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Relevance. Coordination Complexes

Relevance. Coordination complexes, though one may not typically recognize it, have many commercial purposes and contribute to the lives of many in essential ways (even outside of a laboratory setting). They are often a component of household cleaners and are even used in the development of some types of film, both of which are products which people use or encounter daily.

Coordination complexes are formed when metal ions are bonded to some number of ligands, which or molecules or ions, through coordinate covalent bonding. They are important for a few reasons: one, because metal complexes can be used as industrial catalysts which was be used to control reactivity; two, because nearly all metals found on the periodic table make complexes, therefore metal complexes are relevant chemistry-wise for more than half of the elements; third, because complexes are critical for the function of life, as seen in biochemistry. For example, hemoglobin is the primary carrier of oxygen in the bloodstream, is essential to life, and is also an iron complex1(

Background. Coordination compounds have been used in chemistry for a long time; while chemical compositions were known, the chemical natures of complexes were largely guesswork until the late 1880s, when Nobel-Prize Winner Alfred Werner made breakthroughs in coordination chemistry. Werner’s work focused especially on the differentiation between oxidation states and coordination numbers, and he studied in detail the different forms of cobalt complexes and how the arrangements of their ligands altered their colors1. (

Today, it is known that the coordination number (the number of ligands that surround the metal ion) of the complex dictates which type of structure can form. Typically, those with higher coordination numbers can form a higher number of unique complexes-coordination compounds. The field regarding coordination compounds has come a long way since Werner, and is now largely focused on creating more efficient processes; coordination compounds can be used as catalysts, in dyes and pigments, and in analysis and creation of various substances or other compounds. Scientists are continually looking for new uses for coordination complexes. (

This experiment had two parts. The first part was to make four different aquo complexes (metals: Co, Ni, Cu, Zn) and the second part was to make a complicated cobalt complex, pentaamminenitritocobalt (III) chloride; the objective was to determine the absorbencies of all of them. One can infer the coordination number and geometry of aquo-complex compounds in solution from the light-absorption spectra, which was one of the studies conducted in this lab. Due to the mistake which my lab partner and I made in the creation of our coordination complex of cobalt, it will be important to investigate the different parts of this product, its chemical composition, how its UV-vis spectroscopy graph differs from others, and the structure of the created cobalt solution compared to those of the aquo complexes.

Hypothesis. Based on previous work with aquo and cobalt coordination complexes, this lab’s purpose was to expand on previous experimentation with aquo complexes and cobalt complexes over a period of time. By using UV-vis spectroscopy and observing the wavelengths of the different coordination complexes, one can understand a complex’s coordination number (discussed above), the energy differences between d-orbitals in each transition metal, and what factors dictate the interaction of metals and their respective ligands.

The different energies of the d-orbitals result in the varying colors of solution, affecting the maximum absorbencies and their corresponding wavelengths. To create these solutions to determine their wavelengths, a new compound or solution is added to alter the equilibria of the original complex. To experimentally determine the different components that influence metal-ligand bonds, aquo complexes were made using diluted metal solutions (Cobalt, Nickel, Copper, and Zinc) and observed. Because each aquo complex will be different colors, each will have a different absorbance and corresponding wavelength. The absorbencies are the key to determining the differences in metal-ligand interactions in each solution because the different colors of each solution will correspond with the energy differences.

Hypothesis/Description of Current work 

 What does the current work propose based on what is known? 

 Why is this the proposition? How is it grounded? 

 Why is this important to find out? 

Methods and Materials:

Aquo complex. The aquo complexes that were analyzed were Co(H2O)62+, Ni(H2O)62+, Cu(H2O)62+, and Zn(H2O)42+.

The solutions for the aquo complexes were provided and analyzed separately from this lab. However, the process for making a 0.1 M solution of the aquo complexes in this lab would have been as follows: to make the 0.1 M aquo complexes, a certain mass of each salt is calculated dissolved in each 25 mL of DI water each. Measure the mass of each salt using weigh paper and an electronic scale. Measure the 25 mL of DI water using a 10 mL graduated cylinder for maximum accuracy. To make a 0.1 M salt solution in 25 mL of water, 0.0025 moles of each salt is necessary. So, by multiplying the molecular weight (g/mol) of each salt by 0.0025 moles, one can find the grams of salt needed to make a 0.1 M solution.

Co: molecular weight = 58.933 g/mol o 0.0025 mol = 0.147 g Co

Ni: molecular weight = 58.693 g/mol o 0.0025 mol = 0.147 g Ni

Cu: molecular weight = 63.546 g/mol o 0.0025 mol = 0.159 g Cu

Zn: molecular weight = 65.38 g/mol o 0.0025 mol = 0.163 g Zn

Once the desired solution is created, the UV spectrometer would have been used to record the data. For this lab, the solutions were on display at the materials bench (Image 1/2) and UV-vis spectra data is provided in Figures 1, 2, 3, and 4. On display, the cobalt solution was red, the nickel solution was green, the copper solution was light blue, and the zinc solution was colorless.

For the data analysis of the aquo-ion complexes, the given UV-vis spectroscopy data was used to find the maximum wavelength. Additionally, the Beer-Lambert Law was used to find the molar absorptivity constant epsilon (ε) in each of the aquo-ion complexes.

Cobalt complex. The cobalt complex which was created from this method is [Co(NH3)5ONO]Cl2. If not for a fault in procedure, the complex which would have been created after sitting for a week would be the more thermodynamically stable [Co(NH3)5NO2]Cl2, which would also cause a color change. For the creation of the original compound, the method is following:

For this experiment, a clean and dry beaker (100 mL) was used, with a magnetic stir bar and stir plate. Underneath a vent hood, solid ammonium chloride (2.5 g) was be dissolved in 15 M aqueous ammonia (15 mL) inside the beaker, with the magnetic stir bar stirring the solution. The experiment remained under the vent hood for the entirety of the experiment. The solid ammonium chloride was measured using weigh paper and an electronic scale, and the 15 M aqueous ammonia was measured using a graduated cylinder (10 mL) for maximum accuracy.

Powdered hexaaquacobalt (II) chloride (5 g) was added to the beaker in small increments, with no clumps. The hexaaquacobalt (II) chloride was measured using weigh paper and an electronic scale. While the powdered hexaaquacobalt (II) chloride was a magenta color, it transformed the aqueous ammonia-ammonium chloride mixture into an orange-brown slurry. 30% Hydrogen peroxide (4 mL), H2O2, was measured and added using a clean and dry dropper pipette (1 mL) slowly. This caused the solution in the beaker to bubble violently, producing visible gas and heat. The addition of H2O2 also resulted in a color change, from brown to red-brown.

Upon the completion of the solution’s bubbling, the precipitates that had formed were filtered out of the solution. This was done using a Buchner funnel and filter paper of the corresponding size, collecting the filtered solution in a clean and dry Erlenmeyer Flask (250 mL). The filter paper was discarded in the contaminated waste bin. Once the beaker (100 mL) was cleaned, the filtrate was transferred from the flask back into the beaker. The filtrate was then cooled to 10ËšC using a cold-water bath, which was contained in another beaker (400 mL). It was ensured that no tap water accidentally splashed into the filtrate.

While the filtrate was cooling, 3 M HCl (5-10 mL) was added incrementally using a dropper until the pH of the solution became slightly more acidic than or exactly 7. The solution had to be stirred after each addition of HCl using the magnetic stir bar and stir plate. The pH had to be measured indirectly after adding drops of HCl at a time, incrementally diluting a small amount of the filtrate/HCl solution (<0.1 mL) in water (~5 mL) and testing the pH of that water until the pH paper read neutral. The addition of the HCl changed the consistency of the solution slightly.

Continuing stirring the solution using the stir bar and magnetic stir plate, sodium nitrite (2.5 g) was added. This caused a color change in the solution and gas was released. The solution was cooled for about an hour in the ice bath, which precipitated the solid. The ice bath was created in a beaker (400 mL) using a 50/50 mixture of ice pellets and tap water. Finally, the precipitate was filtered again using fresh filter paper, a cleaned Buchner funnel, and a clean and dry filtering flask (250 mL). Keeping the filter paper this time, the precipitate was washed with cold ethanol (20 mL). The ethanol was measured using a clean and dry graduated cylinder (50 mL). The vacuum technique was used for five minutes to dry the precipitate on the filter paper.

The precipitate that was created was an orange/brown color when dried (Image 3). So that its properties could be evaluated with the UV-vis spectrometer, two different solutions were to be made using the precipitate. This is where my lab partner and I failed to complete the procedure, as we created these new solutions using the solution that was left in the filtering flask. Therefore, our UV-vis spectroscopy graph is different than it would otherwise be since our solution was more dilute from the ethanol wash and not as pure as the solid precipitate which came from it.

If the experiment were to be completed correctly, the cobalt precipitate (0.1 g) would have been dissolved in deionized water (50 mL) and put in a small, clean, and dry beaker (50 mL), covered with a Kimwipe and secured with a rubber band. The second solution would be made of a mixture of the cobalt precipitate (0.1 g) and a premade solution of 0.03 M sodium nitrite and 0.15 M ammonium chloride (50 mL). This solution would similarly be placed in a small, clean, and dry beaker, covered with. Kimwipe, and secured with a rubber band. The DI water (50 mL) and pre-made solution (50 mL) would be measured in a graduated cylinder (50 mL). The solid precipitate would be measured using weigh paper and an electronic scale. Both of these cobalt solutions would be measured using UV-vis spectroscopy to find the corresponding wavelength and maximum absorbance of each solution (Figure 5/6). The maximum wavelength for the DI water solution within the visible spectrum was 444.91 nm. The maximum wavelength for the premade solution within the visible spectrum was 453.449 nm.

After a week, the solutions should have changed chemical properties. Since this experiment was not performed correctly, the final cobalt/premade solution complex turned a light color and has some precipitate in the bottom. It can be hypothesized that this color change occurred because the solution was diluted initially because we did not use the solid precipitate, and because of the added ethanol wash. The solutions are once again measured using the UV-vis spectroscopy system to measure the corresponding wavelength and maximum absorbance of the final solution (Figure 7 Parts 1 & 2). The liquid and precipitate of this product solution was separated and measured separately.

For data analysis, the graphs of the UV spectrum were used to visualize where the maximum wavelength occurred for the complex before and after it sat for a week. Using the UV-vis spectroscopy, a graph was obtained and from that graph conclusions can be drawn which are discussed in the results.

Figures 1, 2, 3, and 4 depict the UV-vis spectrum for each of the aquo solutions that were given. UV-vis spectroscopy data, in this case, depicts the absorption spectroscopy in the UV region, or the visible ranges; this affects what color the chemical solution is perceived as. More importantly, in this range of the electromagnetic spectrum, electrons become excited and transition from their ground state to their excited states; in other words, they undergo an electron transition. Therefore, the graphs of the respective aquo complexes obtain their peaks when the electrons become the most excited and undergo transitions; when they do this, they emit visible light.

Figure 1. Cobalt aquo complex. Highest peak of absorbance (.448) occurs at a wavelength of 507.893 nm.

Figure 2. Nickel aquo complex. Highest peak of absorbance (.253) occurs at a wavelength of 706.662 nm.

Figure 3. Copper aquo complex. Highest peak of absorbance (1.131) occurs at a wavelength of 811.992 nm.

Figure 4. Zinc aquo complex. Highest peak of absorbance (.114) occurs at a wavelength of 388.078 nm, outside of the visible spectrum.

The different d-electron configurations of each of the aquo complexes influence how much photon energy the solution absorbs, and what color light the solution will emit. Therefore, the maximum wavelength of the graphs correlates with the number of unpaired electrons the compound has at that concentration of light. The colors of the different complexes are shown below:

The molar absorptivity appeared to be correlated with the number of unpaired electrons in the aquo complex energy diagrams (Figures 8-11). In cobalt, the anti-bonding orbital had 1 unpaired electron. In nickel, the most stable of the aquo complexes when it came to unpaired electrons, there was 1 unpaired electron in both of the antibonding orbitals. In copper, there was one pair of electrons in the anti-bonding orbital and one unpaired electron in the same orbital. From this, it is assumed that the stability of the electron orbitals (how likely they are to want to absorb photoelectrons) was directly correlated with molar absorptivity. Nickel, which is the least stable, had the lowest molar absorptivity constant (2.53 L g-1cm-1). Cobalt, which is unstable but requires more electrons to fill the antibonding orbital, had the second-lowest molar absorptivity constant (4.48 L g-1cm-1). Finally copper, which only needs one electron to fill its antibonding shell and is therefore unstable, had the highest molar absorptivity constant (11.31 L g-1cm-1).

The maximum wavelength for each metal complex also corresponded to the energy diagrams. The cobalt, nickel, and copper energy diagrams corresponded with an octahedral structure, to match the chemical structure of their respective complexes. Specifically, it correlated with the energy gap between the d-orbitals of the diagram. They had a direct relationship; the bigger the wavelength, the smaller the energy gap between the d-orbitals. The energy gap corresponded to the energy required to move between the energy levels of d-orbitals. Cobalt, which had the smallest wavelength (507.893 nm), therefore absorbed more energetic photoelectrons to fill their orbitals because it required the most energy to become ionized. And, as a result, it had the largest energy gap between the d-orbitals (Figure 8).

Along those same lines, nickel had the second smallest energy gap (Figure 9) on account of it having the second smallest wavelength (706.662 nm). The copper solution, which had the highest wavelength (811.992), required the least amount of energy, and therefore had the smallest energy gap (Figure 10).

Since the zinc complexes have a different structure than the other three, its energy diagrams correspond with its structure: tetrahedral. Additionally, zinc’s maximum wavelength takes place outside the visible range; this occurs because it contains all 10 electrons in the d-orbital energy level (Figure 11). The electrons cannot become excited in this case and cannot move between energy levels, therefore, its maximum wavelength is not relevant.

Cobalt Complex synthesis. Due to a mistake in the making of the synthesized cobalt complex that has been explained above, a final mass for the cobalt complex was impossible to obtain. To demonstrate the calculations that would be done for the percent yield of the synthesis of the cobalt complex, the yield for a cobalt chloride aquo complex is:

___ g Co〖Cl〗_2*6H_2 O* (1 mol Co〖Cl〗_2∙6H_2 O)/(237.83 g)*(1 mol complex)/(1 mol Co〖Cl〗_2*6H_2 O)*(260.83 g complex)/(1 mol complex)=____ g complex=theoretical yield

Using this theoretical yield, the masses from both week 1 and week 2 would be used to calculate the percent yields for each week and each isotope of the cobalt complex:

(_.__ g [Co(NH_3 )_5 ONO]Cl_2 experimental)/(____ g [Co(NH_3 )_5 ONO]Cl_2 theoretical )*100%=_.__%yield

(_.__g [Co(NH_3 )_5 NO_2 ]Cl_2 experimental)/(____ g [Co(NH_3 )_5 NO_2 ]Cl_2 theoretical)*100%=_.__% yield

In this experiment, the final mass of the compound should be used at the end of week 1, because the mass could change easily if left too long. Between weeks one and two, liquid can evaporate from the beaker, which might cause a difference in the mass. Alternatively, some of the original precipitate could get stuck to the weigh paper or the side of the beaker, which would reduce the mass that was weighed during week 2, therefore reducing the percent yield.

The percent yield from the earlier question would most likely not be exactly 100% due to human error. The measurements throughout the experiment could not have as precise as they were intended to be. Furthermore, in filtering the solution, it is certainly possible that some of the would-be product got stuck or was not cleaned off properly, especially when it was vacuum-filtered. It is also possible that while the cobalt solution was in the ice bath and precipitate was forming, that the solution was not submerged long enough so not all of the possible precipitate formed.

The color of the actual product can be seen in Image 3. Due to the mistake in the lab technique, the color of the "product" solution did not change very much over the weeks. It started immediately after the synthesis as a light orange color and ended a similar color. The largest difference was that in the immediate post-synthesis solution had no precipitate, and the solution after a few weeks did have small, wispy precipitates. When the precipitates were dissolved in DI water (50 mL), however, the solution was a bright yellow. The chemical reasons behind the lack of change in the color of the product would be because the solution that was used was a dilute liquid, so was not the correct product to most effectively form ligands with the premade solution. Figure 7 Parts 1 & 2 depicts the two separate UV-vis spectroscopies for these product solutions (Figure 7 Part 2 for the precipitate in DI water, Figure 7 Part 1 for the leftover solution in the beaker).