Lab Report Ex 3 Fe(III) Oxalate Complex PDF

Preview

Summary

Lab report for experiment 3...

Description

1 General Chemistry 1 Laboratory Report Name:______Ivan Chan____________________ Date of experiment______2/15/19_______

Lab day____Friday_____ TA:_______Fang________

3. Preparation and Analysis of an Fe (III)-oxalate Complex Purposes of the experiment and brief overview of the experimental method(s) The overarching purpose of this experiment is to prepare and purify a complex of Fe3+ and oxalate and determine its composition. Strong Lewis acid FeCl3 and aqueous K2C2O4 are reacted together, then the techniques of recrystallization and filtration were used to purify the iron (III) oxalate crystals formed. In order to quantitatively analyze the complex and determine its composition and empirical formula, application and verification of the Beer-Lambert Law through spectrophotometry was used to determine the percent composition of Fe (III), while a redox titration with permanganate in acidic solution was used to determine the percent composition of oxalate. In order to apply the Beer-Lambert Law, Fe (III) was reduced to Fe (II) using hydroxylamine through the equation:

This resulting Fe (II) solution was complexed with o-phen to create a bright red solution more suitable for spectrophotometry. The redox titration of oxalate with permanganate in acidic solution followed the equation: The empirical formula of the Fe(III)-oxalate complex was then determined through hydrate composition and charge balance.

2

Results and Calculations – Pay attention to sigfigs and units throughout. Qualitative Tests: For each test, write briefly what you did, and your observations. Write net ionic equations for the reactions of interest in all these tests. See example below. If NH 3 (aq) is the source of hydroxide ions, you should include its hydrolysis as well. Example (Test 1b): FeCl3 (aq) + NH3 (aq) resulted in a reddish orange gelatinous precipitate. Net ionic equations of interest: NH3 (aq) + H2O ()  NH4+ (aq) + OH− (aq) Fe3+ (aq) + 3 OH− (aq)  Fe(OH)3 (s, red/bron, gelatinous) Known ions Test 1a & 1b 1a. Sulfosalicylic acid + FeCl3 resulted in a deep red wine color. Net ionic equation:

1b. FeCl3 (aq) + NH3 (aq) resulted in a reddish orange gelatinous precipitate. Net ionic equation(s): NH3 (aq) + H2O ()  NH4+ (aq) + OH− (aq) Fe3+ (aq) + 3 OH− (aq)  Fe(OH)3 (s)

Flame Tests (net ionic equation is not critical) Na+: NaCl (s) + H2O () + heat resulted in an orange flame K+: KCl (s) + H2O () + heat resulted in a lavender flame

Test 3a & 3b 3a. CaC2O4 (aq) + Ca(NO3)2 (aq) resulted in a foggy white precipitate. Net ionic equation: C2O42- (aq) + Ca2+(aq)  CaC2O4 (s, foggy white precipitate) 3b. KMnO4 (aq) + H2SO4 (aq) + CaC2O4 (aq) + heat resulted in a colorless solution. Net ionic equation: 2MnO4- (aq) + 5H2C2O4 (aq)  6H+ (aq) + 10CO2 (g) + 2Mn2+ (aq) + H2O ()

3

Qualitative tests on your Fe-oxalate complex – In the spaces below, write what you did and observations you made in each test, and your explanation / reasoning for the observation(s). Write relevant balanced net ionic reactions for your observations. Test 1a & 1b 1a. Fe3+-C2O4 complex (aq) + sulfosalicylic acid (aq) resulted in a deep red-orange color. Net ionic equation:

1b. Fe3+-C2O4 complex (aq) + NH3 (aq) resulted in a rusty orange precipitate. Net ionic equation(s): NH3 (aq) + H2O ()  NH4+ (aq) + OH− (aq) Fe3+ (aq) + 3 OH− (aq)  Fe(OH)3 (s) Flame test Fe3+-C2O4 complex (aq) + heat resulted in a slight lavender color. Test 3a & 3b 3a. Fe3+-C2O4 complex (aq) + Ca(NO3)2 (aq) resulted in a foggy white precipitate. Net ionic equation: C2O42- (aq) + Ca2+(aq)  CaC2O4 (s, foggy white precipitate) 3b. KMnO4 (aq) + H2SO4 (aq) + Fe3+-C2O4 complex (aq) + heat resulted in a colorless solution. Net ionic equation: 2MnO4- (aq) + 5H2C2O4 (aq)  6H+ (aq) + 10CO2 (g) + 2Mn2+ (aq) + H2O ()

Quantitative Analysis Part 1 [Fe2+] in the standard solution used to prepare the calibrating solutions = 1.78 x 10-4 mol L-1 Volume of 10 % (mass / volume) sodium acetate solution used in each of the solutions prepared for absorbance measurements = 1.00 mL Volume of 10 % hydroxylamine hydrochloride solution used in each of the solutions prepared for absorbance measurements

= 1.00 mL

Volume of 0.1 % o-phen solution used in each of the solutions prepared for absorbance measurements

= 1.00 mL

Total (final) volume of each calibrating solution = 10.00 mL

4 Visible spectrum of Fe(o-phen)32+ complex Plot absorbance on the vertical axis versus wavelength (range: 400 to 700 nm) for all the four calibrating solutions. Plot all four spectra on the same graph. Mark on the graph the wavelength that you chose for the analysis; this will be around 510 nm. Be sure to include a caption to this graph, and all other graphs as well; see the manual for details. 2.5

2

1.99

Absorbance

1.5 1.28 1 0.78 0.5 0.31 0 400

450

500

550

600

650

700

750

-0.5

Wavelength (nm)

Figure 1. The visible spectrum of Fe(o-phen)32+ complex solutions at different concentrations is pictured above. The absorbance of four solutions of varying concentrations was captured through spectrophotometry and plotted versus wavelength in nanometers. The absorbance at 510 nm labeled for each solution acts as a reference point for comparison and later verification of the Beer-Lambert Law.

5

Verification of Beer-Lambert law and determination of the value of ε Table 2. Absorbance versus concentration [Fe3+] in the calibrating Absorbance at # Volume of Fe3+ standard solution / mol L-1 510 nm solution used / mL 2 1.00 mL 1.78 x 10-5 .312 3 3.00 mL 5.34 x 10-5 .784 -5 4 5.00 mL 8.90 x 10 1.278 5 7.00 mL 1.25 x 10-4 1.988 In the space below, show one sample calculation of [Fe2+] in the calibrating solution for solution # 4 in Table 2 above. To calculate [Fe2+] in the calibrating solution, you need to use the concentration of Fe in, and the volume of, the standard solution used, and the final volume of the calibrating solution prepared. Note: In parts 1 & 2 (Day 2), we start with Fe 3+, but then reduce it to Fe2+, and it does not matter how we refer to the Fe. Sample calculation of [Fe2+] in solution: Solution #2: (1.78 x 10-4 M)(.00100 L) = 1.78 x 10-7 mol Fe2+ / (.0100 L) solution = 1.78 x 10-5 M Solution #3: (1.78 x 10-4 M)(.00300 L) = 5.34 x 10-7 mol Fe2+ / (.0100 L) solution = 5.34 x 10-5 M

2.5

2

Absorbance

f(x) = 15463.35 x − 0.01 R² = 0.99 1.5

1

0.5

0 0.000000 0.000020 0.000040 0.000060 0.000080 0.000100 0.000120 0.000140

Concentration (mol/L)

Figure 2. The linear relationship between absorbance and Fe 2+ concentration, in mol/L, is depicted. The line of best fit further underscores the direct proportionality of the two variables with an R2 correlation value of 0.9901. Slope of the ‘best’ straight line of absorbance versus [Fe 2+] in the calibrating solution = 15463 L/mol Path length for the cuvettes used in absorbance measurements = 1.00 cm Molar absorptivity ε at 510 nm for Fe(o-phen)32+ (show work and units)

6 =

15463 L/mol =¿ 15463 L mol-1 cm-1 1.00cm

Day 2, Part 2. Mass Percent of Fe in the Fe(III)-oxalate complex.

Mass of complex used = .0518 g Measured absorbance at 510 nm of the solution prepared using 1 mL of the stock solution of the complex = 1.393 Measuredabsorbance = 9.01 x 10-5 mol L-1 εat 510 nmxpat h lengt h 1.393 / ((15463 L mol-1 cm-1)(1.00 cm)) = 9.01 x 10-5 mol L-1

[Fe] in the 10 mL prepared solution =

[Fe] in the 100 mL stock solution prepared =

[ Fe ] in the 10 mL prepared solution x 10.00mL 1.00 mL

=9.01 x 10-4 mol L-1 (9.01 x 10-5 mol L-1)(.01000 L) / .00100 L = 9.01 x 10-4 M [ Fe ] from above x 0.100 L x 55.847 g mo l-1 x 100 = mass of complex used ((9.01 x 10-4 mol L-1)(.100L)(55.847 g mol-1)) / .0518 g = .0971 x 100 = 9.71%

Mass % of Fe in the complex =

Day 2, Part 3. Mass percent of oxalate in the Fe(III)-oxalate complex [MnO4−] = .025 mol L-1

Mass of complex used = .2114 g Initial buret reading: 8.05 mL

Final buret reading: 28.21 mL

Volume of KMnO4 solution used in the titration = 20.16 mL Moles of oxalate in .2114 g (from above) of complex = [MnO4-] x Volume in L of MnO4- used in 5 the titration x 2 = .025 mol L-1 x .02016 L x (5 mol oxalate / 2 mol MnO4-) = 0.00126 mol Then, mass % of oxalate in the complex =

Moles of oxalate xmolar mass of oxalate mass of complex used

x 100

7 = (0.00126 mol x 88.020 g / mol) / (.2114 g) = .525 x 100 = 52.5%

Using the mass percentages of Fe and oxalate determined above, calculate the mole ratio between Fe and oxalate in the complex you synthesized. Be sure to include all relevant work. Assume 100 g of iron(III)-oxalate complex. 52.5 g oxalate / (88.020 g / mol) = .596 mol 9.71 g iron(III) / (55.847 g / mol) = .174 mol. .596 mol / .174 mol = 3.42 ~ 3.33 x 3 = 10 .174 mol / .174 mol = 1 x 3 = 3 = Fe3(C2O4)10 Charge balance and hydrate composition to determine complete complex: Charge balance: -11, so need 11 K+ to balance out (11/3)(.174 mol) = .638 mol K+ = K11Fe3(C2O4)10 Since it is a hydrate: Assumption of 100 g of iron(III)-oxalate complex. 100 g – (52.5 g + 9.71 g) = 37.79 g – (.638 mol)(18.0135 g / mol K+) = 26.3 g H2O 26.3 g H2O / (18.02 g / mol) = 1.46 mol H2O, or approximately 1.50 mol H2O = K11Fe3(C2O4)10 · 1.50H2O = K22Fe6(C2O4)20 · 3H2O Calculated Empirical Formula: K22Fe6(C2O4)20 · 3H2O Using an assumption of 100 g of iron(III)-oxalate complex to make mass percent conversions easier, the mole ratio calculated between oxalate and Fe(III) was determined to be 3.42:1 respectively, which can approximate to 3.33:1. Since a whole number ratio is desired for the empirical formula, the entire ratio can be multiplied by 3 to achieve a whole number 10:3 ratio. Utilizing this, a basic partial empirical formula can be formed: Fe3(C2O4)10. However, this formula is not complete. First, the charge of the complex must be balanced. Since Fe has a 3+ charge and C2O4 has a -2 charge, the current formula has an overall charge of -11. In order to preserve the oxalate:Fe(III) mole ratio but also balance out the charge, another ion must be present in the formula. According to the flame test done, the iron(III)-oxalate complex displayed positive results for K+, burning in the signature lavender color of potassium. Thus, K+ can be used to balance the -11 charge, creating the partial empirical formula K11Fe3(C2O4)10. Next, since the complex is a hydrate, the moles of water must be calculated. With

8 an assumption of 100 g of potassium-iron(III)-oxalate complex, the mass composition of the three known ions are calculated and then subtracted from the total mass to obtain the mass of H2O present in 100 g of complex. Dividing by the molar mass resulted in the moles of H2O present, which was approximately 1.50 mol. Therefore, the empirical formula is now K11Fe3(C2O4)10 · 1.50H2O; multiplying by 2 in order to obtain whole number coefficients results in a final determined empirical formula of K22Fe6(C2O4)20 · 3H2O.

One source of laboratory error is that since the blank was used to calibrate the spectrophotometer in two different occasions, stray fingerprints or smudges on the blank cuvette may have made the second calibration with a higher absorbance than the first, making the absorbance of the complex solution higher than it was supposed to be. Although this did not affect the verification of the Beer-Lambert Law, it may have affected the calculated mass percent. This would explain the deviation from the expected Fe(II) mass percent of 11%. As the mass percent was also used to calculate the mole ratio, the empirical formula would also have been affected. Accounting for the greater than actual absorbance, a mole ratio of approximately 3:1 oxalate to iron(III) could have been obtained, rather than the 3.33:1 ratio actually obtained. Factoring in charge balance and hydrate composition, the empirical formula would instead be K3Fe(C2O4)3 · 3H2O.

Calculation of % Yield Moles of Fe in FeCl3 used (concentration x volume) = .0090 mol = (.0060 L)(1.5 mol L-1) = .0090 mol Moles of C2O42- in the 4.8 g of K2C2O4 you used = .02893 mol (4.8098 g K2C2O4)(1 mol / 166.22 g)(1 mol C2O42- / 1 mol K2C2O4) = .02893 mol

Figure out the limiting reagent based on the mole ratio found between Fe and oxalate. Show all relevant work. Mole ratio (oxalate:iron(III)) = 20:6 Expected moles of products produced using all of C2O42- moles: .02893 mol / 20 = .001447 mol Expected moles of products produced using all of Fe3+ moles: .0090 mol / 6 = .0015 mol Thus, C2O42- is the limiting reactant according to our calculated empirical formula. Based on the limiting reagent and the empirical formula, calculate the number of moles of compound expected; show work below Moles of K22Fe6(C2O4)20 · 3H2O expected = (1 / 20)(.02893 mol) = .001447 mol

Mass of beaker + complex = 31.3145 g Mass of beaker = 28.6168 g So, mass of complex obtained = 2.6977 g *Note: some of the complex was used for the qualitative tests

9 Molar mass of complex (based on the empirical formula you deduced) = 3009.66 g mol-1 Number of moles of compound actually obtained = 2.6977 g / 3009.66 g mol-1 = .00089634 mol

molesofcomplexobtained molesofcomplex exp ectedbasedonstoic hiometry mol = .6195 x 100% = 61.95% % yield =

x 100 = .00089634 mol / .001447

Questions. Be sure to first list each question, followed by your answer to that question. Be brief and to the point. Brevity and relevance are more important than volume. Use separate paragraphs if and when appropriate. Brevity and relevance are far more important than volume. 1. Look up ferric oxalate complex online, and draw the structure of this complex. You may copy and paste. If you do, be sure to acknowledge. What is the geometry of the complex, tetrahedral, trigonal bipyramidal, or octahedral?

Source: https://en.wikipedia.org/wiki/Potassium_ferrioxalate The geometry of the complex is octahedral. 2. Were qualitative tests 1a & b, and 3a & b positive on your complex? Why or why not? The qualitative tests 1a and 1b were used to determine the presence of the Fe 3+ ion in the solution through a red color change or rusty orange precipitate produced by reactions with sulfosalicylic acid and ammonia. The qualitative tests 3a and 3b were used to determine the presence of the C2O42- ion through the formation of a foggy white precipitate and a colorless change, respectively. As my complex reaction experienced these results, the presence of both Fe 3+ and C2O42= was confirmed. 3. In the permanganate titration of your complex, you were instructed to first add the permanganate solution to the solution of your complex in the Eflask, and then heat the solution in the Eflask before continuing the titration. What was the purpose of heating? What would have happened if you missed to heat? The reaction occurs extremely slowly without heat. By adding heat, the average kinetic energy of the reactants is increased, and thus the reaction rate is increased. If heat was not added, the titration would have taken incredibly long to complete.

10

11

12...