Exp14_prelab - Experiment 14 Determination of an...

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Unformatted text preview: Experiment 14 Determination of an Equilibrium Constant in Aqueous Solution Learning Objectives Understand what is meant by the term chemical equilibria Know how to determine an equilibrium constant Kc of a chemical reaction Understand the mathematical relationship between absorbance and concentration using the Beer-Lambert Law Many Chemical Reactions are Reversible… … meaning that the reaction can proceed from reactants to products and then back, or in other words, A+B C and A+B C (more commonly depicted as A+B C) A reaction is said to be in equilibrium when the rate of the reaction in the forward direction is equal to the rate in the reverse direction. Chemical Equilibrium Consider the reaction below N2O4(g) 2NO2(g) Kinetically, Rate of the forward reaction, Ratef = kf[N2O4] Rate of reverse reaction, Rater = kr[NO2]2 At equilibrium, kf[N2O4] = kr[NO2]2 Rearranging this we get, [NO2]2/[N2O4] = kf/kr = a constant This constant that we have solved for is called the equilibrium constant, Kc Equilibrium Constant In more complex reactions, exemplified by aA + bB cC + dD The equilibrium constant Kc = [C]c[D]d/[A]a[B]b Where all concentrations are equilibrium concentrations Sample Calculation – Equilibrium Concentration Consider the reaction, Fe3+(aq) + SCN-(aq) FeNCS2+(aq) If the equilibrium concentration of Fe3+(aq) is 1.00x10-4M, in a solution that was initially 2.06x10-3M Fe3+(aq), and 3.00x10-3 M SCN(aq), determine the equilibrium concentrations for SCN-(aq) and FeNCS2+(aq). Sample Calculation – Equilibrium Concentration Fe3+(aq) Initial Concentration (M) 2.06x10-3M + SCN-(aq) 3.00x10-3 M FeNCS2+(aq) 0 Concentration (M) Equilibrium Concentration (M) 1.00x10-4M First calculate the Concentration for Fe3+(aq). 2.06x10-3M - 1.00x10-4M = 1.96x10-3M Since this reaction is 1:1, we also know that Concentration for SCN- (aq) and FeNCS2+(aq) is 1.96x10-3M Sample Calculation – Equilibrium Concentration Fe3+(aq) Initial Concentration (M) Concentration (M) Equilibrium Concentration (M) + SCN-(aq) FeNCS2+(aq) 2.06x10-3M 3.00x10-3 M 0 -1.96x10-3M -1.96x10-3M +1.96x10-3M 1.00x10-4M Now, since we know how much the concentration of FeNCS2+(aq) increased we simply add Concentration to the initial concentration. Likewise we can subtract the Concentration from the initial concentration of SCN-. Equilibrium Concentration SCN- = 3.00x10-3 M - 1.96x10-3M = 1.04x10-3M Equilibrium Concentration FeNCS2+ = 0.00M + 1.96x10-3M = 1.96x10-3M Sample Calculation – Equilibrium Constant Fe3+(aq) Initial Concentration (M) Concentration (M) Equilibrium Concentration (M) + SCN-(aq) FeNCS2+(aq) 2.06x10-3M 3.00x10-3 M 0 -1.96x10-3M -1.96x10-3M +1.96x10-3M 1.00x10-4M 1.04x10-3M 1.96x10-3M Now, since we know the equilibrium concentrations we can use them to determine the equilbrium constant. K eq [FeNCS2+ ] [1.96x10 -3M] = = = 18800 3+ -4 -3 [Fe ][SCN ] [1.00x10 M][1.04x10 M] NOTE: These are just sample concentrations, the equilibrium constant for this reaction is not 18800. Determining Equilibrium Concentrations Indirectly Using Colorimetry While not applicable to all reactions, some reactions can be monitored by colorimetry Colorimetry measures the amount of light at a given wavelength that is transmitted through a solution In solution some light is absorbed by solute molecules, while the rest is transmit. The amount absorbed is usually related to the amount of solute present in solution. By measuring the amount transmit, the amount absorbed by the solution is also known. The amount of absorbed light is then used to determine the concentration of the solute in the solution. The mathematical relationship used to determine concentration in colorimetry is called the Beer-Lambert Law The Beer-Lambert Law A= bc A= Absorbance = -log(I/Io), (no units) I = light intensity of sample Io = light intensity for the blank solution = Molar extinction coefficient – how strongly a solution absorbs at a particular wavelength of light, (M-1cm-1) b = path length of light through the solution in this case b=1cm c = concentration in molarity (mol/L) Constructing a Standard Calibration Curve It is important to construct a calibration curve in order to experimentally determine the value for the molar extinction coefficient To do this, the absorbance of known concentrations of a sample are analyzed and plot on a graph Since A= bc, and A is plot against C, the slope of the line is equal to b, the molar extinction coefficient multiplied by the path length Using a Standard Calibration Curve By using the value for b, as well as absorbance determined experimentally we can determine the concentration. What is the concentration of solution whose absorbance is 0.62? We know that b = 5M-1 from the calibration curve and A= 0.62 If we solve for concentration, we get C=A/ b, or C=0.62/ 5M-1= 0.12 Experimental Guidelines An accurate standard curve is important, make sure to determine the concentrations of each of your solutions very carefully Make sure that the colorimeter is set up properly otherwise your samples may need to be analyzed multiple times Summary You will generate a standard curve for one component of a reaction that you will be observing. Using the Beer-Lambert Law, the absorbance of your solution, and the standard curve you made, you will be able to determine the concentration of one component in the reaction solution. Equilibrium concentrations of the remaining reaction components can be determined by using the equilibrium concentration determined using the Beer-Lambert Law, and the initial concentrations of the other reaction components. The equilibrium constant can then be calculating using the equilibrium concentrations that you determined mathematically ...
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This note was uploaded on 02/06/2010 for the course CHEM 102L taught by Professor N/a during the Spring '07 term at UNC.

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