[RaInBowDaSh1488]

I realize that mastering TLC is a fairly basic skill for organic chemists. In reality, I consider myself somewhat competent in the practice. A recent reaction really has me stumped though. Here is the scenario.

I am synthesizing 1-chloro-3-phenylpropane from the corresponding acid via thionyl chloride. I've done many chlorinations with thionyl chloride and this one seemed to be proceeding as expected. TLCs I've taken have been a bit confusing though. I have a reagent grade sample of the compound I am synthesizing, so I am using it as a reference in my TLCs.

The strange thing is, the pure sample produces two distinct spots on the TLC. The spot with the lower RF corresponds exactly to a spot in the lane containing material from my reaction flask. The odd thing is, the RF seems a bit low for a rather non-polar halide. The solvent system I am using is 8:2 hexanes/EtoAc.

I originally thought I might have produced an aldehyde instead of the halide, but I have confirmed this is not true.

Is there any reason a compound such as 1-chloro-3-phenylpropane might be expected to produce two spots on TLC? I really can't understand why this is happening. The product is very new, I was actually the first to open it. Additionally, I have taken every precaution to avoid contamination in the process of running the TLC. Every time I have used clean vials, fresh capillaries, and very new ethyl acetate to dilute it.

I realize this question is a bit "out there" but any input would be appreciated. I'll be running an NMR tomorrow to ID what was produced, but I am just baffled by the behavior of my reference compound!

[camptzak]

is it possible you did not convert all the alcohol to chloride? The alcohol would account or the low Rf. either way let me know, I am interested.

[RaInBowDaSh1488]

is it possible you did not convert all the alcohol to chloride? The alcohol would account or the low Rf. either way let me know, I am interested.

Thanks for your interest. The reaction did not go to completion after refluxing for roughly four hours.There was likely some of the alcohol remaining in the flask. Frankly, I've had bad luck getting chlorinations with thionyl chloride to go to completion. Usually I do synthesize enough of the desired product to separate and be used for further reactions.

I am very intersted in what you have said about the alcohol lowering the RF. Are you saying that the presence of the alcohol maybe lowering the RF of whatever is being produced? If that is the case, I have a lot of food for thought!

The thing that bugs me the most, is the fact that the a compound that is supposedly "pure" is showing two spots. I realize this can be the case for certain compounds, but I wouldn't expect this sort of behavior from 1-chloro-3-phenylpropane.

[Babcock_Hall]

I realize that mastering TLC is a fairly basic skill for organic chemists. In reality, I consider myself somewhat competent in the practice. A recent reaction really has me stumped though. Here is the scenario.

I am synthesizing 1-chloro-3-phenylpropane from the corresponding acid via thionyl chloride.

(bolding mine) Is the starting material an acid or an alcohol?  When a compound is highly reactive, it can decompose on silica.  I doubt that this is the case here, but it is worth considering.

[camptzak]

I'm just saying that if the alcohol was not all converted to chloride, you have both the chloride and alcohol. two spots on a tlc usually means two different compounds. The alcohol, as you know, would have a lower Rf than the chloride. try using 1.5 equivalents of the thionyl chloride or more.


oh and also, repeating Mr. Babcock hall's question. was your starting material an alcohol or an acid?

[subu]

Did you consider checking purity of your reference compound (reagent grade commercial one)?
If you are done with NMR of what you procduced, how different is this compared to the reference?

[TwistedConf]

The thing that bugs me the most, is the fact that the a compound that is supposedly "pure" is showing two spots.

It probably isn't pure. You need to, as suggested above, use another technique or two if you really want to analyze what you bought.

As an aside-- can you comment on why you're going through all this to prepare a small molecule that's (in theory) readily commercially available.

[RaInBowDaSh1488]

The compound being synthesized is obviously very cheap and available. As I mentioned, we have the exact compound in stock.

The reason I was conducting this chlorination was to test some thionyl chloride that had been sitting for quite a while. Other people had run reactions with it and we considered it possible that it had decomposed, or was no longer useable.

The starting material was an alcohol. Here is the strange thing, by playing around with the TLC solvent system, I was able to produce results that were more inline with what I would have expected. Under a different combination of solvents the chloride I was using as a reference only formed one spot. The spot in question had a reasonable RF and corresponded with the spot that I believe represents the equivalent chloride I have synthesized.

This is very strange to me! I admit, I often have a very hard time finding a suitable solvent system to achieve good resolution in my TLCs. It just seems rather hard to predict which solvent system is appropriate in a given situation. Also, I hate to waste tons of TLC plates trying to find a good solution.

The NMR of the commercially supplied reagent was perfectly normal. I haven't yet had the chance to take an NMR of my product yet.