BACKGROUND Most chemical reactions are reversible. They will proceed forward to a point where the products they have formed begin to collide with one another and reform the original reactants. When the reaction begins, and only reactants are present, the forward reaction occurs quickly. As products begin to form, the reverse reaction accelerates until the two reactions, forward and reverse, are occurring at the same rate. At equilibrium, the concentrations of reactants and products become constant. A ratio of products to reactants may be used to describe the equilibrium position. This ratio, the equilibrium expression, is written as the product of the product concentrations divided by the product of the reactant concentrations with each species raised to a power equal to its coefficient in the balanced equation. Pure solids and liquids are not included in equilibrium expressions because their concentrations do not change. The value of this ratio is called the equilibrium constant, K eq. In this experiment, you will determine the value of K c for the reaction between iron (III) ions and thiocyanate ions, SCN. Fe 3+ (aq) + SCN (aq) FeSCN 2+ (aq) iron (III) thiocyanate thiocyanatoiron (III) When Fe 3+ and SCN are combined, equilibrium is established between these two ions and the FeSCN 2+ ion. In order to calculate K c for the reaction, it is necessary to know the concentrations of all ions at equilibrium: [FeSCN 2+ ] eq, [SCN ] eq, and [Fe 3+ ] eq. You will prepare four equilibrium systems containing different concentrations of these three ions. The equilibrium concentrations of the three ions will then be experimentally determined. These values will be substituted into the equilibrium constant expression to see if K c is indeed constant. In order to determine [FeSCN 2+ ] eq, you will use the Colorimeter shown in at right. The FeSCN 2+ ion produces solutions with a red color. Because the red solutions absorb blue light very well, the blue LED setting on the Colorimeter is used. The Colorimeter measures the amount of blue light absorbed by the colored solutions (absorbance, A). By comparing the absorbance of each equilibrium system, A eq, to the absorbance of a standard solution, A std, you can determine [FeSCN 2+ ] eq. The standard solution has a known FeSCN 2+ concentration. To prepare the standard solution, a very large concentration of Fe 3+ will be added to a small initial concentration of SCN (hereafter referred to as [SCN ] i. The [Fe 3+ ] in the standard solution is 100 times larger than [Fe 3+ ] in the equilibrium mixtures. According to LeChatelier s principle, this high concentration forces the reaction far to the right, using up nearly 100% of the SCN ions. According to the balanced equation, for every mole of SCN reacted, one mole of FeSCN 2+ is produced. Thus [FeSCN 2+ ] std is assumed to be equal to [SCN ] i. Assuming [FeSCN 2+ ] and absorbance are related directly (Beer s law), the concentration of FeSCN 2+ for any of the equilibrium systems can be found by: [FeSCN 2+ ] eq = Knowing the [FeSCN 2+ ] eq allows you to determine the concentrations of the other two ions at equilibrium. For each mole of FeSCN 2+ ions produced, one less mole of Fe 3+ ions will be found in the solution (see the 1:1 ratio of coefficients in the equation above). The [Fe 3+ ] can be determined by: [Fe 3+ ] eq = [Fe 3+ ] i [FeSCN 2+ ] eq Because one mole of SCN - is used up for each mole of FeSCN 2+ ions produced, [SCN - ] eq can be determined by: [SCN ] eq = [SCN ] i [FeSCN 2+ ] eq Knowing the values of [Fe 3+ ] eq, [SCN ] eq, and [FeSCN 2+ ] eq, you can now calculate the value of K c, the equilibrium constant. A eq A std x [FeSCN 2+ ] std Safety The Fe(NO 3 ) 3 solution is prepared with nitric acid (HNO 3 ) as the solvent. Nitric acid is highly corrosive to skins and eyes. Potassium thiocyanate is toxic by ingestion and inhalation. It reacts violently with strong oxidants and decomposes to form toxic substances when heated strongly. Spills should be neutralized and cleaned up immediately. Any solutions that contact skin should be rinsed off with plenty of water. Goggles and aprons must be worn.
Using the GoDirect Colorimeter 1. Plug the GoDirect colorimeter into an available USB port on the laptop. 2. Log into the laptop computer and open the Vernier Graphical Analysis application from the desktop. 3. Calibrate the colorimeter. a. Prepare a blank by filling an empty cuvette 3/4 full with distilled water. Seal the cuvette with a lid. Working with Cuvettes All cuvettes should be wiped clean and dry on the outside with a tissue before inserting into the colorimeter. Handle cuvettes only by the top edge of the ribbed sides. All solutions should be free of bubbles. Tap on table top to remove bubbles. Always position the cuvette with its reference mark facing toward the white reference mark at the top of the cuvette slot on the Colorimeter. b. Place the blank in the cuvette slot of the Colorimeter and close the lid. c. Press the < or > button on the Colorimeter to set the wavelength to 470 nm (Blue). Then calibrate by pressing the CAL button on the Colorimeter. When the LED stops flashing, the calibration is complete. d. Empty the water from the cuvette and set aside to dry. 4. Click on the upper right corner Untitled. a. Select New Experiment from the menu. b. Select Sensor Data Collection c. Click on Graph button in upper right corner. d. Select Meter e. Collect data.
Below is a calculation synopsis if you need additional explanation beyond class instruction. 1. Start each RICE table with the following equation on the REACTION line: Fe 3+ (aq) + SCN (aq) FeSCN 2+ (aq) 2. Calculate the initial concentration of Fe 3+, based on the dilution that results from adding KSCN solution and water to the original 0.0020 M Fe(NO 3 ) 3 solution. See the small table in the procedure for the volume of each substance used in Trials 1 4. Calculate [Fe 3+ ] i using the equation: [Fe 3+ ] i = Fe(NO 3) 3 x (0.0020M) total ml This value will be the same for all four test tubes and is on the INITIAL line of your RICE table. 3. Calculate the initial concentration of SCN, based on its dilution by Fe(NO 3 ) 3 and water: [SCN ] i = KSCN total ml x (0.0020M) For example, in Test Tube 1, [SCN ] i = (2 ml / 10 ml)(.0020 M) =.00040 M. Calculate this for the other three test tubes. These values will be different for each trial and is also in the INITIAL line of your RICE table. 4. [FeSCN 2+ ] eq is calculated using the formula: [FeSCN 2+ ] eq = A eq A std x [FeSCN 2+ ] std Where A eq and A std are the absorbance values for the equilibrium and standard test tubes, respectively, and [FeSCN 2+ ] std = (1/10)(0.0020) = 0.00020 M. Calculate [FeSCN 2+ ] eq for each of the four trials. This is the equilibrium concentration of [FeSCN 2+ ] eq to be put on the EQUILIBRIUM line of your RICE table and will be different for every trial. 5. With all that data substituted into the RICE tables you should now be able to solve for each of the equilibrium concentrations and for K c as well. You will do this four different times for four different RICE tables.
I. PURPOSE To determine the equilibrium constant K c for an equilibrium system using spectrophotometry. II. MATERIALS 1. 0.200 M Fe(NO 3 ) 3 2. 0.0020 M Fe(NO 3 ) 3 3. 0.0020 M KSCN 4. Culture cups (5) 5. Thermometer 6. Lens paper 7. Cuvettes (5) III. PROCEDURES A. Prepare the solutions for testing. 1. Fill the appropriately labeled cups with 0.0020M Fe(NO 3 ) 3 and 0.0020M KSCN and water from the stock bottles. 2. Line up the five cups labeled 1-5. Utilizing the table below as your guide, make the four solutions using the labeled syringes provided. Mix each solution thoroughly by gently stirring the cups. Test Solution # 0.0200 M Fe(NO 3 ) 3 0.0020 M KSCN 1 5.0 2.0 3.0 2 5.0 3.0 2.0 3 5.0 4.0 1.0 4 5.0 5.0 0 Distilled Water 3. Prepare a standard solution of FeSCN 2+ in cup #5 by mixing 18mL of 0.200M Fe(NO 3 ) 3 with 2mL of 0.0020M KSCN. B. Calibrate the colorimeter. Follow the guidelines for using a colorimeter listed in the background section. 1. Prepare a blank by filling an empty cuvette 3/4 full with distilled water. Seal the cuvette with a lid. 2. Place the blank in the cuvette slot of the Colorimeter and close the lid. 3. Press the < or > button on the Colorimeter to set the wavelength to 470 nm (Blue). Then calibrate by pressing the CAL button on the Colorimeter. When the LED stops flashing, the calibration is complete. 4. Empty the water from the cuvette and set aside to dry. C. Collect absorbance data. 1. Using the solution in cup #1, rinse a cuvette twice with ~1 ml amounts and then fill it 3/4 full. Place the cap on the cuvette and wipe the outside with a tissue, place it in the colorimeter, and close the lid. 2. When the value displayed on the screen has stabilized, record the absorbance value in your data table. 3. Repeat Steps C.1 & C.2 with each of the remaining solutions (#2-#5). IV. PRE-LAB QUESTIONS 1. How is an equilibrium constant calculated? What does its magnitude imply about the extent of reaction completion? 2. Describe the dilution formula M 1 V 1 = M 2 V 2. Why does it work mathematically? 3. Using the Table in Procedure 2 and M 1 V 1 = M 2 V 2, calculate the initial concentration of Fe(NO 3 ) 3 used in Trials 1 4. Fill in values in the Data Table. 4. Using the Table in Procedure 2 and M 1 V 1 = M 2 V 2, calculate the initial concentration of KSCN used in Trials 1 4. Fill in values in the Data Table. V. DATA & CALCULATIONS A. DATA [Fe(NO 3 ) 3 ] [KSCN] Absorbance Temperature Data Table Trial 1 Trial 2 Trial 3 Trial 4 Trial 5
B. CALCULATIONS 1. Write the K c expression for the reaction. 2. Construct a RICE table for each trial in order to calculate its K c. 3. Calculate K c for each trial. Show ALL work. 4. Complete the following table using the information from Calculations 2 & 3. K c expression: K c = K c value: Calculation Table Trial 1 Trial 2 Trial 3 Trial 4 Average of K c values: K c = at C 5. Determine the percent error based on the accepted value of 162. Was your constant a constant? VI. POST-LAB QUESTIONS 1. Why was the temperature of the solutions taken in Part A? How would variations in the temperature affect the data? 2. A student fails to wipe their fingerprints off the cuvette before they put it into the colorimeter correctly. What if any effect will this have on the measured absorbance? Explain. 3. What does the magnitude of the K c calculated in this lab imply about this reaction? 4. What is the limiting reactant for this laboratory exercise? (See mixture table in the Procedure 2 for guidance.) 5. What is the name and formula of the colored species being measured in this laboratory exercise? 6. Why is it important to fill the cuvette to about 3 4 full? 7. The wavelength setting you are told to use in this lab is not chosen arbitrarily. A previous experiment was conducted to determine the best wavelength setting. The experiment used a spectrophotometer and collected absorbance data for a given sample of FeSCN 2+ at 10 nm wavelength increments starting at 360 nm and ending at 620 nm. The graph of the data is shown at right. Explain how this graph is used to determine the wavelength setting for this experiment. VII. CONCLUSION