Authors: D. Afzal, H. D. Ervin, A. E. Moody, H. D. Wohlers, and J. M. McCormick*
Last Update: December 29, 2006
In this experiment you will be synthesizing three coordination compounds and analyzing them. Coordination compounds (also known as complex ions and complexes) are formed by the reaction of a Lewis acid (usually a transition metal) with a Lewis base, which is known as a ligand. What is unique about coordination compounds is that they are formed from chemical species that have an independent existence and that this association is often readily reversible (i. e., there is an equilibrium between the solvated metal ion and the ligand). For example, NiCl2 reacts with NH3 in aqueous solution to form the compound Ni(NH3)6Cl2 which contains the complex ion [Ni(NH3)6]2+. This process is easily reversed (by the addition of H+) to give back the starting materials. This type of behavior was thought to be very peculiar by chemists in the 1800's. They were familiar with compounds like CO2, which although it could be made from C and O2, does not act like it is some loose association of C and O2. It wasn’t until the ground-breaking work of Werner (for which he won the Nobel Prize in chemistry) in the late 19th and early 20th centuries that chemists began to understand these compounds. Werner’s work was greatly expanded on in the 20th century especially after it was discovered that coordination chemistry was relevant to the understanding the role of metal ions in biological systems.
You will be preparing complexes of Co3+ with ethylenediamine, NH2CH2CH2NH2 (abbreviated: en), and of Fe3+ with the oxalate ion, C2O42- (abbreviated: ox). Ethylenediamine and oxalate are examples of bidentate ligands, which means that they have two different atoms that can donate electron pairs to a metal ion. Ethylenediamine does this through its nitrogen atoms, while oxalate donates electron pairs from two of its four oxygens.
In all of these complexes the metal ion is directly bonded to six other atoms in what is called an octahedral geometry (if we connected the six atoms, the resulting solid would be an octahedron, and hence the name of this geometry). Although it may not be obvious, there may be many ways in which the same six atoms can be arranged around a central atom in an octahedral geometry, and each of these different arrangements may give rise to compounds with the same chemical formula, but have different arrangements of their atoms (isomers). There are many types of isomers: compounds in which the actual connections between atoms (bonds) are different are called constitutional isomers, while compounds where the bonds are the same, but the atoms are arranged differently are called geometric isomers. Geometric isomers are further classified as enantiomers (two compounds that are mirror images of each other) and diastereomers (geometric isomers that are not mirror images).
Because en is a bidentate ligand, [Co(en)2Cl2]+ exists as three geometric isomers; one pair of enantiomers and their diastereomer. The isomer where the chlorides are situated on either side of the Co3+ (180° from each other) is called the trans isomer (Fig. 1), while the isomer where the chlorides are next to each other in the octahedron is the cis isomer. In addition, there are two different ways in which we can put two Cl atoms cis to one another, and these are enantiomers (Fig. 2). The [Fe(ox)3]3- exists as two enantiomers (which have no diastereomer); one in which the three oxalates form a right-handed propeller, and in the other they form a left-handed propeller (both are shown in Fig. 3).
Figure 1. Structure of trans-[Co(en)2Cl2]+ redrawn from the Cambridge Crystal Structure Database entry CENCOS using the Mercury molecular visualization software package.
Figure 2. Structures of the two cis-[Co(en)2Cl2]+ enantiomers redrawn from the Cambridge Crystal Structure Database entries CENCOC and CLECOC using the Mercury molecular visualization software package.
Figure 3. Structures of the [Fe(ox)3]3- enantiomers redrawn using the Mercury molecular visualization software package from data collected by Professor Russell Baughman at Truman State University on crystals prepared by CHEM 121 students.
In this experiment, you will synthesize trans-[Co(en)2Cl2]+ and then convert it to the cis-[Co(en)2Cl2]+. We will not attempt to resolve (separate) the cis enantiomers. If we did, we would find that their properties are almost identical, except that would rotate polarized light differently. No matter how hard we try, we could not resolve the [Fe(ox)3]3- enantiomers. This is because there is a fast pathway that interconverts the two enantiomers (the complex is said to be labile), which does not exist for the cobalt complexes (which are said to be inert).
The reactions used to synthesize the complexes are shown below. These are obviously greatly simplified as no other reaction products are shown. Note that in the synthesis of trans-[Co(en)2Cl2]Cl·HCl we start with Co2+ and end up with the metal in its +3 oxidation state, while in the synthesis of K3[Fe(ox)3]·3H2O the metal is in its +3 oxidation state throughout the reaction.
The analysis of the K3[Fe(ox)3]·3H2O will use the redox reactions given below for the analysis of the oxalate and the iron, respectively. To save time and sample, you will first perform the oxalate analysis and then treat the titration solution with Zn to convert all of the Fe3+ present to Fe2+. Titrating this solution with the standard MnO4- will allow you to determine the amount of iron present in the original compound.
For this exercise it is critical that you read the procedure carefully and that you and your lab partner plan what you will do before coming to the laboratory. If you want to finish the first week’s work in the allotted time, you must work efficiently.
Each pair of students will need to check out the following items from the stockroom: magnetic stir bar, beaker tongs, Büchner funnel and filter flask.
All work of the first week’s work must be carried out in the hood.
Synthesis of trans-Dichlorobis(ethylenediamine)cobalt(III) Chloride Hydrochloride2,3
There will be two stirring hotplates in each hood. In this part of the first week’s work you will use one hotplate ONLY for heating, and the other ONLY for stirring. In the second part of the first week, you will use the hotplate used for stirring in the synthesis of the iron-oxalate compound. There is enough room on each plate for two beakers or two evaporating dishes; work with your hood-mates to assure that everything is done efficiently and safely.
Fill a 600-mL beaker approximately half full with water, and place it on a hot plate in the hood. Place your wire gauze on top of the beaker and heat to boiling.
Weigh out approximately 1 g of cobalt(II) chloride hexahydrate, CoCl2·6H2O, measured to the nearest milligram. Place the cobalt salt in your evaporating dish and add 2.5 ml of distilled water. Gently swirl the dish until the cobalt has dissolved.
Once the cobalt chloride has completely dissolved in the water, add 4 mL of 10% ethylenediamine. Place your magnetic stir bar in the solution and set the evaporating dish on a stirring hotplate (NOT the hotplate with your water bath). Set the stir control knob so that the stir bar spins but does not splatter solution on the side of the evaporating dish (a setting of "4" on most stir plates is a good starting point). Stir for 10 min.
After 10 minutes, slowly and carefully add dropwise 1.6 mL of 10% H2O2 to the solution while it is still stirring. CAUTION! 10% H2O2 causes severe skin burns. Gloves are recommended and you should wash your hands after using this solution. Once the H2O2 addition is complete, reduce the stir speed (to "3" on most of the laboratory stir plate) and stir for an additional 15 min.
Next, CAREFULLY add 3 mL of concentrated HCl. CAUTION! HCl is VERY corrosive and will cause severe burns. Immediately wash off any HCl that comes in contact with skin with copious amounts of water. Turn off the stir plate and remove the magnetic stir bar using the magnetic stir bar retriever, as your instructor will demonstrate.
Place your evaporating dish on top of the wire gauze on the beaker of boiling water. Reduce the solution’s volume until there is a thick layer of dark green crystals and very little liquid left. This will take between 30 and 45 min. Be sure to monitor the level of boiling water in your water bath, you will need to gradually add water to the water bath during this time to keep it from boiling dry. This is a good point to start the synthesis of the iron complex.
While you are waiting for the volume to be reduced, set up an ice bath. Obtain ~5 mL of methanol and put it in a small beaker or test tube in your ice bath. You will use the ice cold methanol to rinse your crystals in the final step of this synthesis.
Carefully remove the evaporating dish with the beaker tongs and set it on the bench top in the hood to cool. While your solution is cooling, set up the vacuum filtration apparatus (see the CHEM 120 Alum experiment).
Place a piece of filter paper in the Büchner funnel to collect your crystals. Carefully wet the filter paper with a couple drops of methanol BEFORE you add your crystals.
With the aspirator on, transfer the contents of your evaporating dish into the Büchner funnel using a spatula. Add about 2 mL of ice-cold methanol to your evaporating dish, swirl and pour this into the Büchner funnel. This will help remove all of your contents from the evaporating dish. Carefully pour ~5 ml of ice cold methanol over your crystals and filter. Repeat if necessary.
Carefully transfer your crystals to a watch glass, as your instructor will demonstrate. It is not necessary to remove the filter paper at this time. Cover the watch glass with a paper towel and place it in your drawer to dry until next week.
Synthesis of Potassium Tris(oxalato)ferrate(III) trihydrate4,5
In your clean, dry 8 inch test tube dissolve 1.6 g FeCl3·6H2O (measured precisely) in 4 mL of distilled water. It may be convenient to gently heat the solution in a hot water bath on a hot plate to speed the dissolution of larger chunks of the ferric chloride.
Precisely and accurately weight out between 6.0 and 6.5 g K2C2O4·H2O and place it in a clean, dry 50-mL beaker. Add 10 mL of distilled water and gently heat on a hot plate with stirring until all of the potassium oxalate has dissolved. It is not necessary to boil the solution.
Once the oxalate salt has dissolved, quickly and carefully pour the hot oxalate-containing solution into the beaker containing the iron solution. Swirl to mix. Cover and allow the reaction mixture to cool slowly to room temperature. Once the test tube is cool to the touch transfer it to an ice bath and continue cooling.
Green crystals of the product may form during the initial cooling to room temperature, but they might not form until the ice cooling step. If crystals do not form after 30 min of cooling, try gently scratching the bottom of the beaker with a stirring rod or add a seed crystal (a crystal of the product obtained from another preparation).
Decant the solution above the crystals and discard. Recrystallize the crude product by adding approximately 5 to 8 mL of hot distilled water to the crystals and heating gently to affect total dissolution. A dark, insoluble residue may remain after the product has dissolved. If this occurs, decant the green solution containing the product into another clean, dry beaker and discard the residue.
Cover the beaker with a watch glass and set it aside to cool slowly. When the beaker is cool to the touch, transfer it to an ice bath and continue to cool. Green crystals should form within about 20 min. If they do not, consult your instructor for assistance.
Collect the crystals by vacuum filtration. Wash the crystals twice with 2 mL of ice-cold distilled water and then with two 3-mL portions of acetone. CAUTION! Acetone is flammable, so do all these manipulations in the hood. Dry the product on the filter and then transfer it to a watch glass. Cover the watch glass with a paper towel and place it in your drawer to dry.
Before beginning any work this week obtain the mass of the dry trans-dichlorobis(ethylenediamine)cobalt(III) chloride hydrochloride and the dry potassium tris(oxalato)ferrate(III) trihydrate.
Synthesis of cis-Dichlorobis(ethylenediamine)cobalt(III) Chloride
Place approximately 50 mg of the trans-dichlorobis(ethylenediamine)cobalt(III) chloride hydrochloride on your watch glass. Add 1 or 2 drops of saturated aqueous NaHCO3 solution. This will neutralize the excess HCl, which interferes with the conversion to the cis- complex. Use your spatula to gently stir the solution. You may observe gas evolution and the solution may appear cloudy, but the solution will quickly turn from green to purple. Allow the solution to stand for about 10 min. During this time set up a hot water bath on a stirring hot plate. Place the watch glass on the water bath and heat the solution to dryness. The solid tends to scatter, so exercise care when removing the purple solid from the watch glass.
Absorbance Spectrum of cis- and trans-Dichlorobis(ethylenediamine)cobalt(III) Chloride
Place approximately 50 mg of trans-dichlorobis(ethylenediamine)cobalt(III) chloride hydrochloride in a 6-inch test tube. Add approximately 2 mL of water and 1 drop of 6 M HCl; shake to dissolve. Once all of the solid has dissolved, add another 1 mL of water and mix thoroughly.
In a different test tube prepare a sample of the cis-dichlorobis(ethylenediamine)cobalt(III) chloride following the same procedure as for the trans- isomer, except do not add the 6 M HCl.
Set up the Ocean Optics spectrometer (click here to review set up and operation of the spectrometer). Record the absorbance spectrum of each cobalt complex in the wavelength range 450-900 nm. Remove any bubbles by gently tapping with your finger. Under absolutely no circumstances are you to tap a cuvette on a table top. Rinse the cell with distilled water between samples. If the maximum absorbance in the 450 to 900 nm range is greater than 1.00, dilute your sample with distilled water until it is. A helpful hint: the absorbance is directly proportional to the concentration (Beers’ Law). Another helpful hint: do any dilution in small steps by adding a little distilled water and then checking the absorbance. Transfer the spectrum of each complex to Excel and save to your Y: drive.
Determination of Oxalate in Potassium Tris(oxalato)ferrate(III) Trihydrate
Obtain approximately 75 mL of standardized 0.01 M KMnO4 (record the actual concentration from the bottle's label). CAUTION! Potassium permanganate is a strong oxidizing agent! Immediately wash any spilled on your skin with copious amounts of water. Wash your buret several times, as your instructor will demonstrate, with no more than 2 mL of the permanganate solution. Each time allow the solution to drain through the buret's tip into a waste beaker. After the final wash fill the buret to near the 0.00-mL mark such that the upper part of the meniscus is at or below the 0.00-mL mark. Make sure that there are no bubbles in the buret tip. Run some of the solution through to clear any bubbles (gentle taping might help dislodge any recalcitrant bubbles). Refill the buret, if necessary. Record the initial volume of solution in the buret. Because the MnO4- ion is so strongly colored, make all your volume readings from the top of the meniscus.
Accurately weigh out about 0.10 g of your K3[Fe(ox)3]·3H2O. Place it in an Erlenmeyer flask and add about 30 mL of distilled water and 5 mL of 6 M H2SO4. Swirl the flask gently to dissolve the solid.
Place the flask on a stirring hotplate in the hood and carefully heat the solution to approximately 60 °C. Do not boil! Remove the flask from the hotplate (a folded paper towel makes an excellent hot pad). This titration must be performed fairly quickly so that the solution does not cool too much. Be sure that you swirl the flask as you add the permanganate solution. Titrate the sample with the permanganate solution until the slightest purple color persists for at least 30 sec. Record the volume of permanganate used.
If you have enough sample, repeat the titration at least three times or until the data for three determinations agree with each other within 2%.
Determination of Iron in Potassium Tris(oxalato)ferrate(III) Trihydrate (optional)
Take your faintly-purple sample from the oxalate analysis above, and set in on a hotplate in the hood (there may be a small amount of brown precipitate forming at this point, do not be concerned). Gently heat it until it almost boils and add approximately 100 mg of Zn to the hot solution. Cover the flask with a watch glass and continue heating until the yellow color (from Fe3+) disappears (Fe2+ is colorless in solution). While you are waiting, set up a gravity filtration with fluted filter paper (ask your instructor to demonstrate how to flute filter paper, if you are unsure).
Quickly filter the hot, colorless solution into another Erlenmeyer flask, again using a folded paper towel to handle the hot flask. It is important that this filtration be done quickly to minimize the amount of Fe2+ that is reoxidized to Fe3+ by O2 in the air. Rinse the funnel with several small (approximately 5 mL total volume) portions of distilled water. Titrate this solution with the permanganate solution in your buret until the first faint trace of a persistent pink color. Record the volume of permanganate used.
Repeat this procedure after each oxalate determination (at least three times).
Results and Analysis
Week 1
After you have determined the mass of your dry trans-[Co(en)2Cl2]Cl·HCl and K3[Fe(ox)3]·3H2O, calculate a percent yield for each compound. In each case assume that the metal salt is the limiting reagent.
Week 2
In Excel® prepare a graph of absorbance (on a 0 to 1 scale) as a function of wavelength (from 450 to 900 nm) for trans-[Co(en)2Cl2]Cl·HCl and for cis-[Co(en)2Cl2]Cl, where both spectra are shown on the graph. Be sure to clearly indicate which spectrum corresponds to which compound. Print out two copies (one for the original page and one for the duplicate page) and glue them in your notebook (hint: copy and paste into Word® first). Describe the differences you observed between the two spectra.
Calculate the average percent oxalate by mass in your K3[Fe(ox)3]·3H2O and share this with the class. From the class data calculate the average percent oxalate by mass in K3[Fe(ox)3]·3H2O, the standard deviation and the confidence interval at the 95% confidence limit. Be sure to perform a Q-test on the class data first to exclude any suspect point. Calculate the true percent oxalate by mass and determine a percent error for both your data and the class data.
If you performed the iron analysis, calculate the average percent iron by mass in your K3[Fe(ox)3]·3H2O. Share this with the class, and calculate the average percent iron by mass in K3[Fe(ox)3]·3H2O, the standard deviation and the confidence interval at the 95% confidence limit from the class data. Perform a Q-test on the class data first to exclude any suspect points. Calculate the true percent iron by mass and determine a percent error for both your data and the class data.
Conclusions
This experiment contains elements of both a synthesis exercise and a measurement exercise. Your discussion should address whether you think you made the target compounds and whether the analyses support that conclusion. For your discussion of the cobalt complexes, assume that both compounds have the same chemical formula and discuss what it means for them to have different absorbance spectra.
Summary of Results
Your Summary of Results table should contain the percent yield of trans-[Co(en)2Cl2]Cl·HCl and K3[Fe(ox)3]·3H2O and the statistical (average, standard deviation and confidence interval) results for the K3[Fe(ox)3]·3H2O analyses.
2. Szafran, Z.; Pike, R. M. and Singh, M. M. Microscale Inorganic Chemistry: a Comprehensive Laboratory Experience; John Wiley and Sons: New York, 1991.
3. Bailar, J. C., Jr. Inorg. Synth. 1946, 2, 222-225.
4. Beran, J. A. Laboratory Manual for Principles of General Chemistry, 5th Ed.; John Wiley and Sons: New York, 1994.
5. Marcus, S.; Sienko, M. J. and Plane, R. A. Experimental General Chemistry; McGraw-Hill: New York, 1988.
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