After about two weeks of stressing over my progress in my integrated experimental chemistry lab course at university, I've come to learn that I'm actually ahead of the curve. As such, I can finally appreciate what I'm doing and realize just how much fun I'm having! To anyone interested in chemistry, this is an example of what you can expect in a 300-level lab course!

Before we get to the pictures I attached, I should tell you where it all began. Everyone in the lab started with a glass bottle full of a mixture of salts, various soluble and insoluble impurities, and an unknown carboxylic acid (a hydrocarbon with a carboxyl group at one end, meaning there is a carbon atom that is double-bonded to one oxygen atom and single bonded to another oxygen atom which is in turn bonded to a hydrogen atom. I've attached a chart with some examples.).

Our primary goal was to isolate the unknown acid in order to discern its identity. To do this, the first thing we had to do was remove the insoluble impurities by dissolving the mixture in hot deionized (DI) water (water so pure, that if you were to drink it, it would leech the vital minerals from your very cells!). After dissolving what could be dissolved, suction filtration was used to isolate the soluble filtrate from the insoluble impurities. This left me, in particular, with a light-blue filtrate, the light blue color coming from the soluble impurities in the sample (a pure carboxylic acid + salt solution would be clear/colorless).

The next step was to isolate the carboxylic acid and salts from the soluble impurities. This step is a bit less obvious and where actual chemistry comes into play. In order to decolorize, or preferentially remove soluble colored impurities, a small pinch (0.5% of the original weight of the sample) of activated charcoal was mixed into the heated solution and then filtered. The charcoal adsorbed (bonded to, as opposed to took into) the colored soluble impurities, keeping them from passing through the filter and leaving me with a clear and colorless filtrate.

The final step towards isolating the acid involves a process for which, if it does have a name, I don't know what it's called The general idea is that heated solution is poured into a mixture of ice and hydrochloric acid, and then cooled in an ice bath. If the mixture is acidic enough, then an acid slurry should form on top of the solution. After filtration, you're left with a liquid filtrate (DI water, salts, and what soluble impurities the activated charcoal didn't adsorb) and a solid cake of "crude" acid.

Congratulations, you've isolated your acid! However, we're not out of the woods just yet. The reason it's called "crude" is because, while you might think you've removed all the impurities, you haven't. The process from before isolated the acid in a "crude" manner, and we want the sample to be as pure as possible. Therefore, in order to turn all of our "crude" acid into "pure" acid, recrystallization must be performed. This involves dissolving your acid in a heated solvent and then allowing that solvent to cool. Which solvent you use entirely depends on which one works best for your acid during your small and medium batch trials. For this lab, we tested four different solvents: water (Generally the weakest of the four), hexane, toluene (The one that worked best for me), and ethanol (The strongest of the four, I was able to dissolve 20 mg in barely 0.05 mL of the stuff. The reason I went with toluene, however, was that the acid refused to recrystallize in the ethanol even after cooling, seeding, and scratching the sides. Thankfully, this was only a small batch and only about 20 mg of a bulk amount of >7 g of "crude" acid was lost.). Something to note about solvents is that the best ones for the job often have similar structures to their solutes. As such, since toluene is aromatic, so must the acid! (An aromatic molecule, in a nutshell, has a structure that is focused around one or more hexagonal benzene rings, such as the attached picture of toluene.)

This brings us to the needle-like crystals you see forming in the Erlenmeyer flask. After dissolving my bulk "crude" acid into heated toluene and allowing it to slowly cool, the acid began to precipitate out and recrystallize while any impurities that might be left remained in solution. I've since left it to continue recrystallizing over the weekend to allow for the maximum yield, but soon to come will be filtering and drying the pure acid as well comparing the melting temperatures of crude vs recrystallized vs double-recrystallized acids.

While I'm still an undergraduate, I've been a chemistry major for five years and I would love to hear any questions you might have!