Continuing Preparation for my Research at Ole Miss

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Working with Acutely Toxic Chemicals is Always a Fun Time...


The first part of the drama that has been unfolding with my research with Dr. Wadkins has been due to a nasty little chemical called mercury(II) perchlorate trihydrate. This inorganic mercury salt is going to be our source of mercury ions in solution for when we titrate it with the DNA solution to observe the effects of heavy metal ion intercalation of DNA base pairs and subsequent secondary structure formation. Of course everything in life is hard, right? And nothing ever goes to plan. At least I think that's a safe motto to live by especially in scientific research. So what was going wrong?

Mercury(II) Perchlorate is Hygroscopic

What does it mean to be hygroscopic? It essentially means that it absorbs water like crazy! When we opened the bottle of the mercury salt, we discovered there was a thick aqueous layer of water on top of the actual salt and that the salt was nearly cemented to the bottom of the bottle. I tried my best to get it free, but with the compound being so toxic, I didn't want to scrape up any fumes. 

I went back to Dr. Wadkins who suggested just using whatever mercury perchlorate was already dissolved in that aqueous layer; he suspected that it was quite concentrated. The plan of action was to measure the volume of the aqueous layer and weigh it on the balance in the fume hood. So, I went back into the lab and did this, very carefully I might add, in fact on the verge of paranoia. However, paranoia is exactly the mindset you want when working with dangerous chemicals in my opinion. 

So I measured out the volume and it turned out to be nearly exactly 5mL. Phew! That makes my math a lot easier. What I did after this was weigh an empty graduated cylinder, and then place the 5 mLs of mercury perchlorate into it and get the weight in grams. The difference in the weights of course is the weight of the fluid that was added to the graduated cylinder.

Then I poured 5mLs of de-ionized water (A.K.A. distilled water) into the graduated cylinder to get the weight of 5mLs of just pure water. The difference in weight from the 5mLs of mercury perchlorate and 5mLs of straight pure water indicated the presence of lots of mercury perchlorate in the solution. Subtracting the weight of the pure water from the weight of the water + mercury perchlorate solution gives an approximation (not going to say it's exact, just an estimate) of the grams of perchlorate in the solution.

From here it was easy to calculate the number of moles in the 5mLs. Then, after that, calculating the concentration was just as easy. Concentration (in units of Molarity) equals moles divided by volume.
Once I did this, my estimated concentration came out to be 3.15M.

Now What?

That's insanely concentrated compared to what we need! So what I'll do next is dilute it down to something on the order of 1mM or 10mM. I'll place it in the MOPS pH buffer I made last week to keep the pH of the solution steady. From there we might be ready to start the mercury perchlorate titrations!

I hope you enjoyed the read! Stay tuned for the next step of the experiment!

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