure 1. Top loading electronic balance
Balance Rules and Instructions
Figure 1 illustrates one type of top loading electronic balance, which is one type of balance used in CHM 151L. Refer to this figure when following the steps and precautions for using the balance listed below:
1. Never pour or transfer chemicals over the balance. Spilled chemicals can damage the balances, which are very expensive to repair or replace. Never weigh warm or hot objects; if you can feel any heat, the weighing will not be accurate. Always use a container such as a vial, beaker, flask, or watch glass to weigh a solid or liquid chemical on the balance to protect the balance pan.
2. Make sure your hands are clean and dry before you handle containers or objects that are to be weighed. The outside of these containers or objects must also be clean and dry. Clean up any spills on the balance pan or lab bench around the balance immediately.
3. First open or remove the draft lid or cover (if there is one) and check to make sure that the balance pan is clean. If the pan is dirty, have your TA show you how to clean it and gently place it back on the balance.
4. Close or put the cover back on the balance and zero the balance by pressing the "T" or "on/tare" button. Wait 5-10 seconds for the weight display to stabilize. (If the object to be weighed is so large that the draft lid can't be used, do this step without the draft lid in place.)
5. Open or remove the draft lid and place the object to be weighed on the balance pan. Then close or place the draft lid back on the balance. (As long as it does not touch the object to be weighed, leave the lid off if it does touch the object.) After 5-10 seconds the weight display will stabilize and you can record the mass to ±0.001 g.
6. Never unplug the balance.
Weighing Solids and Liquids
Since using the top loading electronic balance is so much easier than using the old triple beam balances or 1 mg mechanical analytical balances, very few errors are made measuring the mass of an object. Most errors are made when trying to measure the mass of solid or liquid transferred from one container to another (weighing by difference). The following are some helpful hints to keep in mind when weighing by difference.
1. Be very careful to avoid spilling material outside the target container.
2. If you are weighing the container that the material is being transferred from, do not use a spatula to transfer the material, but gently tap the container to slowly transfer the material into a new container.
3. Make sure the outside of the container is clean and dry before you weigh it for the first time and then touch it as little as possible until after the final weighing.
4. Set containers to be weighed on clean surfaces only.
5. Always cool containers or samples to room temperature before you weigh them.
6. It is sometimes helpful to preweigh the sample before it is transferred.
Some of these hints will be more important in future experiments (especially Experiment. 4, 6 and 7)
Part A - Balance Calibration Check
All of the balances are regularly checked for correct calibration by the stockroom. To insure that you are using the balance correctly and that it is properly calibrated, the mass of a calibration weight will be measured. Calibration weights are provided in the balance room in the small wooden boxes. Make sure your hands are clean and dry before you touch the weights. (Normally calibration weights are not touched with your hands.)
Select a calibration weight and record the "Known Mass" for the weight (stamped on the weight) on the report sheet found at the end of this experiment. Measure the mass on a balance and record it on the report sheet. If the measured and known mass differ by more than 0.01 g reread the instructions for using the balance and measure the mass of the calibration weight again. If you get the same results again, see your instructor or TA.
Part B - The Density of a Metal Samples and Fermentation
Obtain an Aluminum (Al) metal sample from the reagent bench. You must return this metal sample to the box on the reagent bench as soon as you are done using it or by the end of the current lab period at the latest.
Make sure the metal sample is clean and dry. Weigh the metal sample on the balance and record the mass.
Note: Water is purified using various techniques to varying levels of purity measured by electrical conductivity. For this lab course we will use tap water for general washing and water purified by reverse osmosis (RO) and UV to kill bacteria and other microbial impurities. There are RO taps by side sinks in each lab. Fill your plastic water bottle with this water for use with your experiments. There will also be a rinse tub of RO water to give cleaned glassware a final rinse. Water may also be purified by distillation, ion exchange, carbon filtration, and other steps to even higher levels of purity. Pure water may be referred to as RO, distilled, or deionized. Use RO water in your experiments when pure water is called for.
Fill a 25 mL graduated cylinder approximately halfway with RO water. Read the volume of water accurately by estimating to the nearest 0.1 mL (See Figure 2). Record the result on the report sheet. Then carefully lower the aluminum metal sample into the graduated cylinder without losing any water. Be sure that the water completely covers the sample and that the meniscus is still within the volume scale of the graduated cylinder. If not, start the procedure over using more or less water. Read the total volume of the contents in the cylinder and record it on the report sheet. Repeat this procedure with one piece of the copper sample.
Figure 2. Always read the bottom of the
Densities of Metals meniscus. This reading is 36.5 mL.
|
Symbol |
Metal |
Density(g/mL) |
|
Al |
Aluminum |
2.7 |
|
Cr |
Chromium |
7.2 |
|
Fe |
Iron |
7.9 |
|
Cu |
Copper |
8.9 |
|
Ni |
Nickel |
8.9 |
|
Ag |
Silver |
10.5 |
|
Pb |
Lead |
11.3 |
|
Hg |
Mercury |
13.9 |
|
Au |
Gold |
19.3 |
Record the data for each trial directly on your report sheet. Immediately return the metal samples to the reagent bench (do not store them in your bin or workstation drawer).
Calculate the volume of each metal sample. Finally, calculate the density in g/mL of each sample. Compare the densities you measured with those listed on the previous page. The density for Al should be very close to the listed value, but the density for Cu may differ from the listed value by a large margin. Explain why the value for Cu is poorer than that for Al. What measurement limits the precision or accuracy? Now modify the procedure for measuring the density of Cu so that the density can be measured more accurately. Write your procedure on your report sheet and use the procedure to measure the density of your copper sample again. A very common source of error in the experiment is from inaccurate reading of the graduated cylinder. Take the time to make sure that it is read correctly.
Fermentation of Apple Juice to Cider – First Week
The data for the fermentation will be recorded on the back of the report sheet. First clean and dry a 125-mL Erlenmeyer flask. Weigh the flask. To the empty, clean flask add 5-15 grains (TA will assign exact amounts for each student) of yeast and reweigh the flask. Be sure there are at least 5 grains. Now add 10, 15, or 20 mL of room temperature apple juice using a 25 mL graduated cylinder. Obtain the mass of the flask containing the juice and yeast. Once this is done, place a balloon over the top of the flask, making sure the balloon stretches approximately 1 inch down the neck of the flask. Now place a twist tie or rubber band tightly around the neck of the flask to help seal the balloon on the flask. Let the apparatus sit for 1 week undisturbed. You TA may have you form groups for the fermentation experiment.
The Density of an Unknown Metal Bar
Now go to the chemistry stockroom and check out a metal bar with an unknown density. Record the bar number on your report sheet. Follow the procedure used for Al to determine the density of the unknown and do two trials. Immediately return this bar to the stockroom when you are done. Do not put it in your locker. Be sure to use the calculation check to make sure your density calculations are correct and print a copy of it to attach to the report sheet. In addition to recording the final results for the unknown on the report sheet also record them in the space provided at the end of the experiment in case the report sheet is lost.
Part C - Volume Measurement and Calibration:
You are now familiar with the graduated cylinder. In this part of the experiment the use of the volumetric pipet, the Mohr pipet, bottle top dispenser, and pipettor will be mastered. Use the report sheet at the end of this experiment to record your results. Before you begin, view the videos on using this equipment if you have not done so.
CAUTION: Do not use your mouth to draw liquid into the pipet. Use the pipet bulb and tip provided.
Calibration of Volumetric Pipets
You will begin by calibrating your volumetric pipets with RO water. First, be sure your pipets are clean (review figure 3 before using your pipets). The pipet bulb should have a plastic pipet tip or adaptor that makes the pipet easier to use. If the bulb draws the liquid up slowly make sure the plastic tip or adaptor and bulb make a good seal and perhaps use scissors to cut off about 1/16 of an inch of the plastic tip. Take some time now to practice using your pipets and make sure they are clean. Start by using pure water practice making 2 or 3 transfers with the 10 mL pipet. Measure the temperature of the pure water and record it on the report sheet. Using the correct size pipet, transfer the volume called for into an empty dry vial with cap that has been weighed to the nearest milligram. Cap the vial to prevent evaporation and weigh.
Figure 3. Using a Volumetric Pipet.
Blowup of steps 2, 4 and 6.
(1) The rubber bulb in your locker should have a pipet adapter or a disposable plastic pipet tip attached (if not, see your TA). The pipet is filled by placing a depressed rubber bulb assembly in the top and slowly releasing the bulb, thereby drawing liquid up into the pipet above the graduation line. Be careful not to draw the liquid into the bulb, as the bulb will contaminate the liquid.
(2) Remove the rubber bulb and immediately place your index finger on the top of the pipet.
(3) Wipe the outside of the pipet with a paper towel.
(4) While resting the tip of the pipet on the top inside portion of the container holding the solution, carefully raise your finger to release enough liquid so that the bottom of the meniscus is on the graduation mark.
(5) You are now ready to dispense a specific quantity of liquid; touch the inside upper wall of the target container with the tip of the pipet as you drain it. Hold the tip to the wall 5-10 sec. after the liquid has drained out. Note: If any droplets are found to be adhering to the inner walls of the pipet, it is dirty and should be cleaned. Try cleaning the pipet using hot soapy water. Draw the soapy water up into your pipet several times using your rubber pipet bulb. Then rinse the pipet several times with tap water and finally once with RO water. If the pipet is still dirty take it to the chemistry stockroom where it can be exchanged for a clean one. Be sure to view the video on the pipet.
The volume of water delivered is then calculated by dividing the mass of the pure water by the density.
The density of water can be obtained from the table below using the temperature of the water. Use the density for the temperature closest to the temperature you measured and record it on your report sheet. Calculate the error by subtracting the volume you tried to deliver (size of volumetric pipet used) from the volume calculated above using the mass and density.
The volume for each pipet should agree with the pipet size within ±0.05 mL. If the error in volume is off more than ± 0.05 mL, you have either made a pipetting error, your pipet is dirty or damaged, or possibly the calibration of your pipet is off. In any case, get some help from your TA and repeat the calibration. Record the calibrations for the volumetric pipets at the end of this experiment.
Density of Water
|
|
Temperature |
Density |
Temperature |
Density |
|
|
(Degrees Celsius) |
(g/mL) |
(Degrees Celsius) |
(g/mL) |
|
|
14 |
0.9993 |
25 |
0.9971 |
|
|
15 |
0.9991 |
26 |
0.9968 |
|
|
16 |
0.9990 |
27 |
0.9965 |
|
|
17 |
0.9988 |
28 |
0.9963 |
|
|
18 |
0.9986 |
29 |
0.9960 |
|
|
19 |
0.9984 |
30 |
0.9957 |
|
|
20 |
0.9982 |
31 |
0.9954 |
|
|
21 |
0.9980 |
32 |
0.9950 |
|
|
22 |
0.9978 |
33 |
0.9947 |
|
|
23 |
0.9976 |
34 |
0.9944 |
|
|
24 |
0.9973 |
35 |
0.9941 |
Calibration of the Mohr Pipet
The Mohr pipet and the buret (the buret will be used in experiment 6) are graduated from the top down instead of from the bottom up like the graduated cylinder. This often creates confusion when reading the Mohr pipet or buret but this makes them valuable for transferring variable amounts of liquid with precision.
Figure 4; illustrates a portion of the scale of a Mohr pipet or a buret. Reading the scale from the top down, the bottom edge of the meniscus is resting on 6.63 mL and not 7.37 mL. To deliver a specific volume, fill the Mohr pipet to above the zero mark and drain such that the bottom of the liquid (meniscus) is on zero or below and note the volume. Then, deliver the required amount of solution and note the new volume on the scale. The volume delivered is the difference between the first and second readings. Practice using the Mohr pipet and make sure it is clean.
To become familiar with the use and the calibration of the Mohr pipet, use nearly the same procedure as was done for the volumetric pipet. Use 10.00 mL of RO water to complete the calibration. Because of the design of most of our Mohr pipets, you do not completely drain them to transfer 10 mL. Instead, if you start at 0, the pipet is drained from the 0 to the 10 mL calibration marks on the graduated scale. A few of our Mohr pipets are calibrated all the way to the tip of the pipet. In this case the graduations indicate that the pipet should be completely drained to deliver 10.00 mL. Both this type of mohr pipet and our volumetric pipets are “To Deliver” or “TD” type pipet. This means that the tiny amount of solution left in the tip after a transfer is not blown out!
Comparison of Glassware Accuracy: Graduated Cylinder and Beaker
The terms "accuracy", "precision" and "range" are defined in the Introduction. The choice of glassware to measure volume is dependent on the goal of the experimental procedure, the volume size, the precision required for the final answer, and the time it takes to perform the manipulation. The measurement of volume is a tradeoff between speed and precision. More precise measurements are more time consuming. If a procedure called for 10.00 mL for one part of the experiment and 10 mL for another part, different types of glassware would be used to measure the volumes.
Let's investigate the accuracy to which differing glassware can measure volume. Measure out 10 mL of RO water with your 10 mL graduated cylinder and then using your 50 mL beaker. Weigh the water and calculate the volume transferred or "calculated volume" using the density of water as was done for your pipets. Record your results on the report sheet for experiment one.
For each type of glassware, calculate the volume delivered using the density of water and determine the error in volume delivered (calculated volume delivered minus desired volume). The error in these measurements will give an indication as to just how accurate each type of glassware is (a smaller error indicates higher accuracy). Compare these errors to the precision listed for similar glassware in the table on page 10 in the significant figures section of the Introduction. Complete and print the calculation check for this part of the experiment.
Bottle Top Dispenser and Pipettors
Finally, calibrate other equipment that is used measure and transfer liquids. As you did above, calibrate a bottle top dispenser that has been set to dispense a 10.0 mL volume of pure water. Bottle top dispensers are used to quickly dispense liquids with some, but not exceptional precision. The following steps should be followed when using a bottle top dispenser:
1. Make sure the dispenser is set for the right volume. See your TA if it is not.
2. Make sure the dispenser has been primed and checked. See your TA if you are unsure about this.
3. Bring the container you want to transfer the sample into and hold it under the tip of the dispenser.
4. Pull the dispenser head up slowly until it stops.
5. Slowly press the dispenser head down with the container under the tip until it stops.
6. Check to make sure the volume transferred is about what you wanted.
Record your data for the bottle top dispenser on your report sheet.
Now calibrate the 5mL pipettor from your work station. We are going to modify the calibration procedure slightly but first let's review the steps for using pipettors. Using an automatic pipettor makes transferring precise volumes of liquids easy, but the procedure must be followed exactly. Always calibrate the pipettor with pure water to check the pipettor and your technique before using one. The following are general steps for using a pipettor:
1. Select the correct pipettor that best matches the volume to be transferred. Never go above or below the working volume range for a pipettor (1-5mL for 5mL pipettor).
2. Dial in the volume wanted and snuggly place the correct tip on the pipettor
3. Depress the plunger button to the first stop, put the end of the tip in liquid, and slowly let the plunger button up to draw liquid into the tip. Wait for 5-10 seconds (keep the end of the tip in liquid during this time and keep the pipettor vertical at all times!).
4. Point tip into container you want to transfer to and press the plunger button all the way down (past the first stop) and wait. This procedure may need to be modified for very viscous or dense liquids.
5. Remove the tip before setting the pipettor down (always keep tip pointed down until it is removed from pipettor).
To calibrate the pipettor we will do several trials. Place about 40-50 mL of pure water in a 50 or 100 mL beaker. Bring the pipettor, the beaker with pure water, and another beaker. Follow these steps to calibrate the pipettor:
1. Set the 5mL pipettor on its maximum volume (5.00mL) and place a tip on it.
2. Place the empty beaker on the balance and tare it to zero.
3. Take the beaker off the balance, pipet the maximum volume of the pipettor into the beaker and immediately place it back on the balance and record the mass on your report sheet.
4. Tare the beaker on the balance to zero, remove it from the balance and pipet the same volume into it.
5. Place it back on the balance and record the mass.