Determination of an Empirical Formula of an Oxalato Compex Salt
From ChemEd Collaborative
Determination of an Empirical Formula of an Oxalato Compex Salt by Laura E. Slocum October 2008
BACKGROUND
This lab comes from several of my colleagues and I will be copying most of my colleagues’ work so do NOT cite me as the author of this lab. The original lab came from NSF funded Summer Project in Chemistry at Hope College. I do NOT have the original author’s names for each of the pieces of the projects from those Summer Projects. As I can gather them, I will share them with you. Here is the information that you should use in your procedure citation to start –
Determination of an Empirical Formula of an Iron Oxalato Complex Salt
Hope College
Holland, Michigan
I also do NOT have the year that it was published, but I know that Dr. Brink, who was the author for the synthesis procedure retired in 1996 and it was before that. Since, there have been several adaptations of the synthesis procedure since that time, we will use 1996 for the year of publication for this procedure for now.
In a previous experiment (Synthesis of Potassium Iron Oxalate) you prepared a crystalline product having the formula KxFe y(C2O4)z·nH2O. In these following experiments beginning with this one you will determine the percentage composition of the compound and its simplest formula. Finally the percent yield will be calculated.
Percent Water –
The iron oxalato complex salt is one of a number of solid chemicals that are classified as hydrates. A hydrate contains water chemically bound in the solid state so that it is present in the compound in stoichiometric amounts. Familiar examples of hydrates are gypsum, CaSO4•2H2O; plaster of paris, CaSO4•1/2H2O; and alum, KAl(SO4)2•12H2O. The water of hydration of many hydrates can be removed as a gas by heating the hydrate to a temperature above 100°C for a period of time. The equation below gives the reaction involved when barium chloride dihydrate, BaCl2•2H2O, is heated above 100°C.
The solid remaining after the water has been expelled is called the anhydrous compound, in the above example, anhydrous barium chloride (BaCl2).
The percentage of water of hydration in KxFe y(C2O4)z•nH2O will be determined in this experiment by heating a weighed sample of the hydrate in an open container in an oven until all of the water of hydration has been driven off.
The loss in weight is equal to the mass of the water of hydration.
Percent Oxalate –
The percent oxalate, C2O42– in KxFe y(C2O4)z•nH2O will be determined by titrating a solution containing a known mass of the crystals with the standardized 0.010 M KMnO4 prepared and standardized in this experiment too. [I will prepare the KMnO</nowiki>4 solution, but you will standardize it.]
The mass and percent of C2O42– in the sample can be determined by measuring the volume, V, of KMnO4 of molarity M, required to completely oxidize C2O42-. KMnO4 oxidizes C2O42– in acid solution to give CO2(g) as a product.
Standardization of Potassium Permanganate –
Potassium permanganate, KMnO4, is a relatively inexpensive, intensely colored compound commonly used in laboratories as a powerful oxidizing agent. In this experiment a 0.010 M KMnO4 solution will be prepared by dissolving a weighed quantity of solid KMnO4 in water to give 500 mL of solution. The mole ratio of C2O42– to MnO4– in a titration reaction involving the oxidation of C2O42– by MnO4– will be calculated and the stoichiometry of the reaction determined by completing and balancing equation 1.
Sodium oxalate, Na2C2O4, is the source of the oxalate ions in the titration. The actual molar concentration of the KMnO4 solution will be calculated from the titration data. Combining the titration data with the solution preparation data will do an estimate of the purity of the KMnO4 used to prepare the solution.
Then you will use the standardized KMnO4 solution to determine the percent iron, Fe, and percent oxalate, C2O42-, in the compound, KxFe y(C2O4)z·nH2O, that you synthesized earlier. The actual molarity of the standardized KMnO4 will be used in these calculations.
- Preparation of 0.010 KMnO4: In principle, if one dissolves a weighed quantity of KMnO4 (Molar Mass = 158.04 g/ mole) in an accurately measured volume, say 500.0 mL, the molarity of the solution can be simply calculated by taking the ratio of moles of KMnO4 over liters of solution. However, in practice this is not a good method for an accurate determination of molarity since the KMnO4 purchased from suppliers is contaminated with varying amounts of impurities. KMnO4 is particularly difficult to prepare and preserve in pure form because it is such a strong oxidizing agent that it reacts readily with a great variety of materials including dust in the air. However, the purity of the solid KMnO4 used to prepare the 0.010 M solution in this experiment is good enough to permit the determination of the C2O42– to MnO4– mole ratio in the equation above using the data from the preparation of the KMnO4 solution and the titration results of C2O42– by MnO4–.
- Completing and Balancing Equation 1: When a KMnO4 solution is used as an oxidizing agent in solution, the MnO4– is reduced to any one of several possible manganese containing species depending on the conditions of the reaction. Possible reduction products are MnO4–, MnO2, or Mn2+. Partially balanced equations 2, 3, and 4 give the C2O42– to MnO4– mole ratio for each of the three possible reduction products when MnO4– oxidizes C2O42– in solution.
C2O42– + 2 MnO4– → 2 CO2 + 2 MnO42– (2)
3 C2O42– + 2 MnO4– → 6 CO2 + 2 MnO2 (3)
5 C2O42– + 2 MnO4– → 10 CO2 + 2 Mn2+ (4)
The actual C2O42– to MnO4– mole ratio is experimentally determined by measuring the volume of KMnO4 solution of known concentration required to react with a solution containing a known number of moles of C2O42–, determined from the mass of Na2C2O4 (134.00 g/mol) used in the titration. The experimental mole ratio of C2O42– /MnO4– obtained by this method will not be an exact ratio of small whole numbers but it should be close to either 1/2 = 0.50, 3/2 = 1.50, or 5/2 = 2.50 depending on whether the stoichiometry corresponds to that of equation 2, 3 or 4.
- Determination of Molarity of the KMnO4 by Titrimetry: The value of M, the molar concentration of KMnO4 solution used to determine the stoichiometry of the reaction, was determined from the solution preparation data and is only approximate because of the unknown purity of the solid KMnO4 used to prepare the solution. However, once the C2O42– /MnO4– mol ratio has been established this stoichiometric relationship, rounded off to the nearest whole number ratio (1/2, 3/2, or 5/2), can be used to calculate an exact value of M.
MATERIALS and EQUIPMENT
| Balance | 0.010 M Potassium Permanganate, KMnO4 |
| 250-mL Beakers | Sodium Oxalate, Na2C2O4 |
| 250-mL Beaker | 6 M Sulfuric Acid, H2SO4 |
| 3-250-mL Erlenmeyer Flasks | 3 M Hydrochloric Acid, HCl |
| Dropping Pipet | 85% Phosphoric Acid, H3PO4 |
| Buret | |
| Short Stem Funnel | |
| Crucibles | |
| Dessicator |
SAFETY
Be sure to wear safety goggles as directed by your teacher.
KMnO4 stains the skin, clothing and other items a brown color. Immediately after contact wash affected area with soap and water. Brown stains on the skin will wear off within a few days. Brown stains on clothing are permanent.
3 M HCl is used to clean glassware in this experiment. Treat 3 M HCl with respect; it is a very strong acid. Handle H3PO4, H2SO4 and KMnO4 solutions with care. Sponge up spills immediately using large quantities of water. If the acid spills on skin or clothing rinse affected areas immediately with large amounts of water. If necessary, neutralize the acid with NaHCO3 and wash until all biting sensation ceases. Serious acid burns can result if you ignore or simply endure the biting of the acid. Similarly, handle 6 M H2SO4 with care and clean up spills in the same way with copious rinsing.
Discard all waste solutions in the designated container provided by your teacher.
PROCEDURE
PERCENT WATER –
Notes on Experimental Procedure:
- You will need to know the mass of crystals before heating and the mass loss upon heating. The mass of KxFe y(C2O4)z·nH2O can be calculated by subtracting the mass of the empty crucible from the mass of the crucible plus crystals before heating. The mass of water of hydration can be calculated by subtracting the mass of the crucible plus sample after heating from the mass of the crucible plus sample before heating. A convenient format for entering the experimental data in your research notebook is
Time and date of weighing
crucible empty (without lid) _________
after 1st heating _________
after 2nd _________
(Leave a couple of extra lines to use in case the mass changes by more than 0.002 g between weighings.)
crucible + crystals _________
crucible + residue
after 1st heating _________
after 2nd heating _________
(Continue to dry, cool and weigh until 2 masses agree within 0.002 g.)
- When crucibles are stored in a desiccator or placed in the oven to dry, each crucible should be placed in a small beaker to guard against the crucibles tipping over or becoming contaminated on the outside. Always have crucibles open while they are being heated and while they are cooling to room temperature. The lid may be placed on the crucible for long time storage at room temperature. After they have been dried in the oven, the crucibles should not be touched by fingers until the experiment is completed. Handle them with crucible tongs or a loop of folded paper or a Kimwipe.
- The most frequent cause of erratic balance readings is failure to have objects at room temperature when they are being weighed. Convection currents in a closed balance compartment can have a surprising effect.
- A small inaccuracy in the calibration of a balance is canceled out in the final results when the mass of the empty container and the mass of the container plus the substance of interest are measured using the same balance. Using the same balance for successive weighing when attempting to attain constant mass is recommended for the same reason.
Percent Water Procedure:
Before beginning this experiment, the crystals should have been allowed to dry at room temperature in your locked drawer, weighed and stored in a brown bottle with cap.
Prepare a tabular format in your laboratory notebook for recording the masses that you will be determining on the analytical balance and the date and time that you do the weighings (note 1).
Mark the crucible with pencil (not with ballpoint nor with Sharpie marker) to distinguish one from the other; then place them (without cover) in a 50 mL beaker. Store the covers of the crucibles safely in your drawer. They will not be used until after the final weighing. Obtain the mass of the empty crucible using the 0.001 g balance (these are located at the back of the lab).
Place the empty crucible and beaker assembly in the oven at 110-120°C to dry for at least 30 minutes. Remember to record both the time and the temperature of the oven in your notebook each time you place something in the oven and each time you remove it from the oven. Be sure to include in the beaker, a slip of paper with your name on it. Do not write on the large beaker.
Cool the crucible to room temperature in a desiccator (note 2). This will require at least 30 to 45 minutes (note 3). Weigh the crucible on the analytical balance to the nearest 0.001 g. Record the mass, time and date in the prepared format in your notebook (note 1).
Return the crucible to the oven for an additional 30 minutes of heating, followed by cooling, and weighing as above using the same balance (note 4). Repeat this procedure until the crucible has a constant mass. The mass should not change by more than 0.002 g between weighings.
Pre-weigh about 1.0 g of your crystals into the crucible using the .001 g electronic balance. This mass does not need to be recorded in your notebook. Weigh the crucible containing the crystals to the nearest 0.001 g using the same analytical balance that you used in determining the mass of the empty crucible (note 4). Record this mass in your laboratory notebook.
Place the crucible in the small beaker and heat for two hours at 110-120°C to drive off the water(s) of hydration. Do not leave the crystals in the oven for more than 2.5 hours, or they may begin to decompose and turn brown. Plan this part out appropriately with your lab partner in advance.
After cooling the samples in the desiccator for at least 45 minutes, again weigh each crucible plus sample to the nearest 0.001 g using the same analytical balance that you used prior to heating the crystals. Record in notebook.
Return the crucibles to the oven for an additional 30 minutes of heating, followed by cooling, and weighing to ascertain whether more water has been driven off. Repeat this procedure until crucible plus contents have a constant mass (masses agree within 0.002 g). Record results each time.
Calculate the loss in mass upon heating for the sample and attribute the mass loss to the water of hydration. Determine the percent water of hydration for the sample.
After the final weighing place the cover on the crucible and carefully wrap it in paper toweling to keep out the light. Store the crucible plus their contents for possible use in a later experiment in case your supply of unheated crystals is depleted before you have all of the determinations of % composition completed.
PREPARATION and STANDARDIZATION of 0.010 M Potassium Permanganate Solution –
Notes on Experimental Procedure:
- KMnO4 readily reacts with organic (carbon containing) materials including dust as well as small fibers of paper or cloth. Since it is difficult to rinse glass completely free of organic soap and detergents, the KMnO4 solution will be contaminated by these substances if it comes in contact with glassware which has been cleaned by soaps or detergents. Therefore, avoid using these materials to clean any glassware that comes into contact with KMnO4. Rinse glassware that has been exposed to detergents with a small amount of 3 M HCl followed by at least three tap-water rinses before allowing KMnO4 to contact it. The chloride ion must be completely rinsed from the glassware before it is used with KMnO4 solution because any remaining Cl - will react with the permanganate to form Cl2 and MnO2.
- Brown stains on bottle or stopper are likely to be MnO2 left over from previous storage of KMnO4. Since MnO4 catalyzes the decomposition of KMnO4, it is necessary to remove MnO2 stains to prevent decomposition of KMnO4 during storage.
- Since 0.010 M KMnO4 is so intensely colored it is almost impossible to see the meniscus. Read the mark on the volumetric flask or buret from the top surface of the solution.
- Exposure to light increases the rate of decomposition of KMnO4 solutions.
- The reaction of C2O42- with MnO4 - is slow at room temperature. The purpose of increasing the temperature is to increase the rate of the reaction.
- Since the flask will be too hot to hold with your fingers, use a folded Handi-Wipe or paper towel as a pot-holder.
- The KMnO4 acts as its own indicator here. As long as there is unreacted C2O42- present in solution the MnO4 - will be decolorized by the reaction. However, when the last C2O42- has reacted, the MnO4 - will impart its color to the solution. This coloration will eventually fade because the MnO4 - reacts slowly with trace reducing agents in the water to form colorless reduction products.
- It is important to save as much of your standardized KMnO4 as possible for use in future experiments. Do not discard the KMnO4 remaining in the buret at the end of each titration. Simply add more solution to fill the buret and proceed with the next sample. At the end of your work each day, discard the KMnO4 remaining in the buret and/or beakers. Do not return it to your stock solution because it may introduce contamination. However, in case you are about to run out of standardized solution, it may be worth the risk of contamination to save the solution remaining in the buret.
- A quick method of finding the deviation in parts per thousand (ppt) between samples is to calculate the ratio of g sample to mL titrant for each sample. Divide the largest ratio by the smallest ratio and multiply the answer by 1,000. The amount by which the result differs from 1000 is the deviation between the two samples in ppt.
Example:
| Sample | | | |
| 1 | | | |
| 2 | | | |
| 3 | | | |
0.003225/0.003192 = 1.011; 1.011 x 1000 = 1011; and 1011 - 1000 = 11
Therefore samples 1 and 2 vary by 11 ppt. and sample 3 lies between them.
PREPARATION and STANDARDIZATION of 0.010 M Potassium Permanganate Solution –
Preparation and Standardization of 0.010 M Potassium Permanganate Procedure:
Preparation of 0.010 M KMnO4: I will be preparing the solution for you!!!
Thoroughly rinse a 500 mL reagent bottle with polyethylene or glass stopper with tap water several times. DO NOT USE SOAP OR DETERGENT TO CLEAN THIS BOTTLE (note 1).
If the bottle is not sparkling clean after rinsing, or if it has a yellowish tinge, pour about 10 mL of 3 M HCl into the bottle and carefully tilt and rotate it so that the HCl comes in contact with the stained glass surface; clean the stopper with 3 M HCl if it is stained (note 2). The HCl should dissolve any MnO2 stains. You may ignore any stains that do not respond to HCl treatment.
Thoroughly rinse the bottle and stopper with water at least three times. Invert the bottle in a clean dry beaker (to prevent the bottle from tipping) and allow it to drain dry. Allow the stopper to air dry. Do not use a paper towel or any cloth that leaves lint (note 1).
Calculate the mass of KMnO4 required to prepare 500 mL of 0.010 M solution.
Weigh out between 0.78 and 0.80 g of KMnO4 to the nearest milligram in a clean dry 50 mL beaker using the mg electronic top-loader balance.
Transfer the KMnO4 into a clean beaker and dissolve in distilled water. Transfer the solution using a short stem funnel into a clean (not necessarily dry) 500.0 mL volumetric flask. Rinse the beaker several times with very small amounts of water, adding each rinse to the volumetric flask via the funnel until all KMnO4 has been transferred to the flask. As soon as all of the KMnO4 has been transferred to the volumetric flask carefully add water to bring the solution level to the calibration mark on the flask neck. Use a dropping pipet to add just enough water to bring the liquid level to the mark (note 3).
Holding the stopper securely in place, thoroughly mix the flask contents again by alternately inverting and returning the flask to an upright position at least 20 times.
Transfer the flask contents to the dry 500 mL reagent bottle by means of a dry short stem funnel. If the reagent bottle and funnel are not dry, rinse them with three very small (about 5 to 10 mL) portions of the solution before transferring the bulk of the solution to the bottle. You can afford to discard the rinse portions.
Label the reagent bottle with your name and date and keep the bottle stoppered and out of the light when the KMnO4 is not being used (note 4). Store the bottle of KMnO4 solution in your drawer.
The Standardization of 0.010 M KMnO4: This is YOUR part!!!
Prepare a tabular format in your laboratory notebook for recording the data that will be collected. You will need to record the mass of each sample of Na2C2O4 as well as mL initial and mL final of KMnO4 solution used to titrate each sample.
Pre-weigh between 0.12 and 0.13 g Na2C2O4 into a small beaker using a mg electronic top-loader balance.
Take the beaker with pre-weighed Na2C2O4 and a clean, dry 50 mL beaker to the analytical balance. Place the empty 50 mL beaker on the analytical balance. Press the bar once to tare it. Set the 50 mL beaker on your lab notebook and carefully pour the pre-weighed Na2C2O4 into it. Return the 50 mL beaker plus Na2C2O4 to the balance and record the mass of Na2C2O4 in the tabular format that you have prepared in your notebook. For best results handle the 50 mL beaker with a paper loop or with crucible tongs between the two weighings. This prevents contamination from fingers from affecting the mass of the sample.
The Na2C2O4 must now be quantitatively transferred to a 250 mL Erlenmeyer flask. The flask should be clean, but it need not be dry. The quantitative transfer is accomplished as follows: Pour 15 mL of water into the 50 mL beaker to begin dissolving the Na2C2O4. Carefully transfer the mixture of water and Na2C2O4 into the flask. Rinse the 50 mL beaker with 3 more 15 mL portions of water, transferring each to the flask. Finally pour 6 mL of 6 M sulfuric acid into the 50 mL beaker and transfer it to the flask. Inspect the beaker to make sure that the transfer is complete.
Heat the solution in the flask to just below the boiling point and begin the titration at this temperature (note 5). Avoid using a thermometer as a stirring rod.
While getting the solution up to temperature, rinse a clean buret 3 times with a few mL of your 0.010 M KMnO4 solution. Then fill the buret with the 0.010 M KMnO4, expel any air trapped in the tip, and take an initial buret reading (note 3).
Slowly titrate the hot Na2C2O4 solution with permanganate, swirling the solution constantly, being careful to wait until the solution is colorless before adding further amounts of permanganate (note 6).
Titrate to the first appearance of a faint pink that persists for about 30 seconds (note 7). The temperature should be above 60°C at the equivalence point. If the temperature is less than 60°C, heat it to about 70°C and add more KMnO4 solution, if necessary, to reach the equivalence point (Note 8).
Repeat the procedure with two more samples of Na2C2O4, weighed and titrated in a similar manner. The goal is to get 3 samples within a precision of 12 parts per thousand so that the relative average deviation from the mean will be < 6 ppt. (Note 9).
Do not let KMnO4 stand in the buret for long periods of time. As soon as all of your titrations are finished, thoroughly rinse the buret with water until all traces of KMnO4 are removed. Clean the buret with 3 M HCl if it appeals to have yellow or brown stains. Follow this by thoroughly rinsing with water.
Save your standardized KMnO4 solution for use in future experiments.
PERCENT OXALATE –
Notes on Experimental Procedure:
- The oxidation state of Fe in KxFey(C2O4)z·n H2O is 3+, normally the highest value for Fe. Thus the KMnO4 does not oxidize the Fe in this experiment. However, the presence of Fe3+ imparts a yellow color to the solution. The H3PO4 forms a colorless phosphate complex with Fe3+ making it easier to detect the color change which occurs at the equivalence point.
- The reaction of MnO4 - with C2O42- is rather slow at room temperature. The purpose of heating the solution is to increase the rate of the reaction.
Percent Oxalate Procedure:
Using a tabular format, prepare a form in your laboratory notebook for recording the experimental data.
Weigh to the nearest 0.001 g two samples, each of about 0.125 g (between 0.120 and 0.130 g) of crystals, into two marked 50 mL beakers. Use the same procedure that you used when you standardized the KMnO4 solution. Pre-weigh the samples in 50 mL beakers on the mg electronic top-loader balance; take the pre-weighed samples along with 2 additional clean, dry 50-mL beakers to the analytical balance; and determine the sample masses to the nearest 0.001 g.
Quantitatively transfer the crystals to marked 250 mL Erlenmeyer flasks using four 15 mL portions of water followed by 6 mL of 6 M H2SO4. Add 1 mL of 85% (conc.) H3PO4 to each sample (note 1).
Heat one of the two solutions to just below the boiling point (note 2).
While the solution is heating, rinse a clean buret with three small portions of your standardized KMnO4 solution. Fill the buret with your KMnO4 solution, expel air bubbles from the buret tip and take an initial buret reading. Make sure that the initial buret reading is not exactly 0.00 mL.
Remove the heat source and titrate the solution with the KMnO4.
The equivalence point is detected by observing the solution turn a very light pink color that persists for at least 30 seconds. The temperature should be above 60°C at the equivalence point. If it is not, heat it to about 70-80°C and add more KMnO4 solution, if necessary, to reach the equivalence point. Take a final buret reading.
Heat the second solution and repeat the titration procedure with this sample. Repeat with a third sample if necessary until you have 2 samples that agree within 11 parts per thousand.
Calculate the percent oxalate in KxFey(C2O4)z·n H2O.
DATA TABLES
Prepare an appropriate Data Table for each part of this lab. There should be THREE data tables. I have given some ideas below for parts of this lab –
Preparation of Potassium Permanganate
| Mass of KMnO4 used | _____________________
|
| Volume of solution prepared | _____________________
|
Standardization of Potassium Permanganate
| Run Number | 1 | 2 | 3 |
| Mass of Na2C2O4 | _________ | _________ | _________ |
| Initial Buret Volume | _________ | _________ | _________ |
| Final Buret Volume | _________ | _________ | _________ |
Percent Oxalate
| Experiment Number | 1 | 2 |
| Mass of Sample | _________ | _________ |
| Initial Buret Volume | _________ | _________ |
| Final Buret Volume | _________ | _________ |
RESULTS
Percent Water
| Mass of Sample | _____________________
|
| Mass of Water Lost | _____________________
|
| Percent Water in Sample | _____________________
|
Preparation of Potassium Permanganate
Apparent molar concentration based on the mass of KMnO4 _________________
Standardization of Potassium Permanganate
| Run Number | 1 | 2 | 3 |
| Moles of C2O42- | _________ | _________ | _________ |
| Volume of KMnO4 | _________ | _________ | _________ |
| Approximate moles of KMnO4 | _________ | _________ | _________ |
| use apparent M (above) for this calculation | |||
| Approx. mole ratio C2O42-/ MnO4 - | _________ | _________ | _________ |
| Average mole ratio to 3 sig. figs. | _________ | ||
| Ratio of moles C2O42-/ MnO4 - rounded to small whole numbers. | |||
| Compare with equations 2, 3 and 4 | |||
| Complete and balance equation 1: | |||
| ________MnO4– + ________C2O42– + ________ → ________CO2 + ________ + ________ | |||
| Accurate molar concentration of KMnO4 | _________ | _________ | _________
|
| Approx. mole ratio C2O42-/ MnO4 - | _________ | ||
Percent Oxalate
| Experiment Number | 1 | 2 | |
| mL KMnO4 required | _____________________ | _____________________ | |
| Moles KMnO4 added | _____________________ | _____________________ | |
| Mass of C2O42– in sample | _____________________ | _____________________ | |
| Percent C2O42– in sample | _____________________ | _____________________ | |
| Average percent C2O42– | _____________________ | ||
POST-LAB QUESTIONS
- Look up ALL possible chemical formulas for KxFey(C2O4)z·n H2O and list them Cite the source where you found each chemical formula.
- Based upon the information you have obtained so far, which known formula for KxFey(C2O4)z·n H2O most “matches” to your data. EXPLAIN.
- How did your apparent concentration of KMnO4 compare with the concentration you obtained upon standardization? What could have contributed to any differences?
REFERENCES
Brink, Irwin J., The Synthesis of an Iron Oxalato Complex Salt, NSF Summer Project in Chemistry, Hope
College, 1996.
Holley, Kathleen, The Synthesis of a Complex Iron Salt, The Oakridge School, Arlington, TX, 2006.
Sweeney-Hammond, Kathy, E-mail conversations and Handouts, Maret School, Washington, DC,
2005-2008.
