Be familiar with your glassware Beaker (Conical flask/ erlenmeyer flask, Glass rod with policeman, Short stem funnel, Watch glass, Filter paper, Activated carbon, Hot plates, Stand and ring support)

Be familiar with your glassware

 Beaker



 










Conical flask/ erlenmeyer flask

 










 Glass rod with policeman

 











 Short stem funnel

 







  

Watch glass



 









Other things you will need for this experiment





Filter paper


 








 Activated carbon
 




 








 Hot plates

 

 







 Stand and ring support

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RECRYSTALLIZATION

Recrystallization is a technique used to purify organic solids. That is, you will separate molecules of impurity from molecules of desired materials

A Recrystallization solvent should have the following properties:
1.    Does not dissolve the compound to be purified when cold
2.    Does dissolve the compound then hot (near boiling)
3.    Has a relatively low bp for easy evaporation from the purified compound
4.    Does not react with the compound being purified,
5.    The cold solvent will keep impurities dissolved.

So- you’ll start with impure crystals, dissolve them in a hot solvent and the get the crystals back in the end in a pure state
There are several steps to Recrystallization. They are:

1. Choose a solvent that does not dissolve the compound when cold but does so readily when near boiling
2. Add hot solvent until the compound you are trying to purify just dissolves; do not add too much solvent. Use a minimal amount
3. Add decolorizing carbon (charcoal) to remove colored impurities. The step is skipped most often. Do not add the charcoal to a boiling solution because it will bump and boil over. Add a small amount of charcoal because you will lose good compound as well as colored impurities by adsorption to the charcoal.
4. FILTER BY GRAVITY WHILE HOT (gravity filtration). This is the step that makes the recrystallization a success. It should never be omitted from a recrystallization. Remember this step for future recrystallizations. This step removes the charcoal and any undisclosed impurities. Use fluted filter paper. This provides a large surface area allowed for fast filtration. Recall: when you want the filtrate, use gravity filtration.
5. Cool the filtrate from #4. Cool first in air and then later in an ice/water bath, ask your TA what to do if crystals don’t form.
6. Perform suction filtration to isolate the compound, dissolved impurities stay in this filtrate. Recall: when you want the solid, use suction filtration.
7. Dry the crystals in a safe place in your drawer. Do not bottle them as they need to dry. Do keep them safe.

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Procedure for recrystallization

Procedure for recrystallization-
1.    Place 1.5 to 2.0 g of impure unknown in a 250 mL Erlenmeyer flask and dissolve in the minimum amount of boiling water using hot plate. You should know the exact mass of the amount of material you are using. There may be impurities insoluble in water and dark oil many form in the hot solution. If oil forms this is called “oiling-out” and usually happens when a compound melts near the boiling point of the solvent being used. The oil will eventually go in to solution although a small amount of additional solvent may need to be added.
(Note: You should keep more water boiling in a separate flask than you think you will need- about 250 mL should do. Also, when performing the recrystallization, add the hot solvent.
(water) to the solute and not the solute to the solvent.)

2.    The boiling solution is allowed to cool slightly, and a small amount (about 0.25 – 0.5g) of charcoal s added. Estimate this amount by using the tip of scoopula. If the charcoal is added to the boiling solution, it usually causes the solution to boil over the top over the flask. Never adda solid substance to a solution at or near its boiling point.

3.    The solution is heated again it boiling, and is then filtered by gravity through a pre-warmed funnel and fluted filter paper while hot. This is called hot filtration. This is a critical step in recrystallization when there are insoluble impurities present such as charcoal. See below on how to flute filter paper. The filter paper doesn’t need to be perfect but it should be folded so as to increase the surface area and allow for rapid filtration. (Picture should be added)

4.    The clear filtrate is then allowed to cool slowly to room temperature. Continue cooling on an ice/water bath. After the cooled solution has stood in ice-water for about 30 minutes, crystals will form. If crystals don’t form, try scratching the side of the glass container below the solvent level with a stirring rod. This creates an uneven surface on which crystal formation may occur. You may also try adding a seed crystal (your TA may have one) to induce crystallization. If these fail you may have to concentrate your solution as you probably used more than a minimal amount of solvent.

5.    Collect the crystals by suction filtration. Be sure you use a clean filtration flask. Clamp the filtration flask and turn the water aspiration on full. Wet the filter paper with the solvent you are using for recrystallization. After the filtration is complete, break vacuum by taking the hose off the filtering flask BEFORE turning off the water. Your TA may demonstrate. If you believe your aspirator is not doing the job, tell your TA so that it may be replaced for next time.

6.    Dry the crystals on a watch glass or filter paper in your drawer until next week. Make sure they are stored safely but allow them to air dry. When they are dry, weigh carefully and record their melting point. From the melting point and mixed melting point determine the identity of your unknown. Please see the section on melting pint and mixed melting point. SAVE YOUR ENTIRE SAMPLE. You will run an Infrared Spectrum of this material. Proton NMR spectra will be provided. The filtrate is waste and should be placed in the aqueous acidic waste container.

Report the total weight of pure compound obtained and calculate the % recovery.

% recovery = amount obtained after recrystallization/amount of crude X100

Table 1 :  Unknown list. (Structures are provided on the next page.)

Compound name                                                        MP(in degrees of Celsius)
acetanilide                                                                                113-115
p-nitrobenzaldehyde                                                                106-108
benzoic acid                                                                             122-123
acetylsalicylic acid                                                                   138-140
4-dimethylaminobenzaldehyde                                                 73-75
4-hydroxy-3-methoxybenzaldehyde (also known as vanillin)    81-83
salicylic acid                                                                             159-160
p-aminoacetanilide                                                                   162-163
o-toluic acid                                                                             103-105
p-toluic acid                                                                             180-182
m-toluic acid                                                                            108-110
p-toluic acid                                                                             180-182
3-nitroaniline                                                                           112-114
o-aminobenzoic acid                                                               144-148
m-aminobenzoic acid                                                              178-180
t-cinnamic acid                                                                        134-135
o-nitrobenzoic acid                                                                  146-148
m-nitrobenzoic acid                                                                 139-141
p-nitrobenzoic acid                                                                  237-240

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Five major steps in the recrystallization process ( Dissolving the solute in the solvent Performing a gravity filtration, if necessary Collecting the solute crystals by vacuum filtration Drying the resulting crystals Answers to common questions about the experiment)

• Dissolving the solute in the solvent

• Performing a gravity filtration, if necessary

• Collecting the solute crystals by vacuum filtration

• Drying the resulting crystals

• Answers to common questions about the experiment

1.    Dissolving the solute in the solvent

• Add a minimum amount of boiling solvent to the beaker that contains the impure sample. Minimum amount of solvent will make sure that once your sample is dissolved, it will have a high concentration. High concentration is a must if you need t get crystals from it. If your solution is to dilute, you will have problems in getting a good yield. Always keep control over the amount of solvent you add.
• Heat the beaker containing the solute and continue adding boiling solvent incrementally until all of the solute has been dissolved. If additional solvent can be added with no appreciable change in the amount of solute present, the particulate matter is probably insoluble impurities.
• If the compound has colored impurities, you can add activated charcoal. DO NOT add activated charcoal to a boiling solution.

WHY?
Activated carbon has large number of small pores and air is trapped in it. If you add them in to boiling solvent, all that air is going to expand at once and your solvent will boil off.

2.    Hot Gravity Filtration

* This step is optional if there is no visible particulate matter and the solution is the expected color (most organic compounds are white or light yellow). If you added activated charcoal, you must do gravity filtration.

* Fluted filter paper gives increased surface area for filtering. It can speed up the filtering process because of the increased surface area. A stemless funnel is used for hot gravity filtration. A funnel with a stem is prone to premature recrystallization inside the stem because the filtrate can cool as it passes through the stem. At these cooler temperatures, crystals are likely to form.
     


* If the funnel was properly heated before filtration, all of the solution will have passed through and no crystals will have formed on the paper or in the funnel. If crystals have formed, pouring a small amount of boiling solvent through the funnel will dissolve these. If the solution is still discolored after using activated carbon and filtering, either the color is from the compound and will not go away or you need to repeat the step with the addition of activated carbon.
* The solution should be allowed to cool slowly to room temperature. Gradual cooling is conducive to the formation of large, well-defined crystals.

1.    Vacuum Filtration

When you are asked to draw any kind of a diagram, make sure you label each and every thing on your diagram. This is a point where lot of students lose marks!

* You have to agitate the crystals with a fire polished glass-stirring rod before pouring the mother-liquor along with the crystals through the Buchner funnel. Apply the maximum amount of suction possible using the aspirator.
* Some crystals may have been left behind in the beaker; there are two ways to affect a quantitative transfer of all of this material. Either use a portion of the filtrate to rinse the beaker or use a rubber policeman on the end of your stirring rod to scrape the remaining crystals into the Buchner funnel.
* When the crystals have been collected and washed, allow the aspirator to run for several minutes so that the crystals have an opportunity to dry.

2.    Drying the Crystals

* When the crystals have been dried as much as possible in the Buchner funnel, use a scoopula to remove them to a beaker or crystallizing dish together with the filter paper. DO NOT try to remove the crystals until they dry completely. This will ensure that the crystals are not contaminated by filter paper fibers as they dry.
* Spreading the crystals out in a beaker or a crystallizing dish will provide for the most efficient drying as the crystals will have a maximum of exposed surface area.
* When the crystals are dried, the purity of the sample can be measured by performing a melting point determination.

Melting point apparatus
  

3.    What to do if crystals don't form

* If crystals don't form upon slow cooling of the solution to room temperature there are a variety of procedures you can perform to stimulate their growth. First, the solution should be cooled in an ice bath. Slow cooling of the solution leads to slow formation of crystals and the slower crystals form, the more pure they are. Rate of crystallization slows as temperature decreases so cooling with an ice bath should only be used until crystals begin to form; after they do, the solution should be allowed to warm to room temperature so crystal formation occurs more slowly. If no crystals form even after the solution has been cooled in an ice bath, take a fire polished stirring rod and etch (scratch) the glass of your beaker. The small pieces of glass that are etched off of the beaker serve as nuclei for crystal formation. If crystals still do not form, take a small amount of your solution and spread it on a watch glass. After the solvent evaporates, the crystals that are left behind can serve as seeds for further crystallization. Both these methods of nucleation (i.e. etching and seed crystals) cause very rapid crystallization, which can lead to the formation of impure crystals.
* Crystals will not form if there is a large excess of solvent. If no crystals form with the methods already discussed, a portion of the solvent may need to be removed. This can be accomplished by heating the solution for a period of time in order to evaporate some solvent. The new, concentrated solution, should be cooled, and the previously mentioned methods to stimulate crystallization should again be attempted.
* Another potential problem in recrystallization is that the solute sometimes comes out of solution in the form of impure oil instead of forming purified crystals. This usually happens when the boiling point of the solvent is higher than the melting point of the compound, but this is not the only scenario in which this problem presents itself. If this begins to happen, cooling the solution will not stimulate crystallization, it will make the problem worse. If an oil begins to form, heat the solution until the oil portion dissolves and let the whole solution cool. As the oil begins to form again, stir the solution vigorously to break up the oil. The tiny beads of oil that result from this shaking may act as the nuclei for new crystal formation.

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Recrystallization

Theory

The principle behind recrystallization is that the amount of solute that can be dissolved by a solvent increases with temperature. In recrystallization, a solution is created by dissolving a solute in a solvent at or near its boiling point. At this high temperature, the solute has a greatly increased solubility in the solvent, so a much smaller quantity of hot solvent is needed than when the solvent is at room temperature. When the solution is later cooled, after filtering out insoluble impurities, the amount of solute that remains dissolved drops precipitously. At the cooler temperature, the solution is saturated at a much lower concentration of solute. The solute that can no longer be held in solution forms purified crystals of solute, which can later be collected.

Recrystallization works only when the proper solvent is used. The solute must be relatively insoluble in the solvent at room temperature but much more soluble in the solvent at higher temperature. At the same time, impurities that are present must either be soluble in the solvent at room temperature or insoluble in the solvent at a high temperature. For example, if you wanted to purify a sample of Compound X which is contaminated by a small amount of Compound Y, an appropriate solvent would be one in which all of Compound Y dissolved at room temperature because the impurities will stay in solution and pass through filter paper, leaving only pure crystals behind. Also appropriate would be a solvent in which the impurities are insoluble at a high temperature because they will remain solid in the boiling solvent and can then be filtered out. When dealing with unknowns, you will need to test which solvent will work best for you. According to the adage "Like dissolves like," a solvent that has a similar polarity to the solute being dissolved will usually dissolve the substance very well. In general, a very polar solute will easily be dissolved in a polar solvent and will be fairly insoluble in a non-polar solvent. Frequently, having a solvent with slightly different polarity characteristics than the solute is best because if the polarity of the two is too closely matched, the solute will likely be at least partially dissolved at room temperature.

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Spectrophotometric Determination of Iron in a Vitamin Tablet

Objective: In this experiment, you will use the quantitative technique of
spectrophotometry to determine the mass of iron contained in a commercially available
vitamin tablet.

Introduction
Iron and the Human Body

The element which is found in the fourth greatest abundance in the earth’s crust is also encountered everywhere in biological systems. These two facts about iron are not unrelated as nature has found it logical to adapt this plentiful element to perform a variety of functions. The fact that iron has two readily accessible oxidation states, Fe(II) and Fe(III), also contributes to its usefulness as an active component of proteins. The average adult human body contains 4-6 g of iron. In human beings, the majority of iron present is found in the blood in a protein called hemoglobin. The function of this protein is to transport oxygen from the lungs to the various tissues in the body where it is used to produce energy. One of the byproducts of this metabolism, carbon dioxide, is then transported back to the lungs by hemoglobin. Both the oxygen and carbon dioxide molecules bind to the iron ion present in hemoglobin during transport. Humans obtain the iron necessary for the formation of hemoglobin from their diet in foods such as meat and leafy, green vegetables. When the dietary intake is deficient in iron, a condition called anemia results. Someone who is anemic exhibits a lack of energy and often unusually pale skin tone (the red color of blood is also a result of the presence of iron in hemoglobin). Dietary supplements of iron in the form of vitamin tablets can be administered to help alleviate this condition. The tablets you will examine contain iron in the form of ferrous fumarate, [Fe(II)(C4H2O42–)].

Spectrophotometric Analysis

One of the most common techniques used in the quantitative analysis of samples for a specific chemical substance is called spectrophotometry. In spectrophotometry, the amount of electromagnetic radiation absorbed by a sample is measured with an instrument called a spectrophotometer, and this absorbance is related to the concentration of the species for which the sample is being analyzed. When a source of visible radiation is used, the relationship between absorbance, A, and concentration, c, is known as Beer’s law:
A = εcl Eq. 1

In this equation, the concentration is expressed in molarity. Molarity is represented by a
capital, M :

molarity = M = number of moles of solute in the solution/ volume of solution in liters  Eq. 2
Thus, the number of moles of solute present in a solution can be calculated in the
following manner:
number of moles solute = (molarity)•(volume of solution in liters) Eq. 3
The other variables in Beer’s law are: ε, a constant that is a characteristic of the absorbing species called the molar absorptivity, and l, the length of the absorption cell that contains the solution. You should read all of Section V. in the Techniques section of the lab manual/notebook to understand fully the instrumentation and principles involved in measuring the absorbance of a solution. Because the molar absorptivity value is not always a known quantity, a calibration curve is often constructed from a series of standard solutions. A standard solution is one in which the concentration of the species being analyzed is known. The absorbances of several standard solutions are measured, and these values are plotted as a function of the concentrations of the solutions. If the absorbing species behaves according to Beer’s law, such a plot should produce a straight line (see Eq. 1). The absorbance of a solution of unknown concentration can then be measured with the spectrophotometer, and this value can be used in conjunction with the calibration curve to determine the concentration of this solution.

 

Analyzing for iron spectrophotometrically 

In this experiment, you will spectrophotometrically analyze a commercially available vitamin tablet in order to determine the quantity of iron that is contained in the tablet. To do this, the iron will first be converted into a form that absorbs radiation in the visible region. This conversion will be done by reacting the iron with an organic compound called 1,10-phenanthroline. The structure of this molecule is shown in Figure 1. It reacts with Fe(II) to form an ionic compound that contains one iron ion and three 1,10-phenanthroline molecules (the structure of this compound is also shown, in Figure 2). The color of this compound in solution is bright red-orange; consequently, it absorbs light very strongly in the visible region at a wavelength of 508 nanometers. You will react 1,10-phenanthroline with a vitamin tablet that has been dissolved in hydrochloric acid. In order for this reaction to occur, the iron ions that are present must be in the Fe(II) oxidation state. Since Fe(II) is easily oxidized to Fe(III) in the presence of acid and water, a reducing agent called hydroquinone is added to the solution. In addition, it is important that the acidity of the solution is carefully controlled or compound 1 will not form; the presence of sodium citrate in solution will neutralize some of the acid and maintain the proper pH. The resulting solution will be diluted to an appropriate concentration and its absorbance measured. A spectrophotometric calibration curve will then be constructed from a series of standard solutions which contain known concentrations of the iron–1,10-phenanthroline compound (1). The concentration of iron in your sample solution, and thus the amount of iron contained in the vitamin tablet, cansubsequently be calculated.

Procedure

Preparing the “original” iron solution
Grind a vitamin tablet with a mortar and pestle (it does not have to be finely ground). Obtain 25 mL of 6 M hydrochloric acid in a 150 mL beaker and place the ground vitamin tablet in this solution. Swirl the beaker and wait a few minutes for evidence of a reaction. After the initial reaction has subsided, place the beaker on a hot plate in a fume hood, cover with a watchglass and heat to boiling. Boil the mixture for 15 minutes. You may need to add more distilled water during the boiling period if the volume falls below about 15 mL. Remove the beaker from the hot plate and rinse the bottom of the watchglass using a wash bottle, catching the rinse water in the reaction beaker. Filter the solution while still warm by gravity filtration (see TECH I.G) directly into a 100 mL volumetric flask. Rest the funnel in a wooden funnel holder. While the solution is filtering, heat some distilled water in a 250 mL beaker on a hot plate. When the filtration is completed, use small amounts of this hot water to rinse out the beaker, pouring the rinse water on the solid residue contained in the filter paper such that it filters into the 100 mL volumetric flask. Then use small amounts of the hot water to wash the residue two more times. Be careful not to overshoot the mark on the volumetric flask with these rinses. Turn off the hot plate. Allow the solution in the volumetric flask to cool to room temperature, then dilute to the mark with distilled water and mix. Be sure to add the final amount of water to the volumetric flask with a dropper so that you do not overshoot the calibration mark! This solution in the volumetric flask is what we will call the “original” iron solution. Pour this solution into a clean, dry 250 mL Erlenmeyer flask, label it with the words “Original Iron Solution” and set aside. Rinse the volumetric flask three times with small portions of distilled water, discarding the rinses into the sink.

The first dilution
Obtain a 5 mL volumetric pipet and rinse it two or three times with small portions of the solution contained in the Erlenmeyer flask (see TECH II.D). Pipet 5 mL of the iron solution into the volumetric flask, dilute to the mark with water and mix. Label the volumetric flask "Flask A". [Set aside the Original Iron Solution in case you make a mistake during the dilution procedures. When you are finished with the experiment, pour this solution into the Laboratory Byproducts jar labeled Iron + HCl.] Determining the amount of sodium citrate required to maintain the proper pH Obtain a 10 mL volumetric pipet and rinse it two or three times with small portions of the solution contained in Flask A (the volumetric flask). Label a clean, dry 125 mL Erlenmeyer flask "Flask B", then pipet 10 mL of the solution in Flask A into Flask B. Obtain about 8 mL of sodium citrate solution in a graduated cylinder and a strip of indicator paper. Add the sodium citrate solution dropwise to the solution in Flask B, counting the drops as you add them. After you have added 10 drops, test the solution with the indicator paper. The paper should turn yellow-green, indicating that a pH between 3 and 4 has been reached (a color chart will be available for comparison purposes). If a pH lower than 3 is indicated, continue to add sodium citrate solution dropwise, checking the solution with the indicator paper after every four or five drops until the proper color is reached. If a pH higher than 4 is indicated, start over with a new 10 mL aliquot of the solution from Flask A, and add less than 10 drops of sodium citrate before testing the pH. Record the total number of drops of sodium citrate that were required to reach a pH of 3 to 4 in your notebook.

The second dilution and conversion of iron to the iron-1,10-phenanthroline compound Label a clean, dry 250 mL Erlenmeyer flask "Flask C". Pour the contents of Flask A into Flask C and set aside. Rinse the volumetric flask (Flask A) with distilled water (discard the rinses in the sink). Pipet 10 mL of the solution in Flask C into the volumetric flask. Add the same number of drops of sodium citrate which were required to properly neutralize the solution (the number recorded in your notebook). Then add 2 mL of hydroquinone solution and 3 mL of 1,10-phenanthroline solution to the volumetric flask and swirl. Dilute to the mark with distilled water and mix thoroughly. Let this solution stand for at least 15 minutes before measuring the absorbance.

Preparing the calibration curve and measuring the absorbance of the sample solution
Set the wavelength at 508 nanometers. Note! Never put acetone into a cuvet as it will “cloud” the plastic. Each group assigned to a spectrophotometer must obtain the absorbance readings of the standard solutions provided using their instrument. The standard solutions contain the iron-1,10-phenanthroline compound, 1, in water at the following concentrations: 0.500 x 10-5 M; 1.00 x 10-5 M; 2.00 x 10-5 M; 3.00 x 10-5 M; 5.00 x 10-5 M. Use distilled water to "zero" the instrument before measuring the standard solutions. You need to make a new “blank” solution to “re-zero” the spectrophotometer before measuring the sample solution. Place one drop of sodium citrate solution, one drop of hydroquinone solution and one drop of 1,10-phenanthroline solution into a cuvet which has been thoroughly rinsed with water. Dilute almost to the top of the cuvet with distilled water, and stir gently with a small stirring rod. Use this solution to zero the instrument, then obtain the absorbance reading of your sample solution (the contents of the volumetric flask), rinsing the cuvet first with small portions of the solution. Pour the final sample solution (in the volumetric flask) and the standard solutions into the Laboratory Byproducts jar labeled Iron-1,10-phenanthroline. Discard the Original Iron Solution into the Laboratory Byproducts jar labeled Iron + HCl.

Calculations

1. Plotting and using the calibration curve

• Use the program Excel to plot the calibration curve.
• Follow the instructions to draw the “best fit” line through the data points.
•  Use the LINEST function to determine the slope and y-intercept. Follow the instructions in TECH IV to find the concentration, in molarity, of your sample solution using Excel.
•  Print the plot and data table, and hand them in with your report.

2. Determining the amount of iron in the tablet

• Calculate the concentration of iron in your original iron solution (the solution in the volumetric flask which was obtained after the boiling and filtration steps) in molarity. Remember, there are 2 dilution factors!  for a discussion of dilution calculations. In this case, the unknown quantity is the initial molarity.
•  Knowing the concentration of your “original” iron solution, you can now calculate the number of milligrams of iron in the vitamin tablet (see Eq. 3).

Questions

1. After filtering the hot solution of the vitamin tablet in hydrochloric acid, a student does not rinse the beaker with hot water. Would the omission of this step result in an absorbance reading for the final solution that was too high or too low? Explain. What will be the effect on the value reported for the amount of iron in the vitamin tablet?
2. Is your calibration curve a good example of Beer’s Law? Why or why not? Calculate the value for molar absorptivity, ε, based on the calibration curve that you prepared. Assume the length of the cell, l, is 1.00 cm. Be sure to include the units!
3. List the important sources of error in this experiment and what effect each would have on the results. Discuss in particular any errors that you may have made.

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Synthesis of Alum from Scrap Aluminum

Objective: In this experiment, you will be converting the aluminum metal from a
beverage can into the chemical compound potassium aluminum sulfate, KAl(SO4)2•12 H2O, commonly referred to as alum.

Introduction
It had been a good year for the O’Keefe farm with just the right amount of sun and rain. In fact, the cucumbers had grown so well that Richard was afraid he wouldn’t be able to sell them fast enough. Faced with the prospect of a warehouse full of rotten cukes, he contemplated the idea of preserving them and selling them as “O’Keefe’s Farm Fresh Pickles”. He discussed the plan with his wife, Diane, who agreed it was a good idea and said she had a delicious pickling recipe that had been handed down to her from her mother. “I recall one of the ingredients was alum. I think it helps to keep the pickles firm and crisp,“ Diane said. “It just so happens that I know of a recipe for making alum from aluminum cans,” Richard said. “We can use some of those soda pop cans that have been piling up, and that should save us some money.” “How can you convert aluminum metal into an aluminum salt?” Diane asked. Richard explained the chemistry behind the process: One of the interesting properties of aluminum is that it is amphoteric, meaning it will dissolve in both strong, aqueous acids and strong, aqueous bases In both cases, the formation of hydrogen gas is observed:
 

2 Al (s) + 6 H+ (aq) = 2 Al3+ (aq) + 3 H2 (g)

2 Al (s) + 6 H2O (l) + 2 OH– (aq) = 2 Al(OH)4– (aq) + 3 H2 (g)

The first reaction that needs to be carried out is that of aluminum and potassium
hydroxide, KOH:

2 Al (s) + 2 KOH (aq) + 6 H2O (l) = 2 Al(OH)4– (aq) + 2 K+ (aq) + 3 H2 (g)

Adding sulfuric acid, H2SO4, to the resulting solution will cause the compound Al(OH)3
to precipitate; however, Al(OH)3 is also amphoteric and will re-dissolve when more acid
is added:

2 K+ (aq) + 2 Al(OH)4– (aq) + H2SO4 (aq) =2 K+ (aq) + 2 Al(OH)3 (s) + 2 H2O(l) + SO42–(aq)

2 Al(OH)3(s) + 3 H2SO4 (aq) = 2 Al3+ (aq) + 3 SO42–(aq) + 6 H2O(l)

Crystals of the double salt KAl(SO4)2⋅12 H2O (s), or alum, will form upon cooling this
final solution since the solubility of alum in water decreases as the temperature is
lowered:

K+ (aq) + Al3+ (aq) + 2 SO4 2– (aq) + 12 H2O(l) = KAl(SO4)2•12 H2O (s)

“But how will we know that everything really worked ok and that it’s really alum that you produced? I think we need to implement some quality control measures,” Diane observed. “In order to confirm that your synthesis of alum resulted in the desired product, we need to perform a qualitative analysis of the compound." Qualitative analysis is the name given to the process whereby the identities of the elements or substances present in a sample are determined through the use of simple chemical tests. In this case, chemical reactions will be performed with an alum sample that will confirm the presence of K+, Al3+ and SO42–.
In the first test, alum will be reacted with barium chloride in an aqueous solution,
which will result in the following chemical reaction:

KAl(SO4)2 (aq) + 2 BaCl2 (aq) = 2 BaSO4 (s) + KAlCl4 (aq)

The precipitation of solid material, BaSO4, from a solution of alum upon reaction with
barium chloride is a positive test for the presence of sulfate ion, SO42–. The second test is performed to confirm the presence of potassium and is a flame test. Potassium is volatilized at the very high temperature of a flame (about 1000°C) at which point it imparts a bluish-purple color to the flame. After a few seconds in the flame, the sulfur dioxide will be driven off of your alum sample, and the solid material remaining will consist of aluminum oxides. The third test confirms the presence of aluminum ion and involves its reaction with potassium hydroxide. A wispy, gelatinous precipitate of Al(OH)3 will form upon addition of a small amount of KOH to the aqueous alum solution. Further addition of KOH will cause the precipitate to re-dissolve.
“We should also make sure that we’re really saving some money. We’ll need to
calculate the percent yield of the process once we’ve made the alum,” Diane said.
When carrying out chemical reactions, it is of interest to know what percentage of
the starting material is recovered in the form of the desired product. It is possible to
calculate this percent yield in the following manner:

1. Calculate the number of moles of reactant
2. From the stoichiometry of the reaction, determine the expected number of moles of product. A review of the above reactions reveals that there is a 1:1 relationship between the aluminum containing reactant and the aluminum containing product in all cases. Therefore, one mole of aluminum metal should produce one mole of alum.
3. From the expected number of moles of product, calculate the expected mass of product or the theoretical yield.
4. Calculate the percent yield by dividing your actual yield by the theoretical yield and multiplying by 100.

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Identifying an Unknown Compound - Tests and Observations

Tests and Observations

The following are tests and observations that you should make on your unknown
and on the standards that you characterize:
Observations of the physical appearance of the compound.


Solubility of the compound

 When a solid compound dissolves in a particular liquid, we say that the compound is soluble in that particular solvent. If a compound does not dissolve in a particular solvent, we say that it is insoluble. If some of a compound dissolves in the solvent, we say it is slightly soluble. In quantitative terms, the following ranges for number of milligrams of solid which dissolve
per milliliter of solvent can be used to define the terms:
>30 mg/mL soluble (a significant amount dissolves)
<10 mg/mL insoluble (no detectable amount dissolves)
10-30 mg/mL slightly soluble (a moderate amount dissolves)
These limits are, however, approximate and the solubility of a solid compound is usually determined quickly and easily in the lab in the following manner: Place a "spatula-tip" amount of the compound in a small test tube. You can look at the samples of 0.010 g of solid that are provided in the lab, to get an idea of how much should be used. Add about 10 drops of the solvent to the test tube and agitate the contents by shaking the bottom of the test tube back and forth for a minute or so. Observe whether or not


Density

The density of a compound is an intrinsic property, which means that it is a characteristic property of the compound that can be used to help identify it. As you know from Experiment 1, density can be determined by measuring the mass and volume of a sample and dividing:
d = m/V, where d = density in g/mL; m = mass in g; V = volume in mL
With your teammates, design a procedure for determining the density of a
solid compound. Check this procedure with your TA before beginning.

Acidity/Basicity


Acids are compounds which:
-taste sour
-react with active metals (such as zinc and iron) to dissolve the metal and
 produce hydrogen gas
-react with bases to form water and ionic compounds called salts
-can donate a hydrogen ion

Bases are compounds which:
-taste bitter
-feel slippery or soapy on the skin
-react with acids to form water and a salt
-can accept a hydrogen ion

Both acids and bases are corrosive materials that should never be tasted in the laboratory and should not come into contact with your skin. They have an additional property wherein they will react with certain dyes known as acidbase indicators. A very common acid-base indicator is called litmus. An acid will react with litmus to produce a red color, while the reaction of a base with litmus produces a blue color. Some compounds are neither acidic nor basic and are referred to as neutral compounds. These compounds will not change the color of neutral litmus paper. Thus, a simple and safe test for determining the acidity or basicity of a solid compound is to dissolve it in water and test the resulting solution with litmus paper. Note: Make sure that you have "neutral" litmus paper in your drawer, not "red" or "blue". If you do not have the correct paper, trade it in at the stockroom. The proper technique for testing a solution with indicator paper is to remove a drop of the solution with a stirring rod and place it on a piece of the paper. Do not place the paper directly into the solution being tested. Be sure to rinse off the stirring rod with water before placing it in the next solution you test. Before testing any of the compounds, place a drop of water, a neutral compound, on a piece of neutral litmus paper. The color of the paper in water provides you with a reference against which you can compare the color of the paper that results from other tests. In addition to the litmus test, and especially if the compound is insoluble.

In water, it is useful to observe its behavior in acidic and basic solutions. If the compound is acidic, it will react with a solution of a base, such as sodium hydroxide, NaOH. If it is basic, it will react with a solution of an acid, such as hydrochloric acid, HCl. Evidence that a reaction has occurred can include dissolution of the solid, evolution of a gas (bubbles) or generation of heat. A neutral compound may or may not react with an acid or base. Compounds that have both acidic and basic properties also exist and are called amphoteric.

To test how a compound interacts with an acid or a base, place a “spatulatip” amount in a test tube. Add about 10 drops of HCl solution or 10 drops of NaOH solution to the test tube, and agitate the contents of the tube by shaking the bottom of the test tube back and forth for a minute or so. Record your observations. Using these tests and observations, develop an experimental procedure that will allow you to identify and characterize your unknown compound. Do not use more than 2 g of your unknown compound for the tests. Return the unused portion in the test tube to your TA. Do not contaminate the contents of this test tube. The following resources will be provided for you:

Standards
baking soda
aspirin
corn starch
fertilizer ingredient
caffeine
chalk
caustic lime (corrosive!)
Epsom salt
Borax

Solvents
acetone
ethanol
hexane
don't forget water!
10% aqueous solution of hydrochloric acid, HCl (50 mL per group)
10% aqueous solution of sodium hydroxide, NaOH (50 mL per group)

Note: a 10% solution contains 10 g of the compound dissolved in 90 g of the solvent (water, in this case).
Equipment: 10 mL graduated cylinder with a precision of ±0.02 mL

Precautions:
Hydrochloric acid, HCl, and sodium hydroxide, NaOH, are both corrosive
chemicals. If either of these solutions comes into contact with your skin, rinse it with copious amounts of water. Some of the unknowns are also corrosive. Always assume that an unknown compound is toxic and potentially dangerous and use the proper precautions. Acetone, hexane and ethanol (organic solvents) are all flammable. Keep away from open flames and heat sources!

Questions
1. Describe the procedure that your team followed in performing this experiment and
indicate what your role was in carrying out this procedure (what tasks you
performed and which observations were yours).
2. Write the code number that was on your unknown. What is the identity of your
unknown? What results prompted you to arrive at this conclusion? List the
characteristics of your unknown.
3. Discuss sources of error associated with the procedure your team used for
determining the density of a solid compound and the effects these errors might
have on the results

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Identifying an Unknown Compound

Objective: In this experiment, you will identify an unknown, solid, white compound by
comparing the results of careful observations and tests that you make on the unknown to observations and tests made on a series of standard compounds. You will work with a team to characterize the standards.






Introduction
When Jason’s great uncle Frederick passed away, his parents asked him to help out by sorting through the items that Uncle Fred had stored away in his attic. In a dusty corner behind an old chair, Jason found an antique apothecary chest with nine stoppered bottles inside. Each bottle contained a white powder, however, all the labels had fallen off of the bottles. Jason found the labels in a small pile at the bottom of the chest. Excited about his find, Jason immediately made plans to take the apothecary chest to “The Antiques Roadshow”, which was due to visit nearby Providence in three weeks, to have it appraised. He felt that it would make a more impressive showing if the labels were actually affixed to the bottles. Being a perfectionist, he wanted to put the correct labels on the correct bottles, but all of the powders looked so similar that their appearance did not provide a clue as to their identities. The names on the labels were relatively common materials that he could find in his own home or easily obtain in a pharmacy, grocery or hardware store: baking soda, aspirin, corn starch, an ingredient found in fertilizer, caffeine, chalk, caustic lime, Epsom salt and borax. He decided to perform some tests to identify each of the nine substances.
The experiment In today's experiment, you and the members of your group will perform the same tests that Jason has to carry out in order to solve his dilemma. The nine "known" compounds, or standards, will be provided for you. Work in a group of three or four people to fully characterize these standard compounds. Each student should characterize two or three standards. In addition, each student will be given a test tube that contains a solid, white compound that is one of the nine substances found in Uncle Fred's chest. You are responsible for identifying and fully characterizing this unknown compound. Record the code number of your test tube in your lab notebook.

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