Preparation and Analysis of Alum1
Authors: D. L. McCurdy, V. M. Pultz and J.
Last Update: August 19, 2013
One of chemistry’s goals is to be able to transform any
set of substances (the reactants) to another set of substances (the products)
through a chemical reaction. As we have discussed in class, there are rules,
such as the Law of Conservation of Mass, by which chemical reactions occur, and
it took chemists a long time to understand these basic rules. Even though we
know a great deal about chemical reactions, chemists are still finding new
chemical reactions and new ways of assembling atoms into molecules and molecules
into more elaborate structures. In this and the next laboratory exercise you
will learn some of the basics of how chemists carry out chemical reactions and
how they characterize the chemical substances involved in these reactions.
To fully describe a chemical reaction one needs to know
the identities of both the products and the reactants, and the proportions in
which the reactants combine and the products form. While it may seem a trivial
exercise to identify the reactants, this is not always the case. Needless to
say, identification of the reactants in a complex reaction mixture can be very
difficult, and so we will only work with chemical reactions where the reactants
The description of a chemical reaction consists of a
series of steps: 1) carrying out the reaction, 2) isolating the product(s), 3)
purifying the product(s), 4) and characterizing the product(s) and determining
its(their) purity. The isolation and purification of the products are based on
physical properties such as the ability to form crystals, boiling point, melting
point, solubility, etc. Characterization of the products may be either
quantitative or qualitative. In a quantitative characterization, the chemical
formula and the structure (i. e., how the atoms are connected) are determined.
The former is usually accomplished using elemental analysis, mass spectroscopy,
X-ray crystallography or some spectroscopic method. Sometimes it is sufficient
to show only that certain ions or elements are present in a sample, and in this
case a chemist will perform a qualitative test. Qualitative tests often use
chemical reactions that result in a visible change (formation of an insoluble
solid, a color change, or evolution of a gas) as a way to quickly show whether a
particular chemical species is present or not.
Once the chemical reaction’s products are fully
characterized, and the balanced chemical reaction is known, we can compute a
theoretical and a percent yield. We do these final characterizations of the
reaction because it is important to know how efficiently the reaction converts
reactants to products. Chemists are always trying to strike a balance between
the cost of the reactants, the value of the products, the time a reaction
requires and the cost of any unwanted by-products that must be handled as
hazardous waste. A reaction, even though it gives a valuable product, may be
unusable because it has a low yield, takes too much time or generates too much
In this experiment you will prepare and characterize alum
(potassium aluminum sulfate dodecahydrate, KAl(SO4)2·12 H2O).
The first step in this synthesis, which you will perform during Week 1, is to
react metallic aluminum with a concentrated solution of potassium hydroxide
(KOH) to form the potassium salt of the tetrahydroxoaluminate complex ion,
[Al(OH)4]-. The balanced chemical equation for this
oxidation-reduction reaction is
The second step of the procedure is to convert the KAl(OH)4
to alum by addition of sulfuric acid (H2SO4) in an
acid-base reaction. Under the experimental conditions, the alum has a limited
solubility in water, and so it precipitates from the solution. The balanced
chemical reaction that occurs in this step is
The overall balanced chemical reaction for the conversion
of aluminum to alum, shown below, can be obtained by adding together the
balanced chemical equation for each step (Help
The second and third weeks of this exercise will be
devoted to characterizing the alum. Alum is an ionic compound, which means its
melting and boiling points are likely to be too high to be measured
conveniently. Also, most spectroscopic methods would not yield useful
information. Therefore, we will rely on chemical means to show that we did, in
fact, form alum in our reaction. This procedure duplicates how chemists
characterized chemical reactions until the late 20th century, and in
some cases chemical means of characterization are still the only methods
In Week 2 you will perform qualitative tests to
demonstrate the presence of K+, and sulfate ion (SO42-)
in the alum. You will also perform a quantitative determination to determine the
percent water by mass in alum.
The qualitative test for sulfate uses the insolubility of
barium sulfate (BaSO4). When an aqueous solution of a barium salt
(usually BaCl2) is mixed with an aqueous solution containing sulfate,
a white precipitate of insoluble BaSO4 forms according to the net
Ba2+ (aq) + SO42- (aq) → BaSO4 (s).
A positive test for SO42-, therefore, is the observation
of a white precipitate when an aqueous BaCl2 solution is mixed with
the aqueous test solution.
When placed in a flame many elements give the flame a
distinctive color; an effect that can be used to determine both which elements,
and how much of each one, is present in a sample. Potassium produces a
distinctive lavender flame that we can use as a qualitative test for the
presence of potassium. Potassium's flame is often difficult to see because
sodium, which is often present as an impurity, has an intense yellow flame that
masks other colors. The potassium flame can be seen in the presence of sodium by
viewing the flame through a dark blue cobalt-glass filter, which absorbs the
yellow light from Na, but allows the light from K to pass. When placed in a
flame, aluminum does not change the flame’s color, and so a visual flame test
cannot be used to show the presence of Al.
Alum is a hydrate, which means that it is a compound that
has water molecules trapped within the solid. Hydrates will release some, or
all, of their “waters of hydration” upon heating. If the chemical reaction
between Al and KOH does produce alum as a product, we would expect that heating
the product should result in a decrease in the sample’s weight corresponding to
the loss of 12 water molecules per formula unit of alum. Thus, if one knows the
starting mass of alum, and the amount (mass, and therefore number of moles) of
anhydrous alum remaining after all of the water has been driven off, one can
calculate the amount of water that was present in the alum (by the Law of
Conservation of Mass). A comparison of the experimentally determined waters of
hydration and the number expected from the chemical formula can then be used as
evidence for the formation of the desired product. The process by which the
waters of hydration are driven off is described by the chemical equation shown
below, where the "Δ" written above the arrow indicates that heat was applied to
A quantitative analysis for Al3+ will be made
in Week 3. Normally, Al3+ is colorless, which means that it does not
absorb light in the visible portion of the spectrum. So, we will add a dye
called aluminon that will react with the Al3+ in solution to give a
colored complex ion. For a sufficiently dilute solution, the amount of light
absorbed by a chromophore (a chemical species that absorbs light) present
in the solution is given by Beer’s Law, A = ε•b•C, where
A is the
absorbance (how much light the sample absorbs compared to a solution that
does not contain the chromophore), ε is the molar absorptivity (also
known as an extinction coefficient; ε
depends on the compound and the wavelength of light), b is the pathlength
(how much sample the light must pass through) and C is the concentration of the chromophore (More
Info). According to Beer’s Law the intensity of the color is linearly
dependent on the amount of aluminon-Al3+ complex present. So if we
knew ε for the complex ion
formed between Al3+ and aluminon, we could make a single absorbance
measurement and know the [Al3+] in a solution, and therefore, how
much Al was in the original alum sample. Unfortunately, this is neither a
precise nor accurate way to make this determination. It is imprecise because it
is only a single measurement, and it is inaccurate because 1) we don’t know the
stoichiometry for the reaction between aluminon and Al3+ and 2) the
commercially available dye is not pure (ε cannot be determined). So, we need a
way to increase the method’s precision and to overcome the problem with
The problems with the colorimetric method are solved by
using a calibration curve, which gives the relationship between absorbance and
concentration. A calibration curve is constructed by preparing samples with
known concentrations of the analyte (in our case, Al3+) and then
measuring the absorbance of these samples. If Beer’s Law holds, a calibration
curve is a straight line, for which we can obtain an equation from a regression
analysis. Now if we measure the absorbance of a sample containing an unknown
amount of analyte, it becomes a simple matter of substituting this value into
the equation for our calibration curve and solving for the concentration.
Because more than one measurement was used to construct the calibration curve,
we improve our precision. A calibration curve also improves accuracy because
only the analyte's concentration changes (everything else, such as the
stoichiometry between aluminon and Al3+ and the dye’s purity, is
When you set up your laboratory notebook for this exercise
treat each week of the exercise as a separate experiment. So, each week will
have its own title, statement of purpose, etc. Note that some of your results
will actually be determined during a subsequent week. Be sure to carefully read
the experimental procedure and be aware that there are a number of potential
hazards. Also, there are several places in this exercise where you will be
waiting for something to happen. You can substantially shorten your time in lab
by working on another section of that week’s exercise during these times. Also,
be sure that you have completed all of the calculations for a given week’s work
before coming to laboratory. If you do not come to laboratory prepared, you will
not be able to complete the Week 2 and Week 3 exercises in the allotted
Week 1 (
Synthesis of Potassium Aluminum Sulfate Dodecahydrate
Obtain a piece of aluminum foil weighing about 0.5 g and
weigh it precisely (to the nearest 0.001g). Cut the weighed foil into many small
pieces. The smaller the pieces the faster the reaction will go because of the
increased surface area exposed to the KOH solution.
Place the small pieces of aluminum in a 100-mL beaker. Add
enough hot water to a Styrofoam cup so that, when the 100-mL beaker is placed
inside, the beaker is completely surrounded by water, but the water does not
spill out of the cup or into the beaker. If water from the hot water bath spills
into the beaker there will be a drastic decrease in the yield of alum.
Place the 100-mL beaker containing the aluminum into the
hot water in the Styrofoam cup and transfer everything to the hood. Slowly and
carefully add 25 mL of the 1.4 M KOH solution to the aluminum. CAUTION!
No open flames can be present in lab while the reaction between KOH and Al is
taking place. Stir the solution with your glass stirring rod and cover it with a
watch glass. Repeat the stirring every few minutes until all of the aluminum
dissolves. If the reaction slows down, replace the water in the bath with fresh
hot water. If the reaction becomes too vigorous, remove the beaker from the
water bath until the reaction subsides. CAUTION! Avoid inhaling the gas
evolved during this reaction. The gas is not toxic at this concentration, but a
fine mist of the corrosive KOH solution is formed by the gas evolution.
When the aluminum has completely dissolved (do not be
concerned if the solution appears cloudy or contains black specks), gravity
filter the reaction mixture into a 50-mL beaker through fluted filter paper (the
instructor will demonstrate). Dispose of the used filter paper in the laboratory
garbage. CAUTION! The filter paper will be wet with the corrosive KOH
solution. So, wash your hands after handling the wet filter paper.
Obtain approximately 5 mL of 9 M H2SO4
in your 10 mL graduated cylinder. Use a plastic pipet to slowly and carefully
add the H2SO4 solution to the 50-mL beaker containing your
filtered solution. Do not dip the pipet into the filtered solution!
Continue adding the H2SO4 solution until no more
precipitate forms. This should require no more than about 5 mL of the H2SO4
solution. Do not add too much H2SO4, or your yield will
Info). After the H2SO4 addition, carefully stir the
new mixture with your stirring rod and record your observations. CAUTION!
The H2SO4 solution is very corrosive and the reaction
between the H2SO4 and KOH is very exothermic (gives off
Prepare a cold bath. Place the 50-mL beaker with the
filtered reaction solution in the ice bath. Do not introduce any of the water
from the ice bath into the beaker. Also place a test tube containing 15 mL of
95% ethanol in the ice bath. The ethanol solution will be used to wash residual
H2SO4 from the alum crystals.
After a crop of crystals has formed, set up a vacuum
filtration apparatus (Help
Me). Do not under any circumstances push the rubber tubing more than 1/4” on
to the side-arm of the filter flask and do not leave the tubing attached to the
flask while the flask is unclamped.
While the vacuum is on, carefully remove some of the
supernatant (the solution above a solid) from your crystals using a pipette
and wet the filter paper. This will help the paper adhere to the filter and
prevents leaks. CAUTION! The solution is corrosive. Remove the 50-mL
beaker from the ice bath, swirl it gently to suspend the crystals and pour it
into the Büchner funnel. Use your glass stir rod to remove any crystals that
adhere to the side of the beaker. Once the aqueous solution has been filtered
completely (leaving the crystals on the filter paper), place 2 - 3 mL (the
plastic pipets hold about 3 mL) of the cold ethanol solution in the 50-mL
beaker. Use your glass stirring rod to loosen any remaining solid that clings to
the side. Swirl to suspend any crystals remaining in the beaker, and pour the
suspension into the filter. Once the ethanol has been filtered away, repeat this
washing several times. After the last ethanol wash, allow the vacuum to run for
a minute or two to draw air through the crystals to help them dry.
After the ethanol solution has stopped draining from the
funnel, inspect the product. If it looks dry, gently prod it with your metal
spatula. If it is dry enough to remove from the filter, the solid will not be
very sticky and will have the consistency of fine sand. Break the vacuum by
removing the vacuum hose from the side-arm of the filter flask, and then turn
off the aspirator. Transfer the solid and the filter paper from the funnel to a
pre-weighed watch glass with the help of your metal spatula, as your
instructor will demonstrate. Carefully scrape any alum that adheres to the side
of the Büchner funnel onto the watch glass.
If the alum is dry, the filter paper will separate from
the crystals and you can remove the filter paper. Gently scrape any
crystals adhering to the filter paper onto the watch glass. If the alum is still
too wet, leave the filter paper and remove it next week.
Obtain the mass of the wet alum. You will need to have
about 2 g of wet alum (3 g if the mass includes the filter paper) so that you
will have enough for the next two weeks. If you don’t have enough, collect the
second crop of crystals and/or redo the synthesis. Keep the crystals from
different crops and syntheses separate. Cover the container holding the crystals
with a piece of paper towel, and place it in your drawer to dry.
You may notice that more crystals formed in the filter
flask during the washings. This second crop of crystals may also be collected,
but if you choose to collect these crystals, they should be kept separate from
the main crop. It is always good laboratory practice to keep different crops of
crystals separate until the identity and purity of each crop is determined
(second crops almost always contain more impurities than the first crop and the
time needed to purify them sometimes far outweighs the additional yield).
Collect the second crystal crop by vacuum filtration; wash with several small
portions of the cold ethanol solution and dry, as described above.
Before doing anything else in the laboratory obtain the
mass of each crop of alum to the nearest milligram (three decimal places). Make
observations on the crystalline product (color, texture, etc.), and record all
of your observations in your laboratory notebook. Share your results with your
Qualitative Chemical Tests
Perform the following qualitative tests for SO42-
and potassium on your sample. If you collected a second crop of alum crystals,
you should perform the sulfate and potassium qualitative tests on both the first
and second crops (are your two crops qualitatively the same?).
Place a few crystals of your alum in a 6” test tube. Add
distilled water dropwise while stirring until the alum dissolves. Add one drop
of 0.5 M BaCl2 (barium chloride). Record your results. Does alum
Potassium Flame Test
The instructor will demonstrate the proper techniques for
using the Bunsen burner and heating the needle. In the hood, heat the provided
needle in the flame to remove impurities. Once the needle is clean, carefully
scoop up a small amount of alum on the end of the hot needle. Place the alum in
the flame and heat it until the crystals begin to melt and the solid glows. Note
the color of the flame. If your flame is bright yellow (indicating the presence
of sodium), try cleaning your needle again, or use the cobalt glass filter. Does
this sample contain potassium?
Quantitative Determination of Waters of Hydration
Before beginning this section be sure that your alum
sample is powdered and that you have weighed your alum sample!
Set up a ring stand, ring clamp and porcelain triangle, as
your instructor will demonstrate.
Clean your crucible by placing a few drops of 1 M NH3
solution in the empty crucible and scrubbing with a paper towel. CAUTION!
This ammonia solution has a strong odor and is corrosive. Rinse the crucible
with distilled water and place the empty crucible on the porcelain triangle
supported by a ring and ring stand.
With majority of the flame remaining below the bottom of
the crucible, heat the crucible until its bottom glows a dull red. After heating
for five minutes, remove the flame and let the crucible cool to room temperature
on the triangle. CAUTION! Do not touch the crucible with your hand. It is
extremely hot and will remain hot for several minutes. Remember that a hot
crucible looks exactly the same as a cool crucible. When cooled, you can move
the crucible to the bench top using the crucible tongs. Do not set a hot
crucible on the bench top, because the temperature differential may cause the
crucible to shatter. Once you have cleaned the crucible, it is important that
you handle it only with the crucible tongs. This prevents burns and will
eliminate a systematic error caused by the weight of your fingerprints.
Weigh the cooled crucible (and its cover) to the nearest
milligram (three decimal places) and record this mass in your notebook. If the
balance does not show three decimal places, notify the instructor. Place about
1.0 g of your alum sample in the crucible. Obtain the mass of the crucible, its
cover, and the alum to the nearest milligram and record this in your notebook.
Return the crucible to the porcelain triangle and set the
cover slightly ajar so that the water vapor can escape. For the first few
minutes gently heat (only the light blue portion of the flame touches the
crucible) the crucible by holding the Bunsen burner off to the side. Take care!
The water can violently leave the alum at this point, if it is heated too
strongly. Move the Bunsen burner such that the tip of the inner blue cone is
approximately 3 cm below the crucible. Heat until the crucible glows red and
continue heating for 10 minutes. If at any time you observe white smoke being
given off, or smell an acrid odor, discontinue heating immediately (the sulfate
is being decomposed to SO2).
Remove the heat and completely cover the crucible with the
lid. Cool the crucible to room temperature on the triangle (this takes about ten
minutes). Weigh the cooled crucible (including its cover and the contents) to
the nearest milligram (three decimal places). Using the tongs, move the crucible
and contents back to the triangle and repeat the heating step for 10 minutes.
When this heating step is over, cover the crucible and allow it to cool on the
triangle to room temperature, and then reweigh the crucible, cover, and its
contents. Record this second mass in your notebook. If the second mass is
within a 50 mg of the mass after the first heating, then you have driven off all
of the water. If the masses are not within 50 mg, then repeat the heating
procedure until two subsequent masses agree.
Once you have made your final weighing, invert the
crucible and the anhydrous alum should fall out. If it does not, add some water
from a squirt bottle and use your metal spatula gently to dislodge it. The
anhydrous alum may be disposed of in the trash or in the sink with plenty of
water. Rinse the crucible with distilled water and dry it before returning it to
Before coming to the laboratory you must have completed
the following: 1) prepare a table, like Table 1, in your notebook's Results
section in which to write your data for the calibration curve, 2) set up the
calculations to calculate the [Al3+] in Table 1 (the number of
significant figures in each volume is given in Table 1 and in Table 2), 3)
prepare an Excel spreadsheet to graph the calibration curve (save on your Y:
drive or a flash drive), and 4) familiarize yourself with the
instrument before laboratory; your instructor will review spectrometer
operations before you begin work (click
here for operating instructions).
Volume of Al3+ Stock Solution Used
of Solution (mL)
at 525 nm
Table 1. Example of a table that could be used to
present the data for the calibration curve.
Your instructor will demonstrate how to prepare solutions
using volumetric glassware and will review the protocols for using the balances.
Colorimetric Determination of Aluminum2,3
Preparation of the Aluminum Stock Solution
Precisely weigh out (to the nearest milligram) about 0.1 g
AlCl3·6H2O using an analytical
balance. Quantitatively transfer this solid to a 100-mL volumetric flask
(assume the flask's volume is 100.0 mL). Add about 10 mL of distilled water and
swirl to dissolve the AlCl3·6H2O.
If the solid does not dissolve, carefully add small amounts of distilled water,
swirling between each addition, until it does. Add distilled water to bring the
level of the solution in the flask to the mark on the neck (this procedure is
called "diluting to the mark"). Mix thoroughly by stoppering the flask, and
then inverting and shaking the flask. Repeat if necessary.
Pipet 3.00 mL of the aluminum solution that you just made
into a 25-mL volumetric flask. Dilute to the mark and mix thoroughly. This is
the aluminum stock solution that you will use to construct the calibration
Construction of the Calibration Curve
Number five 50-mL volumetric flasks 1 to 5. Do not add
any of the aluminum stock solution to flask 1. To flask 2 add 1.00 mL of the
aluminum stock solution; to flask 3 add 2.00 mL; add 3.00 mL to flask 4 and 5.00
mL to flask 5. These measurements must be precise, and so you must use
To each flask then add 20 mL of the acetate buffer
solution and 5 mL of the aluminon solution (in that order!) and swirl gently to
mix. These volume measurements do not need to be highly precise. So, you can use
your 50-mL and 10-mL graduated cylinders here. Dilute all the solutions to the
mark by adding distilled water and mix thoroughly. Allow the solutions to sit
for 20 min while monitoring the solutions’ colors. Note any changes in your
spectrometer's operating instructions to ready the instrument for use.
Fill the cuvette with the buffer solution to use as a blank (IMPORTANT!
you must use the same cuvette for both the blank and for your samples). Remove any bubbles by
gently tapping the cuvette with your finger. Under absolutely no circumstances are
you to tap a cuvette on a table top. Do not handle the
cuvette by the clear window (your fingerprints will cause an error in the
measurement). Before placing the cuvette into the spectrometer, be sure to
thoroughly wipe the clear sides of the cuvette with a Kim-Wipe (do not
use a paper towel). When placing the cuvette in the spectrometer, be sure
that clear sides are aligned with the light beam and that the cuvette is placed
in the spectrometer the same way every time. The major sources of error when
using these spectrometers come from poor technique, and you can avoid these by
following these guidelines every time you make a measurement with the
A representative spectrum of a solution
of the Al3+:aluminon complex is shown in Figure 1. The spectrum
should exhibit a broad peak near 525 nm. If the shape of your spectrum looks
dramatically different than that in Figure 1, consult your instructor.
Measure the absorbance at 525 nm (More
Info) for solutions 1 through 5. Graph the absorbance at 525 nm as a
function of [Al3+] in Excel (Help
Me) and perform a linear regression of the data by inserting a trendline (Help
Me) on the graph. Show your graph to your instructor; once he or she has
approved it, you may proceed to the next section.
Figure 1. Representative absorbance spectrum for a
dilute solution of the Al3+/aluminon complex in acetate buffer.
Determination of Aluminum in Alum
Precisely (to three decimal places) weigh out about 0.2 g
of your alum. Quantitatively transfer, as your instructor will demonstrate, to
a small beaker and add distilled water to bring the volume to about 15 mL.
Place the beaker on a hot plate in the hood, cover with a watch glass and heat
to boiling, stirring occasionally with your glass stirring rod. After stirring
rinse the glass rod into the beaker with a small amount of distilled water from
your wash bottle. While the mixture is heating, clean and dry (exterior only)
your volumetric flasks. Also prepare for a gravity filtration directly into the
100-mL volumetric flask using a long-stemmed glass funnel.
Remove the beaker from the hot plate just as the solution
starts to boil. CAUTION! The beaker, watch glass and the hot plate's top
are all hot. Your instructor will demonstrate the safe way to remove the beaker
from the hot plate. Immediately pour the hot solution into the funnel. As the
solution is being filtered rinse the beaker, the bottom of the watch glass and
the stirring rod each with several small washes of distilled water into the
funnel. When the solution is completely filtered, remove the funnel from the
volumetric and rinse the volumetric's neck with several small portions of
distilled water. The volumetric should now be cool to the touch, but if it is
not, wait until it is. Dilute the solution in the flask to the mark. Transfer
3.00 mL of this solution to the 25-mL volumetric flask and dilute as before.
Pipet 3.00 mL of the alum solution that you just prepared
into a 50-mL volumetric. Add 20 mL of acetate buffer and 5 mL of aluminon
solution using graduated cylinders and then dilute to the mark with distilled
water. Wait 20 min and measure the absorbance at 525 nm, as you did for the
other solutions. Record this value in your notebook.
Results and Analysis
From the amounts of the reactants that you actually used,
calculate the theoretical yield of alum (Help
Calculate the percent yield of alum from the theoretical
yield you determined last week and the amount of alum that you actually
obtained. Share your percent yield with your classmates.
Determine the percent water by weight in alum and the
number of waters of hydration in the alum. Share these numbers with the rest of
the class, as instructed. Perform a Q-test (Help
Me) on the class data, and discard the discordant datum, if there is one.
From the class data, calculate the average percent water by weight in alum, the
standard deviation (Help
Me) of the data and determine the confidence limits (Help
Me) at 95% confidence. Based on the known formula for alum, determine the
expected value of the percent water by weight in the sample. Calculate a percent
error for the class average and for your result. Record all data and calculated
results in your notebook. You may do the calculations in Excel, and if you do,
you will need to paste copies of your output in your laboratory notebook.
Calculate the concentration of the aluminum stock solution
(the solution that you had after the second dilution) and the concentration of
aluminum in each of the solutions that you prepared from the stock solution.
Write these values in your table (Table 1, above) in the Results section
of your notebook. In your calculations assume that the volumes of the flasks and
pipets are as shown in Table 2.
Table 2. Nominal volumes of the volumetric
glassware used in this exercise.
From the equation for the best-fit line for the absorbance
at 525 nm as a function of [Al3+] determine the percent aluminum by
weight in alum (Help
Me) and share your results with the class. Perform a Q-test (Help
Me) on the class data and then calculate the average percent Al by weight in
alum, the standard deviation (Help
Me) of the data and finally find the confidence limits (Help
Me) at 95% confidence. Determine what the true percent Al by weight is for
alum and then calculate a percent error for the class average and for your
result. Record your calculations in your notebook, as you did for the Week 2
calculations, and include any spreadsheet output.
The first week of this exercise was a synthesis.
Therefore, your Discussion of Conclusions section for
this week should follow the
synthesis outline. Note that you will not be able to discuss your results
for Week 1 until after you have obtained the mass of your product and done the
qualitative tests on it. It is advisable to reserve two or three pages in your
notebook for the Week 1 Discussion of Conclusions and Error Analysis when
you prepare for Week 2. One important question that you will need to address in
your Discussion of Conclusions section for Week 1 is why is your
percent yield of alum less 100% with specific references to what you did and
Weeks 2 and 3 are both
measurement exercises. In the Week 2 Discussion of Conclusions and Error
Analysis you should include a brief discussion of the qualitative test
results. In both Week 2 and Week 3 you gather evidence for the identity and
purity of your alum. So, you must include a short discussion of whether your
quantitative results support your purported synthesis of alum. Although a
propagation of error analysis is possible, we won't perform one here. However,
you should be able to identify where your major sources of uncertainty are and
qualitatively discuss how they affected your results.
Summary of Results
Use Table 3 to report your results for Week 1.
Mass of Al Used (g)
of Alum (g)
Mass of Alum
Yield of Alum
Table 3. Summary table for the first week.
Summarize your results for Week 2 using Tables 4 and 5.
In the second column of Table 4, write either "positive", or "negative", as
appropriate. Don't forget to report the confidence interval on the class data
in Table 5.
|Test for potassium:
|Test for sulfate:
Table 4. Summary table for the qualitative tests.
of Alum (g)
Mass of Anhydrous
Mass in Alum
% Water by Mass
Number of Waters
Table 5. Summary table for quantitative
determination of water in alum.
Table 6 should be used to summarize the results for the
third week's work. Remember to include the confidence interval on the class
average % Al by mass in alum.
Mass of Alum Used (g)
Slope of Calibration Curve
Intercept of Calibration Curve (AU)
Absorbance of Alum Solution (AU)
Mass in Alum
% Al by Mass
Table 6. Summary table for Week 3.
- 1. Click here to download this file in PDF format
(link not yet active).
- 2. Smith, W. H.; Sager, E. E. and Siewers, I. J.
Anal. Chem. 1949, 21, 1334-1338.
- 3. Marczenko, Z. Spectrophotometric Determination
of Elements; Ellis Horwood Ltd.: Chichester, England, 1976, p. 116-117.