Week 7: Quenching Experiments

During this last week of research funded by the Charles Center, I have been attempting to hone my accuracy regarding the aforementioned quenching experiments.  During the first part of the week, I performed triethylamine additions to a fluorescein solution in order to replicate the baseline experiment done in a previous paper.  I succeeded; I created a Stern-Volmer plot from the fluorimeter’s data, which plots relative intensity versus catalyst (or, in this case, triethylamine) concentration times a constant.  This plot had the same quenching constant, the slope of the linear fit of the data points, as found by previous lab mates.  Later in the week, I have been trying to master catalyst additions, which have proven to be more difficult. The quenching constant can give us insight into the quenching pathway of the system, which is why I will keep working to verify the findings by other lab mates.  This will give us insight into the physical properties of the system.  While I am finished with my Charles Center research, I am continuing to work on these quenching experiments and other aspects to my initial project over the next three weeks.  Other than quenching, I set up more crystallizations for Monday, hoping to grow some illusive crystals that have been difficult to grow all summer.

Week 6: Crystallization of Complex

Last week, I mentioned my attempts to synthesize a pure product from my coordination reaction.  This week, I toyed with many different solvents in order to find the perfect combination for crystallization via slow diffusion.  This technique works by dissolving the product of a reaction in a certain solvent, and then slowly adding a second solvent on top of it.  The two solvents must be miscible, but the second solvent must not dissolve the complex.  When the two progressively mix together, the product then crashes out in a crystal formation.  This week was used to find solvents that did not dissolve my complex.  The two that do not are pentanes and hexanes. I then needed to test the different solvents with pentanes and hexanes.  After much trial and error, I tried an acetone/pentanes and acetone/hexanes mix.  I am hoping to discover crystals in my vial when I go in for my last week supported by the Charles Center.

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Week 5: Generation of Hydrogen

In my previous post, I touched on the efficiency of the iron catalyst I have been studying, but I also mentioned that I need to ensure that my catalyst was actually generating hydrogen. This week, I introduced my complex into an environment containing a sacrificial donor of electrons (triethylamine), a proton source (water), a chromophore (fluorescein), and a solvent (ethanol).  Six different test tubes were prepared.  They were then illuminated in a controlled environment.  This experiment was purely used to see if any hydrogen was generated at all.  So, instead of taking many measurements of hydrogen gas in the test tube every few hours, which can be useful to measure the robustness of a catalyst as well as measure reaction rates, I took one after 24 hours of illumination.  The experiment was successful; I observed hydrogen generation from each test tube by collecting samples and analyzing them via gas chromatography.  This proves that my catalyst is, in fact, producing hydrogen and at an efficient rate.  However, the catalyst I used was impure, and the turnover number (moles of hydrogen per mole of catalyst) was relatively low.  The next step is finding ways to ensure crystallization of my catalyst after synthesis.  While I have made strides in increasing my ligand yield, it is still less than half of what I should be synthesizing.

Beginning of Week 4: A Change of Catalyst

Since starting my research, I have found out quite a bit about my original complex.  The cobalt version of my complex doesn’t seem to form when a complexation reaction is performed with the ligand and cobalt precursor.  However, in order to verify that the ligand that I synthesized was, in fact, my desired product, I was advised to perform a metallation with iron and run electrochemical tests.  This would allow me to compare it with a previous person’s research of the same complex.  If the results were similar, then I could confidently say that I had the correct ligand.  So, I performed the metallation, and ran the electrochemical tests, received some interesting results.  The electrochemical technique I used is called cyclic voltammetry.  It is effective in observing if a complex can reduce hydrogen efficiently.  A potential is introduced to a solution containing the catalyst.  In this experiment, the potential is brought from a positive to a negative potential, and then back again.  I am then able to  observe the change in current as the potential is varied.  While performing an acid addition experiment, I anticipated seeing a redox couple (observed by two peaks nearly at the equialent potential, reduction on the potential being brought out negative and the oxidation on being brought back positive) and a catalytic peak.  Where I expected to see one catalytic peak, I saw many peaks.  It is possible that this was observed because the test was carried to a more negative potential than needed.  My professor proposed running my experiment to a less negative potential so that we could see the one reduction needed to ensure accurate results. So, I ran it in a relatively small potential range, and observed on my plot that there was in fact one catalytic peak in this range, and a redox couple.  Since the catalytic peak is at such a cathodic potential, it seems that the complex is extremely efficient.  However, in order to make sure that the peak does, in fact, correspond to a reduction event, I will run photochemical experiments of my complex in order to fully observe it (hopefully) generating hydrogen in a system that should allow it to do so.

Abstract: Hydrogen Generation via Transition Metal complexes

Artificial Photosynthesis is a promising mechanism through which clean, renewable energy can be generated.  Photosynthesis allows a plant to synthesize glucose to use as fuel.  Instead of generating glucose as a final product, artificial photosynthesis generates hydrogen. The mechanism involves a chromophore, a semiconductor, and an electrocatalyst (Fig. 1).  When sunlight hits a chromophore, an electron becomes excited.  This electron is then transferred to a semiconductor, which then transfers the electron to an electrocatalyst.  The electrocatalyst reduces the protons naturally found in water to generate hydrogen gas.   Over the summer, I will be researching a cobalt complex that will act as an electrocatalyst.  Cobalt is earth abundant and has low cost, which will ultimately lower the cost of this mechanism.  The cobalt complex itself is inspired by a similar iron complex previously synthesized in my research lab.  Specifically, I will be running organic syntheses in order to create the ligand so that I may later complex it with cobalt.  This will then yield the electrocatalyst.  Once synthesized, I will then analyze it  using electrochemical experiments to observe its stability and efficiency.  Hydrogen can be introduced into a fuel cell to generate energy, releasing water vapor as waste.  Making this process cost effective and efficient may eventually turn it into a promising alternative energy source.

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