Synthesis of chloropentaaminecobalt(lll) chloride Essay Example

Synthesis of chloropentaaminecobalt(lll) chloride Paper Insert the tube into the MSB and take a mental average of the fluctuating reading. 5. Repeat the previous step three times. 6. If the MSB displays a negative number, the coordination complex is diamagnetic and no further steps are necessary. 7. Determine how many unpaired electrons reside on the cobalt. E. Interpreting Infrared Spectrum 1. Assign the bands of the infrared spectrum of the complex. Results A. Synthesizing Carbonatotetraamminecobalt (III) nitrate: [Co(NH3)4CO3]NO3 The first attempt to dissolve the ammonium carbonate in 30. 0 mL of water failed because some chunks still remained. It was important to grind the ammonium carbonate down to a fine powder because the dissolution process took place very slowly and bigger chunks would not have dissolved. After all of the 10. 012 g of ammonium carbonate was dissolved, the liquid remained clear. The addition of 7. 56 g of cobalt nitrate to 15. 0 mL of water was very fluent and took on a color between dark purple and maroon. When the ammonium carbonate and cobalt nitrate solutions were added together, the new mixture was a blood red. After the mixture was stirred, however, it darkened and became nearly black. The 4. We will write a custom essay sample on Synthesis of chloropentaaminecobalt(lll) chloride specifically for you for only $16.38 $13.9/page Order now We will write a custom essay sample on Synthesis of chloropentaaminecobalt(lll) chloride specifically for you FOR ONLY $16.38 $13.9/page Hire Writer We will write a custom essay sample on Synthesis of chloropentaaminecobalt(lll) chloride specifically for you FOR ONLY $16.38 $13.9/page Hire Writer 0 mL of 30% H2O2 solution which was slowly added caused the mixture to bubble, sizzle and gain heat, while staying black. While the solution, which was poured into a large crucible, was concentrated over a Bunsen burner, evaporation occurred very slowly. During evaporation, the solution bubbled a little, gas was evolved, and the dissolution of 2. 49 g of ammonium carbonate in the hot solution was visible due to tiny bubbles. After the evaporation was complete and the volume had dropped to 48 mL, the solution had the same appearance. However, after suction filtering the solution, the color became a lighter purple. After the cold water bath, it looked somewhat powdery on the bottom. Filtering a second time resulted in only a pink/purple powder which appeared somewhat crystalline. When the solid was mixed with a scupula during the 10 minute drying process, it looked a lot more like crystals. The final mass of the crystals was 3. 039 g. The balanced equation for the synthesis of carbonatotetraamminecobalt (III) nitrate is shown below as well as the calculations for determining percent yield. 2Co(NO3)2 + 6NH3 + 2(NH4)2CO3 + H2O2 2[Co(NH3)4CO3]NO3 + 2NH4NO3 + 2H2O Theoretical Yield: Density of NH3 (liquid) = . 628 g/ml Density of 30% H2O2 = 1. 11 g/ml Moles of NH3 = (30 ml)(. 628 g/ml)/(17. 031 g/mol) = 1. 20 mol Moles of H2O2 = (4. 0 ml)(1. 11 g/ml)/(34. 015 g/mol) = . 131 mol Moles of [Co(H2O)6](NO3)2 = . 025 mol Moles of (NH4)2CO3 = . 105 mol Limiting Reactant = [Co(H2O)6](NO3)2 (. 025 mol of [Co(H2O)6](NO3)2)(2 mol of [Co(NH3)4CO3]NO3 / 2 mol [Co(H2O)6](NO3)2) = .025 mol of [Co(NH3)4CO3]NO3 (. 025 mol of [Co(NH3)4CO3]NO3)(249. 066g) = 6. 22 g Percent Yield (Actual Yield/Theoretical Yield)(100) = (3. 039 g/6. 22 g)(100) = 48. 86 % Yield B. Measuring Absorbance Spectroscopy. Two distinct peaks of intensity were noticeable for the absorbance spectrum of the cobalt solution between 350 to 650 nm. Below are the calculations for determining the mass of crystals required to create a solution of the right concentration to produce a spectrum of absorbance at approximately 0. 6. A = ? lC 0. 6 = (100 M-1cm-1)(1 cm)(C mol/L) 0. 006 M = C mol/L mol = . 0006 .0006 mol = (mass)/(249. 037 g/mol) mass = 0. 149 g [Co(NH3)4CO3]NO3 The mass of [Co(NH3)4CO3]NO3 added to 100 mL of water was 0. 149 g. Initially, the crystals were dissolved in a beaker to ease stirring. After dissolving, the solution turned purple. The solution was poured into an Erlenmeyer flask and filled with more water to the line on the flask. Below are the calculations for determining the extinction coefficient at each lambda max. The graph for the absorption spectrum is attached. .579 = (? max)(1 cm)(. 006 M) ?max = 96. 5 .519 = (? max)(1 cm)(. 006 M) ?max = 86. 5 C. Measuring conductivity Below are the calculations for preparing a . 001 M solution of the cobalt complex in 100 mL of water. (. 001 M/1000 mL) = (. 0001 mol/100 mL) .0001 mol = (mass/249. 037 g/mol) mass = .025 g. Below are the calculations for finding k, the correction factor. k = literary value/actual value Literary value = 1384 i seimens Actual value of KCl = 1420 i seimens k = (1384/1420) = . 975 Below are the calculations for finding the conductance of the cobalt solution. Probe reading of cobalt solution = 90 i seimens [((1000)(90 i seimens))(. 975)/. 002M](1x10e-6) = 43. 88 i seimens D. Measuring Magnetic Susceptibility Average empty tube reading = -004 Mass of empty tube = . 884 g Mass of tube with crystals = 1. 022 g Height of chemical in capillary rod = 3. 25 cm. Average filled tube reading = -004 (diamagnetic) Temperature of the room = 18. 6 ? C No calculations required. E. Interpreting Infrared Spectrum [Co(NH3)4CO3]NO3 IR spectrum: spikes at frequencies(cm-1) of 280, 500, 830, 1290, 1380, 1600, and 3300. [Co(NH3)5Cl]Cl2 IR spectrum: spikes at frequencies(cm-1) of 830, 1290, 1550, 3200. NaNO3 IR spectrum: spikes at frequencies(cm-1) of 830 and 1380. Discussion A problem occurred during the conductivity measurements of the cobalt complex. The expected values for conductance were between 118 and 131, but the actual calculated value was 43. 88 microseimens. Clearly, the problem arose during the probe reading in the cobalt solution. It is possible that the probes reading was skewed because it made contact with the glass wall of the beaker. This problem would have definitely lowered the reading relative to what it should have been. The absorbance spectrum of the cobalt complex on the wavelength interval 350-650 nm displayed two distinct peaks, meaning that both ions of the coordination compound were separated. This suggests strongly that the coordination compound was indeed synthesized correctly. The absorption intensities at both lambda maxes were relatively close, around 0. 6. These absorption values were used in the Beers Law equation to solve for the molar absorbance, also called the extinction coefficient. Both molar absorbance values are fairly close to 100, but deviations of 4 and 14 imply that some impurities still remain in the compound, possibly left over from the evaporation process. The coordination compound characterization via magnetic susceptibility was rather simple data requiring no calculations. The reason for this is that the magnetic susceptibility reading of the crystal-filled glass tube was equal to the reading of the empty tube, and that both were negative values. Negative values signify that the coordination compound is diamagnetic, which means that no unpaired electrons are present. The reason for the pairing of electrons can be explained with crystal field splitting. When electrons are introduced to d-orbitals, a change in energy occurs. Electrons will move to orbitals with the lowest possible energy. The ideal behavior for electrons after occupying the dxy, dx2, and dyz orbitals is to pair with the electrons already in these same orbitals. This is true only if ? E, or the energy difference from these orbitals to the dx2-y2 and dz2 orbitals, is greater than the energy cost for pairing with electrons in the lower orbitals. In the case of the cobalt complex, the spin pairing energy is much less than ? E, causing the electrons to pair in the lower energy orbitals. This pairing makes the complex a low-spin complex, implying that it is in fact diamagnetic. The goal of the infrared spectroscopy analysis was to prove that the cobalt coordination compound which was created during this experiment was actually created instead of a different compound with similar properties and bonds. [Co(NH3)5Cl]Cl2 and NaNO3 were available to compare with the infrared spectrum of [Co(NH3)5CO3]NO3. Clearly, [Co(NH3)5CO3]NO3 cannot be the same compound as NaNO3 because they only share one infrared band. Both coordination compounds have a band at 3300 cm-1 which corresponds to a N-H stretch, one at 1600 cm-1 corresponding to N-H bending, one at 1300 cm-1 corresponding to N-H symmetric bending, and one at 830 cm-1 corresponding to N-H bending. The most important difference between these coordination compounds lies in the common spike of [Co(NH3)5CO3]NO3 and NaNO3 at 1380 cm-1. Because these two compounds share this spike while the Cl coordination compound does not, the 1380 cm-1 peak must represent an N-O bond, which is the most significant spike on the NaNO3 spectrum as would be expected. During the synthesis of carbonatotetraamminecobalt (III) nitrate, several factors could have contributed to the relatively low quantity of crystals formed in terms of the percent yield. After calculations, it was concluded that 6. 22 grams of crystals should have been produced. However, only 3. 039 grams was actually produced. One of these factors that may have contributed to the low percent error of 48. 86% is the short evaporation stage. Since time was a factor during the experiment, the solution may not have been given enough time to evaporate any impurities. While the final volume of the solution was verified to be 48 mL, under the designated 50 mL, more impurities could have evaporated. If the solution had sat above a Bunsen burner for a greater length of time, it would have undoubtedly become more concentrated. Another factor which possibly contributed to a low percent yield is measurement errors. Although a good level of confidence can be felt about whether or not the right amounts were added, there is always room for error. A simple miscalculation or slight deviation in a measurement could have thrown off the remainder of the synthesis. Furthermore, a very obvious source of error can be found regarding the filtering system. The crystal product is quite soluble in water for the fact that its ions can be easily separated. For this reason, ice cold water was used whenever contact had to be made with the crystals; the low temperature of the water prevents the dissolution of the crystals to some extent. In the filter, some of the crystals could have dissolved and fallen through, which would have resulted in a poor percent yield. In general, this experiment ran very smoothly and achieved its purpose of providing detailed information regarding the properties and synthesis processes of a transition metal coordination compound. To improve the results of percent yield and perhaps to improve the results in many areas of characterization, the evaporation stage should be lengthened to facilitate the removal of remaining impurities.