Southern California University of Health Empty Graduated Cylinder Lab Report

General Chemistry ISouthern California University of Health Sciences
GENERAL
CHEMISTRY
I
Laboratory Manual
Revised Spring 2021
General Chemistry I
Southern California University of Health Sciences
Contributions
Many individuals made significant contributions to this manual, especially the laboratory instructors and
students in general chemistry, and the technical staff of the Chemistry Stockroom at SCU.
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Table of Contents
Table of Contents ………………………………………………………………………………………………………………………………………………..3
Contributions ………………………………………………………………………………………………………………………………………………………2
Message to the Students ……………………………………………………………………………………………………………………………………….5
Laboratory Safety and Procedures ……………………………………………………………………………………………………………………….6
Laboratory Attire ………………………………………………………………………………………………………………………………………………..6
Safety Equipment ………………………………………………………………………………………………………………………………………………..6
Chemical Safety …………………………………………………………………………………………………………………………………………………..9
Handling Chemical Reagents ……………………………………………………………………………………………………………………………..10
Experiment 1: Mass Volume and Significant Figures …………………………………………………………………………………………..11
Experiment Procedure ………………………………………………………………………………………………………………………………………15
Prelab …………………………………………………………………………………………………………………………………………………………….18
Report…………………………………………………………………………………………………………………………………………………………….19
Postlab ……………………………………………………………………………………………………………………………………………………………21
Experiment 2: Formula of a Hydrate ………………………………………………………………………………………………………………….24
Experiment Procedure ………………………………………………………………………………………………………………………………………28
Prelab …………………………………………………………………………………………………………………………………………………………….29
Report…………………………………………………………………………………………………………………………………………………………….30
Postlab ……………………………………………………………………………………………………………………………………………………………33
Experiment 3: Reaction Stoichiometry………………………………………………………………………………………………………………..34
Experiment Procedure ………………………………………………………………………………………………………………………………………37
Prelab …………………………………………………………………………………………………………………………………………………………….38
Report…………………………………………………………………………………………………………………………………………………………….39
Postlab ……………………………………………………………………………………………………………………………………………………………42
Experiment 4: Beer’s Law Spectrophotometric Determination of Copper …………………………………………………………..43
Experiment Procedure ………………………………………………………………………………………………………………………………………46
Prelab …………………………………………………………………………………………………………………………………………………………….48
Report…………………………………………………………………………………………………………………………………………………………….49
Experiment 5: Analysis of Unknown Solutions ……………………………………………………………………………………………………51
Experiment Procedure ………………………………………………………………………………………………………………………………………54
Prelab …………………………………………………………………………………………………………………………………………………………….55
Report…………………………………………………………………………………………………………………………………………………………….56
Experiment 6: Calorimetry and Specific Heat Capacity ………………………………………………………………………………………58
Experiment Procedure ………………………………………………………………………………………………………………………………………62
Prelab …………………………………………………………………………………………………………………………………………………………….63
Report…………………………………………………………………………………………………………………………………………………………….64
Postlab ……………………………………………………………………………………………………………………………………………………………66
Experiment 7: Molecular Structure and Properties ……………………………………………………………………………………………..67
Experiment Procedure ………………………………………………………………………………………………………………………………………74
Prelab …………………………………………………………………………………………………………………………………………………………….76
Report…………………………………………………………………………………………………………………………………………………………….77
Experiment 8: Charles’ Law ………………………………………………………………………………………………………………………………79
Experiment Procedure ………………………………………………………………………………………………………………………………………81
Prelab …………………………………………………………………………………………………………………………………………………………….83
Report…………………………………………………………………………………………………………………………………………………………….84
Postlab ……………………………………………………………………………………………………………………………………………………………86
Worksheets ………………………………………………………………………………………………………………………………………………………..87
Worksheet A …………………………………………………………………………………………………………………………………………………..88
Worksheet B……………………………………………………………………………………………………………………………………………………89
Worksheet C……………………………………………………………………………………………………………………………………………………91
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Worksheet D …………………………………………………………………………………………………………………………………………………..94
Conversions…………………………………………………………………………………………………………………………………………………….97
Stoichiometry ………………………………………………………………………………………………………………………………………………..101
Thermochemistry …………………………………………………………………………………………………………………………………………..108
Periodic Table ………………………………………………………………………………………………………………………………………………….112
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Dear Students
The laboratory exercises in this manual are designed to help you develop a sense of how chemistry (and
science in general) is accomplished. In addition to learning many essential laboratory techniques, you will learn
how to plan experiments, acquire data, record observations, and properly present your findings. You will also
learn procedures for working safely in the laboratory.
To ensure your success in the laboratory, it is very important that you come to class prepared. Prior to the
laboratory period, you should always read through the description of the assigned exercise found in this manual.
You will also be expected to review short sections in your lecture text, in order to better understand the
principles of the exercise. Each activity will also have a short set of pre-laboratory questions. Because time is
limited in the laboratory, you must plan your course of action in advance. Students who fail to demonstrate the
proper preparation before coming to laboratory will not be allowed to work in the laboratory.
At the beginning of each laboratory period your instructor will usually present a short pre-laboratory lecture.
This discussion will be used to highlight safety provisions, demonstrate proper techniques and equipment use,
and discuss any questions you may have. Come to the laboratory prepared to ask questions about any part of the
experiment that you do not understand.
For many of the laboratory exercises you will work with other students in a team. Every student is expected
to contribute equally to the completion of the laboratory activity. You may be responsible for maintaining an
individual laboratory notebook, and when unknowns are analyzed, each student will have a different unknown
and will report separate results.
For most chemists, the laboratory is where the excitement of chemistry occurs. Hopefully, you too will
discover some of this excitement as you explore these activities. Your comments and suggestions are an
important contributor to this endeavor and are always welcome.
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Laboratory Safety and Procedures
There are many hazards associated with work in the chemistry laboratory. However, when the proper
procedures are followed these hazards can be adequately controlled. The key to safe laboratory work is proper
preparation. Always come to laboratory with a thorough understanding of what steps you will follow in a
particular experiment. If something is not clear, ask your laboratory instructor for more information before you
begin work. You should never change the experimental procedure unless you have discussed this with your
laboratory instructor. The following sections provide basic information about laboratory safety and common
laboratory procedures.
Laboratory Attire

Safety goggles must be worn at all times in the laboratory. While in the laboratory, you must wear
approved splash-proof safety goggles, even if you personally are not doing any experimental work. Many of
the injuries that result from flying glass or splashed chemicals originate at a considerable distance from the
victim and can be avoided through the use of safety glasses. Failure to wear safety goggles over your eyes
may be grounds for dismissal from the laboratory.

Contact lenses should not be worn in the laboratory. The concern with contact lenses is that chemical
splashes might get trapped behind the contacts and cause more damage before the chemicals get flushed
from the eyes. If you wear contact lenses, inform your laboratory instructor and follow his/her
recommendations.

Shoes must be worn at all times in the laboratory. Open-toe shoes or sandals should not be worn because
they do not offer adequate protection against chemical spills or broken glass.

A laboratory coat and long pants are highly recommended. It is advisable to wear a laboratory coat or
apron and long pants to prevent damage to your clothing or skin from corrosive chemicals.
Safety Equipment

Learn the location of the safety shower, eyewash, fire extinguisher, and first-aid kit. Your laboratory
instructor will point out the location and operation of the laboratory safety equipment available in the lab.

Use of a Fire Extinguisher. Know the location of the fire extinguisher(s) in your laboratory. A simple
PASS method can help you remember the use of the extinguisher in the event of a fire:
1. Pull the pin on the extinguisher.
2. Aim the hose nozzle low toward the base of the fire.
3. Squeeze the handle to release the extinguishing agent.
4. Sweep the nozzle from side to side at the base of the flames until extinguished.
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Proper Laboratory Procedures

In case of an accident, notify your laboratory instructor immediately. Even if you believe that a cut or
chemical in contact with your skin does not need attention, you should inform your laboratory instructor
immediately.

Do not eat or drink in the laboratory. Eating and drinking (even water) are prohibited in the laboratory,
in order to avoid the accidental ingestion of toxic materials. Food or water that you might have with you
must be contained in a closed backpack/bag, and cannot be placed on the laboratory bench.

Keep your laboratory bench workspace clean and neat. An unorganized work area leads to confusion,
whereas experiments performed in a neatly kept area will proceed more rapidly, successfully, and safely.
You must leave your work area clean when you leave the laboratory.
Any spilled chemicals must be cleaned up immediately. If in doubt as to the proper method of cleaning up a
spilled chemical, ask your laboratory instructor. Special care should be taken not to spill chemicals on the
balance pans, and any spills that do occur should be cleaned right away.
If acids or bases are spilled, first neutralize them with solid sodium bicarbonate (a container is located at
each sink), then remove the solid residue and clean the area with a wet paper towel.

Dispose of chemicals properly. Before the beginning of each experiment your laboratory instructor will
instruct you in the proper disposal of each chemical you will be using.
Chemical collection bottles will be provided for those chemicals which should not be flushed down the sink.
Make sure you dispose of each chemical in the proper bottle. Placing a chemical in the wrong bottle may
result in undesirable chemical reactions. When in doubt, ask!

Dispose of broken glass properly. Broken glass should be discarded only in the specially marked glass
disposal containers. If a piece of glassware breaks, use the dustpan and brush to collect all of the broken
pieces. Avoid handling the broken glass directly.

Always use clean glassware for each experiment. If you wash dirty equipment at the end of each
laboratory period, it will be easier to clean and will be available when you next need it. Clean glassware by
(using the designated glassware detergent bottles located at each sink) first preparing a flask of detergent
solution made by mixing a small amount of liquid detergent in warm water. Use this solution to clean your
dirty glassware, and then rinse the detergent away. A final rinsing should always be made with a stream of
deionized water from your plastic wash bottle. Never dry glassware with compressed air, as it is often oily.

Laboratory equipment should be returned to its proper location. All common equipment (e.g., clamps,
ring stands, burets, etc.) used for an experiment must be returned at the end of the laboratory period to the
location from which it was taken. Any special equipment that you obtain from the stockroom for an
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experiment must be returned to the stockroom at the end of the laboratory period.

Balances are to be kept clean at all times. The analytical balances and top loading balances in the
laboratory are very useful and convenient tools, however, proper care of the balances is necessary if they are
to function reliably. Chemical spills are the biggest problems. To prevent corrosion, it is very important that
any chemicals spilled on a balance are cleaned up immediately.
Do not place chemicals directly on the balance pans. Instead, first place a small container on the pan, zero
the balance, then weigh your material. If you cannot tap out the chemical conveniently into the container,
remove the container from the balance pan area to make the transfer.
Do not use weighing paper to hold your sample. It is too easy to spill chemicals in the balance pan area
when transferring to weighing paper.

Never pipet any liquid directly by mouth. Always use a safety pipet bulb when pipetting. The possibility
of getting a toxic chemical into your mouth greatly outweighs any inconvenience involved with the use of
the pipet bulb.

Chemicals should be heated with great care. When heating a substance in a test tube, use a test tube
holder and be careful not to direct the tube toward yourself or your neighbor. A suddenly formed bubble of
vapor may eject the liquid contents violently and dangerously. Do not heat small test tubes directly in a
flame because the probability of the contents being ejected increases substantially.
Never apply the direct heat of a flame to glassware such as volumetric flasks, burets, graduated cylinders,
bottles, and thermometers. These objects break spectacularly when heated too much. Volumetric glassware
can be permanently distorted from the calibrated volumes by such heating.
Avoid heating any object too quickly. Apply the flame intermittently at first, especially to control the rate of
a chemical reaction.
During any procedure involving the heating of glassware, avoid burns by allowing adequate time for the
glassware to cool before touching. It takes 20–30 minutes for heated glassware to cool to room temperature.
Use the appropriate type of tongs when moving hot objects. Test tube holders should be used for test tubes
only. Larger tongs should be used to carry flasks or other heavier objects.
Never place hot glassware directly on the laboratory bench. Use a wire mesh or other suitable insulator to
support the hot glassware.
Never use an open flame and flammable liquids at the same time. If you are using a flammable liquid, make
certain there are no open flames in the area and preferably carry out your experiment in the fume hood.
Rinsing or soaking the burned area in cold water generally best treats burns. As with any accident, notify
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your laboratory instructor immediately.
Chemical Safety

All chemicals should be considered potentially toxic and flammable. A complete description of the
hazards associated with a chemical is found on its Material Safety Data Sheet, or MSDS. The MSDS for
each chemical used in your experiments is on file in the Chemistry Stockroom and may be examined upon
request. Your instructor will discuss the toxicity and other dangerous properties of any materials you are
using for which special precautions are necessary.

Never taste a chemical. Never smell a chemical unless instructed to do so. Nearly all chemicals are
poisonous to the human body to some extent. If you are asked to smell a laboratory chemical, gently fan the
vapors toward your nose and smell cautiously.

Any chemical in contact with the skin should be washed off immediately. Immediately wash off any
chemical with copious quantities of water. If the area of contact is large, use the safety shower.
If the eyes are splashed with a liquid, immediately flush them using the eyewash. Continue flushing the eyes
for 15–20 minutes to remove the chemical. Notify your laboratory instructor if you spill any chemical on
yourself. Many toxic organic compounds, which are not corrosive, are absorbed through the skin with no
immediately visible symptoms. Be sure to wash your hands before leaving the laboratory.

Use a fume hood for all experiments involving hazardous gases or vapors. Never purposely allow
fumes or volatile liquids to escape into the open room. Inhalation of more than a slight dose of such fumes
should be reported to your laboratory instructor and you should get to fresh air immediately. The sash on the
fume hood should be positioned as indicated, to ensure that fumes do not escape into the room accidentally.

Never pour water into acid; slowly add acid to water with constant stirring. “Do what you oughta,
pour acid into water.” Heat is liberated in the dilution process. If water is poured into acid, steam may form
with explosive violence, causing splattering of the acid. By pouring the acid slowly into water, you reduce
the risk of a dangerous chemical splatter.
Handling Chemical Reagents

Read the label carefully before using a reagent. Some experiments involve the use of substances with
similar names and formulas. Use of the wrong chemical could result in a serious injury. It is critical that you
take the time to select the proper reagents for a particular experiment.

Do not take reagent bottles to your work area. Use test tubes or small beakers to obtain chemicals from
the dispensing area. Label your container with a laboratory marker before dispensing the material. Do not
take any more material than is required.
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Do not contaminate reagents. Do not insert your own pipets, droppers, or spatulas into the reagent bottles.
To avoid contamination of the materials, pour them from the bottles into your container.
Never return unused chemicals to the reagent bottles. An entire supply of a reagent can be contaminated this
way.
Always return the proper cap to the proper bottle and ensure that it is secured on the bottle.

Pour from a reagent bottle so the liquid flows from the opening on the side opposite the label. Pouring
from the side opposite the label ensures that any chemical that spills will not obscure or destroy the reagent
label. Make it a habit to hold the label when pouring from a reagent bottle.
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Experiment 1: Mass, Volume, and Significant Figures
Learning Objectives:
As you work through this exercise you will learn how to:

Use balances to properly weigh samples.

Make volume measurements with volumetric glassware.

Report the proper number of significant figures for an experimental result.
Introduction
Note: You may wish to review significant figures, the units for mass and volume, and the powers-of-ten
prefixes used to scale units in your textbook before starting this experiment.
Mass Measurements
One of the most common measurements made in any laboratory involves determining the mass of an object.
Modern balances make this process very simple. However, large errors can result if the proper technique is not
followed. Part of this exercise will give you practice weighing objects with an analytical balance and a toploading balance. An analytical balance is used to weigh small samples (typically much less than 100 g) to the
nearest tenth of a milligram (± 0.0001 g), whereas a top-loading balance is used for heavier objects or to get
approximate masses, and usually reports the mass to the nearest 0.1 g or 0.01 g.
The reliability of a delicate balance depends upon how it is treated by the user. Avoid placing chemicals directly
on the balance as this can lead to corrosion. For long balance life always:
1. Weigh objects in a small container. Liquids should be weighed in a closed container.
2. Immediately clean up any spills (solid or liquid) on, in, or near the balance.
3. Avoid damaging the balance by carefully placing objects to be weighed on the balance pan. Most
balances have an upper limit of mass that can be placed on them. Most analytical balances can’t weigh
more than 100g without potentially damaging the balance.
Volume Measurements
Another important measurement in the laboratory is that of volume. Many reagents are dispensed as
aqueous solutions (substances dissolved in water) and the volume used is the means of knowing how much of a
substance you are working with. To measure volumes, you will use equipment such as graduated cylinders,
transfer pipettes, burets, and volumetric flasks. Each piece of equipment is used for a specific purpose and the
error associated with a volume measured by each of these devices varies, as noted in Table 1. Each of these
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pieces of glassware has a certain amount of error that is associated with them as no manufacturing process is
perfect. The error that a piece of glassware has will inform you on how many significant figures are appropriate.
For example, a 10-mL transfer pipette will have 2 significant figures past the decimal place.
Table 1. Errors in Volume Measurements
Volumetric Glassware in Lab Expected Error (mL)
50-mL buret
± 0.03
10-mL transfer pipette
± 0.06
50-mL graduated cylinder
> ± 0.5
50-mL volumetric flask
± 0.05
Note that burets, pipettes, and volumetric flasks tend to deliver volumes much more precisely (to the nearest
few hundredths of a milliliter) compared to the graduated cylinder. A graduated cylinder is used, much like a
top-loading balance, for larger samples or for approximate volumes.
All pieces of volumetric glassware have calibration marks (typically a white or red line marked on the
glassware) that correspond to specific volume levels. If you are filling a transfer pipette or a volumetric flask,
the liquid level is adjusted until the bottom of the meniscus is just touching the calibration mark. For burets and
graduated cylinders, the liquid level may be between calibration marks. In all cases, carefully read the liquid
level at the bottom of the meniscus making sure your line of sight is horizontal (Figure 1). The magnifying lens
in your locker might make it easier to locate the meniscus level.
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24
Figure 1. Reading a meniscus – keep your eye level with the meniscus and always estimate one more significant figure than indicated by
the engravings or markings on the scale. A piece of paper behind the meniscus sometimes helps improves contrast.
All volumetric glassware must be clean before use. Cleanliness can be checked by filling the pipette or
buret with water and allowing the liquid to drain. No water droplets should be observed on the inner walls. After
cleaning and rinsing with water, volumetric glassware should not be dried. Instead, rinse pipettes, burets, and
graduated cylinders with a small amount of the liquid to be dispensed/measured before use. Your
volumetric pipette is calibrated to deliver (TD) the indicated amount of liquid by gravity drain only. As the
pipette drains, hold the tip of the pipette to the inner wall of the collecting vessel. When the flow of liquid from
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the pipette is complete, a small amount of liquid will remain in the tip. A TD pipette is calibrated taking this
extra liquid into account. Do not shake or blow out this liquid.
The volume of a given amount of liquid will usually increase with an increase in temperature. For example,
the volume occupied by 1000 mL of water at 15 °C will occupy 1002 mL at 25 °C. For this reason pipettes,
burets, and volumetric flasks are usually calibrated at a specific temperature (20 °C). Using such glassware at
room temperature (about 25 °C) will produce an error in the measured volume of about 1 part per thousand
(0.1%). This error is generally negligible for our experiments.
Density
Density is a physical property of a substance and can be used as an aid to its identification. Density (d) is
defined as the ratio of the mass (m) of a substance to the volume (V) occupied by that mass:
Density 
mass
m

volume V
Thus, density is the conversion factor between mass and volume. Often, the units used for density are grams per
cubic centimeter, g/cm3. Since 1 cm3 = 1 mL, g/cm3 is equivalent to g/mL. While the mass of a sample does not
change, the volume occupied by the sample varies with the pressure and temperature to which it is subjected.
Density, therefore, will also vary with pressure and temperature. The density of gases is affected by temperature
and pressure more than that of liquids, while solids are affected the least. For most considerations, the effect of
pressure on the density of liquids and solids is negligible. In this experiment you will use different tools to
determine the density of solid and liquid samples, and then report your calculated results to the proper number
of significant figures and with the appropriate units.
Significant Figures
Significant figures refer to all digits in a measurement that are certain, plus one uncertain digit.
Consequently, significant figures provide information about the quality of the measuring tool and the resulting
measurement. A greater number of significant figures indicates a more certain measurement. In this experiment,
you will explore how the significant figures in a calculated result depend upon the tools used to make the
measurements.
Rules for Significant Figures
There are ‘rules’ for guiding how many significant figures you should quote in an answer when performing
any calculation. These depend on what type of calculation you are performing. In general, write down as many
digits as you can all the way through to the end of the calculation and then perform rounding to bring your
answer to the correct number of significant figures. In this way you will avoid so-called ‘rounding errors’ that
result in the accumulation of many small errors through a multi-step calculation.
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Rules for Addition and Subtraction
When adding or subtracting numbers, report the answer with the same number of significant figures after
the decimal point as the least certain value.
For example, 13.01 + 423.8236734 = 436.8336734 = 436.83 because 13.01 has two significant figures after
the decimal point and so should the answer.
Rules for Multiplication and Division
When multiplying or dividing numbers, report the answer with the same number of significant figures as the
least certain value.
For example, 13.01 x 423.8236734 = 5513.94599093 = 5514 because 13.01 has four significant figures and
so should the answer.
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Experiment 1: Experimental Procedure
Weighing
1.
Obtain a metal strip of known mass and record its identification number. Determine its mass on an
analytical balance and record the result. The mass you obtain should agree with the posted mass within ±
0.5 mg (± 0.0005 g). If the mass differs by more than ± 0.5 mg, consult with your laboratory instructor.
Density of a Solid (white stir bar)
1.
Determine the density of a solid (stir bar) using a top-loading balance and a graduated cylinder for
volume displacement
Obtain a solid and record its weight using a top loading balance. Add approximately 25-35 mL of
water (this does not have to be exact) to a 50-mL graduated cylinder. Read the liquid level and record the
volume (to the proper number of significant figures). Carefully lower the stir bar into the graduated
cylinder. Remove any air bubbles that may appear and read and record the final water level. From this
volume change and the known mass, calculate the sample density to the proper number of significant
figures.
2.
Determine the density of a solid (stir bar) using an analytical balance and a burette for volume
displacement
Thoroughly dry your stir bar and weigh it on an analytical balance. Add approximately 20–30 mL of
water to a 50-mL burette (once again, this does not have to be exactly between 20 and 30. Read the liquid
level and record the volume. Carefully lower the solid sample into the burette. Remove any air bubbles that
may appear. Read and record the final water level. From this volume change and the measured mass,
calculate the sample density to the proper number of significant figures.
Calibrating the Volume of the 10-mL Volumetric Pipette
As mentioned in the introduction section, no piece of glassware is perfectly made in the manufacturing process
and your pipette has some error in it. Check the pipet to see if there is a +/- notation next to the volume
measurement, this is the expected error. In this section we will find out exactly how much volume your pipette
actually gives out by calculating density using temperature and measuring the mass of deionized water.
1.
Clean your 10.00 mL volumetric pipette with DI water. If it does not drain cleanly, see your instructor. Fill
a large beaker with deionized water and measure the water temperature. Estimate the temperature to the
nearest 0.1 °C.
2.
Weigh a clean, dry 50-mL Erlenmeyer flask, and a well-fitting rubber stopper on the analytical balance.
Next, remove the stopper, carefully pipette 10.00 mL of water into the flask, then replace the stopper and
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weigh the flask and contents. Calculate the mass of water delivered by taking the difference between the
two measurements.
3.
From the data listed in Table 2 below, find the density of water at your measured temperature. Estimate the
density by interpolation if your temperature is not one of the values listed on the table. For example, if your
estimated temperature is 24.4 °C, the density is given by:
 4

Density at 24.4 C  Density at 24.0 C    (Density at 25.0 C  Density at 24.0 C)
 10

 4

 0.9972995 g / mL    (0.9970479 g / mL  0.9972995 g / mL)
 10

 0.9971989 g / mL
Note: To do this calculation, you start with the whole number below your temperature and you add a
correction value to it. In this case, since our temperature was 24.4°C we use the density at 24.0°. Next, we
add our correction factor which has a fractional term in it. Since we are at 24.4°C, we use 4/10 as our
fraction. If it was 24.7, we would use 7/10. Finally, we subtract the two densities that are closest to our
value. Since our temperature is 24.4°C, the two closest values we have on the table are 25.0°C and 24.0°C.
If our value was 26.7°C, we would do D27.0°C – D26.0°C.
Temperature (°C) Density (g/mL) Temperature (°C) Density (g/mL)
10.0
0.9997026
24.0
0.9972995
12.0
0.9995004
25.0
0.9970479
13.0
0.9993801
26.0
0.9967867
14.0
0.9992474
27.0
0.9965162
15.0
0.9991026
28.0
0.9962365
16.0
0.9989460
29.0
0.9959478
17.0
0.9987779
30.0
0.9956502
18.0
0.9985986
31.0
0.9953440
19.0
0.9984082
32.0
0.9950292
20.0
0.9982071
33.0
0.9947060
21.0
0.9979955
34.0
0.9943745
22.0
0.9977735
35.0
0.9940349
23.0
0.9975415
40.0
0.9922187
Table 2: Density of Water at Various Temperatures
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Convert the mass of delivered water to milliliters using the density. To do this, use the mass you measured
and the density you just calculated to solve for the volume (Remember d=m/v and we are solving for v). Be
sure to use all provided significant figures for the density values.
5.
Repeat this process several more times until consistent results are obtained, then average similar values. The
calculated volume of water should be within 0.05 mL of the expected volume. This volume is the volume
your specific pipette is calibrated “to deliver” and will be used in the calculation of your unknown.
Density of an Unknown Liquid
1. Obtain an unknown liquid sample from your instructor and record its number and appearence in your
notebook. Weigh a clean, completely dry 50-mL Erlenmeyer flask with a well-fitting rubber stopper on an
analytical balance.
2. Remove the stopper, carefully pipette your unknown into the flask using your calibrated 10 mL pipet,
replace the stopper and weigh the flask and contents. Solve for density using the calibrated volume and
mass of liquid. Repeat this process 2 more times, taking the difference between the mass prior to each
subsequent addition and after. Each liquid mass should be approximately 10 g.
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Experiment 1 Mass, Volume, and Significant Figures: Pre-Lab Exercise
Name
Date
Laboratory Instructor
Lab Day & Time
1.
When rinsing your pipette to clean it, should you do your final rinse using water? If not, what should you
use?
2.
What is the difference between an analytical balance and a top loading balance?
3.
If a 0.050 L sample weighs 59.6 g, what is the sample density to the correct number of significant figures?
4.
The density of a sample usually ____________ (increases, decreases) as its temperature is raised.
5.
A student adds 10.00 mL of an unknown sample to a 25.0000 g flask three times to get masses of 36.0100
g, 47.0211 g, and 58.0305 g. Calculate the density of this sample. Take the mass difference for each
addition, i.e (47.0211 g – 36.0100 g) so the masses are each approximately 11 grams.
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Name ______________________________________
Lab Report_______/
Laboratory Instructor __________________________
Date_______________
Report – Experiment 1 Mass Volume and Significant Figures
Weighing Measurements of Known Metal
Mass of known metal strip trial 1___________________
Mass of known metal strip trial 2 __________________
Top loading balance/graduated cylinder
Mass of stir bar: _________________
Volume of stir bar:_________________
Density ____________________
Analytical balance/burette
Mass of stir bar: _________________
Volume of stir bar:_________________
Density ____________________
Volume Calibration of 10 mL Volumetric Pipet
Temperature of water _____________
Density of water ____________________________
Trial 1: Mass of water _____________
Calculated volume of water ___________________
Trial 2: Mass of water _____________
Calculated volume of water ___________________
Average volume __________________
Percent error _______________________________
Accepted volume is 10.00 mL, measured is the calibrated volume
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Liquid Unknown Results
Unknown #:________
Each mass should be close to 10 g
Trial 1:
Mass _____________
Density ___________________
Trial 2:
Mass _____________
Density ___________________
Trial 3:
Mass _____________
Density ___________________
Average density of liquid unknown __________________
20
General Chemistry I
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Experiment 1 Mass, Volume, and Significant Figures: Post-Lab Exercise
Name
Date
Laboratory Instructor
Lab Day & Time
1. If, unknown to you, your pipette was incorrectly calibrated so that it transferred less than 10.00 mL of
solution, the density you calculated for the liquid would tend to be ____________ (larger, smaller) than the
correct value. Explain.
2. An empty graduated cylinder has a mass of 42.91 g. When filled with 40.00 milliliters of an unknown
liquid, it has a mass of 103.26 g. Calculate the density of the unknown liquid.
3. A block has a length of 3.00 cm, a width of 4.00 cm, and a height of 7.00 cm. The mass of this solid object
was found to be 1.249 kg. Find the density of this substance in g/cm3.
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General Chemistry I
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4. A sample of solid gold has a density of 19.3 g/cm3 and a volume of 2.0 mL. What is the mass of the sample
of gold?
5. Use the figure below to answer the following questions.
a) Use the provided data to determine the density of this piece of rebar.
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6. Perform the following conversions. Make sure your answer has the correct number of significant digits.
a) 5.00 mm3 to cm3
b) 66.3 m3 to mL
c) 71.63 g/cm3 to mg/dm3
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General Chemistry I
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Experiment 2: Formula of a Hydrate
Learning Outcomes:
As you work through this exercise you will learn how to:

Determine the relative amount of water in a hydrated compound

Use heat to cause dehydration of a metal hydrate salt

Find the molar equivalence of anhydrous salt to moles of water
Note: You may wish to review your textbook before attempting this lab.
Introduction
Many ionic compounds incorporate a fixed number of water molecules into their crystal structures.
Water has a polar structure: it has partially charged positive and negative parts within each molecule. This
gives it a strong attraction toward ions. The ions in some salts attract and form strong bonds with water
molecules. These salts, when they have absorbed water, are called hydrates. Anhydrous salts are salts that can
form hydrates but which have had all the water driven off, usually by heat. Hydrated salts are characterized by
the number of moles of water molecules per mole of salt. The so-called water of hydration of nickel (II)
chloride (NiCl2) is six moles of H2O for every one mole of NiCl2. The hydration reaction is shown below. The
hydrate in this reaction is called nickel (II) chloride hexahydrate.
The formula of this hydrate shows the molar amount of water incorporated into the crystal matrix. For
most hydrates the amount of water included in the formula is only important when trying to measure molar
amounts of the salt. You need to know the true formula weight (molar mass) in order to measure out the mass
needed to give a certain number of moles. The chemical importance of the water of hydration is minimal since
it can be driven off by heat. From the reaction above, the molar mass of nickel (II) chloride hexahydrate
(NiCl2•6H2O) is 237.69 g/mol, not 129.60 g/mol. The molar mass of 129.60 g/mol corresponds to the
anhydrous salt (NiCl2).
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General Chemistry I
1-mono
6- hexa
Southern California University of Health Sciences
2-di
7- hepta
3-tri
4- tetra
5- penta
8- octa
9- nona
10- deca
Greek Prefixes: used to indicate the molar amount of water in hydrated salt.
Another example is Iron (III) Chloride, FeCl3, which crystallizes with six water molecules weakly
bound to the Fe3+ ion in an octahedral arrangement:
Figure 1. Iron (III) Chloride
Formulas for hydrates are written using a dot convention: a dot is used to separate the formula of the
salt from the formula of the water of hydration. A numerical coefficient gives the molar amount of water
included in the hydrate. Hydrates are named using prefixes for the word hydrate (at right). For example,
CuCl2•2H2O is copper (II) chloride dihydrate and CuSO4•5H2O is copper (II) sulfate pentahydrate. One
key point: the dot is not a multiplication sign. When calculating the molar mass you add the molar mass of
water (multiplied by the coefficient).
An everyday example of hydration is concrete. Concrete is made by mixing Portland cement with water
and aggregate materials. The aggregate materials are the gravel and sand that add strength to the final concrete.
The Portland cement is a mixture of calcium silicates, calcium aluminate, calcium aluminoferrite and gypsum.
All of these chemicals absorb water by hydration. This means that concrete does not ‘dry’ in a conventional
sense. Instead the water mixed with the concrete combines chemically with the materials in the cement and
the resulting hydrates form a strong matrix that holds the concrete together and makes it strong.
Another interesting example of the value of hydration is the incorporation of hydrated building materials
(such as concrete, gypsum wall board, and plaster). The building materials will not rise above the 100°C boiling
point of water until all of the water of hydration has been driven off. This can help keep damage to a minimum
until a fire can be put out. In the construction business this is known as passive fire protection.
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In this experiment, you will be instructed to determine the mass of a sample of an unknown hydrate by
“difference” using a pre-weighed crucible and cover as the container. The substance will be “dehydrated” by heat
and weighed again. The loss of mass represents the mass of water in the original sample, which may be expressed
as percentage of water of hydration, using the following relationship:
In this experiment you will be directed to “heat to constant mass.” Your purpose is to heat the
substance until all of the water is driven off. After the first heating, cooling, and weighing, you cannot tell if all
water has been removed or if some still remains. You therefore must repeat the heating, cooling, and weighing
procedure. If the same mass is reached after the second heating, you may assume that all water was removed
the first time. If mass is lost in the second heating, you may be sure that all water was not removed in the first
heating, and you are still unsure whether all water was driven off in the second heating. Another heating is
therefore required. The heating, cooling, and weighing sequence is repeated until three successive identical
weighings are recorded. Duplicate weighing within 0.02g is satisfactory for this experiment.
The first step to finding the formula for a hydrate is to record the mass of the hydrate. After heating the
hydrate, the mass is determined for the anhydrate that remains. The mass of the water that was present is
calculated by finding the difference between the mass of the hydrate and the mass of the anhydrate. The mass
of the water and the mass of the anhydrate are each converted to moles using their respective molar masses.
From this a whole number ratio can be determined.
Example 1
Data Table
Mass of Hydrate (CaCl2 • xH2O)
4.72 g
Mass of Anhydrate (CaCl2)
3.56 g
Mass of Water
1.16 g
A calcium chloride hydrate has a mass of 4.72g. After heating for several minutes the mass of the
anhydrate is found to be 3.56g. Use this information to determine the formula of the hydrate.
mass of hydrate – mass of anhydrate = mass of
1. Find the mass of the water driven off
water
4.72g – 3.56g = 1.16g
3.56𝑔 𝐶𝑎𝐶𝑙2 1 𝑚𝑜𝑙 𝐶𝑎𝐶𝑙2
×
1
110.98
2. Convert the mass of the anhydrate to moles
= 0.0321 𝑚𝑜𝑙 𝐶𝑎𝐶𝑙2
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General Chemistry I
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1.16𝑔 𝐻2 𝑂 1 𝑚𝑜𝑙 𝐻2 𝑂
×
= 0.0644 𝑚𝑜𝑙 𝐻2 𝑂
1
18.00 𝑔
3. Convert the mass of the water to moles
0.0644 𝑚𝑜𝑙 𝐻2 𝑂
0.0321 𝑚𝑜𝑙 𝐶𝑎𝐶𝑙2
4. Find the ratio of moles H2O to moles CaCl2
=
2 𝑚𝑜𝑙 𝐻2 𝑂
1 𝑚𝑜𝑙 𝐶𝑎𝐶𝑙2
Since the compound contains 2 moles of water for every 1 mole of anhydrate, the formula is
CaCl2
• 2H2O.
Example 2
A 140.5-g sample of NiSO4•XH2O is heated until there is no further decrease in mass. The mass of the
anhydrous salt is 77.500 g. Find the number of water molecules in the formula of this hydrate of nickel (II) sulfate
and calculate the percentage of water.
Reaction:
NiSO 4 •XH 2 O  NiSO 4 + XH 2 O
77.500 g NiSO4 ×
1mol NiSO4
154.76 g NiSO4
= 0.50078 mol NiSO4
(140.5 g NiSO4•XH2O – 77.500 g of NiSO4) = 63.0 g H2O
1 𝑚𝑜𝑙 𝐻 O
2O
63.0 g 𝐻2 O × 18.00 𝑔 𝐻2
= 3.50 mol 𝐻2 O
3.50 mol H O
2
Mole ratio = 0.50078 mol NiSO
= 7 moles of H2O per 1 mole of anhydrous salt
4
Therefore,
the
hydrate
formula
is
NiSO4
•
27
7H2O,
nickel
(II)
sulfate
heptahydrate.
General Chemistry I
Southern California University of Health Sciences
Experiment 2: Experimental Procedure
Observation of Heating a Metal Hydrate
1.
Add about a tablespoon of copper sulfate pentahydrate, CuSO4•5H2O into an evaporating dish.
2. Using a hot plate set low, heat the evaporating dish to remove the water. Increase the heat slowly as
the water is driven off.
3. Continue heating until the contents are completely dry (about 5 minutes).
4. Record observations.
5. After the evaporating dish has cooled to room temperature, add a few drops of water.
6. Record your observations.
Determintation of Waters of Hydration
1.
Obtain a metal ring clamp and place the clay triangle on top
2. Place clean crucible with the cover on the triangle. Light the Bunsen burner and heat for a couple
of minutes to make certain the container is thoroughly dry. Turn off the burner and cool the
container inside a desiccator for five minutes.
3. Weigh the empty crucible and record its mass. Place ~2 g of the solid unknown hydrate crystals in
the crucible and weigh again.
4. Place the crucible and cover on the triangle again. Heat the container and its contents vigorously
with the top slightly open (see figure 1) for 15 minutes. Remove the lid if water is collecting on the
inner surface.
5. Remove and allow the crucible to cool (to the touch) inside the desiccator for 10 minutes.
6. Weigh the crucible with its contents.
7. Return the crucible to the clay triangle and heat the “dry” sample again following steps 4-7.
Repeat this until a constant mass is obtained between weighings. Do not discard the sample until
checking with your instructor.
Figure 1: Unknown set up
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Experiment 2 Formula of a Hydrate: Pre-Lab Exercise
Name
Date
Laboratory Instructor
Lab Day & Time
1.
What is the purpose of the desiccator?
2.
What two components are present in a hydrate?
3.
Write the formulas for the following hydrates:
Barium phosphate dihydrate
Copper (II) sulfate pentahydrate
Cobalt (II) chloride hexahydrate
Magnesium sulfate heptahydrate
Nickel (II) sulfate hexahydrate
4.
A hydrate of magnesium sulfate has a mass of 13.52 g. This sample is heated until no water remains. The
MgSO4 anhydrate has a mass of 6.60 g. Find the formula and name of the hydrate
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Name _______________________________________
Lab Report _______/
Laboratory Instructor __________________________
Date_______________
Report Form: Experiment 2 Formula of a Hydrate
PART I
a) Write your observations of the copper hydrate when it is heated in the evaporating dish.
b) Write your observations of when water is added to the copper anhydrous salt.
PART II
Unknown Number______________
Color_______________
Mass of empty crucible
Mass of crucible and
hydrated unknown
Mass of hydrated
unknown
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After 1st Heating
After 2nd Heating
Mass after heating
(crucible and unknown)
Mass of water lost by
hydrated salt (Step 2 step 4)
Show calculations
Mass of anhydrous salt
Show calculations
Mass percentage of water
in hydrated unknown
(Show calculations)
Formula of anhydrous
salt (check with
instructor)
Moles of anhydrous salt
Show calculations
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After 3rd Heating
General Chemistry I
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Formula Mass of Water
Moles of water in
hydrated salt (show
calculations)
Ratio of moles of water
to moles of anhydrous
salt
Empirical formal of
hydrated salt
Name of Hydrated salt
SHOW ALL CALCULATIONS FOR ALL TRIALS
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Experiment 2 Formula of a Hydrate: Post-Lab Exercise
Name
Date
Laboratory Instructor
Lab Day & Time
1.
What is the percentage of water found in Na2S•10H2O?
2.
A 6.0 g sample of a hydrate of BaCl2 was heated, and only 4.4g of the anhydrous salt remained.
What percentage of water was in the hydrate?
3.
A 5.0 g sample of Cu(NO3)2•xH2O is heated, and 3.9 g of the anhydrous salt remains. What is the
value of x?
4.
A hydrate is determined to be 45.43% water and 54.57% CoCl2. Find the chemical formula and
name for this hydrate. (*Hint – assume that there is 100 g total of hydrate compound.)
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Experiment 3: Reaction Stoichiometry
Learning Outcomes:
As you work through this exercise you will learn how to:

Determine the mass of a product from an acid/base reaction

Determine molar relationships from chemical formula and reactions

Calculate therotical yield and percent error
Introduction:
When a chemical reaction occurs, reactants are converted to products. The quantitative relationship
between the reactants and products is termed stoichiometry. Because matter is not created or destroyed in a
chemical reaction, the number of atoms of each element should be balanced on each side of the reaction. To do
this, we use stoichiometric coefficients. Stoichiometric coefficients are the numbers placed in front of atoms,
ions, or molecules in a chemical equation. Stoichiometric coefficients establish the mole ratio between reactants
and products in a chemical reaction. For example, in this reaction:
N2(g) + 3 H2(g) → 2 NH3(g)
The mole ratios between the molecules are:
1 mol N2 : 3 mol H2 : 2 mol NH3
The mole ratio means that 1 mole of N2 will react with 3 moles of H2 to form 2 moles of NH3. The mole
ratio is used to compare the amount of reactants needed for the reaction to occur, the amount of products formed
in the reaction, or to convert between amounts of reactants and products. The stoichiometric coefficient and
resulting mol:mol ratios are essential to solving quantitative problems based on chemical reactions.
Example #1– How many moles of NH3(g) can be produced from 2.62 moles of N2(g)?
2.62 𝑚𝑜𝑙 𝑁2 ×
2 𝑚𝑜𝑙 𝑁𝐻3
1 𝑚𝑜𝑙 𝑁2
= 𝟓.𝟐𝟒 𝒎𝒐𝒍 𝑵𝑯𝟑
While the mol:mol ratio provides information on different reactants or products in a chemical reaction,
to calculate measureable data in the laboratory, the units must be in grams. To convert between moles and grams
the molar mass of the substance is used. For a general stoichiometric calculation, the figure below can be used
as a guide. In this example, ‘A’ represents the product or reactant we know something about and ‘B’ represents
the product or reactant we want to know something about.
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grams A
moles B
moles A
grams B
Figure 1 – The Relationship between Stoichiometric Coefficients
Example #2 – How many grams of NH3 can be produced from the reaction of 1.32 grams of H2?
𝑁2(𝑔) + 3𝐻2(𝑔) → 2𝑁𝐻3(𝑔)
Grams H2 → Moles H2 → Moles NH3 → Grams NH3
1.32 𝑔 𝐻2 ×
1 𝑚𝑜𝑙 𝐻2
2.02 𝑔 𝐻2
×
2 𝑚𝑜𝑙 𝑁𝐻3
3 𝑚𝑜𝑙 𝐻2
×
17.04 𝑔 𝑁𝐻3
1 𝑚𝑜𝑙 𝑁𝐻3
= 𝟕. 𝟒𝟐 𝒈 𝑵𝑯𝟑
In most chemical reactions, one of the reactants will run out first. For example, if you wanted to make
peanut butter and jelly sandwiches, you may go to the store and buy a jar of peanut butter, a jar of jelly and a
loaf of bread. To assemble the sandwiches (the product), you would use the following ‘reaction’:
2 𝐵𝑟𝑒𝑎𝑑 𝑆𝑙𝑖𝑐𝑒𝑠 + 1 𝑆𝑐𝑜𝑜𝑝 𝑃𝑒𝑎𝑛𝑢𝑡 𝐵𝑢𝑡𝑡𝑒𝑟 + 1 𝑆𝑐𝑜𝑜𝑝 𝐽𝑒𝑙𝑙𝑦 → 1 𝑃𝑒𝑎𝑛𝑢𝑡 𝐵𝑢𝑡𝑡𝑒𝑟 & 𝐽𝑒𝑙𝑙𝑦 𝑆𝑎𝑛𝑑𝑤𝑖𝑐ℎ
How could you predict how many sandwiches you could make? It’s unlikely that the bread, the peanut
butter, and the jelly will run out at exactly the same time. The component that runs out first determines how
many sandwiches you can make. If you run out of bread, you can no longer make sandwiches. The same is true
in chemical reactions. In chemistry, the reactant that runs out first is called the limiting reactant. Because the
limiting reactant determines the maximum amount of product that can form, the amount of product formed when
all of the limiting reactant is converted to product is the theoretical yield. In our sandwich example, the bread is
the limiting reactant. If a loaf consists of 36 slices of bread, the theoretical yield would be 18 peanut butter and
jelly sandwiches. Any reactant that is not the limiting reactant is referred to as an excess reactant. The limiting
and excess reactants for a chemical reaction can be identified using stoichiometry. The maximum amount of
product expected from each reactant is calculated. The reactant yielding the least amount of product is the
limiting reactant. The yield produced by the limiting reactant is the theoretical yield, the maximum amount of
product that could be formed from a reaction.
Example #3 – Upon reaction of 2.42 g of nitrogen gas and 1.44 g of hydrogen gas, 0.165 grams of ammonia
was formed. What is the limiting reactant and theoretical yield?
2.42 𝑔 𝑁2 ×
1 𝑚𝑜𝑙 𝑁2
2 𝑚𝑜𝑙 𝑁𝐻3
×
= 0.0173 𝑚𝑜𝑙 𝑁𝐻3
28.02 𝑔 𝑁2
1 𝑚𝑜𝑙 𝑁2
1.44 𝑔 𝐻2 ×
1 𝑚𝑜𝑙 𝐻2
2 𝑚𝑜𝑙 𝑁𝐻3
×
= 0.475 𝑚𝑜𝑙 𝑁𝐻3
2.02 𝑔 𝐻2
3 𝑚𝑜𝑙 𝐻2
The reactant that produces the smallest amount of moles of product (NH3) is N2. The amount of product
35
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formed from the limiting reactant is the theoretical yield. Therefore, the theoretical yield is:
0.0173 𝑚𝑜𝑙 𝑁𝐻3 ×
17.04 𝑔 𝑁𝐻3
= 0.295 𝑔 𝑁𝐻3
1 𝑚𝑜𝑙 𝑁𝐻3
Notice that the theoretical yield (0.295 g NH3) is larger than the amount of product obtained (0.165 g NH3).
Although it is possible to form product from all of the limiting reactant, chemical reactions do not
always go to completion. Often some of the reactant remains at the end of the reaction, or some of the reactant is
consumed by other chemical reactions. To describe how efficiently the limiting reactant is converted to product
in a reaction, we use percent yield. The percent yield is:
% 𝑦𝑖𝑒𝑙𝑑 =
𝑎𝑐𝑡𝑢𝑎𝑙 𝑦𝑖𝑒𝑙𝑑
× 100%
𝑡ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝑦𝑖𝑒𝑙𝑑
The percent yield for a chemical reaction is typically high when the amount of experimental errors are
low. If a yield of over 100% is found, the product most likely contains a contaminant, such as water.
Example #4 – The reaction in Example #3 yielded 0.165 g of NH3 in the laboratory. In a laboratory experiment,
the actual yield will be the mass of product formed from the reaction. What is the percent yield for the reaction?
% 𝑦𝑖𝑒𝑙𝑑 =
0.165 𝑔 𝑁𝐻3
= 55.9%
0.295 𝑔 𝑁𝐻3
In today’s experiment, the stoichiometry of the following double-displacement, gas-forming, reactions will be
examined.
NaHCO3 (s) + HCl(aq) → NaCl (aq) + CO2 (g) + H2O(l)
Na2CO3 (s) + 2 HCl → 2NaCl (aq) + CO2 + H2O (l)
In the first reaction, there is a 1:1 mol ratio between NaHCO3 and NaCl. In the second reaction, there is
a 1:2 mol ratio between Na2CO3 and NaCl. You will perform both reactions, and determine the amount of NaCl
formed by each reaction. In these reactions, the reactants NaHCO3 and Na2CO3 are the limiting reactants, as the
HCl is used in excess. The solid NaCl remaining after the experiment is the actual yield. You will use the
amount of reactant and product in each reaction to calculate the mol:mol ratios. Then, you will use
stoichiometry to determine the theoretical yield for both of the reactions, and then compare the theoretical yields
with the actual yields obtained from the reactions. Finally, the percent yield of both reactions will be
determined.
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Experiment 3: Experimental Procedure
1. Obtain an evaporating dish and glass stirring rod. Make sure that both are clean and dry.
2. Weigh the empty, dry evaporating dish on the top-loading balance and record the mass.
3. Weigh out 0.3-0.4 g of NaHCO3 using a weigh boat. Transfer the NaHCO3 to the evaporating dish.
Reweigh the evaporating dish + NaHCO3 and record the mass.
4. Using the small bottles of 6M HCl found in the fume-hood, add the HCl very slowly (drop by drop) to
the NaHCO3 in the evaporating dish. After every 4-5 drops of HCl, carefully mix the reactants with a
stirring rod. You will know that the reaction is occurring, as the mixture will be bubbling.
5. Continue adding HCl until the bubbling has stopped and all of the NaHCO3 has dissolved. Be sure to
rinse the stir rod as well to remove any residue.
6. Using a hot plate set to 4, heat the solution to remove the water and excess HCl from the product.
(Using a higher heat setting will result in ‘popping’, which will cause product loss.)
7. Continue heating until the contents are completely dry (about twenty minutes).
8. Allow the evaporating dish to cool to room temperature for ten minutes.
9. Record the mass of the evaporating dish and the product of the reaction (NaCl).
10. The waste from the reaction can go down the sink. Clean your evaporating dish and dry it.
11. Repeat steps 1-10 using Na2CO3.
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Experiment 3 Reaction Stoichiometry: Pre-Lab Exercise
Name
Date
Laboratory Instructor
Lab Day & Time_____________________
1. What is the highest heat setting you should use for the hot plates? Why?
2. How do you know the reaction is occurring? How do you know when the reaction is finished?
3. What is the name of the product at the end of both reactions? Is the mass of that you measure considered
the actual yield or the theoretical yield?
4. What is the percent yield of a reaction that produces 12.5 g CF2Cl2 from 32.9 g of CCl4 and excess HF?
The reaction is as follows:
CCl4 (l) + 2 HF(aq) → CF2Cl2 (g) + 2 HCl(aq)
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Name _______________________________________
Lab Report _______/
Laboratory Instructor __________________________
Date_______________
Report Form: Experiment 3 Reaction Stoichiometry
Sodium Bicarbonate NaHCO3 Reaction Data
Mass of evaporating dish
Mass of evaporating dish + NaHCO3
Mass of NaHCO3
Mass of evaporating dish + NaCl
Mass of NaCl (actual yield)
Sodium Carbonate Na2CO3 Reaction Data
Mass of evaporating dish
Mass of evaporating dish + Na2CO3
Mass of Na2CO3
Mass of evaporating dish + NaCl
Mass of NaCl (actual yield)
Data Analysis:
Please provide sample calculations for each step of the calculation.
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Part A: NaHCO3 Reaction
1. Convert the mass of NaHCO3 used and mass of NaCl obtained to moles.
a. NaHCO3
b. NaCl
2. Determine the theoretical yield NaCl in the NaHCO3 reaction. Show your work.
3. Determine the percent yield for the NaHCO3 reaction. Show your work.
4. Divide your answers from Question 1 by the lower mole value to determine the simplest mole-to-mole ratio
between NaHCO3 and NaCl. Calculate your answer and then round them to the nearest whole number.
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Part B: Na2CO3 Reaction:
1. Convert the mass of Na2CO3 used and mass of NaCl obtained to moles.
a. Na2CO3
b. NaCl
3. Determine the theoretical yield and percent yield of NaCl in the Na2CO3 reaction. Show your work for each
step.
4. Divide your answers from Question 1 by the lower mole value to determine the simplest mole-to-mole ratio
between Na2CO3 and NaCl. Calculate your answer and then round them to the nearest whole number.
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Experiment 3 Reaction Stoichiometry: Post-Lab Exercise
Name
Date
Laboratory Instructor
Lab Day & Time
1. A student performed a reaction between NaHCO3 and HCl and obtained a percent yield of 108%. What
experimental error(s) could have cause this to occur?
2. A 1.274 g sample of copper (II) sulfate was dissolved in water and allowed to react with excess zinc
metal. The reaction produced 0.392 g of copper metal.
a) Write the balanced equation for this reaction, including all phases.
b) What is the limiting reactant in this reaction?
c) What is the theoretical yield for this reaction?
d) What is the percent yield for this reaction?
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Experiment 4: Beer’s Law Spectrophotometric Determination of
Copper
Learning Outcomes:
As you work through this exercise you will learn how to:

Dissolve a sample and dilute to a specific volume.

Prepare a dilution series covering a specific concentration range.

Perform absorbance measurements using a spectrophotometer.

Construct a calibration curve and interpolate using the equation for this curve.
Note: Review dilution, molarity calculations, and Beer’s law in your lecture notes and/or textbook
(Chapter 4) before attempting this lab.
Introduction
.
Since 1983 the composition of the U.S. one-cent coin has changed from being primarily copper to a coin
that has a zinc core with a thin copper coating. This change reflected an increase in the price of copper metal.
In this experiment your team will determine the mass percent of copper in a post-1983 penny. The results will
be correlated to the wear of the penny and compared with the expected composition. In order to measure the
amount of copper present, the penny will first be dissolved in 6.0 M nitric acid to produce Cu(NO3)2 (aq). The
reaction is as follows:
Cu (s) + 4HNO3 (aq) → Cu(NO3)2 (aq) + NO2 (g) + H2O (l)
Safety Note: The reddish-brown nitrogen dioxide gas is toxic, so the reaction must be done in a fume
A blue-green aqueous solution is formed due to the presence of the Cu (II) ion. The intensity of the bluegreen color is proportional to the concentration of the Cu (II) ion. Since the solution is colored, it means that the
solution significantly absorbs some wavelengths of visible light; in this case, the red end of the visible spectrum.
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If there was no significant absorption of visible light, the solution would be colorless. The zinc and nitrate ions
are colorless and so do not interfere with the measurement of Cu (II) by spectrophotometry.
Experimentally, absorbance measurements are made using a spectrophotometer. This instrument consists of
a light source, a diffraction grating to select a narrow wavelength region of light from the source, and a device to
detect the light passing through the sample. For monitoring the absorbance of the Cu (II) ion in solution, light
with a wavelength (λ) of 800 nm will be used. The extent to which light is absorbed by a substance in solution
depends on the concentration of the substance – the higher the concentration, the larger the absorbance. Imagine
directing a beam of light through a solution held in a tube, as shown in Figure 1.
Figure 1. A beam of light directed through a sample is affected by the thickness of the sample, the concentration of the absorber,
and its molar absorptivity
When an absorbing solution is placed in the tube, the intensity of the light passing through decreases because the
sample absorbs some of the light, thus reducing the amount that exits the tube. As the concentration of the
sample increases, even more light is absorbed leading to a corresponding decrease in the amount of transmitted
light. Mathematically
𝐴 = − log
𝐼
𝐼0
where I0 and I are the intensities of incident (starting intensity) and transmitted light, respectively. This
relationship is described by Beer’s law
𝐴 = ∈ 𝑙𝑐
Where A is called “absorbance,” “absorbance units,” or “AU,” and has no actual units; l is the thickness of
thes ample (called the pathlength) through which the light passes (in cm); c is the concentration of the lightabsorbing substance (M); and Ɛ is the molar absorptivity, a measure of how well a substance absorbs at a given
wavelength (M-1 cm-1). Notice that Beer’s law (A = Ɛ・l ・c) has the same mathematical form as the equation
for a straight line (y = m・x + b) if b = 0. This means that a plot of absorbance (A) versus concentration (c)
should be linear with slope (Ɛ ・l ) and pass through the origin (corresponding to b = 0).
In this experiment, a series of standard solutions with known copper (II) nitrate concentrations will be
prepared by diluting a stock solution with DI water (note: since the solutions will be prepared with DI water as
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the solvent, the inside of the flasks do not necessarily need to be dry). Once the solutions are made, the
absorbance will be measured using the same spectrophotometer that will be used to measure the dissolved penny
solution. A calibration curve (graph of absorbance vs. concentration) will be plotted using these data points and
fit with a linear trendline. This is the purpose of a standard series – it gives us a mathematical relationship
between two variables which can be used to determine an unknown value. In this case, the equation of this line
reveals how the absorbance varies with Cu (II) concentration in your experiment and can be used to determine
the copper concentration in an unknown (e.g., penny) solution. Using the absorbance value measured for the
penny solution, the trendline equation can be used to calculate the concentration that corresponds to this
measured absorbance by plugging the absorbance into the y of the equation. (Make sure you understand this
concept… if it’s unclear, ask your instructor!)
Knowing the concentration (found from using the trendline of the calibration curve) and volume of the
solution, the moles of copper (and thus grams of copper) can be determined. As a reminder, concentration is
typically given in the units of Molarity which is moles/Liters. If you wanted to calculate mass% from your
concentration, you would do a calculation similar to the following:
𝑣𝑜𝑙𝑢𝑚𝑒 𝑖𝑛 𝐿 𝑜𝑓 𝐶𝑢 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 𝑥
𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝐶𝑢
𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝑜𝑓 𝐶𝑢
𝑥
= 𝑚𝑎𝑠𝑠 𝑜𝑓 𝐶𝑢
𝐿 𝑜𝑓 𝐶𝑢 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛
𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝐶𝑢
The mass percent of copper in the penny is then obtained by comparing the mass of the copper to the mass
of the penny.
𝑚𝑎𝑠𝑠 𝑜𝑓 𝐶𝑢
𝑥 100 = 𝑚𝑎𝑠𝑠 𝑝𝑒𝑟𝑐𝑒𝑛𝑡
𝑚𝑎𝑠𝑠 𝑜𝑓 𝑝𝑒𝑛𝑛𝑦
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Experiment 4: Experimental Procedure
Caution: Dissolve the penny in nitric acid in the fume hood to avoid release of toxic NO 2 gas into the
laboratory. Nitric acid is corrosive, if it contacts your skin immediately rinse the area with large amounts
of cold water and notify your instructor
Quick Look
1. Dissolve a penny in nitric acid. (Note: use your own post-1983 pennies from home)
2. While the penny is dissolving, create a set of standard solutions of Cu(NO3)2 and calculate their
concentrations (in M or moles/Liters) using M1V1=M2V2
3. Measure the absorbance of the solutions at 800 nm using a spectrophotometer
4. Once the penny has dissolved, dilute it in a 50.00 mL volumetric flask and measure the absorbance of
this solution.
5. Make a Y vs. X linear chart of Absorbance (AU) vs. Concentration (M) from your standard solutions.
6. Calculate the penny copper concentration by plugging the absorbance into Y and solving for X. Convert
this concentration to mass % using the calculations found in the introduction section.
Preparation of an Unknown Solution
1. Select a post-1983 penny. Note the date and surface appearance of the penny. Clean the surface with
soap and water, then dry thoroughly. Weigh the dry penny using an analytical balance and record the
mass (if the mass is decreasing slowly the penny is not dry).
2. In the fume hood, place the penny in a clean, dry 150mL beaker and slowly, carefully add about 30 mL
of 6.0 M HNO3 to cover the penny. After a short time, the reaction becomes very vigorous, releasing
heat and brown fumes (NO2(g)). Be careful of splattering acid – if any gets on your skin or clothing,
wash quickly with copious amounts of cold water and alert the instructor.
Preparation of a Series of Standard Solutions from 0.100 M Cu(NO3)2 Stock Solution
1. Obtain ~ 15mL of 0.100 M Cu(NO3)2 in a 50mL Erlenmeyer flask
2. (This procedure can be carried out while the penny is dissolving.) Using a graduated pipet obtain the
amount necessary of 0.100 M Cu(NO3)2 to make a 10.00 mL solution of 0.020M in a volumetric flask.
Use DI water to make your solutions, Use M1V1=M2V2 equation. Show your calculations to your
instructor.
3. Transfer your solution to a 50mL Erlenmeyer flask and label it with the proper concentration.
4. Repeat steps 2 and 3 to make 10mL of 0.040 M, 0.060 M, and 0.080 M standard solutions of Cu(II)
from the stock solution. Be sure to label each container.
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Finish preparation of unknown penny solution
1. Once the penny is completely dissolved, use a glass funnel and stirring rod to transfer the solution to a
clean 50-mL volumetric flask. Use a few milliliters of deionized water (DI) to rinse the beaker and add
this to the volumetric flask. Carefully dilute to the mark on the flask with deionized water. Stopper the
flask and invert several times to mix the solution thoroughly.
Making Absorbance Measurements
1. Set the wavelength of the spectrophotometer to 800 nm and zero the instrument with DI water as a
‘blank’ in the cuvette.
2. Starting with the least concentrated solution, measure and record the absorbance of each of the four
dilute standard solutions.
3. Measure and record the absorbance of your dissolved penny solution. Use this absorbance, together with
your graph of absorbance versus concentration for the standard solutions, to calculate the concentration
of copper in your penny solution. Refer to the introduction to figure out how to do this.
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Experiment 4 Beer’: Pre-Lab Exercise
Name
Date
Laboratory Instructor
Lab Day & Time
1. Why are we preparing a standard solution series? What will it be used to calculate?
2. How many milliliters of a 1.30 M NaCl solution are needed to prepare 0.50 L of 0.10 M NaCl?
3. Why are we using a volumetric flask in this experiment rather then another type of glassware?
4. When preparing aqueous solutions in a volumetric flask, does it matter if there are a few droplets of
water in the bottom of the flask or does the flask need to be completely dry? Explain.
5. Recall Beer’s Law.
In your experiment a calibration plot of ________________ (y) versus
________________ (x) will be constructed. For the unknown penny solution, ________________
will be measured and ________________ will be calculated using the calibration curve.
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Name _______________________________________
Lab Report _______/
Laboratory Instructor __________________________
Date_______________
Report Form: Experiment 4 Beer’s Law
Dilution Series for Standard Solutions
Solution
Initial volume
Initial
Concentration
Stock
1
2
3
4
Sample Dilution calculation
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Final Volume
Final
Concentration
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Data
Solution
Concentration
Absorbance
1
2
3
4
Penny
—
Calibration Curve (attach EXCEL plot)
Use EXCEL to prepare a calibration curve of absorbance vs. concentration for the 4 dilute copper standard
solutions. Include a point with x = 0, y = 0. Fit these data with a linear trendline and obtain the equation and R 2
(display on the plot and note them here). Report R2 to four decimal places.
Equation_______________________________________________
R2 ____________________________________________________
Determination of Copper in Penny
Show calculations
Date of Penny___________________________________________
Mass of Penny___________________________________________
Concentration of Cu(II) in penny solution_____________________
Moles of Cu(II) in penny solution___________________________
Mass % Cu in penny_____________________________________
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Experiment 5: Unknown Solutions
Learning Outcomes:
As you work through this exercise you will learn how to:

Make careful observations of chemical reactions.

Develop logical testing procedures for performing qualitative analyses.

Describe solubility rules and fundamental acid/base chemistry.
Note: You may want to review your textbook before attempting this experiment.
Introduction
One of the central roles of a chemist involves the analysis of materials. In some instances the analysis is
quantitative, that is, a determination is made of how much of a particular substance is present. In other cases the
analysis is qualitative and a determination of what components are present is made. In this experiment you will
perform a qualitative analysis on a set of unknown solutions. The identification process will involve
observations of solution properties such as pH, color, odor, and the results of mixing two solutions, such as the
formation of a precipitate – the formation of an insoluble solid in your solution. However, before you are able to
assess a set of solutions you must understand some basic solubility rules.
Table 4 shows the list of common ions and their solubilities. The top half of the table shows the ions that
are almost always soluble (dissolving in water and becoming homogeneous solutions) with their notable
exceptions. For example, compounds containing Cl- are always soluble unless they contain Ag+, Hg2+, or Pb2+.
Understanding this, we could predict that KCl would be soluble in water but PbCl 2 would be insoluble and
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precipitate (form a solid) out of solution. This can be shown in a balanced or net ionic equation:
𝐾𝐶𝑙 (𝑎𝑞) + 𝑃𝑏(𝑁𝑂3 )2 (𝑎𝑞) → 𝑃𝑏𝐶𝑙2 (𝑠) + 2𝐾𝑁𝑂3
𝐶𝑙 − (𝑎𝑞) + 𝑃𝑏 2+ → 𝑃𝑏𝐶𝑙2 (𝑠)
Note in the net ionic equation, we only include the ions that are reacting and changing their state of matter; the
ones that stay in the same state of matter (aq) are not included and are called spectator ions. The bottom half of
the table works in the opposite fashion – the anions listed will always precipitate unless matched with one of the
exceptions listed.
Important exceptions
NO3-
None
CH3CO2-
None
Cl-, Br-, I-
Compounds of Ag+, Hg2+ and Pb2+
Soluble compounds
SO42-
Compounds containing NH4+, alkali metal cations, Ca2+,
2-
S
Insoluble compounds
Compounds of Sr2+, Ba2+, Hg2+, and Pb2+
Sr2+ and Ba2+
CO32-
Compounds containing NH4+ and alkali metal cations
PO43-
Compounds containing NH4+ and alkali metal cations
–
OH
Compounds containing alkali metal cations, NH4+, Ca2+,
Sr2+ and Ba2+
Table 1: Solubility guidelines for common ionic compounds in water
As an example to understand how we would use these solubility rules and other qualitative analyses,
suppose you know that your set of unknown solutions consists of the following solutions (in no particular
order):
Fe(NO3)2 (aq)
KOH (aq)
Cu(NO3)2 (aq)
H3PO4 (aq)
H2SO4 (aq)
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The first thing we could do would be to observe the pH of each solution. The pH of a solution is a measure
of the acidity or basicity (ability to donate or receive H+) and can be roughly measured by adding a drop of the
solution to universal pH indicator paper. A red result will indicate an acid (pH < 7.0) and a blue result will indicate a basic solution (pH > 7.0) with stronger acids and bases being darker shades of red and blue. Using this
pH test on our solutions we could see that there are two acids (H2SO4 and H3PO4), one base (KOH), and two
relatively neutral solutions (these solutions may show up as a slightly acidic orange but we will still consider
them neutral). With the pH tests, we have already confirmed the identity of the KOH and have limited what the
other four solutions could be. The next test we could do would be to observe the color of each solution. As the
Cu2+ ion in solution has a distinctive blue color (we learned this in Experiment #3) the identification of the
Cu(NO3)2 solution can be made on the basis of color. As the pH test gave us two relatively neutral solutions and
one of them is a blue solution, we can determine the neutral colorless solution is Fe(NO 3)2. This leaves us with
two solutions that are acidic and colorless and we may need to rely on other tests. In order to identify the
remaining solutions, you may want to observe what happens when pairs of solutions are mixed. Note that if
Cu(NO3)2 solution is mixed with the other four solutions, a precipitate of an insoluble copper salt is expected
with KOH and H3PO4, but not with Fe(NO3)2 or H2SO4 (see Table 6.1). We can see that Cu2+ is not an exception
with the soluble NO3- or SO42- ions and thus won’t precipitate with either of them. It also isn’t an exception to
the insoluble OH- or PO43- ions and thus will precipitate with both of them. Thus, the identity of the H3PO4
solution can be made and differentiated from the other acid – H2SO4.
For the experiment we will be doing today, you will have one of the following three solutions sets, the
possible sets of unknown solutions in this experiment are:
Set A:
AgNO3(aq), Mn(NO3)2(aq), Ba(NO3)2(aq), HCl(aq), NaOH(aq)
Set B:
AgNO3(aq), Ba(NO3)2(aq), HCl(aq), H2SO4(aq), NaOH(aq)
Set C:
AgNO3(aq), Pb(NO3)2(aq), H2SO4(aq), NH4OH(aq), H2O(aq)
The first thing you will need to do is to determine which set you have and unlike our example none of the
solutions will be colored. You will have to determine what test or tests that could be used that would
differentiate the sets from each other. Once you have your set, you must figure out which solution is which but
you don’t have to worry about the possibility of other solutions from other sets in your set.
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Experiment 5: Experimental Procedure
Caution: When taking samples for testing on the spot-plate, be very careful not to contaminate any of the
solutions; doing so can adversely influence your results. Use a separate Pasteur pipette for each solution.
Identifaction of Unknown Solution
1. Record your unknown set number.
2. Determine the pH of each solution; identify which are acidic, basic, or neutral. Confirm with your
instructor.
3. Make a determination about which of the 3 solutions sets you may have
4. Using a spot plate and a pastuer pipet for each solution, react solutions A-E with each other in sequence,
record the observations. Take detailed notes on color, precipitation, and any chemical reactions that
may occur.
5. Using the solubility rules determine the identity of each unknown.
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Experiment 5 Analysis of Unknown Solutions: Pre-Lab Exercise
Name
Date
Laboratory Instructor
Lab Day & Time
1.
a. Write a balanced equation for the reaction that occurs when an aqueous solution of iron (II)
chloride is mixed with an aqueous solution of potassium hydroxide.
b. Write a net ionic equation for the above reaction.
2.
What simple test would you use to identify the possible presence of HCl(aq)? What result do you
expect if the solution is HCl(aq)?
3.
What preliminary test results are you looking for to identify conclusively your unknown solutions
as Set B? Explain. Hint. Mixing is not necessary
4.
Give the formula and name of the acids and bases we will be working with in today’s experiments.
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Name _______________________________________
Lab Report _______/
Laboratory Instructor __________________________
Date_______________
Report Form: Experiment 5 Analysis of Unknown Solutions
Unknown number for set______________
Solutions present in set: ___________________________________________________
Test #
Specific Test Performed
Result
1
2
3
4
5
6
7
8
9
10
Discuss how these tests confirmed or ruled out the identity of each component.
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(in no particular order)
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Indicate the specific identity of each solution:
Solution #
Identity
Balanced Equations
Provide balanced net ionic equations for all mixtures in which a chemical reaction occurred.
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Experiment 6: Calorimetry and Specific Heat Capacity
Learning Outcomes:
As you work through this exercise you will learn how to:

Measure the mass and temperature of water in a calorimeter

Calculate the change in temperature caused by adding a heated substance to water

Calculate the specific heat of an unknown metal
Introduction
Chemists identify substances on the basis of their chemical and physical properties. One physical
property is the amount of energy each gram of a substance will absorb. This property can be measured quite
accurately and is called specific heat (Cp). Specific heat is the amount of energy, measured in joules, needed to
raise the temperature of one gram of a substance one degree Celsius. Often applied to metallic elements, specific
heat can be used as a basis for comparing how different substances absorb and transfer energy. To measure
specific heat in the laboratory, a calorimeter of some kind must be used. A calorimeter is a well-insulated
container that can measure energy changes. The calorimeter is insulated to reduce the loss or gain of energy to
or from the surroundings. Energy always flows from an object of higher temperature to an object of lower
temperature. The heat gained by the cooler substance equals the heat lost by the warmer substance, if we assume
no loss of heat to the surrounding environment. Therefore, Heat lost by substance = heat gained by water
Substance
Specific Heat (in J/g∙ ℃)
Water
4.184 (1.00 cal/g∙ ℃)
Wood
1.76
Aluminum
0.902
Glass
0.84
Iron
0.451
Nickel
0.444
Copper
0.385
Zinc
0.385
Tin
0.222
Lead
0.129
Table 1. Specific Heats of Some Common Substances
Compare the heat capacities of concrete and wood. Because the specific heat of wood is twice as great
as that of concrete, it takes about twice as much heat to raise the temperature of wood versus that of concrete.
This can be verified by comparing the feel of walking on concrete versus walking on wood on a hot, sunny day
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with bare feet. The concrete feels hotter. The sun gives off energy which is absorbed by the concrete and the
wood equally. However, because the wood has a greater specific heat value, it is able to absorb more heat before
its temperature rises, and thus does not feel as hot as the concrete.
General Rule #1 – The greater the specific heat value, the less the temperature will rise when a
given heat energy is absorbed.
Not only does the specific heat value describe how much heat may be absorbed by a substance before its
temperature rises, it also describes the ability of a substance to deliver heat to a cooler object.
General Rule #2 – As the specific heat value decreases, the ability to deliver heat to a cooler object
increases.
For example, imagine holding two hot pieces of metal — X (Cp = 2 J/g∙ ℃) and Y (Cp = 3 J/g∙ ℃ ). If
the hot piece of metal X was held in one hand and the hot piece of metal Y in the other hand, the hand holding
the metal X would get hotter. Because metal X has a specific heat less than metal Y, the metal X sample
transfers heat to a cooler object (your hand) more readily.
Why do different materials possess different specific heat values?
One reason for the variation is that each substance is made up of atoms that have different masses. The
mass of each copper atom is larger than the mass of each aluminum atom, for example. Therefore, a given mass
(such as 58 g of copper) has fewer atoms than the same mass of aluminum. When heat is added to 58 g of
copper, fewer atoms need to be put in motion (remember temperature is related to kinetic energy). Thus, less
heat is needed to increase the kinetic energy of the atoms in the sample, and raise the temperature by 1℃. As a
result, the specific heat value for copper is lower than the specific heat of aluminum. Notice that copper and zinc
have identical specific heat values. This is due to the similar mass of the atoms.
General Rule #3 – The larger the metal atom, the lower its specific heat value.
How is the specific heat of a material determined?
Heat transfer or heat flow always occurs in one direction – from a region of higher temperature to a
region of lower temperature – until some final equilibrium temperature is reached. In this experiment, heat is
transferred from a hot metal sample to a colder water sample. Because each metal has a different specific heat,
each metal will cause the temperature of the water to increase to a different extent. The transfer of energy can be
detected by measuring the resulting temperature change, ∆T, calculated by taking the final temperature minus
the initial temperature, as shown in Equation 1.
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∆T = final temperature – initial temperature = Tf – Ti (Equation 1)
For the hotter object in this scenario (the metal), the amount of heat (q) delivered by the metal (qmetal)
is equal to the mass of the metal (mmetal) multiplied by the specific heat of the metal (Cp metal) multiplied by
the change in temperate of the metal (∆Tmetal). This relationship is given by Equation 2.
qmetal = (mmetal )(Cp metal )(∆Tmetal )
(Equation 2)
For the cooler object in this scenario (the water), the amount of heat absorbed by the water (qwater) is
equal to the mass of the water (mwater) multiplied by the specific heat of the water (Cp water), multiplied by
the temperature change of the water (∆Twater). This relationship is given by Equation 3.
qwater = (mwater )(Cp water )( ∆Twater )
(Equation 3)
By convention, the sign of q is a signal showing the direction of heat transfer. When heat is transferred
out of a material, the sign of q is negative. Conversely, when heat is absorbed by a material, q is positive. The
signs of q, along with the necessary associated temperature changes, are summarized in Table 2.
Direction of Heat
Transfer
Sign of ∆T
Sign of q
Heat is absorbed
(transferred into a
+
+
–
–
Heatmaterial)
is delivered
(transferred out of
a material)
Change in
Temperature of
Material
Temperature
increases
Temperature
decreases
Table 2
According to the Law of Conservation of Energy, the heat delivered by the heated metal, qmetal, must
be equal to the heat absorbed by the water, qwater, and its surroundings. Incorporating the sign convention given
in Table 2 gives Equations 4 and 5.
qmetal = – qwater
(Equation 4)
(mmetal )(Cp metal )(ATmetal ) = – (mwater )(Cp water )( ∆Twater )
(Equation 5)
In this laboratory activity, Equation 5 is used to calculate the specific heat of a heated metal added to a
water sample. For calculation purposes, it is important to realize that when the metal is added to the water, the
final temperature of both materials will be the same. The calculated specific heat value will then be compared to
the known specific heat value given in Table 1.
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To make accurate measurements of the heat transfer and to prevent heat loss to the surroundings, an
insulating device known as a calorimeter is used. A calorimeter is a device used to measure heat flow, where the
heat given off by a material is absorbed by the calorimeter and its contents (often water or another material of
known heat capacity).
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Experiment 6: Experimental Procedure
1. Fill a large beaker about half-full with DI water. Heat the water to a boil using a hot plate. The hot
plate setting should not be above 5.
2. Weigh between 30-40 grams of the unknown metal provided by your instructor. Be sure to use a
weighing dish. Record the exact mass and color of the metal in your notebook.
3. Put the metal in a large test tube, then place the test tube in the boiling water (9 5-100°C) bath for
15 minutes. Be sure the test tube is placed in the water such that the area containing the metal is
completely submerged, but that no water enters the test tube.
4. While waiting for the metal to heat up, weigh the empty calorimeter and record its mass.
5. Pour 50 mL of DI water into the calorimeter and weigh the calorimeter again. Record the mass.
6. Measure the temperature of the water in the calorimeter in degrees Celsius.
7. Determine the temperature of the metal sample. To do this, measure the te mperature of the
boiling water bath (This should be done after the metal has been in the boiling water for 15
minutes). An assumption is made that the temperature of the metal is equal to the temperature
of the water bath.
8. Using a test tube holder, lift the test tube containing the heated metal from the boiling water bath and
quickly, yet carefully, pour the metal into the calorimeter. Make sure no hot water from the outside of
the test tube drips into the calorimeter.
9. Gently stir the water and metal shot in the calorimeter with the accompanying stirring rod.
10. Measure and record the temperature of the water every 30 seconds until the temperature
remains constant for two consecutive readings (this can take up to 3 minutes).
11. Drain the water out of the calorimeter and pour the metal onto paper towels. Pat the metal dry
thoroughly. Dry the calorimeter. Repeat steps 3-10 for 2 additional trials.
62
General Chemistry I
Southern California University of Health Sciences
Experiment 6 Calorimtry and Specific Heat Capacity: Pre-Lab Exercise
Name
Date
Laboratory Instructor
Lab Day & Time
1.
Define specific heat using your own words.
2.
If the hot water in this experiment is at 98 °C, how much heat must be released in order for the
temperature of 100 mL of water to decrease to 32 °C? Is this exothermic or endothermic? (The
specific heat of water is 4.184 J/g ∙˚C.)
3.
If a 5.26 g sample of copper at 258 °C is placed in 125 mL of water at 21.0 °C, how hot will the water
get? Assume no heat loss to the surroundings. (The specific heat of copper is 0.385 J/g·°C.)
4.
What is the specific heat of a substance that absorbs 2500 joules of heat when 100g increases in
temperature from 10 °C to 70 °C?
5.
Why do different materials possess different specific heat values?
63
General Chemistry I
Southern California University of Health Sciences
Name _______________________________________
Lab Report _______/
Laboratory Instructor __________________________
Date_______________
Report Form: Experiment 6 Calorimetry and Specific Heat
Trial 1
Trial 2
Mass of H2O
Mass of Metal
Initial Temperature of
H2O
Final Temperature of H2O
ΔTwater (Final – Initial)
Initial Temperature of
Metal
Final Temperature of
Metal
ΔTmetal (Final – Initial)
64
Trial 3
General Chemistry I
Southern California University of Health Sciences
Calculations
1. Using the data from your lab, calculate the specific heat (Cp) of your metal. Show all work in the space
provided. You must show the calculations for each all trials
mwater × Cp water × ΔTwater = -(mmetal × Cp metal × ΔTmetal)
65
General Chemistry I
Southern California University of Health Sciences
Experiment 6 Calorimetry and Specific Heat Capacity: Post-Lab Exercise
Name
Date
Laboratory Instructor
Lab Day & Time
1. What is the identity of the metal?
2. Explain how you were able to identify the unknown metal. What evidence do you have to support
your claim?
3. Were you correct in your claim? If not, what sources of error may have contributed to your incorrect
findings?
4.
Given more time, we should have completed more trials of this lab. Why would that have been a
good idea?
66
General Chemistry I
Southern California University of Health Sciences
Experiment 7: Molecular Structure and Properties
Learning Outcomes:
As you work through this exercise you will learn how to:

Draw Lewis electron-dot structures for small molecules.

Assign shape and polarity to small molecules.

Determine the nature of intermolecular forces of attraction for molecules.
Introduction
Note: You should review electron-dot structures, VSEPR theory, polar bonds, polar molecules, geometric
isomers and inte…