Question

Can't figure out the concentration of the complex of solns. 10, 11, 12, 13, & 14.

Modern Experimental Chemistry Chemistry 153 The Iron(III) Thiocyanate Complex Purpose of the Experiment To determine the chemical formula of a complex ion and measure its formation equilibrium constant Equipment Spectro Vis spectrophotometer and LabQuest, cuvette, 25-mL buret (3), ring stand, buret clamp (2), 50-mL beaker (3), small plastic beakers Reagents SCN, as KSCN, 0.00200 M solution Fe; as Fe(NO), 0.00200 M solution and a 0.200 M solution HNO, 1.5 M solution and a 0.4 M solution INTRODUCTION The iron(IIl) thiocyanate complex is labile and only exists in solution. Equilibrium concentrations of the complex and the reactants are established if iron(III) ion, Fe", and thiocyanate ion, scN, are initially present at constant temperature. Metal cations exist in aqueous solution in the form of complexes with water molecules. In the case of the iron(LII) cation, with a coordination number of 6, the iron(III) is in the form of the hexaaquairon(III) complex cation. The reaction between Fe" and SCN in aqueous solution is indicated below where n is a whole number between 1 and 6 (inclusive) whose value depends on the ratio of SCN to Fe in the most stable form of the complex Since the iron(II) and thiocyanate ions contribute very little to the absorbance of visible light at a wavelength of 450 nm, it is possible to use the intense absorbance of the deep red iron(III) thiocyanate complex at 450 nm to measure its concentration without interference from the uncomplexed ions. For example, the composition of the complex may be determined by observing the molar ratio of thiocyanate ion to iron(IIl) ion that produces the maximum yield of complex. Since neither ion contributes to the absorbance, if the total concentration of iron(Ill) ion and thiocyanate ion is kept constant while the molar ratios of the two are varied, the composition of the complex will be given by the molar ratio which produces the maximum absorbance After determining the composition of the complex, data will be measured which will allow the construction of a Beer's Law plot so that the concentrations of the complex may be calculated from measured absorbances. In part 2 of the procedure, several solutions of iron(lII) thiocyanate complex will be formed from mixtures of a low concentration of thiocyanate ion The Iron(III) Thiocyanate Complex 101 Exp 16 Chemistry 153 Modern Experimental Chemistry with a large excess of iron(III) ion. Under these conditions, it is reasonable to assume that all of the thiocyanate ion is consumed in forming the iron(III) thiocyanate complex. This allows the direct calculation of the complex concentration from the initial concentrations and volumes of the solutions mixed together. A plot of the absorbance of each of these solutions versus their calculated concentrations will be used to generate a Beer's Law Plot, from which the complex concentration can be determined from the measured absorbance of other complex solutions. In part 3 of the procedure, the absorbance of four more iron(lII) thiocyanate solutions is measured. From the determination of the composition (part 1), the stoichiometry is known and with the initial concentrations, the equilibrium concentrations can be calculated once the complex concentration is determined from the absorbance using the absorptivity from part 2. Overview The composition of the complex is determined. 2. A Beer's Law plot for the iron(III) thiocyanate complex is constructed 3· The formation equilibrium constant, Kf, is calculated from spectral data. 1· Using three burets (attach two buret clamps, at right angles to each other, to a single ring stand), make the following 9 measurements in order to determine the composition of the complex. One partner can carefully prepare the solutions by delivering the reagents from the burets into a small plastic beaker. The other partner can then measure and record the spectrophotometric data. Record absorbance readings to 3 decimal places. Read Instructions for Use of Vernier LabQuest2 and SpectroVis Spectrophotometer DANGER: HNO,CORROSIVE Be meticulous in the measurements and in the recording of data. Use care in preparing the solutions with the correct concentrations. To avoid contamination, burets and beakers must be rinsed thoroughly with distilled water before changing reagents. 1. Determining Compositions V of0.00200 M | V of 0.00200 M Fe (mL) V of 1.5 MM HNO, (mL) Soln # 8.00 2.00 0.00 2.00 7.20 0.80 2.00 6.40 1.60 5.60 4 2.00 4.80 3.20 2.00 4.00 4.00 3.20 4.80 2.00 1.60 2.00 0.00 8.00 102 Exp 16 The Iron(III) Thiocyanate Complex Chemistry 153 Modern Experimental Chemistry Discard the solution from the cuvette after each absorbance measurement. Use a small portion of the new solution to rinse the cuvette before each new measurement. Record the absorbance es for each of the nine solutions in Table I at the top of the "Data and Report Sheet". This e also contains columns for recording calculated values necessary for determining the composition of the iron( III) thiocyanate complex 2. Determining the Molar Absorptivity by a Beer's Law Plot Change the iron(III) solution in the first buret (note: the concentration of iron(III) used in this part is 100 times greater than that used in part 1!) and change the last buret to pure H,O. Measure the absorbance for each of these solutions using the technique of rinsing the cuvette with the appropriate solution before each measurement. Record absorbance values in Table 2 on the DATA AND REPORT SHEET Fe (mL) SCN (ml (mL) 1.00 7.50 1.50 10 1.00 1.50 7.50 1.70 7.50 0.80 0.40 7.50 13 0,00 2.50 7.50 ) For the Beer's Law plo, assume that all ofthe scr is converted to Fe(SCN. NOTE: Measuring the Equilibrium Concentrations and Determining the Formation Constant 3. Measure the absorbance of these last four solutions to complete the required laboratory measurements. Record absorbance values in Table 3 on the DATA AND REPORT SHEET V of 04 M HNO, (mL) V of0.200 M Fe (mL) İ V of O 00200 M | So lni 7.00 1.00 2.00 15 8.00 1.00 1.00 16 6,00 1.00 17 3.00 5.00 1.00 4.00 103 The Iron(III) Thiocyanate Complex Exp 16 Chemistry 153 Modern Experimental Chemistry Calculations 1. Determining Compositions Calculate the initial moles of Fe" and ScN for all nine s total number of moles (the sum of the moles of Fe" p Calculate the mole fractions of thiocyanate ion, zscN, and iron(II) ion, Table 1 solutions and record in Table 1. The plus the moles of SCN-) should be constant. Using the data from these solutions, create a plot of absorbance versus the mole fraction of SCN . The maximurn in this plot will occur at the mole fraction of SCN, χ-present in the most stable complex. If the stable form of the complex has the formula [Fe(SCN, thern from the definition of mole fraction, Determine xfrom your plot and use it to calculate the value of n using the above equation. Remember, n must be an integer. Use this value of n in subsequent calculations. 2. Determining the Molar Absorptivity from a Beer's Law Plot Make a plot of the solution absorbance versus the molar concentration of the iron(ll) thiocyanate complex assuming that all of the SCN is conv as unconnected data points. The concentration may be calculated by to the complex because of the large excess of Superimpose on the plot the straight line passing through the origin that best fits g using EXCEL, this is done by adding a trendline to the plot. In the the data points. If plottin Format Trendline pop-up box, check the three boxes at the bottom: Set Intercept -0.0, Display Equation on chart, and Display R-squared value on chart. The slope of this best-fit line is the molar absorptivity, e, of the complex. Measuring the Equilibrium Concentrations and Determining the Formation Constant 3. Knowledge of the stoichiometry, initial concentrations, and the equilibrium concentration of the iron(III) thiocyanate complex allows the calculation of the formation equilibrium constant. x, [complex] (0.2) where [complex] is the equilibrium concentration of the [Fe(SCN).J complex. This equation for Kr is valid only if all of the concentrations used in it are equilibrium concentrations. The equilibrium concentration of the complex for a given solution can be calculated from Beer's Exp 16 The Iron(III) Thiocyanate Complex 104 Chemistrv 1s2 Chemistry 153 Modern Experimental Chemistry Law by using its measured absorbance along with the value of e from part 2 of the calculations. [complex]-Ale (0.3) Once the value of [Fe(SCN),(3-a)] is known for the solution, the equilibrium concentrations of Fe" and SCN can be calculated easily from the stoichiometry of the complex formation: [Fe]-[Feb-[complex) (0.4) and (0.5) [SCN 1-[SCN-complex] Perform these calculation for solutions 15 through 18 and solutions 2 through 8, and enter the results in the appropriate places in Table 3 of the DATA AND REPORT SHEET 105 Exp 16 The Iron(III) Thiocyanate Complex The Iron(III Thiocyanate Complex DATA AND REPORT SHEET 1. Determining Compositions Soin # Mols Fe" Mols SCN Abs (at ScN 1.6%10:5 | 0.008 | 106 o.093 0.4oo 0.40 16 x 10 0.5oo o.500 103416 I 21 K16 0.184 0a00 o 5 Solution # yielding maximum absorbance above: 2 Value of n in [Fe(SCN) n+1 2. Determining the Molar Absorptivity by a Beer's Law Plot Soln # 12 14 complex (M) Abs (at λ .. 450nm)| 1369 |0.974 | Ο.Sb7|0.185 |-o.co6 Molar absorptivity, a (units)

Answer 1

Since the number of moles of SCN^- is a lot less than the number of moles of Fe³⁺, number of moles of the complex= number of moles of SCN^-

For complexes 10,11,12, 13 and 14, total volume of the solution- 10 mL

Complex number	No. of moles of SCN-	No. of moles of the complex	Volume of the solution	Concentration of the complex
10	3 X 10-6	3 X 10-6	10 mL	3 X 10-4
11	2 X 10-6	2 X 10-6	10 mL	2 X 10-4
12	1.6 X 10-6	1.6 X 10-6	10 mL	1.6 X 10-4
13	8 X 10-7	8 X 10-7	10 mL	8 X 10-5
14	0	0	10 mL	0