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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.

Copyright © Houghton Mifflin Co. All rights reserved.

2 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.


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



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.

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.


Figure 2-6. Filling technique for volumetric pipet.


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).

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.

SafetyPrecautions

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.


Figure 2.7. A volumetric buret. Typically, 50-mL burets are used in the chemistry labs.


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).



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.


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.


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?



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.


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?


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

Copyright © Houghton Mifflin Co. All rights reserved. 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 heightFor 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.

SafetyPrecautions


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.


Unknown liquid sample, regularly shaped metal, irregularly-shaped metal, distilled water


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.


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.


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.


Experiment 3 Density Determinations 13

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.



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.


Experiment 11 Stoichiometric Determinations 15

SafetyPrecautions


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.


Porcelain crucible and cover, crucible tongs, clay triangle, magnesium turnings (or ribbon), pH paper, 6 M HC1

Procedure


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.


16 Experiment 11 Stoichiometric Determinations

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.


Experiment 11 Stoichiometric Determinations 17

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.


18 Experiment 11 Stoichiometric Determinations

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.


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 = PVRT

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 = gM

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.


Experiment 15 Molar Mass of a Volatile Liquid 20

SafetyPrecautions


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.


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


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.


Experiment 15 Molar Mass of a Volatile Liquid 21

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 determinationof the molar mass of a volatile liquid. Most of the flask containing the unknown liquid must be beneaththe 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.


22 Experiment 15 Molar Mass of a Volatile Liquid

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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faculty.gordonstate.edufaculty.gordonstate.edu/cfermin-ennis/CHEM 1211 labs/Chem... · Web viewThe most common apparatus for routine determination of liquid volumes is the graduated