Exp. 25 Lab Report PDF

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2/21/18 Experiment 25. Formula and Formation Constant of a Complex Ion by Colorimetry

Purpose: The purpose of this experiment was to verify the formula of FeSCN^2+ and to determine its formation constant by using a spectrometer. Theory/Principles: In a dilute solution where there is a large amount of Fe3+ present, the Fe3+ will react with SCNto form a complex ion: Fe3+ (aq) + SCN-(aq)  FeSCN2+ (aq) (reaction is reversible). The equilibrium constant for this reaction is written as a formation constant kf: kf= [FeSCN2+ (aq)]/ [Fe3+ (aq)][SCN-(aq)]. In this experiment the solution which contains Fe(SCN)2+ absorbs a blue-green light at 400-500 nm and transmits a light that appears red at 500-700 nm. The color intensity all depends on the concentration of substance which absorbs the light which is called Beer’s Law. To define the light of a given wavelength with transmittance T is given by: T= I/Io where “I” is the intensity of the light transmitted and “Io” is the intensity of the light incident on the sample. Defining absorbance (A) also called optical density as: A= log1/T=logIo/I. Furthermore, Beer’s Law also states that the absorbance is proportional to both molar concentrations and distance that light travels through the solution given in the equation form of: A= e b c. Where e depends on the molecule absorbing light and the wavelength chosen by using a spectrometer to determine the measurement. To determine the concentration of an unknown solution through using Beer’s Law we are given the equation of: Cunknown= Aunknown/Aknown x Cknown, where Aknown is of a known compound, Cknown is a wavelength, and Aunknown measure absorbance of another solution that contains the same compound. Once we obtain our data’s, plot A vs λ, and plot our remaining solutions we prepared at the highest wavelength we should be able to connect the points with a smooth curve and deduced the stoichiometry of the reaction between Fe(III) and SCN-. Next we can calculate the

concentrations of iron(III) thiocyanate from the our solutions in test tubes B2, B3, and B4 by using: [FeSCN2+]= A/Astd [FeSCN2+]std. Finally, we can calculate the values for Kf by their measured absorbance’s using the formula given earlier of Cunknown and [FeSCN2+] that is found in each of these solutions. Experimental Procedures: -

Just changed one step from the procedure which is we used burets instead of pipets for getting our solutions.

Goldwhite, H.; Tikkanen, W. Experiment 25. Formula and Formation Constant of a Complex Ion by Colorimetry, Experiments in General Chemistry, 4th ed.;The McGraw Hill Companies. (151-153) Data Tables/Summary: -

Table #1: Raw data collected during experiment for parts A/B & signed pre-lab

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Table #2: Plot for λ (nm) vs A3 (Absorbance’s)

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Table #3: Plot for Fe(III) mL vs. Absorbance at λ max

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Table #4: Plot for SCN- mL vs. Absorbance at λ max

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Table #5: Data from part B of experiment including calculated Kf values

Results & Discussions: Sample Calculations: Calculating Kf value for B1:B2: Kf= [Fe(SCN-)^2+]/[Fe^3+][SCN-] mol SCN-=1.0 × 10^-3L(1.0×10-3 M)= 1.0 × 10^-6 mol SCN- = mol FeSCN2+ [FeSCN^2+] std = 1.0 × 10-6 mol FeSCN^2+/ 0.010 L = 1.0 × 10-4 M [FeSCN^2+] = A/A std [FeSCN^2+] std [FeSCN^2+]= 0.359/ 0.436 (1.0 × 10^-4 M)= 8.2 × 10^-5 M [Fe^3+] initial = molFe^3+/ V(L) = 0.5 x 10^-3(0.20 mol/L)/ 0.010 L= 1.0 × 10^-2 M [Fe^3+] equilib. = 1.0 × 10^-2 M- (8.2 × 10^-5 M)= 9.91 × 10^-3 M [SCN-] equilib. = 1.0 × 10-4 M- (8.2 × 10^-5 M)= 1.8 × 10^-5 M Kf= 8.2 × 10^-5 M/ (9.91 × 10^-3 M)(1.8 × 10^-5 M) = 459.7 Average Kf values: Average= 459.7 + 157.9 + 10.201/ 3 = 209.3 Logical Explanation: By determining the formula for iron (III) thiocyanate by using a spectrometer to obtain the absorbance’s for our solutions I was able to calculate the formation constants for B2, B3, and B4. As we can see in my graph of “Plot for λ (nm) vs A3 (Absorbance’s)” my highest wavelength

was at 450 nm and with this information we were able to obtain the rest of the absorbance measurements. With the three plots I gave above, they helped to determine the ratio of the reactants that was able to give me an idea of the stoichiometry of the reaction happening. Based off my Kf values we can see that solutions B2 and B3 gave the highest constants while B4 gave the lowest. However, the Kf values are not nearly all the same which can be due to an error of not accurately obtaining the solutions needed for each. Although, my average formation constant was 209.3, showing me that the reaction went to completion because there was a larger amount of Fe3+ than SCN- causing all of SCN- to be used up. Discussion Ques.: 1. The validity of the assumption we made for solution B1 was that thiocyanate ion present was nearly completely complexed because the average Kf value was greater than 1. Therefore, in my solutions, B4 had the lowest Kf value while B2 and B3 had the highest Kf values. Thus, B2 and B3 are valid. 2. It would have been difficult to obtain reliable results if the absorbance’s had been measured at 675 nm because in this experiment we needed to measure the absorbance at the wavelength that absorbs the most light, would give the highest absorbance per unit concentration, and highest sensitivity. For this reason my absorbance’s were measured at 450 nm. 3. a) For an endothermic reaction between two substances in solution based on their heat changes we can know what type of reaction it is. b) For a reaction between two gaseous reactants to produce gaseous products where ∆n≠0, from the equilibrium constant Kp we can know the product formation and equilibrium.

c) For a reaction between two strong electrolytes in solution to form a soluble, nonionized product we wouldn't able to use this because this situation would change the ionic strength. 4. Uncertainty for Buret: ±0.05 mL

∑= (9±0.05)+(0.9±0.05)+(0.5±0.05)+(6±0.05)+(1±0.05)+(1±0.05)+(1.11+±0.05)+(0.9±0.05)/ 81 = 2.97 √2.97= 1.72 Conclusion: By using verifying the formula of FeSCN^2+ and determining its formation constant by using a spectrometer I was about to collect the needed absorbance’s in order to calculate my three Kf values and its average value of 209.3. References: Goldwhite, H.; Tikkanen, W. Experiment 25. Formula and Formation Constant of a Complex Ion by Colorimetry, Experiments in General Chemistry, 4th ed.;The McGraw Hill Companies. (149-154)...