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Coordination Compounds
Introduction - Part I
The majority of compounds studied in the laboratory to date have either been simple ionic salts
such as KCl or covalent molecules such as cyclohexane, C6H12. One way to understand the
difference between these compounds and the coordination compounds to be studied in this
experiment is to use Lewis acid-base terminology. In coordination chemistry, a pair of electrons
from one atom is donated to an electron deficient atom to form a coordinate covalent bond. The
atom donating the electron pair is acting as the Lewis base, and the atom accepting the electron
pair is acting as the Lewis acid. The product of the interaction can be called a Lewis acid-base
adduct or a coordination compound. The reaction shown in Figure 1. between ammonia, NH3,
and boron trifluoride, BF3, to produce BF3NH3 is a simple illustration. The nitrogen atom in the
ammonia donates a lone pair of electrons to the boron atom in the boron trifluoride to make a
coordinate covalent bond. The chemistry of most transition metals is rich and diverse because of
these coordinate covalent interactions.
Figure 1. The reaction between boron trifluoride and ammonia.
Prior to 1893, the formula for a compound known as yellow luteocobaltic chloride was
written CoCl3.6NH3 as if it were made up of two compounds, ammonia (NH3) and cobalt (III)
chloride (CoCl3). Alfred Werner, now known as the Father of Coordination Chemistry, studied a
series of reactions involving this compound, and determined the correct formula.
CoCl3.6NH3 + HCl (concentrated) → no reaction (at 100°C)
CoCl3.6NH3 +3 AgNO3(aq) → Co(NO3)3.6NH3 + 3AgCl
CoCl3.6NH3 + 1 1/2 H2SO4 (concentrated) → 1/2 [Co2(SO4)3.12NH3] + 3 HCl
CoCl3.6NH3 + 3 KOH → 1/2 Co2O3 + 6 NH3 + 3 KCl + 1 1/2 H2O (at 100°C)
CoCl3.6NH3 + 1 1/2 Ag2O + 1 1/2 H2O (moist) → Co(OH)3.6NH3 + 3 AgCl
CoCl3.6NH3 + 3 AgX → CoX3.6NH3 + 3 AgCl
CoCl3.6NH3 + 3 HX → CoX3.6NH3 + 3 HCl
Werner observed a number of facts in these reactions. He noted that six ammonia
molecules always remained in the product for each cobalt atom. and that the ammonia
didn't react. There was no reaction with concentrated hydrochloric acid. The yellow
luteocobaltic chloride did react with sulfuric acid but the acid didn't break the ammonia
away from the cobalt. Only reaction with KOH destroyed whatever interaction was
holding the together the cobalt and the ammonia. However, Werner noted the chloride in
the yellow luteocobaltic chloride reacted with almost anything. The chloride could be
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precipitated by reaction with silver nitrate. The chloride could be replaced with another
halide (X) by reacting the yellow luteocobaltic chloride with either a silver halide (AgX) or
a binary acid (HX). Finally, Werner observed from freezing point depression and osmotic
pressure studies that the compound was ionic and the van’t Hoff factor (i) for the
compound was close to 4. That meant there were four moles of ions present per mole of
compound. Werner concluded the proper formulation of the compound was
[Co(NH3)6]C13.
Hexamminecobalt(III) chloride, [Co(NH3)6]Cl3, is an ionic compound. The cation is the
complex ion hexamminecobalt(III), Co(NH3)63+, and the anion is the chloride ion. The cation has
a charge of 3+ requiring three chloride ions each with a charge of 1- to form a neutral
compound. The complex cation Co(NH3)63+ is a coordination complex. The Co3+ cation is
electron deficient and is surrounded by six ammonia molecules each capable of donating an
electron pair to the metal ion in a coordinate covalent bond as shown in Figure 2. The structure of
the complex is octahedral, with the metal ion at the center of the octahedron and the nitrogen
atoms at the vertices.
Figure 2. The hexamminecobalt(III) ), Co(NH3)63+, anion.
In theory, any species with a lone pair of electrons can behave as a Lewis base and donate a
pair of electrons to a central metal atom to form a coordination complex or compound. The term
ligand is often used to describe any species that exhibits this kind of chemistry and literally
hundreds of ligands are known. Two examples of neutral molecules that can be ligands are
ammonia and water. Anions such as hydroxide (OH1-), fluoride (Fl-), chloride (Cl1-), bromide
(Br1-), iodide (I1-), and cyanide (CN1-) ions can also behave as ligands. The cyanide ion is
different from the other ligands mentioned so far because it can donate an electron pair from
either the carbon atom or nitrogen atom end of the ion. If the electron pair is donated from the C
end of the ion a M-C≡N interaction is formed. If the electron pair is donated from the N end of
the ion a M-N≡C interaction is formed. All the ligands mentioned so far usually participate in
only one coordinate covalent interaction at a time. The term monodentate is regularly used to
describe ligands which donate only one (mono-) pair of electrons at a time to a central metal
atom. There are other ligands capable of donating more than one pair of electrons at a time to a
central metal atom and they are referred to as polydentate ligands.
Polydentate ligands range from those capable of donating simultaneously two pairs of
electrons to a central atom (bidentate) to those capable of donating simultaneously six pairs of
electrons to a central atom (hexadentate). Two examples of bidentate ligands are the neutral
compound ethylenediamine (commonly abbreviated as "en"), NH2CH2CH2NH2, and the oxalate
anion, C2O42-. The ethylenediaminetetraacetate ion (commonly abbreviated as “edta”) is a
hexadentate ion regularly found in food products. The edta complexes metal ions which make
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their way into the food from metal cooking and storage containers. The metal ions cannot react
with the food causing oxidation (spoilage) when complexed by edta.
Polydentate ligands are also referred to as chelating agents and the complexes formed by
the interaction of chelating agents with a central metal atom are referred to as chelates. These
last terms come from the Greek work chela for claw and is appropriate because of the way the
chelating agent claws on to or wraps around a central metal atom.
A complex salt containing iron(III) and the bidentate ligand oxalate will be prepared in part
2 of this experiment. The formula for the compound can be written KxFey(C2O4)z.3H2O. The
synthesis of the compound can be considered as being in two steps. In the first step, insoluble
iron(II) oxalate is formed from the reaction of a soluble iron(II) salt and oxalic acid:
Fe2+(aq) + C2O42- (aq) → FeC2O4(s)
In the second step, the insoluble iron(II) oxalate is oxidized with hydrogen peroxide (H2O2) in the
presence of excess oxalate ion to form the desired iron(III) complex:
K1+ + FeC2O4 + H2O2 + C2O42- → KxFey(C2O4)z.3H2O + OH1(Note: this reaction equation is NOT balanced.)
The compound will be analyzed in part 2 of the experiment to determine the values of x, y,
and z.
Experimental Procedures - Part 1
Synthesis of an Unknown Coordination Compound of Iron(III)
Be sure to record all weights and observations about colors, color changes,
precipitates, etc., in your laboratory notebook. label each section of your notebook to agree
with the different sections of this experiment. This section should be labeled: Part 1 Synthesis of an Unknown Coordination Compound of Iron(III). Other sections should be
similarly labeled.
Weigh out about 10 g of iron(II) ammonium sulfate hexahydrate, Fe(NH4)2(SO4)2.6H2O,
recording the mass of the sample to three places after the decimal point. Add 15 drops of 3M
H2SO4 to about 30 mL of distilled water in a 250 mL beaker. Warm the solution gently on a hot
plate. Slowly add the sample of iron(II) ammonium sulfate hexahydrate until it is completely
dissolved in the solution. While stirring this solution, add 50 mL of 1 M oxalic acid using a
graduated cylinder.
The precipitation reaction between iron(II), Fe2+, and oxalate ion, C2O42-, to form insoluble
iron(II) oxalate should occur. Carefully heat and stir the mixture almost to boiling on the hot
plate. Remove the beaker from the hot plate and let the mixture cool and settle. The liquid layer
above the solid precipitate should be clear. Separate the liquid layer from the iron(II) oxalate
precipitate by first, gently decanting (pour off) as much of the liquid as possible from the solid.
Discard the liquid and save the precipitate. Then use an eyedropper to remove as much
additional liquid as possible from the precipitate. Wash the insoluble precipitate by adding about
30 mL of almost boiling distilled water to the precipitate in the beaker. Stir the mixture of the
hot water and precipitate for 2-3 minutes and then allow the mixture to settle. Repeat the process
of decanting away the wash liquid followed by removing excess wash solution with an
eyedropper. Wash a second time with hot water, again removing as much of the wash liquid as
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possible. (Remember, in this process, the solid precipitate is the desired product, FeC204,
and is to be saved. Discard the wash liquid.)
To the beaker containing the washed FeC2O4 precipitate, add 18 mL of a saturated (approx.
1.8 M) potassium oxalate (K2C2O4) solution using a graduated cylinder. Warm this reaction
beaker and a large beaker (400 or 600 mL) containing approximately 100 mL of water on the hot
plate to 40-45°C. Turn the hot plate to a low setting and place the reaction beaker inside the
large beaker. Monitor the temperature in the reaction beaker to ensure it stays in the range 4045°C. Increase the setting on the hot plate if necessary. Mix 10 mL of concentrated hydrogen
peroxide, H2O2, and 10 mL of water in a 25 or 50 mL graduated cylinder (Careful with the
concentrated H2O2!). Slowly (over a period of 2-3 minutes maximum) and with stirring, add
the hydrogen peroxide slowly to the warm reaction beaker. (Try adding the solution 1-2 mL at a
time. Wait for the fizzing to stop between additions.) The solution may become basic enough
that some rust colored iron(III) hydroxide precipitates from the reaction mixture. Do not be
concerned, it should be redissolved in the next step. Remove the reaction beaker from the larger
beaker, place it directly on the hot plate, and heat it almost to boiling. Quickly add 9 mL of 1 M
oxalic acid, H2C2O4, to the reaction beaker. Continue to add 1 M oxalic acid dropwise and stir
until a clear solution is obtained. (If precipitate is still present after 15 mL of oxalic acid has
been added, filter the solution hot, discard the precipitate and use the resultant clear solution for
the rest of the experiment.) Add 20 mL of 95% ethyl alcohol, C2H5OH, to the clear solution.
This prevents oxalic acid from precipitating from the reaction mixture because it is more soluble
in an alcohol-water solution than in pure water. If some crystallization has already occurred,
warm the solution gently on a hot plate under the hood until all crystals are redissolved.
Caution: alcohol is flammable and no open flames should be present in the laboratory
when it is in use! Clearly label the beaker with your name and cover it with an inverted watch
glass. Store it as instructed until the next laboratory period.
A mass of green crystals should have formed in the beaker by the beginning of the next
laboratory period. Prepare a suction filter as demonstrated by your laboratory instructor. Use a
rubber policeman to transfer as much of the crystals and solution to the filter as possible. Use the
liquid that comes through the filter (called mother liquor or filtrate) to wash any crystals
remaining in the beaker on to the filter. Prepare 15 mL of a 1:1 mixture of ethyl alcohol and
water. Use this alcohol solution to rinse the crystals in the filter, slowly pouring the solution over
all the crystals.
Wash the crystals in the funnel with three 10 mL portions of acetone. Caution: acetone is
flammable and no open flames should be present in the laboratory when it is in use! Then
allow the crystals to dry in the air. This should take no more than 15-20
Clean, dry and weigh a stoppered container. Transfer the crystals to this container when
they are thoroughly dry. Reweigh the container and calculate the mass of product obtained.
This compound is light sensitive. Wrap the storage vial with aluminum foil when it is not in use
and keep it in a drawer as much as possible.
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Introduction - Part 2
As stated earlier, the formula for the compound synthesized in Part 1 of this experiment has
been written KxFey(C2O4)z.3H2O. The goal of part 2 is to determine the values of x, y, and z.
Both oxalate ion and iron(II) ion can be analyzed by oxidation-reduction titration methods using
potassium permanganate.
Part 2 of this experiment will broken down into three steps. In step A, a solution of
potassium permanganate, KMnO4, will be standardized by titration of sodium oxalate, Na2C2O4.
After the concentration of the KMnO4 solution is known, it can be used in step B to determine
the amount of oxalate in a sample of KxFey(C2O4)z.3H2O. In step C, the iron(III) in a sample of
KxFey(C2O4)z.3H2O will be determined by separating it from the oxalate, reducing it to iron(II),
and then titrating it with the potassium permanganate solution.
Step A. Standardization of Potassium Permanganate Solution
Permanganate ion, MnO41-, oxidizes oxalate ion, C2O42- , in acidic solution as shown in this
reaction equation:
5 C2O42- + 2 MnO41- + 16 Hl+ → 10 CO2 + 2 Mn2+ + 8 H2O
The permanganate ion is an intense purple in solution while all other products and reactants are
colorless. Thus, when permanganate ion is delivered from a burette as the titrant into a flask
containing a sample of oxalate ion, the solution in the flask will remain colorless until all the
oxalate ion has been oxidized and there is a slight excess of permanganate present. The first
persistent trace of violet in the reaction flask marks the endpoint of the titration.
In this step, samples of known mass of Na2C2O4 will he titrated with a solution of KMnO4
of unknown concentration. The mass of Na2C2O4 in a sample can be converted to moles using
the molar mass of Na2C2O4. The mole relationships implied in the balanced reaction equation
above can then be used to convert the number of moles of Na2C2O4 in the sample to the number
of moles of KMnO4 titrant required to reach the equivalence point of the titration. Finally the
number of moles of KMnO4 required for reaction and the volume used can he used to calculate
the molarity (M) of the potassium permanganate solution.
Good laboratory technique is required to obtain good results. There is a second, interfering
reaction between permanganate ion and oxalate ion that can lead to unreliable results. If a high
local concentration of permanganate is allowed to collect, the permanganate can be only partially
reduced in insoluble MnO2, instead of Mn2+.
3 C2O42- + 2 MnO41- + 8 Hl+ → 6 CO2 + 2 MnO2↓ + 4 H2O
The reaction flask should be swirled constantly to prevent this build up of extra MnO41- in
the region where the titrant is entering the solution. The appearance of a brownish color due
to suspended manganese(IV) oxide indicates this interfering reaction has occurred. If the
equivalence point has not been reached, the unreacted oxalate ion in the flask should react with
the MnO2 and complete the conversion of the permanganate to manganese(II).
C2O42- + MnO2 + 4 Hl+ → 2 CO2 + Mn2+ + 2 H2O
These last two reactions can be added together in such a way as to cancel the MnO2, and obtain
the desired reaction completely converting MnO41- to Mn2+. This preserves the stoichiometry,
meaning the titration can be saved. However, if the equivalence point has been overshot, the
titration is ruined and must be discarded.
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A properly titrated solution will be faintly pink at the endpoint. The presence of BOTH
pink and brown indicates a potential problem. A slight brown cloudiness may form in a properly
titrated reaction if it is allowed to stand a few minutes after the titration is completed. This does
NOT invalidate the titration.
Experimental Procedures - Part 2
Analysis of an Unknown Coordination Compound of Iron (III)
Step A. Standardization of Potassium Permanganate Solution
Be sure to record all weights and observations about colors, color changes,
precipitates, etc., in your laboratory notebook. Label each section of your notebook to
agree with the different sections of this experiment. This section should be labeled: Part 2 Analysis of an Unknown Coordination Compound of Iron(III), Step A., Standardization of
Potassium Permanganate. Other sections should be similarly labeled.
Note: It is important that all apparatus be clean and distilled water used in all solutions,
because impurities are frequently oxidized by permanganate.
Obtain and clean a polyethylene storage bottle. After rinsing several times with distilled
water, add 1-2 mL of 6 M HCl, and an equal amount of 3% H2O2 solution. Cap and shake the
bottle so the solution comes in contact with all the surfaces. This step is required to clean it of
any MnO2 (dark deposits). Then, rinse twice with 2-5 mL portions of the permanganate stock
solution. Finally, obtain about 500 mL of the permanganate stock solution which is
approximately 0.2 M. Keep this bottle closed as much as possible.
Make sure the polyethylene squeeze bottle containing distilled water is full. A lot of
distilled water will be used rinsing the walls of reaction flasks.
Weigh out a sample of sodium oxalate of approximately 0.12 g in a 250 mL Erlenmeyer
flask. Record all masses to three significant figures after the decimal point. Add 60 mL of 1.0
M H2SO4 to the flask to dissolve the sodium oxalate.
Clean and rinse a burette. Rinse the burette with three 3-5 mL portions of the
permanganate solution before filling to slightly above the zero marking. Open the stopcock fully
to sweep any air bubbles out of the tip of the burette, and allow the level of solution to drop
below the zero marking. Because the solution is colored, readings are taken with respect to the
top of the meniscus rather than the bottom as is done with clear, colorless solutions. Do NOT
reload the burette in an attempt to get the reading to 0.00. Force yourself to read the
burette to two significant figures beyond the decimal point.
Heat the solution of sodium oxalate in the Erlenmeyer flask to 80-90°C on a hot plate.
Remove the thermometer from the flask before titrating. Be sure to rinse the thermometer with
distilled water into the flask to prevent removing any sodium oxalate. Record the initial burette
reading to two significant figures beyond the decimal point and begin to titrate. Do not add the
titrant rapidly (about 5 drops per second is a good rate of addition), and constantly swirl the
solution. Use one hand to control the stopcock while swirling the flask with the other hand. If
titrant is added too quickly, or the flask is not swirled sufficiently, or the flask is not warm
enough, brown MnO2 may form. If that happens, immediately stop adding titrant to prevent
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overshooting the equivalent point. In a proper titration, the purple titrant is decolorized as it falls
into the hot sodium oxalate solution. As the titration continues, the decoloration will take longer
and the rate of addition of titrant should be reduced until it is essentially drop-by-drop.
Eventually the addition of a single drop of titrant will cause a faint pink color to persist in the
solution. (Place the flask on a piece of white paper to help see the end point. Also a flask of
water makes a good color comparison.) Record the final burette reading, again to two significant
figures beyond the decimal point.
Repeat this procedure until consistent results are obtained. It may be necessary to complete
the calculations to check consistency.
The sulfuric acid used to dissolve the sodium oxalate may contain impurities which react
with permanganate ion, and a blank should be run to take these contaminants into account. Add
about 50 mL distilled water to 60 mL of the 1.0 M H2SO4. Be sure to add the water slowly to
avoid overheating and spattering! Warm the flask to 80-90°C on the hot plate and record the
initial burette reading. Add titrant drop-by-drop until a faint pink color persist. This titration
should only take 1-2 drops. This volume will be subtracted from the volume of titrant used to
obtain the corrected volume that will be used in the calculation of the molarity of the KMnO4.
Calculations - Part 3 - Step A
1.
2.
3.
4.
5.
For each titration, convert the mass of sodium oxalate in the flask to moles.
Convert the number of moles of sodium oxalate to potassium permanganate using the mole
relationship implied in the balanced reaction equation.
Calculate the volume of potassium permanganate titrant used in each titration. First
subtract initial from the final burette readings to obtain the volume of titrant added. Then
subtract the volume of titrant used for the blank from the volume of titrant added to the
sample to obtain the corrected volume of titrant used.
Calculate the concentration of potassium permanganate from number moles of
permanganate and the corrected volume.
Average the most consistent results to obtain an average value for the concentration of the
potassium permanganate solution.
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Step B. Analysis of Oxalate in the Unknown Compound
The same reaction used in step A is used in step B. In step B however the concentration of
the titrant is known and the number of moles oxalate is unknown. This is essentially the reverse
of what was done in step A.
Experimental Procedures - Part 2
Analysis of an Unknown Coordination Compound of Iron(III)
Step B. Analysis for Oxalate in the Unknown Compound
Be sure to record all weights and observations about colors, color changes,
precipitates, etc., in your laboratory notebook. label each section of your notebook to agree
with the different sections of this experiment. This section should be labeled: Part 2 Analysis of an Unknown Coordination Compound of iron(III), Step B., Analysis for
Oxalate in the Unknown Compound. Other sections should be similarly labeled.
Weigh at least two 0.10-0.15 g samples of the complex salt in separate, labeled 250 mL
Erlenmeyer flasks. Dissolve each sample in 10 mL of 3 M H2SO4. Warm the solution gently to
dissolve the sample if necessary. Dilute each sample to a volume of about 75 mL using distilled
water. Heat the sample to 80-90°C and titrate as in the standardization procedure. The endpoint
color change may be somewhat different than before because of the presence of the iron(III) ions.
Repeat the procedure until consistent results are obtained. It may be necessary to quickly
complete the calculations to check for consistency.
A blank should be run on a sample prepared identically as above without any complex salt.
Calculations - Part 2 - Step B
1.
2.
3.
4.
5.
For each titration, calculate the number of moles of permanganate that reacted using the
average concentration of the permanganate solution and the corrected volume of titrant
delivered.
Convert the number of moles of permanganate to oxalate using the mole relationship
implied in the balanced equation for the titration reaction.
Convert the number of moles of oxalate ion, C2O42-, to the mass of oxalate ion in the
sample using the molar mass of the oxalate ion.
Calculate the percent oxalate in each sample from the mass of oxalate in the sample and the
total mass of the sample.
Average the most consistent results to obtain an average value for the percent oxalate in the
complex salt.
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Step C. Analysis of Iron(III) in the Unknown Compound
There are two problems which must be addressed before the amount of iron in the sample
can be determined by titration with permanganate. First, the oxalate in the sample must be
removed before the iron in the complex is titrated because, as demonstrated in steps A and B,
oxalate reacts with permanganate. Second, the iron(III) in the complex must be reduced to
iron(II) before it can be oxidized by permanganate. That is why the iron(III) in the complex did
not interfere with the determination of oxalate in step B; it is iron(II) which reacts with
permanganate, but iron(III).
The first problem can be overcome by taking advantage of the same reaction used to
analyze for oxalate.
5 C2O42- + 2 MnO41- + 16 Hl+ → 10 CO2 + 2 Mn2+ + 8 H2O
Addition of an excess of permanganate oxidizes the oxalate to carbon dioxide gas which evolves
from the solution. That leads to a solution containing Fe3+ from the complex, Mn2+ produced
during the destruction of the oxalate, plus excess unreacted MnO41-.
The second problem can be overcome with a good reducing such as tin(II). An excess of
tin(II) can be used to reduce iron(III) to iron(II).
Sn2+ + 2 Fe3+ → Sn4+ + 2 Fe2+
It also reduces the excess MnO41- left over from the destruction of the oxalate to Mn2+. That
leads to a solution containing Fe2+, Mn2+, Sn4+ produced in the reduction of the iron, and excess
Sn2+. Unfortunately the solution is not yet ready for titration!
The oxalate has been removed; it has been converted to carbon dioxide. The iron is in the
proper oxidation state; it has been reduced from iron(III) to iron(II). However, the tin(II) left
over from the reduction of the iron, can also be oxidized by permanganate. If it is left in the
solution, it will interfere with the iron(II)-permanganate titration; it must be converted to another
form that will not interfere
Mercury(II) oxidizes tin(II) to tin(IV).
2 Hg2+ + Sn2+ + 2 Cl1- → Hg2Cl2↓ + Sn4+
Equally important however is the fact the product mercury([) precipitates from solution and
cannot be oxidized by permanganate! The remaining solution contains Fe2+, Mn2+, Sn4+, Hg2+
and insoluble Hg2Cl2. The only species that can be oxidized by MnO41- is the Fe2+.
5 Fe2+ + MnO41- + 8 Hl+ → 5 Fe3+ + Mn2+ + 4 H2O
The laboratory work should go smoothly if you understand the chemistry associated with
each addition. It is strongly recommended that you understand this section before beginning the
experiment!
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Experimental Procedures - Part 2
Analysis of an Unknown Coordination Compound of Iron(II)
Step C. Analysis for Iron(III) in the Unknown Compound
Be sure to record all weights and observations about colors, color changes,
precipitates, etc., in your laboratory notebook. label each section of your notebook to agree
with the different sections of this experiment. This section should be labeled: Part 2 Analysis of an Unknown Coordination Compound of Iron(III), Step C., Analysis for
Iron(III) in the Unknown Compound. Other sections should be similarly labeled.
Weigh at least two samples weighing between 0.5 and 0.7 g of the complex salt into
separate labeled 250 mL Erlenmeyer flasks. Be sure to record the mass to three significant
figures after the decimal point. Add approximately 10 mL of 3M H2SO4 to one of the samples
and warm the flask gently on a hot plate while the sample dissolves.
Add 3% KMnO4 (NOTE: This is NOT the standardized KMnO4 solution, but a more
concentrated solution available in the laboratory.) dropwise to the flask until the first
permanent appearance of pink persists or a brown precipitate forms. What could cause the pink
color? What could the brown precipitate be? Be sure to answer these questions in your
laboratory notebook. You may also observe effervescence of the solution. Be sure to explain
that in your notebook. The oxalate ions should have been removed from the sample by these
additions.
The next two steps are critical! The flask can be rendered useless if the wrong amount of
tin(II) chloride is added. This is no time to be careless or impatient!
Carefully add the tin(II) chloride, SnCl2, solution dropwise to the flask until a clear yellow
solution is obtained. The yellow color is due to the Fe3+ in solution. Any precipitate formed in
the previous step should redissolve. This should have reduced the excess permanganate ion in
the solution to Mn2+.
Carefully bring the solution to a low boil on the hot plate and continue to add the SnCl2
solution dropwise until the yellow color of the Fe3+ ion just disappears, and then add 1 or 2
drops more. This should have reduced the Fe3+ in solution to Fe2+. Only about 40 drops, 2 mL,
of SnCl2 should have been used for both of these reductions,
Cool the solution under the cold water tap or in a cold water bath. Add 10 mL of
mercury(II) chloride, HgCl2, solution front a graduated cylinder to the reaction flask all at once.
A white, silky precipitate of mercury(I) chloride, Hg2Cl2 should form.
If the white precipitate forms, the iron(III) in the complex has successfully been reduced to
iron(II) and all the interfering ions have been converted to non-interfering species. Iron(II)
cations can begin to oxidize to iron(III) in the presence of iron, so proceed without delay.
There is a problem if either no precipitate forms, or a gray precipitate forms. If the
precipitate does NOT form, not enough SnCl2 was added in the previous step. There is either
MnO41- or Fe3+ ion left in the solution, or both! If a gray precipitate results, elemental mercury
has been formed. This can happen if either too much tin(II) chloride was added in the previous
step, or the HgCl2 was added too slowly in this step. If either one of these problems is observed,
the sample cannot be titrated. It is lost and you must begin the determination again with another
sample.
If the white precipitate forms, the sample is almost ready to be titrated. Dilute the solution
with approximately 100 mL of distilled water and 25 mL of Zimmerman-Reinhart reagent. (The
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Zimmerman-Reinhart reagent is necessary for two reasons. It prevents the chloride ion present
from being oxidized by the permanganate titrant. It also contains a reagent that ties up the Fe3+
formed in the titration as a phosphate complex. This makes the oxidation of the Fe2+ to Fe3+ by
MnO41- more complete by shifting the position of the equilibrium toward the formation of
product Fe3+). Titrate the sample slowly with the standard permanganate solution until a faint
pink color persists for more than 5 seconds.
Repeat this entire procedure with fresh samples of complex until consistent results are
obtained. It may necessary to complete calculations quickly to check consistency.
Finally, it will necessary to titrate a blank as in steps A and B. Carry through the above
procedure, omitting the sample. Add one drop of 3% KMnO4 to 10 mL of 3M H2SO4 and only
1-2 drops of SnCl2 as there are no oxalate or iron(III) ions present. Only a few drops of
standardized permanganate should be required to reach the endpoint of the titration. Again, this
volume will be used to correct the volumes of titrant used for contaminants present.
Calculations - Part 2 - Step C
1.
2.
3.
4.
5.
6.
For each titration, calculate the number of moles of permanganate that reacted using the
average concentration of the permanganate solution and the corrected volume of titrant
delivered.
Convert the number of moles of permanganate to iron(II) using the mole relationship
implied in the balanced equation for the titration reaction.
Convert the number of moles of iron(II) titrated to the number of moles of iron(III) in the
sample using the mole relationship implied in the balanced equation for the reduction of the
iron(III) to iron(II).
Convert the number of moles of iron(III) to the mass of iron(III) ion in the sample using the
molar mass of the iron(III) ion.
Calculate the percent iron in each sample from the mass of iron in the sample and the total
mass of the sample.
Average the most consistent results to obtain an average value for the percent iron in the
complex salt.
Page 11
Principles of Chemistry II
Coordination Compoounds
Determination of the Empirical Formula of the Unknown Iron(III) Compound
The following information is needed to determine the empirical formula of the unknown
iron(III) compound: the percent oxalate in the compound, from Part 2 - Step B; the percent iron
in the compound, from Part 2 - Step C; and the generic form of the empirical formula for the
compound given in the Introduction to Parts 1 and 2 of this experiment.
1.
2.
3.
4.
5.
6.
Calculate the number of moles of oxalate in a 100 g sample of the compound using the
percent oxalate determined in Part 2 - Step B and the molar mass of oxalate.
Calculate the number of moles of iron in a 100 g sample of the compound using the percent
iron determined in Part 2 - Step C and the molar mass of iron.
Calculate the ratio of moles of oxalate to moles of iron using the results of computations 1
and 2 above. This should be quite close to an integer or a simple fractions
Use this ratio to construct likely formulae for the compounds. Use potassium ions, K1+
ions, to produce a neutrally charged compound, recalling both the Fe3+ and C2O42- ions are
charged. Finally, include the three waters of hydration as indicated in the generic form of
the empirical formula for the compound.
Using the empirical formula for the compound calculate the theoretical yield for the
synthesis assuming the actual mass of iron(II) ammonium sulfate hexahydrate,
Fe(NH4)2(SO4)2.6H2O, used in the synthesis is the limiting reagent. Note: It is not
necessary to write the balanced equation for the synthesis to complete this calculation.
The mole relationship between the starting material, iron(II) ammonium sulfate
hexahydrate, Fe(NH4)2(SO4)2.6H2O, and the final product, can be inferred from the
number of moles of iron in the formulae for the two compounds. However, it should
also be possible to balance the equation for the synthesis given in the Introduction to
Part 1 of this experiment, now that the formula for the final product is known.
Calculate the percentage yield for your synthesis using the results of the calculation above
and the actual mass of product obtained in Part 1 of the experiment.
Page 12