I pen this from the desk of my hotel suite in North Carolina, almost but not quite ready for bed. Tomorrow brings with it the big day. The one I’ve been waiting for. Tomorrow I board a plane to St. Johns Antigua, and begin my education as a Doctor.
I would be lying if I told you that I wasn’t nervous. Quite the contrary. My anxiety stems from several factors, the deepest being my impending separation from my family and girlfriend. Not to mention the fact that this will be my first time living and providing for myself. However, I am confident that I will do fine, it may just take time.
I figured I would write, the benefits being twofold: It serves as an outlet for my fears, and provides a timely update to those who still read this little site of mine. Nevertheless, this narration is helping in some small way to calm my rattled nerves.
Thanks to the wonderful donation from my Aunt Becky and Uncle Don, I was able to afford a new Canon digital camera. I took a slew of photos on the trip to North Carolina, and have posted them in a small gallery on my facebook. Interested parties can find the link to my profile on the front page of this very site. I plan on continuing to write and photograph my adventure, so there will eventually be new updates. First, I must just get organized tomorrow.
That’s all I have for this evening. Hopefully tomorrow will leave me with a list of good events to recall, and a stable internet connection in my new apartment in which to utilize in writing about them!
For this experiment, I decided to attempt to isolate the scent of pine needles via Soxhlet Extraction.
I prepared fresh pine needles from a tree in the back yard by trimming them with scissors and packing them into the thimble of the Soxhlet Extractor, utilizing cotton on top and bottom of the needle samples. I decided to try an organic solvent, and settled on denatured ethanol. I started with 100mL of the ethanol, and began the extraction process.
About halfway through the process, I found that 100mL was not enough, and added ~100mL extra to the Florence Flask. This allowed the proper siphon action on the side-arm of the Soxhlet.
The final distillate was not what I had expected. I received a yield of forest-green liquid which, upon wafting, smelled vaguely like pine, yet retained far too much of the original ethanol odor to be usable for aroma purposes. If I decide to repeat this experiment, I believe I will try either another solvent other than ethanol, or adjust the quantity of needles that I decide to extract from.
If anyone has successfully done this procedure, or has ideas as to how I might modify the experiment further, please let me know!
The Soxhlet Extraction of Pine Needles.
Close-up of the distillate and the pine needles.
Upper part of the assembly, highlighting the Allihn condenser with water flowing and the Soxhlet Extractor.
A molal solution is a specific number of moles of solute per kilogram of solvent.
In this experiment, I mixed 100mL of a 0.5 Molal solution of Potassium Hexacyanoferrate (III). To begin, I weighed out 16.65g of the chemical. It’s molecular weight is 329.24. Half of that (0.5M) is 164.62g, and 1/10 of that is 16.46. The reason we use 1/10 of 0.5M is because 100mL is 1/10 of 1L, and our desired solution is 0.5M. Now that we have 16.46 as our value, I assumed 99% chemical purity (and it said so on the label!) so I divided 16.46 by .99, yielding 16.63g.
The next part differs slightly. Instead of using the volumetric flask like we did for the molar solution, we simply need to weigh out 100.00g of the solvent, which is water. I weighed out 100.15g (this is close enough.) and dissolved the potassium hexacyanoferrate in it. After the chemical is dissolved, it can be transferred to a storage bottle.
When transferring the solution to a storage bottle, do NOT rinse the beaker to transfer any remaining chemical stuck on the sides. The additional water will increase the weight value of the solvent!
Potassium Hexacyanoferrate (III) dissolved in water, and an empty storage bottle.
For those interested, take the time to check out http://www.elvesandmoreneo.org/. This charity, which I found out about last night, is comprised of wonderful people who take the time to donate bicycles to poor children at Christmas in the inner cities of North-Eastern Ohio! I wanted to give them a shout-out, and hopefully bring them even more traffic.
Keep up the great work!
To demonstrate the mixing of Molar solutions, I prepared a simple 1M solution of copper(II) sulfate pentahydrate, which I commonly use for other experiments.
The Molecular Weight of copper sulfate is 249.7g/M. A 1M solution of copper sulfate would simply be 249.7g of chemical dissolved in 1L of water. Since I like to store my solutions in 100mL glass bottles, I decided to make up 100mL of a 1M copper sulfate solution. 100mL is 1/10th of a L. Therefore, 1/10th of 249.7g would be 24.97g. However, the copper sulfate I used has a purity of 99%, so to account for this, I divided 24.97 by .99, yielding 25.22g for usage. After weighing out 25.22g, I mixed it with 85mL of distilled water in a 150mL beaker and agitated the solution until the solute was completely dissolved. At that point, I transferred the solution to a 100mL volumetric flask, which I then filled to the graduation line with an eye dropper. I then bottled the resulting solution for storage.
A wash bottle of distilled water, a 100mL graduated cylinder, the finished and bottled 1M copper sulfate solution, and a 100mL graduated cylinder.
A hydrate (in inorganic chemistry) is a salt combined with water molecules in a specific ratio. We can determine that ratio by driving off the water molecules (referred to as the water of hydration) via heat, and then performing some simple mathematics to determine the mass relationship.
For this experiment, I utilized Copper(II) sulfate pentahydrate, the name of which suggests that for every molecule of copper sulfate, there are five associated water molecules in the crystalline structure. By adding 5 grams of CuSO4●5H2O to a crucible and driving off the water of hydration via an alcohol lamp, I was left with the white, anhydrous form of CuSO4. By comparing the weight of the hydrated salt and the anhydrous form, I was able to determine the mass of water lost from evaporation. Translating this to terms of moles, I ended up with a final ratio of 4.970 to 1, or when we round, 5 to 1, which would represent the pentahydrate. This value was not exact due to several predictable sources of error (weighing discrepancies, not driving off all water of hydration, etc.).
Heating the Copper sulfate in a crucible over an alcohol lamp.
By utilizing a liquid-liquid solvent extraction technique, we can extract a certain component from a liquid solution. For this experiment, I used a tincture of iodine, which is simply iodine dissolved in ethanol.
I added around 1mL of tincture of iodine to a test tube filled with distilled water. Then, I added around 1mL of an organic solvent. This is important, as the ethanol is not soluble in an organic solvent (being one itself.) However, the iodine WILL dissolve in my solvent, Petroleum Ether Ligroin. By agitating the test tube, it mixes the solvents and the iodine moves from the waste layer to the product layer. It is then as simple as drawing the top iodine layer out with an eye dropper, and placing it on a watch glass to evaporate.
To purify crude Copper sulfate, available commercially as root killer and other garden uses, one can use recrystallization to remove impurities in the chemical composition. By dissolving a quantity of Copper sulfate in boiling water, then allowing it to cool causes the Copper sulfate to precipitate out of solution and form pure crystals, which contain no room for impure material in their lattice formation. The solution can then be poured through a filter funnel to isolate the crystals. A final wash of Acetone serves to remove all the water from the crystals, as Copper sulfate is almost insoluble in acetone.
Purified Copper sulfate crystals resting in a 250mL beaker.
After the recrystallization is complete, we can utilize the leftover impurities from the filtering process to obtain a supply of crude Copper carbonate. First, Sodium carbonate is dissolved in hot water. Once finished, this solution can be added to our waste solution from the first procedure to precipitate out the copper ions as the insoluble carbonate salt, leaving Sodium sulfate in solution. Once this reaction takes place, the new solution can be poured through a filter funnel, leaving behind the water-insoluble Copper carbonate crystals on the filter paper. These can be dried and washed with Acetone, then retained for further experiments. The waste solution from this process contains only Sodium sulfate and a small amount of copper.
Crude Copper carbonate crystals drying in filter paper.
