Project 11: Identification, Properties, and Synthesis of an Unknown Ionic Compound
Chemistry 1011 Laboratory
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information, thoughts, and ideas are my own, except as indicated in the references.
Discussion of Results:
The goal of this project was to identify and verify the identity and composition of an
unknown ionic sample. During the first week, we performed a variety of tests to conclude that
we believed our substance to be Magnesium Sulfate. To explain that conclusion, one must look
at the results of our tests. The first test that we did was a qualitative solubility test in which our
substance dissolved in water but did not dissolve in acetone or toluene. From this test, we
concluded that our substance was polar, most likely ionic, and slightly acidic. These results
narrowed our choices down slightly before going into the next phase of testing.
Our subsequent test was a cation test in which we initially tested for ammonium using a
solution comprised of a solution of our compound and Sodium Hydroxide (NaOH). This test
returned negative, so we proceeded to conduct a flame test to determine which cation was
present in our sample. Because we did not observe a change in color, we concluded that our
substance was a compound that contained magnesium, which narrowed our choices down
significantly.
The next tests that we conducted were the quantitative solubility test and conductivity
test. These tests led us to the conclusion that our substance was highly soluble (1g/0.01L) as well
as conductive in water (with a conductivity of 0.035). These tests were important because the
results of both the conductivity test and solubility test helped confirm that our substance was an
ionic compound. If we had determined that our compound did not conduct electricity or was not
made up of ionic bonds, it would refute some of our other conclusions and our results would no
longer be meaningful. The fact that our substance conducted electricity and was highly soluble in
water helped to confirm our earlier results.
Our final test for the first week was a series of anion tests, once we had narrowed down
the possibilities to magnesium chloride or magnesium sulfate. In these tests, our substance tested
negative for a chloride ion but tested positively for a sulfate ion, which led us to the conclusion
that our substance was magnesium chloride because our substance formed a precipitate when
mixed with a solution of HCl and BaCl2. All of the data from week one can be found on Table 1.
Our goal for week two was to perform a series of reactions with our “unknown”
compound and a substance that we knew to be magnesium sulfate. Our results found that our
sample reacted in the same way as magnesium sulfate in each of the reactions. Both samples that
we tested (the known and unknown, that is) formed precipitates with Sodium Hydroxide
(NaOH), Barium Chloride (BaCl2), and Lead Acetate (Pb(C2H3O2)2), and not with Sodium
Acetate (Na(C2H3O2)) or Hydrochloric Acid (HCl). We chose these solutions to compare how the
compounds reacted with acids, bases, an inorganic compound, and multiple salt solutions.
Compared to the known sample, our sample reacted exactly the same way, which helped us to
confirm that our substance was, in fact, magnesium sulfate because our sample performed the
same as the known sample of magnesium sulfate. This data can be found on Table 2.
Our goal for week three was to determine the percent composition of our sample as well
as the percent composition of a known sample of magnesium sulfate using processes such as
gravimetric analysis and stoichiometry. After dissolving three samples of our unknown
compound into beakers of water, we used Lead Nitrate (Pb(NO3)2) to form a precipitate and used
centrifugation to separate out a precipitate. We then weighed the precipitates (the data for which
can be found on Table 3) and used the following calculations to determine the average percent
composition:
.374 g PbSO4 * [1 mol / 303.26 g] = .0012 mol PbSO4
1:1 ratio → .0012 mol SO4
Mass of SO4 = 32.066 + (15.999 * 4) = 96.069 g/mol
.0012 mol SO4 * [96.069 g / mol] = .115 g SO4
[m / .5] * 100 = 23.1% Sulfate/Target Ion
100 - 23.1 = 76.9% Magnesium/Counter Ion
We then repeated the experiment two additional times with samples of the known
compound, magnesium sulfate, and repeated those calculations:
.347 g PbSO4 * [1 mol / 303.26 g] = .0011 mol PbSO4
1:1 ratio → .0011 mol SO4
Mass of SO4 = 32.066 + (15.999 * 4) = 96.069 g/mol
.0011 mol SO4 * [96.069 g / mol] = .106 g SO4
[m / .5] * 100 = 21.2% Sulfate/Target Ion
100 - 23.1 = 78.8% Magnesium/Counter Ion
These consistent results confirmed, once again, that our unknown compound was, in fact,
magnesium sulfate. Compared to the known sample, our unknown compound reacted similarly in
acid, base, double displacement, and salt solution reactions and had an extremely similar percent
compositions. All of these properties allowed us to conclude that our “unknown” sample was, in
fact, magnesium sulfate.
Scientific Explanation of Claims/Support:
Several claims were made in the results of this experiment. One of the first conclusions
that our group came to was the fact that our substance contained magnesium, based on the results
of the flame test. The “flame test” works by exciting electrons found in ions and metals using
intense heat to produce an emission spectra unique to each substance. By observing the
wavelength of light given off by a given substance, it is possible to determine the presence of a
specific metal within a compound. Because our substance did not produce a color in the flame
test, we were able to eliminate each of the other choices and conclude that our compound must
have magnesium.
Another claim that requires explanation is our qualitative solubility testing. We used
water, acetone, and toluene as solvents because making those solutions would give us a good
idea of the overall structure and behavior of our compound. We used water because it is polar
and it is a universal solvent, we used toluene because it is a strong solvent that is only slightly
polar, and we used acetone because it allowed us to determine whether or not we were dealing
with an organic compound. Because our compound dissolved in water and not in toluene, we
were able to conclude that the substance was polar and because it did not dissolve in acetone, we
concluded that our substance was not an organic compound. Additionally, according to the pH
test that we conducted, our substance was very slightly acidic, which helped to narrow down the
possible choices.
One important aspect of note involving these experiments is that we made sure to do each
of the quantitative experiments with multiple trials to ensure that our results were reasonable and
repeatable. It is easily possible to find an outlier in an experiment and conducting multiple trials
allows for possible errors and makes it easier to find more exact and repeatable values.
Gravimetric analysis is the process of analyzing percent composition by comparing the
masses of precipitates formed by known masses of a compound. Our week three experiment used
gravimetric analysis to compare the masses of precipitates formed to find the percent
composition of the known and unknown compounds. Before separating the precipitate out, we
first had to use a solution that would form a very strong precipitate, in which we used Lead
Nitrate (Pb(NO3)2). We chose lead nitrate because it is soluble in water, it reacts with a
double-replacement reaction with our compound, and because it forms a very strong precipitate
with magnesium sulfate. Once the precipitates stopped forming, we used a centrifuge to separate
the precipitates out from the rest of the solutions, though we also could have used vacuum
filtration. We chose to use centrifugation because though both processes would serve our
purpose, centrifugation was simply easier to set up. Once we had separated the precipitates out of
the solution and scraped them out of the cuvette, we heated the precipitates in an oven to ensure
that any moisture did not affect our results. Once they had been in the oven for around ten
minutes, we weighed the samples and performed calculations to complete our gravimetric
analysis.
Possible Errors:
There were several instances where a potential error could have been made. First, there is
always a possibility of human error contaminating the data. Specifically, any errors in
measurements or calculations are possible, particularly when dealing with extremely small
quantities. In addition, there is certainly a possibility that some cuvettes or beakers may have had
other trace chemicals or particulates left inside that could alter our results. This is the most likely
scenario in which an error could have been made with this project because not all glassware had
been thoroughly cleaned.
Another instance where an error could have been made comes from the lack of
standardization or precise measurement, particularly with the week three experiment. Because
each cuvette may have had slightly different volumes and some volumes of precipitates may
have been more difficult to scrape out than others, it is possible that an error could have been
made in those experiments. To try to prevent or counteract this potential error, we used multiple
trials and averaged the masses to hopefully find an estimate for the exact value for the mass of
the precipitate formed.
References:
1. Burke, John. "Solubility Parameters: Theory and Application." The Book and Paper
Group Annual, vol. 3, 1984,
https://cool.conservation-us.org/coolaic/sg/bpg/annual/v03/bp03-04.html.
2. Clark, Jim. "Flame Tests.", Oct 15, 2019,
https://chem.libretexts.org/Bookshelves/Inorganic_Chemistry/Supplemental_Modules_(I
norganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/1_s-Block_
Elements/Group__1%3A_The_Alkali_Metals/2Reactions_of_the_Group_1_Elements/Fla
me_Tests.
3. Cooper, Melanie M. "Project 11: Identification of an Unknown Ionic Salt." Cooperative
Chemistry: Laboratory Manual. McGraw-Hill Higher Education, 2009.
4. Yoder, Claude. "Gravimetric Analysis.", 2019,
http://www.wiredchemist.com/chemistry/instructional/laboratory-tutorials/gravimetric-an
alysis.