EXPERIMENT
2
The Use of Volumetric
Glassware
Objective
Familiarity with the various instruments used for making physical measurements in the
laboratory is essential to the study of experimental chemistry. In this experiment, you will
investigate the uses and limits of the various types of volumetric glassware.
Introduction
Most of the glassware in your laboratory locker has been marked by the manufacturer to indicate
the volume contained by the glassware when filled to a certain level. The graduations etched or
painted onto the glassware by the manufacturer differ greatly in the precision they indicate,
depending on the type of glassware and its intended use. For example, beakers and Erlenmeyer
flasks are marked with very approximate volumes, which serve merely as a rough guide to the
volume of liquid in the container. Other pieces of glassware, notably burets, pipets, and
graduated cylinders, are marked much more precisely by the manufacturer to indicate volumes. It
is important to distinguish when a precise volume determination is necessary and appropriate for
an experiment and when only a rough determination of volume is needed.
Glassware that is intended to contain or to deliver specific precise volumes is generally
marked by the manufacturer with the letters “TC” (to contain) or “TD” (to deliver). For example,
a flask that has been calibrated by the manufacturer to contain exactly 500 mL of liquid at 20°C
would have the legend “TC 20 500 mL” stamped on the flask. A pipet that is intended to deliver
a precise 10.00 mL sample of liquid at 20°C would be stamped with “TD 20 10 mL.” It is
important not to confuse “TC” and “TD” glassware: such glassware may not be used
interchangeably. The temperature (usually 20°C) is specified with volumetric glassware
since the volume of a liquid changes with temperature, which causes the density of the
liquid to change. While a given pipet will contain or deliver the same volume at any
temperature, the mass (amount of the substance present in that volume) will vary with
temperature.
A. Graduated Cylinders
The most common apparatus for routine determination of liquid volumes is the graduated
cylinder. Although a graduated cylinder does not permit as precise a determination of volume as
do other volumetric devices, for many applications the precision of the graduated cylinder is
sufficient. Figures 2-1 and 2-2 show typical graduated cylinders. In Figure 2-2, notice the plastic
safety ring, which helps to keep the graduated cylinder from breaking if it is tipped over. In
Figure 2-1, compare the difference in graduations shown for the 10-mL and 100-mL cylinders.
Examine the graduated cylinders in your lab locker and determine the smallest graduation of
volume that can be determined with each cylinder.
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Experiment 2 Volumetric Glassware
Figure 2-1. Expanded view of 10-mL and 100-mL
cylinders. Greater precision is possible with the 10-mL
cylinder, since each numbered scale division represents
1 mL.
Figure 2-2. 100-mL graduated cylinder with a plastic
safety ring.
When water (or an aqueous solution) is contained in a narrow glass container such as a graduated
cylinder, the liquid surface is not flat as might be expected. Rather, the liquid surface is curved
downward (see Figure 2-3). This curved surface is called a meniscus, and is caused by an
interaction between the water molecules and the molecules of the glass container wall. When
reading the volume of a liquid that makes a meniscus, hold the graduated cylinder so that the
meniscus is at eye level, and read the liquid level at the bottom of the curved surface.
Figure 2-3. Reading a meniscus. Read the bottom of the
meniscus while holding at eye level.
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Experiment 2 Volumetric Glassware
3
B. Pipets
When a more precise determination of liquid volume is needed than can be provided by a
graduated cylinder, a transfer pipet may be used. Pipets are especially useful if several
measurements of the same volume are needed (such as preparing similar-sized samples of a
liquid substance). Two types of pipet are commonly available, as indicated in Figure 2-4. The
Mohr pipet is calibrated at least at each milliliter and can be used to deliver any size sample (up
to the capacity of the pipet). The volumetric transfer pipet can deliver only one size sample (as
stamped on the barrel of the pipet), but generally it is easier to use and is more reproducible.
Pipets are filled using a rubber safety bulb to supply the suction needed to draw liquid
into the pipet. It is absolutely forbidden to pipet by mouth in the chemistry laboratory. Two
common types of rubber safety bulb are shown in Figure 2-5.
Figure 2-5. Pipet safety bulbs. Never pipet by mouth.
The simple bulb should not actually be placed onto the barrel of the pipet. This would most
likely cause the liquid being measured to be sucked into the bulb itself. Rather, squeeze the bulb,
and just gently press the opening of the bulb against the opening in the barrel of the pipet to
apply the suction force, keeping the tip of the pipet under the surface of the liquid being sampled.
Do not force the pipet into the plastic tip of the safety bulb, or the pipet may break.
Allow the suction to draw liquid into the pipet until the liquid level is 1 or 2 inches above the
calibration mark on the barrel of the pipet. At this point, quickly place your index finger over the
opening at the top of the pipet to prevent the liquid from falling. By gently releasing the pressure
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4
Experiment 2 Volumetric Glassware
of your index finger, the liquid level can be allowed to fall until it reaches the calibration mark of
the pipet. The tip of the pipet may then be inserted into the container that is to receive the sample
and the pressure of the finger removed to allow the liquid to flow from the barrel of the pipet.
(See Figure 2-6.)
To use the more expensive valve-type
bulb (see Figure 2-5b), squeeze the
valve of the bulb marked A, and
simultaneously squeeze the large
portion of the rubber bulb itself to expel
air from the bulb. Press Valve A a
second time, release the pressure on the
bulb, and attach the bulb to the top of
the pipet. Insert the tip of the pipet
under the surface of the liquid to be
measured, and squeeze the valve
marked S on the bulb, which will cause
liquid to being to be sucked into the
pipet. When the liquid level has risen to
an inch or two above the calibration
mark of the pipet, stop squeezing valve
S to stop the suction. Transfer the pipet
to the vessel to receive the liquid and
press valve E to empty the pipet, The
use of this sort of bulb generally
requires considerable practice to
develop proficiency.
Figure 2-6. Filling technique for volumetric pipet.
When using either type of pipet, observe the following rules:
1. The pipet must be scrupulously clean before use: wash with soap and water, rinse with
tap water, and then with distilled water. If the pipet is clean enough for use, the water will
not bead up anywhere on the inside of the barrel.
2. To remove rinse water from the pipet (to prevent dilution of the solution to be measured),
rinse the pipet with several small portions of the solution to be measured, discarding the
rinsings in a waste beaker for disposal. It is not necessary to completely fill the pipet with
the solution for rinsing.
3. The tip of the pipet must be kept under the surface of the liquid being measured out
during the entire time suction s being applied, or air will be sucked into the pipet.
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Experiment 2 Volumetric Glassware
5
4. Allow the pipet to drain for at least a minute when emptying to make certain the full
capacity of the pipet has been delivered. Remove any droplets of liquid adhering to the
tip of the pipet by touching the tip of the pipet to the side of the vessel that is receiving
the sample.
5. If you are using the same pipet to measure out several different liquids, you should rinse
the pipet with distilled water between liquids, and following with a rinse of several small
portions of the next liquid to be measured.
C. Burets
When samples of various sizes must be dispensed or measured precisely, a buret may be used.
The buret consists of a tall, narrow, calibrated glass tube, fitted at the bottom with a valve for
controlling the flow of liquid. The valve is more commonly called a stopcock. (See Figure 2-7).
Figure 2.7. A
volumetric buret.
Typically, 50-mL
burets are used in the
chemistry labs.
Like a pipet, a buret must be scrupulously clean before use.
The precision permitted in reading a buret is 0.02 mL, but if the
buret is not completely clean, this level of precision is not
attainable. To clean the buret, first use soap and water, using a
special long-handled buret brush to scrub the interior of the glass.
Then rinse the buret with tap water, followed by several rinsings
with distilled water.
Before use, the buret should be rinsed with several small
portions of the solution to be used in the buret. The buret should be
tilted and rotated during the rinsings to make sure that all rinse
water is washed from it. Discard the rinsings. After use, the buret
should again be rinsed with distilled water. Many of the reagent
solutions used in burets may attack the glass of the buret if they are
not removed. This would destroy the calibration. To speed up the
cleaning of a buret in future experiments, the buret may be left
filled with distilled water during storage between experiments (if
your locker is large enough to permit this).
A common mistake made by beginning students is to fill the
buret with the reagent solution to be dispensed to be exactly the
0.00 mark. This is not necessary or desirable in most experiments,
and wastes time. The buret should be filled to a level that is
comfortable for you to read (based on your height). The precise
initial liquid level reading of the buret should be taken before the
solution is dispensed and again after the liquid is dispensed. The
readings should be made to the nearest 0.02 mL. The volume of
liquid dispensed is then obtained by simple subtraction of the two
volume readings.
Safety 
Precautions


Wear safety glasses at all times while in the laboratory.
When using a pipet, use a rubber safety bulb to apply the
suction force. Never pipet by mouth.
Rinse the buret carefully do not attempt to admit water
directly to the buret from the cold water tap. Fill a beaker
with tap water, and pour from the beaker into the buret.
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Experiment 2 Volumetric Glassware
Apparatus/Reagents Required
Graduated cylinders, pipets and safety bulb, buret and clamp, beakers, distilled water.
Procedure
Record all data and observations directly in your notebook in ink.
A. The Graduated Cylinder
Your instructor will set up a display of several graduated cylinders filled with different amounts
of colored water. Several sizes of cylinder are available (10-mL, 25-mL, 50-mL, 100-mL).
Examine each cylinder, paying particular attention to the marked scale divisions on the cylinder.
For each graduated cylinder to what fractional unit of volume does the smallest scale mark
correspond?
Read the volume of liquid contained in each graduated cylinder and record. Make your readings
to the level of precision permitted by each of the cylinders.
Check your readings of the liquid levels with the instructor before proceeding, and ask for
assistance if your readings different from those provided by the instructor.
B. The 25-mL Graduated cylinder
Obtain about 100 mL of distilled water in a clean Erlenmeyer flask. Determine and record the
temperature of the distilled water.
Weigh an empty 25-mL graduated cylinder. Then fill the graduated cylinder with distilled water
so that the meniscus of the water level lines up with the 25-mL calibration mark of the cylinder.
Weigh the graduated cylinder (copy all digits seen on weighing scale) and calculate the mass of
water each contains.
Take the temperature of the dispensed water using a thermometer and record. Make sure to read
the thermometer up to 1 decimal place.
Using the Density of Water table from Appendix H to this manual, calculate the volume of water
present in the graduated cylinder from the exact mass of water present.
Compare the calculated volume of water (based on the mass of water) to the observed volumes
of water determined from the calibration marks on the cylinder. Calculate the percentage
difference between the calculated volume and the observed volume from the calibration marks.
C. The 50-mL Beaker
Clean and wipe dry your 50-mL beaker. Weigh the beaker and record its mass (copy all digits
seen on weighing scale).
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Experiment 2 Volumetric Glassware
7
Place distilled water in the 50-mL beaker up to the 25-mL mark. (DO NOT USE ANY
MEASURING DEVICE TO DO THIS.) Since the beaker has insufficient calibration marks,
just eyeball where you think the 25-mL mark would be and place distilled water up to that level.
Weigh the beaker (copy all digits seen on weighing scale) and calculate the mass of water the
beaker contains.
Take the temperature of the dispensed water using a thermometer and record. Make sure to read
the thermometer up to 1 decimal place.
Using the Density of Water table from Appendix H to this manual, calculate the volume of water
present in the graduated cylinder from the exact mass of water present.
Compare the calculated volume of water (based on the mass of water) to the observed volumes
of water determined from the calibration marks on the beaker. Calculate the percentage
difference between the calculated volume and the observed volume from the calibration marks.
Why are the calibration marks on laboratory beakers taken only to be an approximately guide to
volume?
D. The 25.00-mL Pipet
Obtain a 25-mL pipet and rubber safety bulb. Clean the pipet with soap and water. Rinse the
pipet with tap water, and then with small portions of distilled water. Practice filling and
dispensing distilled water from the pipet until you feel comfortable with the technique. Ask your
instructor for assistance if you have any difficulties in the manipulation.
Clean and wipe dry a 150-mL beaker. Weigh the beaker (copy all digits seen on weighing scale)
and record. (Note: This beaker is to be used solely to hold water released from the pipet.)
Obtain about 100 mL of distilled water in a clean Erlenmeyer flask. Determine and record the
temperature of the water. Read thermometer up to 1 decimal place.
Pipet 25 mL of the distilled water from the Erlenmeyer flask and release the water into the clean
beaker you have weighed. Reweigh the beaker containing the 25 mL of water. Determine the
weight of water transferred by the pipet.
Take the temperature of the dispensed water using a thermometer and record. Make sure to read
the thermometer up to 1 decimal place.
Using the Density of Water table from Appendix H to this manual, calculate the volume of water
transferred by the pipet from the mass of water transferred. Compare this calculated volume to
the volume of the pipet as specified by the manufacturer. Calculate % Error. Any significant
difference in these two volumes is an indication that you need additional practice in pipeting.
Consult with your instructor for help.
How does the volume dispensed by the pipet compare to the volumes as determined in part A
using a graduated cylinder or beaker?
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8
Experiment 2 Volumetric Glassware
E. The Buret
Obtain a buret and set it up in a clamp on your lab bench.
Fill the buret with tap water, and check to make sure that there are no leaks from the stopcock
before proceeding. If the stopcock leaks, have the instructor examine the stopcock to make sure
that all the appropriate washers are present. If the stopcock cannot be made leakproof, replace
the buret.
Clean the buret with soap and water, using a long-handled buret brush to scrub the inner surface
of the buret. Rinse all soap from the buret with tap water, being sure to flush water through the
stopcock as well. Rinse the buret with several small portions of distilled water, and then fill the
buret to above the zero mark with distilled water.
Open the stopcock of the buret and allow the distilled water to run from the buret into a beaker or
flash. Examine the buret while the water is running from it. If the buret is clean enough for use,
water will flow in sheets down the inside surface of the buret without being up anywhere. If the
buret is not clean, repeat the scrubbing with soap and water.
Once the buret is clean, refill it with distilled water to a point somewhat below the zero mark.
Determine the precise liquid level in the buret to the nearest hundredth place, for example, 0.02
mL. (Note: The buret must be read up to 2 decimal places.)
With a paper towel, clean and wipe dry a 150-mL beaker. Weigh the beaker (copy all digits seen
on weighing scale) and record.
Place the weighed beaker beneath the stopcock of the buret. Open the stopcock of the buret and
run water into the beaker until approximately 25.00 mL of water have been dispensed. Read the
precise liquid level in the buret to the nearest 0.02 mL (2 decimal places). Calculate the volume
of water that has been dispensed from the buret by subtracting the two liquid levels (Final
Volume – Initial Volume). Reweigh the beaker containing the water dispensed from the buret
(copy all digits seen on weighing scale), and determine the mass of water transferred to the
beaker from the buret.
Take the temperature of the dispensed water using a thermometer and record. Make sure to read
the thermometer up to 1 decimal place.
Use the Density of Water table from Appendix H to calculate the volume of water transferred
from the mass of the water. Calculate % Error. Compare the volume of water transferred (as
determined by reading the buret) with the calculated volume of water (from the mass
determinations). If there is any significant difference between the two volumes, most likely you
need additional practice in the operation and reading of the buret.
How does the volume dispensed by the buret compare to the volumes as determined in Part A
using a graduated cylinder or beaker? How does the volume dispensed by the buret compare to
that dispensed using the pipet in Part B?
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EXPERIMENT
3
Density Determinations
Objective
Density is an important property of matter, and may be used as a method of identification. In this
experiment, you will determine the densities of regularly and irregularly shaped solids as well as
of pure liquids and solutions.
Introduction
The density of a sample of matter represents the mass contained within a unit volume of space in
the sample. For most samples, a unit volume means 1.0 mL. The units of density, therefore, are
quoted in terms of grams per milliliter (g/mL) or grams per cubic centimeter (g/cm3) for most
solid and liquid samples of matter.
Since we seldom deal with exactly 1.0 mL of substance in the general chemistry
laboratory, we usually say that the density of a sample represents the mass of the specific sample
divided by its particular volume.
density =
𝑚𝑎𝑠𝑠
𝑣𝑜𝑙𝑢𝑚𝑒
Because the density does in fact represent a ratio, the mass of any size sample, divided by the
volume of that sample, gives the mass that 1.0 mL of the same sample would possess.
Densities are usually determined and reported at 20°C (around room temperature)
because the volume of a sample, and hence the density, will often vary with temperature. This is
especially true for gases, with smaller (but still often significant) changes for liquids and solids.
References (such as the various chemical handbooks) always specify the temperature at which a
density was measured.
Density is often used as a point of identification in the determination of an unknown
substance. In later experiments, you will study several other physical properties of substances
that are used in the identification of unknown substances. For example, the boiling and melting
points of a given substance are characteristics of that substance and are used routinely in
identification of unknown substances. Suppose an unknown’s boiling and melting points have
been determined but on consulting the literature, it is found that more than one substance has
these boiling and melting points. The density of the unknown might then be used to distinguish
the unknown. It is very unlikely that two substances would have the same boiling point, melting
point, and density.
Density can also be used to determine the concentration of solutions in certain instances.
When a solute is dissolved in a solvent, the density of the solution will be different from that of
the pure solvent itself. Handbooks list detailed information about the densities of solutions as a
function of their composition (typically, in terms of the percent solute in the solution). If a
sample is known to contain only a single solute, the density of the solution could be measured
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9
10
Experiment 3 Density Determinations
experimentally, and then the handbook could be consulted to determine what concentration of
the solute gives rise to the measured solution density.
The determination of the density of certain physiological liquids is often an important
screening tool in medical diagnosis. For example, if the density of urine differs from normal
values, this may indicate a problem with the kidneys secreting substances which should not be
lost from the body. The determination of density (specific gravity) is almost always performed
during a urinalysis.
Several techniques are used for the determination of density. The method used will
depend on the type of sample and on the level of precision desired for the measurement. For
example, devices have been constructed for determinations of the density of urine, that permit a
quick, reliable, routine determination. In general, a density determination will involve the
determination of the mass of the sample with a balance, but the method used to determine the
volume of the sample will differ from situation to situation. Several methods of volume
determination are explored in this experiment.
For solid samples, different methods may be needed for the determination of the volume,
depending on whether or not the solid is regularly shaped. If a solid has a regular shape (e.g.,
cubic, rectangular, cylindrical), the volume of the solid may be determined by geometry:
For a cubic solid, volume = (edge)3
For a rectangular solid, volume = length x width x height
For a cylindrical solid, volume = π x (radius)2 x height
If a solid does not have a regular shape, it may be possible to determine the volume of the
solid from Archimedes’ principle, which states that an insoluble, nonreactive solid will displace
a volume of liquid equal to its own volume. Typically, an irregularly shaped solid is added to a
liquid in a volumetric container (such as a graduated cylinder) and the change in liquid level
determined.
For liquids, very precise values of density may be determined by pipeting an exact
volume of liquid into a sealable weighing bottle (this is especially useful for highly volatile
liquids) and then determining the mass of liquid that was pipeted. A more convenient method for
routine density determinations for liquids is to weigh a particular volume of liquid as contained
in a graduated cylinder.
Safety 
Precautions


Wear safety glasses at all times while in the laboratory.
The unknown liquids may be flammable, and their fumes
may be toxic. Keep the unknown liquids away from open
flames and do not inhale their vapors. Dispose of the
unknown liquids as directed by the instructor.
Dispose of the metal samples only in the container
designated for their collection.
Apparatus/Reagents Required
Unknown liquid sample, regularly shaped metal, irregularly-shaped metal, distilled water
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Experiment 3 Density Determinations
11
Procedure
Record all data and observations directly in you notebook in ink.
A. Determination of the Density of Regularly Shaped Solids
Obtain a regularly shaped solid, and record its identification number (located on the face
of the solid). With a caliper, determine the physical dimensions (e.g., length, width,
height, radius) of the solid to the nearest 0.2 mm. From the physical dimensions, calculate
the volume of the solid.
Determine the mass of the regularly shaped solid (copy all digits seen on weighing scale)
and record. From the mass and volume, calculate the density of the solid.
Calculate your % Error.
B. Determination of the Density of Irregularly Shaped Solids
Obtain a sample of metal pellets (shot) and record its identification number. Weigh a
sample of the metal of approximately 50 g and record the actual mass of the metal (copy
all digits seen on weighing scale).
Add water to your 100-mL graduated cylinder to approximately the 50-mL mark. Record
the exact volume of water in the cylinder to the precision permitted by the calibration
marks of the cylinder.
Pour the metal sample into the graduated cylinder, making sure that none of the pellets
stick to the walls of the cylinder above the water level. Stir/shake the cylinder to make
certain that no air bubbles have been trapped among the metal pellets. (See Figure 3-1.)
Figure 3-1. Measurement of volume by
displacement. A nonsoluble object
displaces a volume of liquid equal to its
own volume.
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12
Experiment 3 Density Determinations
Read the level of the water in the graduated cylinder, again making your determination to
the precision permitted by the calibration marks of the cylinder. Assuming that the metal
sample does not dissolve in or react with water, the change in water levels represents the
volume of the metal pellets.
Calculate the density of the unknown metal pellets. Calculate your % Error.
After blotting them dry with a paper towel, turn in the metal pellets to your instructor (do
not discard).
C. Density of Pure Water
Clean and dry a 25-mL graduated cylinder (a rolled-up paper towel should be used).
Weigh the dry graduated cylinder (copy all digits seen on weighing scale).
Add distilled water to the cylinder so that the water level is above the 20-mL mark but
below the 25-mL mark. Record the temperature of the water in the cylinder up to 1
decimal place.
Reweigh the cylinder.
Record the exact volume of water in the cylinder, to the level of precision permitted by
the calibration marks on the barrel of the cylinder.
Calculate the density of the water. Compare the measured density of the water with the
value listed in the handbook for the temperature of your experiment.
Calculate your % Error.
D. Density of Unknown Solution
Clean and dry a 25-mL graduated cylinder (a rolled-up paper towel should be used).
Weigh the dry graduated cylinder (copy all digits seen on weighing scale).
Add the unknown liquid to the cylinder so that the liquid level is above the 20-mL mark
but below the 25-mL mark.
Reweigh the cylinder.
Record the exact volume of unknown liquid in the cylinder, to the level of precision
permitted by the calibration marks on the barrel of the cylinder.
Calculate the density of the unknown liquid. Compare the measured density of your
unknown liquid with the other density values listed on the data sheet. Identify your
unknown.
Calculate your % Error.
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EXPERIMENT
11 Stoichiometric Determinations
Objective
Stoichiometric measurements are among the most important in chemistry, indicating the
proportions by mass in which various substances react. In this experiment, three examples of
stoichiometric determinations will be investigated. The reactions to be studied are between
magnesium and molecular oxygen, O2 (Choice III).
Choice III. Determination of a Formula
Introduction
Magnesium metal is a moderately reactive elementary substance. At room temperature,
magnesium reacts only very slowly with oxygen and can be kept for long periods of time without
appreciable oxide buildup. At elevated temperatures, however, magnesium will ignite in an
excess of oxygen gas, burning with an intensely white flame and producing magnesium oxide.
Because of the brightness of its flame, magnesium is used in flares and in photographic
flashbulbs.
In this experiment, however, you will be heating magnesium in a closed container called
a crucible, exposing it only gradually to the oxygen of the air. Under these conditions, the
magnesium will undergo a more controlled oxidation, gradually turning from shiny metal to
grayish-white powdered oxide. Because the air also contains a great deal of nitrogen gas, a
portion of the magnesium being heated may be converted to magnesium nitride, Mg3N2, rather
than magnesium oxide. Magnesium nitride will react with water and, on careful heating, is
converted into magnesium oxide
Mg3N2 + 3H2O → 3MgO + 2NH3
The ammonia produced by this reaction can be detected by its odor, which is released on heating
the mixture.
Magnesium is a group IIA metal and its oxide should have the formula MgO. Based on
this formula, magnesium oxide should consist of approximately 60% magnesium by weight. By
comparing the weight of magnesium reacted, and the weight of magnesium oxide that results
from the reaction, this will be confirmed.
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13
14
Experiment 11 Stoichiometric Determinations
Safety 
Precautions

Wear safety glasses at all times while in the laboratory.
Magnesium produces an intensely white flame if ignited,
which may be damaging to the eyes. If the magnesium used
in this experiment ignites in the crucible by accident,
immediately cover the crucible and stop heating. Do not look
directly at the magnesium while it is burning.

When water is added to the crucible to convert magnesium
nitride to magnesium oxide, the contents of the crucible may
spatter when heated. Use only gentle heating in evaporating
the water. Do not heat the crucible strongly until nearly all
the water has been removed.

Use crucible tongs to handle the hot crucible and cover.

Hydrochloric acid is damaging to skin and clothing. If it is
spilled, wash immediately and inform the instructor.
Apparatus/Reagents Required
Porcelain crucible and cover, crucible tongs, clay triangle, magnesium turnings (or ribbon), pH
paper, 6 M HC1
Procedure
Record all data and observations directly in your notebook in ink.
Obtain a crucible and cover and examine. The crucible and cover are extremely fragile and
expensive. Use caution in handling them.
If there is any loose dirt in the crucible, moisten and rub it gently with a paper towel or spatulato
remove the dirt. If dirt remains in the crucible, bring it to the hood, add 5-10 mL of 6 M HC 1
and allow the crucible to stand for 5 minutes. Discard the HC1 and rinse the crucible with water.
If the crucible is not clean at this point, consult with the instructor about other cleaning
techniques, or replace the crucible. After the crucible has been cleaned, use tongs to handle the
crucible and cover.
Set up a clay triangle on a ringstand. Transfer the crucible and cover to the triangle. The crucible
should sit firmly in the triangle (the triangle’s arms can be bent slightly if necessary).
Begin heating the crucible and cover with a small flame to dry them. When the crucible and
cover show no visible droplets of moisture, increase the flame to full intensity, and heat the
crucible and cover for 5 minutes.
Remove the flame, and allow the crucible and cover to cool to room temperature.
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Experiment 11 Stoichiometric Determinations
15
When the crucible and cover have completely cooled, use tongs to move them to a clean, dry
evaporating dish. Do not use a watch glass as it will break upon contact with the hot crucible. Do
not place the crucible directly on the lab bench as it will burn the bench top. Weigh the crucible
and cover. Record. (Copy all digits seen on the weighing scale.)
Return the crucible and cover to the clay triangle. Reheat in the full heat of the burner flame for 5
minutes. Allow the crucible/cover to cool completely to room temperature.
Reweigh the crucible after it has cooled. If the weight this time differs from the earlier weight by
more than 5 mg (0.005 g), reheat the crucible for an additional 5 minutes and reweigh when cool.
Continue the heating/weighing until the weight of the crucible and cover is constant to within 5
mg.
Add approximately 1 teaspoon of magnesium turnings (or about 8 inches of magnesium ribbon
coiled into a spiral) to the crucible. Fold the magnesium ribbon so that it lays flat at the base of
the crucible.
Using tongs, transfer the crucible/cover and magnesium to the balance and weigh them to the
nearest mg (0.001 g).
Set up the crucible on the clay triangle with the cover very slightly ajar. (See Figure 11-1.) With
a very small flame, begin heating the crucible gently.
Figure 11-1. Set-up for oxidation
of magnesium.
If the crucible begins to smoke when heated, immediately cover the magnesium completely and
remove the heat for 2-3 minutes. The smoke consists of the magnesium oxide product and must
not be lost from the crucible. The appearance of smoke indicates that you have lost some of the
desired product.
Continue to heat gently for 5-10 minutes with the cover of the crucible slightly ajar. Remove the
heat and allow the crucible to cool for 1-2 minutes.
Remove the cover and examine the contents of the crucible. If portions of the magnesium still
demonstrate the shiny appearance of the free metal, return the cover and heat with a small flame
for an additional 5 minutes; then re-examine the metal. Continue heating with a small flame until
no shiny metallic pieces are visible.
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16
Experiment 11 Stoichiometric Determinations
When the shiny magnesium metal appears to have been converted fully to the dull gray oxide,
return the cover to its slightly ajar position and heat the crucible with the full heat of the burner
flame for 5 minutes. Then slide the cover to about the half-open position and heat the crucible in
the full heat of the burner flame for an additional 5 minutes.
Remove the heat and allow the crucible and contents to cool completely to room temperature.
Remove the crucible from the clay triangle and set it on a sheet of clean paper on the lab bench.
With a stirring rod, gently break up any large chunks of solid in the crucible. The crushed
magnesium oxide must look like fine, gray powder. If not, heat the crucible for another 5
minutes and cool. Do not set the stirring rod down as some powder will be lost on the lab bench.
Rinse any material that adheres to the stirring rod into the crucible with a few drops of distilled
water. With a dropper, add about 10 drops of distilled water to the crucible, spreading the water
evenly throughout the solid. (Note: If water is added to hot MgO, spattering occurs and more of
the desired product will be lost.)
Return the crucible to the clay triangle and set the cover in the slightly ajar position. With a very
small flame, begin heating the crucible to drive off the water that has been added. Beware of
spattering during the heating. Remove the flame and close the cover of the crucible if spattering
occurs.
As the water is driven off, hold a piece of moistened pH paper (with forceps) in the stream of
steam being expelled from the crucible. Any nitrogen that had reacted with the magnesium is
driven off as ammonia during the heating and should give a basic response with pH paper (you
may also note the odor of ammonia).
When it is certain that all the water has been driven off, slide the cover so that it is in
approximately the half-open position and increase the size of the flame. Heat the crucible and
contents in the full heat of the burner for about 5 minutes.
Allow the crucible and contents to cool complexly to room temperature. When completely cool,
weigh the crucible and contents to the nearest milligram (0.001 g).
Return the crucible to the triangle and heat for another 5 minutes in the full heat of the burner
flame. Allow the crucible to cool complexly to room temperature and reweigh. The two
measurements of the crucible and contents should give weights that agree within 5 mg (0.005 g).
If this agreement is not obtained, heat the crucible for additional 5-minute periods until two
successive weighings agree within 5 mg.
Clean out the crucible, and repeat the determination, if required. Time constraints normally
allow for only one determination per laboratory session.
Calculate the weight of magnesium that was taken as well as the weight of magnesium oxide that
was present after the completion of the reaction. Calculate the percentage of magnesium in the
magnesium oxide from your experimental data. Calculate the mean for your two determinations,
if two determinations were performed.
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Experiment 11 Stoichiometric Determinations
17
Calculate the theoretical percentage of magnesium (by mass) in magnesium oxide, and compare
this to the mean experimental value. Calculate the percent error in your determination.
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EXPERIMENT
15
Molar Mass of a Volatile Liquid
Objective
The molar mass of a volatile liquid will be determined by measuring what mass of vapor of the
liquid is needed to fill a flask of known volume at a particular temperature and pressure.
Introduction
The most common instrument for the determination of molar masses in modern chemical
research is the mass spectrometer. Such an instrument permits very precise determination of
molar mass and also gives a great deal of structural information about the molecule being
analyzed; this is of great help in the identification of new or unknown compounds.
Mass spectrometers, however, are extremely expensive and take a great deal of time and
effort to calibrate and maintain. For this reason, many of the classical methods of molar mass
determination are still widely applied. In this experiment, a common modification of the ideal
gas law will be used in the determination of the molar mass of a liquid that is easily evaporated.
The ideal gas law (PV = nRT) indicates that the observed properties of a gas sample
[pressure(P), volume(V), and temperature(T)] are directly related to the quantity of gas in the
sample (n, moles). For a given container of fixed volume at a particular temperature and
pressure, only one possible quantity of gas can be present in the container:
n=
𝑃𝑉
𝑅𝑇
By careful measurement of the mass of the gas sample under study in the container, the molar
mass of the gas sample can be calculated, since the molar mass, M, merely represents the number
of grams, g, of the volatile substance per mole:
n=
𝑔
𝑀
In this experiment, a small amount of easily volatilized liquid will be placed in a flask of known
volume. The flask will be heated in a boiling water bath and will be equilibrated with
atmospheric pressure. From the volume of the flask used, the temperature of the boiling water
bath, and the atmospheric pressure, the number of moles of gas contained in the flask may be
calculated. From the mass of liquid required to fill the flask with vapor when it is in the boiling
water bath, the molar mass of the liquid may be calculated.
A major assumption is made in this experiment that may affect your results. We assume
that the vapor of the liquid behaves as an ideal gas. Actually, a vapor behaves least like an ideal
gas under conditions near which the vapor would liquefy. The unknown liquids provided in this
experiment have been chosen, however, so that the vapor will approach ideal gas behavior.
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18
Experiment 15 Molar Mass of a Volatile Liquid
Safety 
Precautions

19
Wear safety glasses at all times while in the laboratory.
Assume that the vapors of your liquid unknown are toxic.
Work in an exhaust hood or other well-ventilated area.

The unknown may also be flammable. All heating is to be
performed use a hotplate.

The liquid unknowns may be harmful to skin. Avoid contact,
and was immediately if the liquid is spilled.

A boiling water bath is used to heat the liquid, and there may
be a tendency for the water to splash when the flask
containing the liquid is inserted. Exercise caution.

Use tongs or a towel to protect your hands from hot
glassware.
Apparatus/Reagents Required
250-mL Erlenmeyer flask and 1000-mL beaker, 2 pieces aluminum freezer foil (one larger than
the other), 2 rubber bands, needle or pin, oven (110°C), unknown liquid sample, hotplate
Procedure
Record all data and observations directly in your notebook in ink.
Prepare a 250-mL Erlenmeyer flask by cleaning the flask and then drying it completely. The
flask must be completely dry, since any water present will vaporize under the conditions of the
experiment and will adversely affect the results. An oven may be available for heating the flask
to dryness, or your instructor may describe another technique.
Cut 2 square pieces of thick (freezer) aluminum foil to serve as “inner” and “outer” covers for
the flask. Make sure one piece is larger than the other. Mark the larger piece of foil to identify it
as the”outer” cover. Obtain 2 rubber bands. Mark the second rubber band to indicate it is the
“outer” rubber band.
Prepare a 1000-mL beaker for use as a heating bath for the flask. The beaker must be large
enough for most of the flask to be covered by boiling water when placed in the beaker. Add the
required quantity of tap water (approximately 800-mL) to the beaker. Add 2 small pieces of
boiling chips in the beaker to maintain even boiling of the water. Set the beaker on a hot plate
and heat.
Weigh the dry, empty flask with its inner foil cover to the nearest mg (0.001 g).
Obtain an unknown liquid and record its identification number.
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20
Experiment 15 Molar Mass of a Volatile Liquid
Add 3-4 mL of liquid to the dry Erlenmeyer flask. Cover the flask with the inner foil cover,
making sure that the foil cover is tightly crimped around the rim of the flask and held in place
with the inner rubber band. Fold up the edges of the foil over the rubber band.
Cover the inner foil and inner rubber band with the larger piece of aluminum foil (outer foil) and
hold in place with the second rubber band. Make sure the outer rubber band seals both the inner
foil and inner rubber band.
Punch a single small hole in the foil covers with a needle or pin. Make sure the hole goes through
the 2 foils. If the holes are too big, replace the foils. (Note: The outer foil is needed to keep the
inner foil dry.)
When the water in the beaker begins to boil, carefully immerse the flask containing the unknown
liquid in the boiling water so that most of the flask is covered with the water of the heating bath
(see Figure 15-1). Clamp the neck of the flask to hold the flask in the boiling water. Adjust the
temperature of the hotplate so that the water remains boiling but does not splash from the beaker.
Figure 15-1. Apparatus for determination
of the molar mass of a volatile liquid. Most of the flask
containing the unknown liquid must be beneath
the surface of the boiling water bath.
Watch the unknown liquid carefully. The liquid will begin to evaporate rapidly, and its volume
will decrease. The amount of liquid placed in the flask is much more than will be necessary to
fill the flask with vapor at the boiling water temperature. Excess vapor will be observed escaping
through the pinhole made in the foil cover of the flask.
When it appears that all the unknown liquid has vaporized (approximately 3 minutes), and the
flask is filled with the vapor, continue to heat for 1-2 more minutes. (Total heating must not
exceed 5 minutes; otherwise, the gas in the flask will be displaced by air from the outside.)Then
remove the flask from the boiling water bath; use the clamp on the neck of the flask to protect
your hands from the heat.
Set the flask on the lab bench, remove the clamp, and allow the flask to cool to room
temperature. A small amount of liquid will reappear in the flask as the vapor in the flask cools.
While the flask is cooling, measure and record the exact temperature of the boiling water in the
beaker, as well as the barometric pressure in the laboratory.
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Experiment 15 Molar Mass of a Volatile Liquid
21
When the flask has completely cooled to room temperature, carefully dry the outside of the flask
to remove any droplets of water. Carefully remove the outer aluminum foil and outer rubber
band. Then weigh the flask, foil cover, and condensed vapor to the nearest mg (0.001 g).
Do a second determination using the same flask, foils and rubber bands. There is no need to dry
the flask since excess sample will escape through the pinhole anyhow.
Add another 3-4 mL sample of unknown liquid. Reheat the flask until it is filled with vapor;
cool, and reweigh the flask. The weight of the flask after the second sample of unknown liquid is
vaporized should agree with the first determination within 0.05 g. If it does not, do a third
determination.
When two acceptable determinations of the weight of vapor needed to fill the flask have been
obtained, remove the foil cover from the flask and clean it out.
Fill the flask to the very rim with tap water. Weigh the flask with water, inner foil cover, and
inner rubber band using the weighing scale at the back of the lab. Record the weight to the
nearest 0.1 g. Determine the temperature of the tap water in the flask using a thermometer.
Record the water temperature up to 1 decimal place and determine its density from Appendix H.
Using the density of water in the flask and the mass of water the flask contains, calculate the
exact volume of the flask.
(If no balance with the capacity to weigh the flask when filled with water is available, the
volume of the flask must be approximated by pouring the water in the flask into a 1-L graduated
cylinder and reading the water level in the cylinder.)
Using the volume of the flask (in liters), the temperature of the boiling water bath (in kelvins),
and the barometric pressure (in atmospheres), calculate the number of moles of vapor the flask is
capable of containing. R = 0.08206 L atm/mol K.
Using the weight of unknown vapor contained in the flask, and the number of moles of vapor
present, calculate the molar mass of the unknown liquid.
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