Organic chemistry help

[Lioness]

Basically, I did an ester preparation practical. I refluxed glacial ethanoic acid and 1-butanol to make an ester, purified it, and then distilled it into the separate compounds.

Problem being, I didn't get to finish the distillation, and I'm unsure of what would have happened. The main problem being that ethanoic acid and 1-butanol have almost exactly the same boiling point. (118.1C and 117.2C)

I need to record observations, and I'm not sure what would've been observed. Any help would be really appreciated.

[Strawberries]

What kind of distillation equipment were you using?
Basically, you could have managed to separate the two compounds if you had a good enough equipment (the longer and more complex your fractionating column is, the closer the boiling point of the substances you manage to separate). On a good equipment, the temperature would have risen to 117.something, then stabilized then begin to rise again, then stabilized for the separation on the second compound. It really depends on the conditions, though.

EDIT: in my experience, since 1 degree is a really small difference, the temperature would have shot up directly to 118 and then stayed there, and you wouldn't have managed to separate the two compounds

Also, what was the reagent in excess in the original synthesys? (I assume 1-butanol, but you never know, since I assume this is a didactic experiment). Basically, are you sure that you have only butanol and ester in your flask?

[llamamushroom]

Also, what was the reagent in excess in the original synthesys? (I assume 1-butanol, but you never know, since I assume this is a didactic experiment). Basically, are you sure that you have only butanol and ester in your flask?

I don't think that's important, as the excess reactant(s) would have been removed in the purification process along with the sulfuric acid.

I admit that mine isn't exactly the most comprehensive Chemistry course in the world, but would distillation separate it again? Wouldn't it need to be in the presence of excess water and heat, preferably with dilute H2SO4, to push it back to the alcohol and carboxcylic acid?

In any case, you could probably just say that you didn't manage to separate the two (what with the ludicrously close BP°C, and the likely quality of you gear unless you're doing a university/uni-sponsered course) and mention in your hypothesis "the ester will split into its original reactants", but in your conclusion "the hypothesis was not met". Then, just whack in a bit about possible sources for error (close BP°C, faulty apparatus, etc.) and you're done.

[Strawberries]

I admit that mine isn't exactly the most comprehensive Chemistry course in the world, but would distillation separate it again? Wouldn't it need to be in the presence of excess water and heat, preferably with dilute H2SO4, to push it back to the alcohol and carboxcylic acid?



Yes, I think the ester would be pretty stable. And I'm an idiot for posting too early in the morning without thinking to the purification step :smalltongue:

But now that I'm more awake I checked and what do you know? Butanol and butyl acetate form an azeotrope, boiling point 390.75 K (117,6°C). Basically you wouldn't have managed to separate them anyway.

Source http://eweb.chemeng.ed.ac.uk/chem_eng/azeotrope_bank.html

Edit, not as awake as I thought. It was ethanoic acid, not acetic acid. Disregard what I said, I'm a double idiot. :smallbiggrin:

[Lioness]

Edit, not as awake as I thought. It was ethanoic acid, not acetic acid. Disregard what I said, I'm a double idiot. :smallbiggrin:

Ethanoic acid and acetic acid are the same thing, no?

Anyways, I discovered why I was so confused. We were meant to separate the ethanoic acid and butanol from the pure ester (BP 126), not from each other.

It all makes sense now. Thanks for trying to help though :)

[Strawberries]

Ethanoic acid and acetic acid are the same thing, no?

Anyways, I discovered why I was so confused. We were meant to separate the ethanoic acid and butanol from the pure ester (BP 126), not from each other.

It all makes sense now. Thanks for trying to help though :)

*facepalm*. That's IT. no more posts for me till 12 in the morning if I had a late night the day before. You are, of course, correct. They are the same thing.

With butanol and ester you expect to have a standard distillation. So, initially the temperature rises to 118 (and you are separating butanol and eventually the remaining of acetic acid that hasn't reacted), then it rises again at 126 (that's when you begin to separate the pure ester). As llamamushroom said, heating only isn't enough to break the ester into its component (you need an acid or basic catalist for that and of course at least a little water).

[rogueboy]

As people have said, you will see boiling points at ~117 and 126 (assuming that's the correct boiling point for butanoic acetate [the ester], I couldn't find it in my brief Google searching). Based on what my school does (I've TA'd our organic chemistry labs for 2 years now), I would expect the ethanoic (acetic) acid to be your excess reagent, since it is relatively cheap. If you haven't been told what your limiting reagent is, do the calculation (find the moles of each used): it's probably a key piece of what's being looked for in the report, and it should only take a minute once you have the density and molecular weight (assuming you're starting with volumes of acid and alcohol). This is also the first step in finding your yield, another piece of information I would expect to be in the report.

But yeah, it looks like pretty much everything has been covered already. Feel free to shoot me a PM if you have other OChem questions and I should be able to get back to you pretty quickly.

[llamamushroom]

Anyways, I discovered why I was so confused. We were meant to separate the ethanoic acid and butanol from the pure ester (BP 126), not from each other.

You were supposed to do that via distillation? But, surely you would've been using concentrated sulfuric acid in your initial esterification so as to only have one excess reactant (most likely the ethanoic acid), and then the washing process (distilled water, then a base, then distilled water again, all through a separating funnel until you're left with a not-colourless, and therefore not water, clear solution). Distilling it would only risk losing product.

I have an issue with your experiment, but not how you did it. It confuses me greatly.

[rogueboy]

You were supposed to do that via distillation? But, surely you would've been using concentrated sulfuric acid in your initial esterification so as to only have one excess reactant (most likely the ethanoic acid), and then the washing process (distilled water, then a base, then distilled water again, all through a separating funnel until you're left with a not-colourless, and therefore not water, clear solution). Distilling it would only risk losing product.

I have an issue with your experiment, but not how you did it. It confuses me greatly.

Actually, I would expect that the washing process would remove the majority of the acids (the concentrated catalytic acid as well as the majority of the ethanoic acid), but I would not expect the reaction to go to 100% completion, which means that there will still be some of the butanol in your flask, along with the ester, that can't be removed by simple extraction (chromatography of some sort would do it, but that's a more complicated and involved process). As such, you need a simple, quick way to separate the ester and the alcohol. Distillation (likely a fractional distillation so you get a better separation) is an outstanding way to do this when you have 2 liquids with differing boiling points (like an alcohol and its acetate ester). Yes, you do expect to lose a small portion of your product, but that is always true of any purification method (the exact amount you will lose depends on the internal volume of your distillation apparatus if your product is the lower boiling component, and how long your fractionating column is if it's the higher boiling component [like it is in this case]).

In the organic lab that I TA, we do the same esterification with 1-, 2-, or 3-pentanol. All of them give acetates that, after the wash, are colorless, clear liquids that are then separated from the starting alcohol. Note: students typically have a colored solution when they are done heating the solution, which is likely from contamination on the stir bar or the inside of the flask. The color is removed during the washing process, one that typically involves a water wash, a bicarbonate [basic] rinse, and then a saline [sat. NaCl] rinse to "pre-dry" the organic layer.

As a side note, in a reaction like this where you only have 2 reactants, how do you plan to have 2 excess reagents? This seems to be what you're suggesting with your comment about the reasoning for having concentrated sulfuric acid... the actual motivation here is that the reaction will not proceed without an acid catalyst.

Out of curiosity, how would (or did) you do the purification of a similar ester product if not by distillation?

[llamamushroom]

As a side note, in a reaction like this where you only have 2 reactants, how do you plan to have 2 excess reagents? This seems to be what you're suggesting with your comment about the reasoning for having concentrated sulfuric acid... the actual motivation here is that the reaction will not proceed without an acid catalyst.

Out of curiosity, how would (or did) you do the purification of a similar ester product if not by distillation?

Sorry if my reply was poorly worded. The reason for having the concentrated sulfuric acid is twofold: yes, it is a catalyst, but also quite important is the fact that concentrated sulfuric acid is a dehydrating agent. Thus, the reaction [alcohol] + [corboxylic acid] -><- [ester] + [water] can be driven to completion (by Le Chatelier's Principle) via the removal of water by the acid.

When we did this experiment, we had excess organic acid (the exact one escapes my memory), so all the alcohol was used up. Then, a base wash (to remove the excess acid) and a distilled water wash (to remove the water-soluble contaminants) made it fairly pure... then the bell rang, so we spent the 5 minutes of getting-to-class time debating on how best to remove the rest of them. I believe a non-polar solvent was discussed, but shot down, and our teacher said he'd tell us how to purify it the next lesson. He didn't. Suddenly, I want to find out the answer to that question... (he says, guessing that it was "fractional distillation". Damn! Foiled again by my ancient enemy)

[Lioness]

We had some of everything left over after reflux - we didn't have time to reflux it all the way. Sulphuric acid was used as a catalyst; I'm not sure about using it as a dehydrating agent. We washed it with water, washed it with some sort of carbonate, and then fused calcium chloride. The final distillation was to separate the ester from whatever was left after all that

[rogueboy]

We had some of everything left over after reflux - we didn't have time to reflux it all the way. Sulphuric acid was used as a catalyst; I'm not sure about using it as a dehydrating agent. We washed it with water, washed it with some sort of carbonate, and then fused calcium chloride. The final distillation was to separate the ester from whatever was left after all that

Not refluxing it to 100% doesn't surprise me, either because of time issues or just the equilibrium. I've never actually heard of using the sulfuric acid as a dehydrating agent... I'll have to ask some of my profs about it. The calcium chloride is a drying agent, which will just remove any water that's remained in the organic layer after the extractions (washes).

[rogueboy]

No one I spoke to today (2 organic profs and a lab coworker) could make any sense out of the idea of using H2SO4 as a dehydrating/drying agent ("dehydrating" typically refers to the removal of water, typically a protonated alcohol, from the substrate rather than the removal of water from the solvent). The closest explanation that we (as a group) could come up with was the blunting of the nucleophilicity of water by protonation. However, it should be noted that using sulfuric acid, even concentrated sulfuric acid, does not mean that H2SO4 is the only thing being added: concentrated sulfuric acid is an 18M solution of H2SO4 in water. Also, since sulfuric acid is being used as a catalyst, there simply isn't enough of it in order to "soak up" the water generated by the esterification in the way that a drying agent would.

There are 2 standard ways to remove water during a reaction: use of a drying agent or a Dean-Stark trap. Using a drying agent is simple: you put some drying agent (cesium chloride, calcium chloride, sodium sulfate, magnesium sulfate, molecular sieves, or any of the other common drying agents) in the flask with the rest of your reagents. Note that because you're using an aqueous (water-based) catalyst, I'm not sure how this would work for this reaction. The other option, and one that is more typically used if you have access to the glassware, is the Dean-Stark trap. This is basically a distillation where you allow the water and organic solvent evaporate and condense into the side flask of the trap. This is typically filled with the organic solvent so that when the water and solvent condense it replaces the solvent, preventing your reaction from "running dry" (losing all the solvent). Obviously, this only works with an organic solvent that is less dense than water (fortunately, this is almost all of them). It is possible to do something similar with a more dense organic solvent (such as dichloromethane), but it is more complicated and I don't know what it's typically called.

Hopefully this hasn't scared anyone off from chemistry (it's a lot of text, I know, but I usually get excited when I talk about OChem) and it's been useful and/or informative :smallwink:

[rogueboy]

No one I spoke to today (2 organic profs and a lab coworker) could make any sense out of the idea of using H2SO4 as a dehydrating/drying agent ("dehydrating" typically refers to the removal of water, typically a protonated alcohol, from the substrate rather than the removal of water from the solvent). The closest explanation that we (as a group) could come up with was the blunting of the nucleophilicity of water by protonation. However, it should be noted that using sulfuric acid, even concentrated sulfuric acid, does not mean that H2SO4 is the only thing being added: concentrated sulfuric acid is an 18M solution of H2SO4 in water. Also, since sulfuric acid is being used as a catalyst, there simply isn't enough of it in order to "soak up" the water generated by the esterification in the way that a drying agent would.

There are 2 standard ways to remove water during a reaction: use of a drying agent or a Dean-Stark trap. Using a drying agent is simple: you put some drying agent (cesium chloride, calcium chloride, sodium sulfate, magnesium sulfate, molecular sieves, or any of the other common drying agents) in the flask with the rest of your reagents. Note that because you're using an aqueous (water-based) catalyst, I'm not sure how this would work for this reaction. The other option, and one that is more typically used if you have access to the glassware, is the Dean-Stark trap. This is basically a distillation where you allow the water and organic solvent evaporate and condense into the side flask of the trap. This is typically filled with the organic solvent so that when the water and solvent condense it replaces the solvent, preventing your reaction from "running dry" (losing all the solvent). Obviously, this only works with an organic solvent that is less dense than water (fortunately, this is almost all of them). It is possible to do something similar with a more dense organic solvent (such as dichloromethane), but it is more complicated and I don't know what it's typically called.

Hopefully this hasn't scared anyone off from chemistry (it's a lot of text, I know, but I usually get excited when I talk about OChem) and it's been useful and/or informative :smallwink:

[Strawberries]

It is possible to do something similar with a more dense organic solvent (such as dichloromethane), but it is more complicated and I don't know what it's typically called.


It's still called a Dean-Stark trap. There are two kind, with different arrangements based on the density of the solvent.

Molecular sieves are very used. Unfortunately they work best in "dry" reactions, where you have only traces of water to remove, otherwise they are quickly inactivated. In water-rich reactions, Dean Stark is your best bet (introducing other drying agents can cause unwanted side reactions, wich is why you use them after the reaction is quenched, in the washing phase)

You didn't scare me off :smalltongue:

[llamamushroom]

No-one had heard of using H2SO4 as a dehydrating agent in esterification? Suddenly, I'm beginning to question my teacher, and yet he has a ludicrous amount of experience in Chem, has taught it for a very long time, and is one of the markers for the IB*. I checked my text book (Conquering Chemistry by Roland Smith), and it also said that the sulfuric acid "absorbs the product water and so moves the equilibrium to the right". This is... unsettling, to say the least.

When I talked to my teacher about the purification of our ester, he said he'd never really thought that using distillation was particularly useful. His preferred method, which he used recently to help the ANU chem department in producing pure esters, is to use different concentrations of ethanol in water at low temperatures. This removes the non-polar contaminants without dissolving the ester.

Gah! This thread has raised more questions than it answered (to my mind, at least).
___

*He showed us some of the papers he's marked, some of which come with notes like "[applicant's name withheld by teacher]'s experiment was well thought-out, and there are no flaws in their practical work or results." Which was when he showed us how this person had stated that you get sulfuric and hydrochloric acids in acid rain. He/She apparently mixed up solutions of their chosen acids to a pH of 4.5, except that one of them was carbonic acid. There were quite a few more errors, like when sulfuric acid was suddenly replaced by sulfur dioxide mid-write-up. The only reason they got 40% was because it was formatted very well, and that's one of the criteria.

[Asta Kask]

No-one had heard of using H2SO4 as a dehydrating agent in esterification? Suddenly, I'm beginning to question my teacher, and yet he has a ludicrous amount of experience in Chem, has taught it for a very long time, and is one of the markers for the IB*. I checked my text book (Conquering Chemistry by Roland Smith), and it also said that the sulfuric acid "absorbs the product water and so moves the equilibrium to the right". This is... unsettling, to say the least.

Raises hand - I have. Try pouring concentrated sulfuric acid onto a carbohydrate and see the water disappear... wee - coal sausage!

[rogueboy]

It's still called a Dean-Stark trap. There are two kind, with different arrangements based on the density of the solvent.

I suppose this is what I deserve for trying to talk about something I know very little about (my research is almost entirely dry chemistry, and nothing I do produces water as a byproduct) :D


Molecular sieves are very used. Unfortunately they work best in "dry" reactions, where you have only traces of water to remove, otherwise they are quickly inactivated. In water-rich reactions, Dean Stark is your best bet (introducing other drying agents can cause unwanted side reactions, wich is why you use them after the reaction is quenched, in the washing phase)

You didn't scare me off :smalltongue:

Yay, someone else who can explain this stuff more effectively than me (while I'm busy leaving for the ACS conference, no less)! Glad to see there are other people who know chemistry quite well on these boards, too! But yeah, non-MS drying agents can create side reactions, but that's why you would have to select them carefully (why you would use them, I don't know, but it's an option! :smalltongue:)

And llama, I have no idea what to say to your prof/book's comments about the water... I know I recognize the name Roland Smith (I'm blanking on how, but I definitely know the name - yes, the fact that I'm blanking on this is bugging me...) as a good organic chemist, so now I'm confused as well... My primary argument against its use as a dehydrating agent would be that there is very little of it, since it's a catalyst - even if each molecule of H2SO4 can absorb 10-20 molecules of H2O (a *very* strong dehydrating agent), that would still require 5-10 mol% in order to act as a drying agent, which would mean you need 6-15% as a catalyst... and typically no more than 1-2% is used (this is ignoring the fact that it is 18M, while the water it's dissolved in is ~56M, or already 3x as concentrated).

[llamamushroom]

Yeah, this has been a frankly worrying experience.

So, I hope we managed to help you with your problem, Lioness!

[Lioness]

So, I hope we managed to help you with your problem, Lioness!

Yup, and I've had a rather enjoyable read of the debate following. When I get my results back I'll tell how it went

[Strawberries]

I suppose this is what I deserve for trying to talk about something I know very little about (my research is almost entirely dry chemistry, and nothing I do produces water as a byproduct) :D


Ah ah, I graduated with a thesis in synthetic organic chemistry and, as my professor liked to say, "an organic chemist uses water only to drink :smallwink:"


When I get my results back I'll tell how it went

Good luck! I'll keep my fingers crossed for you!