kilomentor | 15 November, 2010 15:34
Besides the Minimum Stirrable Volume another Critical Volume Parameter of a Proposed Reactor is its Maximum sirrable Volume.
When scaling up to kilolitre or pilot plant size, using fewer pieces of equipment significantly reduces the cost of a process step. This is true because equipment costs money to buy, suffers depreciation upon usage, and is only available for one process step at a time in a plant where several projects are going on in parallel. Furthermore, when exposed to process chemicals, equipment must be washed, rinsed, dried, validated as clean, and in many instances stored for reuse. These things all cost money and take time.
In order to minimize the number of vessels used in a step, the workup of a reaction is preferably conducted in the reactor vessel. In order to calculate the highest possible kilograms of intermediate product that can be produced in a particular reactor in a single run one needs to first determine what point in the protocol requires the largest stirred volume. Then the protocol is rescaled so that at that point this volume is equal to the maximum stirrable volume of the particular reactor. For example, suppose the protocol you are going to follow will produce 45 grams of purified product and at the point of maximum volume in this procedure the combined organic and aqueous solutions are 500 mL. if the maximum stirrable volume for the reactor you propose using in production is 2200 litres then the maximum throughput per run will be (2200/500) X 45 = 198 kg. Now if the target of your project will require you to make 450 kg of this intermediate, you will need at least three runs of this step because the maximum amount two runs can give is 396 kg.
Now if you could modify the protocol so that at the point of maximum volume the volume were only 440 litres, you could make the 450 kg of material in just 2 runs. Providing a margin of safety against shortfalls by planning for 3 runs may be the wisest course of action.
kilomentor | 15 November, 2010 15:31
In a chemical pilot plant or plant one cannot choose a reactor as easily as one chooses a 250, 500,1000 or 5000 mL flask. One has to work with what is available and that may be constrained not just by the equipment within four walls by what other processes are planned to be run at the same time as yours.
A plant reactor can be characterized by two crucially important volumes: the Minimum stirrable Volume and the Maximum stirrable volume. The former will be discussed here and the latter in another blog article.
The minimum stirrable volume is exactly what the name teaches. It is the minimum volume of liquid needed in the reactor so that the turning stirrer paddles effectively stir the reactor contents. To some extent this minimum stirrable volume depends upon what is being stirred and what will happen during the initial stage of the reaction of interest. For example, a homogeneous solution in which an exothermic refluxing occurs at the beginning of the reaction may only need the moving blades to touch the liquid surface because gas evolution and thermal convection are going to move the homogeneous liquid phase around. At the other extreme, if zinc powder, tin granules or magnesium turning need to be swept up off the bottom of the reactor, the immersion of the stirrer blades probably needs to be complete and the stirring rapid.
The reason chemists, who are more accustomed to working in the laboratory tend to forget the minimum stirring volume constraint is that two of the most common laboratory stirrer types are the magnetic stirrer and the crescent bladed overhead mechanical type. The magnetic stir bar on the bottom of the flask so that the minimum stirrable volume is very small while the crescent bladed stirrer is very often set up with its curved edge nestled up against the bottom of the round bottomed flask in which it is installed so that again the minimum stirrable volume is close to zero.
In order to use this parameter the chemist examines the procedure which he intends to use and decides at what points in the process stirring will be required. The chemist estimates what the reactor volume will be throughout the process. The process cannot be run with less material that will allow stirring at the point of lowest point of low volume where stirring is essential. This is important in the very earliest pilot plant runs when one does not want to risk more material than is required to see whether the scale up proceeds without problems.
An example may clarify this. Suppose you want to run a process step for the first time in the 2000 litre reactor which is planned for the production. The starting materials are expensive and you do not want to risk any more materials in this first run than is necessary. The minimum stirrable volume in this 2000L reactor is 200L. Your protocol process will produce 45 gm of the desired intermediate from 40 g of starting material. Twenty grams of insoluble anhydrous potassium carbonate need to be stirred in the solvent during reaction. The point in the process where stirring is essential and the total volume is lowest is right at the beginning when you starting adding a reagent drop-by-drop. At this point you have 40 gm of starting material, 20 gm of potassium carbonate and 300 ml of solvent. Very roughly the total volume at this point of minimum volume is 360 ml.
The minimum amount of starting material you can use without any change in your lab procedure is 200/360 X40 =22.22kg. If you try using less after loading the reactor with the starting material, the potassium carbonate and the organic solvent the slurry will not touch the stirrer blades.
Suppose management is not happy with this. They don’t want to risk more than 10 kg of starting material in the first run. The best answer is to perform the first run in the smallest reactor that can be found that has the same configuration as the reactor planned for the full production. A smaller reactor will typically have a smaller minimum stirrable volume. This one reason that a process is often scaled up in steps.
kilomentor | 12 November, 2010 10:40
Neal G. Anderson in his book, Practical Process Research and Development, breaks down the time required to perform the operation of drying an organic solution with a solid inorganic desiccant such as sodium sulfate. He estimates that, including equipment preparation and equipment cleaning, the unit operation would take two operators 8 hours. Only three hours are actually spent performing the physical operations that parallel the laboratory manipulations. An additional one hour is spent assembling and testing the filter and its associated piping. At the end of the actual filtration, four hours are required to rinse, disassemble, clean, validate and store the equipment ready for reuse. These same times apply to the operation of treating an organic solution with decolorizing charcoal and filtering it off.
Chemists more accustomed to working at a laboratory scale cannot initially imagine how these preparation and cleaning phases can consume so much time. In the lab, we are accustomed to using glassware that is cleaned with a quick acetone rinse, transferring liquids simply by gravity, pouring hot fluids through the air and assessing cleanliness by visual inspection. On-scale however whether the solution is hot or cold, filtration needs to be done in a closed inert system. The general uncomplicated pilot plant filtration compares more closely to working in the laboratory with a toxic liquid or a pyrophoric solid in Schlenk-tube equipment.
With regard to drying organic solvents, once it is grasped that removing water is difficult and costly process, chemists can try to avoid adding water or think about ways to remove it without the complication of additional equipment. Pouring a reaction mixture into water or adding water into an organic reaction mixture is a frequent operation incorporated into a reaction step. This standard approach, which weshould be trying to rethink, uses a humongous excess of water, usually added rapidly.
Why is this done? One reason is that it usually stops the reaction. This is why the operation is often called quenching the reaction. Water is a evaporable, cheap low strength buffer. Since the strongest acid that can exist in a wet solution or an aqueous phase is H3O+, and the strongest base is -OH, many reactions that require either more acidic or more basic conditions are stopped by adding water. Another mechanism also operates when a water quench is performed. The addition of sufficient water to produce two immiscible phases in the reactor quenches by phase separation of reaction participants. Some chemicals are much more soluble in the aqueous phase while others are much more soluble in the organic layer. Separation of reactants, reagents, products, co-products and byproducts into one or the other of water or organic usually reduces the rate of reaction among them to essentially zero. A third mechanism for aqueous quenching is the rapid decomposition of one of the reactants by water.
If we remove a water quenching step in a process, it is still usually necessary to stop the reaction at the correct end point. Whatever that operation is it must be rapid. What can be done depends upon the characteristics of the reaction itself. Whatever the quenching additive used it is very preferably inexpensive.
One option is to employ some water, but not the large volumes that result in a second phase. In the laboratory, the volume of water used in a quench often doubles the total volume in the reaction vessel. When the quench is done by pouring, with vigorous stirring, the reaction mixture onto an ice-water slush, the volume is usually still greater. The large volume is used ‘just to be sure’ because the laboratory reactions are not investigated enough to figure out the actual lowest amount of water that is needed. At scale, the total reactor volume immediately after quenching but before separation of this water is very often the point of maximum reactor volume and it limits the batch size. Using a larger quench volume than necessary immediately lowers the possible throughput for the process step. Increasing the volume at the point of maximum volume in the reactor increases the cost per kilogram of intermediate product from the step.
Furthermore, creating a large aqueous phase saturated with organic substances creates a disposal cost. Water contaminated with a saturation level of organic contaminants cannot be discharged to a municipal sewage treatment system. The water becomes a waste stream and one that cannot be treated inexpensively by burning. Reducing the amount of the water quench phase will greatly reduce the kg of waste produced per kg of product. It is an easy way to make a process step more environmentally friendly.
Replacements for copious amounts of water could be ammonium chloride or ammonium acetate both of which are quite soluble in lower alcohols and for which any residue is volatile. These would satisfy the buffering function. In other situations acetic acid or ammonia might work. Grignard reactions and hydride reactions are often treated with ethyl acetate.
Even if water-organic partitioning in a liquid-liquid extraction is required in the work-up procedure, there can still be an advantage to using a minimum volume of quench. Once the reaction has been stopped it may be possible to reduce the volume of the reaction mixture by distillation before adding the larger portion of the aqueous extraction phase so that the volume/kg product at the point of maximum volume is reduced and the throughput can be increased. For example, an alkylation under strongly basic anhydrous conditions could be quenched with ammonium chloride in methanol and then the volume of the reaction mixture could be reduced by half using vacuum distillation before adding water for an acid-base extraction to purify the product by phase switching. The volume/kg at the point of maximum volume could be reduced by half and the number of runs required to meet a production target cut in half.
Scientists more accustomed to laboratory synthesis would be benefited by recognizing that in a plant setting, , one usually cannot do the equivalent of simply switching from a 500 ml to a 1 liter reaction flask. The chemical reactors available are few in number and cannot be simply interchanged since, for example, some may be glass-lined and others stainless steel, or one may have a different heating/cooling capability from another. Still, increasing the throughput per batch becomes a big deal!
Drying an organic liquid has the meaning here of reducing the water content most often to negligible or alternately to any other practically acceptable level.
Drying is often not necessary. In the laboratory drying organic solutions is conducted routinely. Most often no effort is made to determine whether it is useful or necessary. Drying may be performed because the cut between an aqueous and an immiscible organic phase may not be perfect and there may be a concern about small water droplets in the organic layer. In the lab it is easier to be safe rather than sorry. In the plant discovering thatdrying is not of concern can simplify the process and save costs.
The most widely known and widely practiced method applied in the plant is azeotropic distillation. Azeotropic distillation has the advantage that the overall volume is reduced during the operation. Also because almost all reactors have reflux and simple distillation capabilities, no additional equipment is used and dirtied with product by the operation. The disadvantages are that the operation requires heating and also frequently a vacuum. Water is very often less volatile than the organic solvent used in the reaction and many solvents do not form azeotropes with water. It is not much of an answer to suggest that the process step reaction solvent be switched to one that does form a useful azeotrope because this reduces solvent choice which is one of the most influential variables that can be used in reaction yield optimization. Reaction solvent choice should be kept as free as possible.
Where heating is a concern or where an azeotropic method is not possible, operators at scale have found that they can pass a solution through a fixed bed of molecular sieves to achieve drying. Molecular sieves that have absorbed water can be renewed by the passage of hot dry air once the sieves have been washed to remove external contamination.
Another procedure that can remove water from a solution is to add a reagent that reacts preferentially with water and produces products that are more easily removed than water.
Triethylorthoformate or trimethylorthoacetate for example can consume water and produce an ester and alcohol products. DMF dimethylacetal can react with water to produce dimethylformamide and methanol and remove water. Bis-trimethylsilyl-trifluoracetamide can also react with water to give Bis trimethylsilylether and trifluoracetamide.
Replacing Charcoal Treatment
Charcoal decolorizing treatment has the same time and equipment disadvantages as drying with inorganic salts only more severe. Because charcoal is insoluble and very finely divided in many cases, proper cleaning is harder and the evidence of inadequate cleaning (black particles) is embarrassingly obvious. Methods of purification of crude solids that are less burdensome are preferred but there are occasions when charcoaling cannot be replaced. At scale, producers often have dedicated charcoaling equipment so that most of their equipment is never contacted with charcoal. Sometimes pumping a solution through a fixed bed containing charcoal works as well as adding it into the organic solution.
kilomentor | 13 October, 2010 11:33
For every pure compound there may be one related substance that arises naturally from its route of synthesis that would be particularly difficult to remove. This I call the Impurity from Hell.
In November 2007, I wrote an article for my blog, Kilomentor, that should be of interest to chemists performing stability studies. The article imagined the worst possible impurity that could realistically contaminate a final product. The idea was to consider using a synthetic preparation of such impurity to test the resolving power of the analytical separation technique that is to be used in stability studies. Since thechemist developing the analytic methdology does not know at the outset of a stability investigation what all the impurities will be present as synthetic impurities in the API arising from different syntheses or that will arise in the degradation studies, it is a problem to know when one has achieved an HPLC method that is discriminating enough to separate all potential impurities, almost certainly, without any exception. The separation of a compound selected because of its verisimilitude with the product should be an excellent test. The original blog was titled, Avoiding the Impurity from Hell.
I have exemplified this strategy here by applying it to the drug, pregabalin. This exercise is not a veiled statement that the impurity is present in any company’s product. Even ‘an impurity from hell’ may be removed by the phase shifts inherent in the synthesis itself. What we can say with some confidence is that the impurity could easily be present and would, if present, represent an analytic challenge. Some of the routes to pregabalin start from and these routes do risk the impurity discussed here. Certainly including these diastereomers of ‘this impurity from hell’ as part of a test of an analytic method for pregabalin would be a significant challenge.
As I have discussed in my previous related kilomentor blog, one can hypothesize the structure of the impurity from hell. Usually it is a structural isomer differing in a floppy alkyl side chain and it is a potential synthesis derived impurity based on the route used.
I propose that the impurities from hell in the drug substance pregabalin may be any of the diastereomers of (3RS, 4RS)-3-aminomethyl-4-methylhexanoic acid. They would arise if instead of an isobutyl group a secondary butyl group gets incorporated into the molecule. Pregabalin itself is (3S)-3-aminomethyl-5-methylhexanoic acid. 3-Aminomethyl-4-methylhexanoic acids arise if there is any 2-methyl-butanal impurity in the 3-methylbutanal (isovaleraldehyde) used as starting material in the pregabalin synthesis. The likelihood of having this isomeric aldehyde in the starting isovaleraldehyde depends upon the route of its synthesis. Unfortunately obvious routes of synthesis can be expected not to mitigate the risk. If the isovaleraldehyde is made from isoamyl alcohol, it is important that the isoamyl alcohol be free from 2-methyl-1-butanol. This cannot be guaranteed by distillation. Isoamyl alcohol has bp. 130 C which is the same as for 2-methyl-1-butanol. Once 2-methy-1-butanal gets into a sample of isovaleraldehyde, it would be difficult to remove it since Aldrich reports that isovaleraldehyde has boiling point 90 C and 2-methyl-1-butanal is 90-92 C!
kilomentor | 10 October, 2010 14:10
It has been repeated in the Kilomentor blog like a mantra that a synthetic route with intermediates which can be reversibly extracted into aqueous solution by pH adjustments is very advantageous. If intermediates are neutral and do not meet this need, then on scale they most often will have to be crystallized or distilled for purification. When an intermediate’s molecular weight is too high its boiling point is impractical for large scale distillation. When structures are conformationally very flexible they have a greater tendency to be low melting and difficult to crystallize in good yield.
Triethanolamine is immiscible or very poorly miscible with many organic solvents. The liquid is used in large volumes in scrubbing systems for chemical reactors. Even the 98% pure compound is only about $20/kg from Aldrich. The b.p. is 190-193 C @ 5 tor. Because it is a base it can be readily extracted into aqueous solution with weak acid.
Codistillation with many neutral organic substances is highly likely. Since the triethanolamine is only partially miscible with quite a few popular solvents [benzene (4.2%), toluene (?%), ether (1.6%), n-heptane (<0.1%)], the co-distilled substrates can be back extracted into these lower boiling common solvents, washed with weak aqueous acid to remove the small amount of triethanolamine carried with the substrate and the thus purified product fraction isolated by crystallization or precipitation with or without an antisolvent.
This suggestion needs to be examined with the perspective that direct crystallization of a somewhat impure compound say with 10% of an impurity is unlikely to recover any more than 80% of the actual assay yield of product present. A high yielding codistillation that leaves the 10% of major impurity behind followed by crystallization is much more likely to crystallize the desired intermediate in near quantitative yield since the overall impurity level is now very low.
kilomentor | 25 September, 2010 13:40
Process Intermediates are dried because it is only measurement of the dry weight of the isolated product that can give the percentage yield. It is not necessary to dry the entire amount of the intermediate in order to measure the purity so long as a representative sampling can be done, the purity can be measured for the entire batch.
The process chemist needs to know the weight of dry intermediate so the amount of reagents and solvents for the next step can be correctly calculated.
Drying an intermediate to negligible further weight loss is not a simple task for kilograms of material. Because pumps and ovens are needed for the intermediate from step where isolation is planned, if several process steps are running in parallel drying capacity can be easily exceeded. When drying capacity may become rate limiting for the processing it is important to consider what can be done to reduce as much as possible the starting wet weight and increase the rate of drying. This will reduce the time to complete a run for a particular intermediate and increase the rate at which a campaign can proceed. Drying time and the equipment dedicated is often not considered by laboratory chemists optimizing a procedure for scale up because in the laboratory capacity is almost always more than adequate and drying is done most often over-night when time expended is “free”. As long as the drying is complete when the researcher comes back innext morning, no problem is registered.
Drying proceeds both on the solid filter or in the centrifuge and in the drying oven or other drying equipment. A longer period sitting with air being sucked through a solid cake on the filter will reduce the time required in the drying oven.
The length of time required for drying is a function of the volatility of the liquid being evaporated, the temperature of the incoming ga,s and the pressure in the dryer. Water is among the most difficult solvents to remove. There can be advantages to washing a high boiling anti-solvent with a low boiling anti-solvent before attempting drying because this will reduce the drying time.
Small crystals occlude more solvent than large ones.
When convection air drying is being used a more elevated inlet air temperature can substantially reduce the drying time. Even when a product has some instability from heating, the inlet temperature can be higher at the beginning of drying because the chemical substance is protected by the cooling from solvent evaporation. The temperature inside the solid mass is lower than the inlet air temperature.
Tray drying is very irregular. To get a reasonable estimate of the residual solvent, portions of the sample must be taken from different locations and combined.
Drying is usually much better understood by engineers than by chemists. If you are a chemist get help; it can save real money.
kilomentor | 08 August, 2010 10:14
Kilomentor has already written about this subject. See the blog, Making a Good Recrystallization Better.
Kilomentor consistently prefers processe that utilize iintermediates that can be protonated or deprotonated at pHs accessible in water and which therefore can undergo phase switches as part of purifications. Such routes are more rugged to the extent that they are less depend upon purification by crystallization of intermediates whose physical properties at the time the route is designed, are unknown. Nevertheless, even with ones best efforts to use routes where the intermediates are acidic or basic, once the molecular weight of intermediates exceeds what can be practically distilled, crystallization is the predominant isolation and purification method for neutral-unionizable intermediates.
Crystallization is of such importance that it is taught early on in (what there is of) laboratory training in universities. The disadvantage of this is that treatment is elementary and laboratory scientists never seem to get around to more sophisticated discussions about it, but learn what more they can by experience, both good and bad.
The crystallization / recrystallization procedures are probably the most important in the laboratory, yet the basic principles that underlie the operations are not widely understood by the average practitioner and rule-of-thumb insights are not passed on.
The most common mistake relating to crystallization is try to crystallize without first applying other methods of purification. The deleterious effect by impurities upon the rate and completeness of crystal formation for the main product is so pronounced that crystallization of crude products should never be attempted until other methods of purification have been applied This is particularly true because these pre-treatments are often very simple. Although crystallization has a high probability for successfully removing impurities, many synthetic chemists do not recognize that they pay an unnecessarily heavy penalty by recrystallizing crude products because the loss from incomplete sluggish crystallization is avoidable. Consideration should be given to first:
· distilling in vacuum
· co-distilling with another solvent
· steam distillation with or without the presence of salt
· superheated vacuum steam distillation;
· exhaustive digestion with a poor solvent
· extraction in a continuous extraction apparatus;
· passage through a plug of a solid adsorbant;
· or acid base extraction (to remove acidic or basic impurities even when the major compound is neutral)
· treatment with derivatizing reagent to trap identified (or guessed impurities)
· treatment with ionic exchange resins
· treatment with scavenger resins
Water and other solvent residues count as impurities that reduce the rate and completeness of crystallizations. Of course these solvents also play a part by affecting the actual crystalline substance that is obtained since they form solvates and hydrates. Thus adding water during the work up means that a thorough drying and pumping down needs to be done before crystallization. This in turn leads to the question whether the traditional pouring into water is always a good idea.
kilomentor | 03 July, 2010 15:11
Aldehydes and Ketones are among the most common and best understood functional groups in organic chemistry; however, they can be problematic as intermediates in large scale process chemistry because they are neither markedly acidic nor basic and so cannot be extracted as salts into aqueous solution to achieve purification by phase shifting. During the route planning stage, the synthesis creator cannot readily guess whether these carbonyl intermediates with be crystalline or not. Aldehydes and ketones that are neutral and have no other extraction handle thus are potentially isolation and purification problems. They can turn out to be oils or low melting solids.
Low molecular weight aldehydes and ketones are most often purified by fractional distillation either at atmospheric pressure or under reduced pressure. When they have boiling points in the neighbourhood of 200 C, steam distillation can provide a partial fractionation, but steam distillation is almost always unacceptable because of the very high large point of maximum volume inherent in the procedure. Compounds such as 7-tridecanone [m.p. 30-32.5 C; b.p. 264 C]; 2-pentadecanone [m.p. 7-41 C; b.p. 293 C]; or 2-heptadecanone [m.p. 47-51 C] are representative of these in-between type substances. So although one cannot say with certainty that an intermediate molecular weigh,t neutral carbonyl compounds is going to difficult to separate/purify it is good to have some precautionary potential patches in mind.
Although oximes derivatives of carbonyl compounds are not completely dependably solids, the likelihood that the oxime is a recrystallizable is more than for the carbonyl itself and increases as the number of carbons increases. Shriner, Fuson and Curtin in their classic manual, The Systematic Identification of Organic Compounds, A Laboratory Manual Wiley 1964 report that out of 63 liquid ketones, 44 had solid oximes. Out of 44 liquid aldehydes, 34 had solid oximes, some of which were separated into syn and anti forms.
If the oxime is not a solid, then the possibility for making the oxime hydrochloride adds an additional opportunity to get one’s hands on a crystallizable solid that can be easily converted back into the original carbonyl. It is not routine for synthetic chemists to think of oximes as substances that can be converted into addition salts, because we more typically think of oximes as being reactive with acids to give Beckmann rearrangement products, but in fact the oxime nitrogen is reasonably basic and can produce acid addition salts with mineral and other strong organic acids. These salts can solidify and provide a means of phase shifting (from liquid or solution) to solid that can provide a basis for purification. In Organic Syntheses Coll. Vo. V pg. 266 2-chloro-cyclooctanone oxime in trichloroethylene solution was converted into an oxime hydrochloride by blowing in hydrogen chloride gas. When the solvent was removed the oil solidified to give oxime hydrochloride in 100% crude yield. It seems likely that all that is required to provide an isolable salt is the addition of a strong acid to an anhydrous medium with the oxime. Another non-solid strong acid that could be considered is the liquid acid, dichloroacetic acid.
Whether it is the oxime or the oxime addition salt that is isolated, whether it is as a crystallized or precipitated solid; obtaining the solid provides the opportunity foradequate purification. Then, either the oxime or oxime addition salt can be converted back to the carbonyl in high yield by a variety of well documented treatments. Successful application of this strategy would be one more demonstration of the concept that it is not always the protocol with the fewest identifiably steps, or the fewest chemical reagents, but the one that is simplest to execute and most rugged that is best suited for scale-up and cost minimization.
kilomentor | 10 June, 2010 18:12
Intuitively it seems sensible that to obtain the highest melting derivative it might be sensible to derivatize with the highest melting derivatizing agent. Although there is no firm basis for such a prediction, there is some enhanced likelihood of this being so. The derivative will share some partial structure with the derivatizing agent and if features of that agent contribute significantly to the high melting property and if those features are preserved in the derivative’s structure, some of the high melting characteristic could be anticipated to be retained.
The highest melting pharmaceutical salt former for reaction with basic substances is orotic acid. Its melting point is around 325 C. Its structure contains multiple hydrogen bond donors and acceptors that most likely contribute the melting character and these hydrogen-bonding sites are retained in the pharmaceutical salt formed.
Despite this the orotic acid salts are rarely used.
The alkali salts of orotic acid are poorly soluble. Orotic acid is also poorly soluble in water. Thus the sodium salts can be precipitated by adding a solution of the N,N-dimethylethanol salt of orotic acid in 80% aqueous ethanol. This might be tried as a method for removing an alkali salt from an alcoholic solution of an organic compound. The alkaline cation as an alkaline salt would be precipitated with orotic acid. The residual orotic acid would also be essentially insoluble. N,N-dimethylethanol would be left in solution but this could be removed by extraction with a neutral organic solvent.
Orotic acid or a salt of orotic acid might be expected to form a complex with triphenyl phosphine oxide which is a good hydrogen bond acceptor. Orotic acid has an imide NH which is a good hydrogen bond donor.
Orotic acid may be a good substance to neutralizes aqueous base since both the alkaline salt and the free acid are essentially insoluble.
The unusual solubility properties of orotic acid and its metal salts make it worth bearing in mind when trying to isolate organic bases particularly because, since it is pharmaceutically acceptable, trace residues are not too critical.
kilomentor | 03 May, 2010 17:02
It is the flavour f the month for Organic Research & Development Process Chemists to modify laboratory synthesis steps, which are environmentally troublesome, with green equivalents. Although this is often a priority pressure in the opposite direction can arise because, to meet deadlines for specified quantities, it is more straight forward to keep the original chemistry. This problem can come up rather frequently when the reaction ‘sinning’ environmentally is a classic such as the chromic acid oxidation.
Chromium(VI) has been a popular oxidant in the laboratory for 70 years. It has been used in many complex reagent combinations. It has appeared in many of the volumes of Fiesers Reagents for Organic Synthesis. Twenty three procedures using a chromium (VI) appear in Coll. Vol. I-IV of Organic Syntheses.
Unfortunately all the variants of this oxidant present problems for a scaled up procedure. Some: (Sarrett and Snatze) can lead to fire when the oxidant is being prepared . Some variants use undesirable solvents (benzene, diethyl ether, HMPT, pyridine, acetic acid and methylene chloride). What all share is that chromium is a heavy metal that cannot be discarded in waste-water. Compounding these difficulties, the work-up of many of these oxidations produce inorganic gases, and large volumes of waste extraction solvents. Traces of residual chromium not removed quickly in the isolation are known to lead to by-products during the solvent concentration steps in f the isolation protocols causing products which ought to be white, to be yellowish.
Chromium (VI) in many variants have trivial names
Jones Reagent H2CrO4/ sulfuric acid/ sodium dichromate/ acetone
Brown-Garg H2CrO4/(ether or benzene)/water
Kiliani Reagent H2CrO4 /H2SO4/ water/ acetic acid
Chromic anhydride CrO3/ water/ acetic acid
Fieser Reagent CrO3/ acetic acid
Sarrett/Ratcliffe Reagent CrO3/ pyridine
Cornforth Reagent CrO3/ pyridine/ water
Thiele Reagent CrO3 /acetic anhydride/ H2SO4
Snatze Reagent CrO3/ DMF
and with no special name
CrO3/ acetic anhydride/ acetic acid
The reagents can be used with co-catalysts like mercuric acetate, ceric ammonium nitrate, manganous nitrate, oxalic acid (F.& F. Vol. 9 pg. 114) and special effects come with controlled amounts of water in the reagent.
kilomentor | 05 April, 2010 19:26
Compounds that contain double bonds often suffer from the presence of isomers in which the double bond is in a different location. These contaminants are usually very difficult to remove. Hexachlorocyclopentadiene is a reagent that reacts readily with a double bond that is not conjugated to any electron withdrawing group. This is precisely the situation that is most likely to be difficult to separate. Hexachlorocyclopentadiene is reactive enough that it adds twice to naphthalene molecules once at each of the two double bonds in an unsubstituted ring. This use in fact is its best known application where it acts as a protecting group for an unsubstituted ring in a naphthalene.
There is also a literature report that ordinary alkenes react at different rates with hexachlorocyclopentadiene. 1-Octene has an activation energy of 20 kcal and 4-methyl-1-cyclohexene has an activation energy of 24 kcal. These exemplify Diels-Alder reaction occuring with unactivated alkenes. Although hexachlorocyclopentadiene is expensive when obtained from Aldrich, it is still probably available somewhere in bulk inexpensively. The disposal of the chemical must be handled carefully since it is an environmental toxin.
The adducts of this reactive dienophile are reported to be split back into their precursors by heating with phosphorus pentoxide and distilling the non-chlorinated adduct. This is an inconvenient, inefficient and damaging procedure that cannot promote widespread use. Using the adducts to remove an unwanted contaminant by reaction however would not have this disadvantage since reaction with an impurity would not need to be reversed. An excess of the hexachlorocyclopentadiene could be used and when the reaction had proceeded to the most beneficial extent, as judged by in process analysis, the excess could be quenched with maleic anhydride followed by mild base extraction of the hydrolyzed adduct as sodium salt. This would leave the hexachlorocyclopentadiene adduct of a more reactive alkene impurity and the desired less reactive product. This should prove a much simpler separation.
The titration of a mixture of reactive alkene, hexachlorocyclopentadiene and adduct with bromine gives the residual amount of the olefin because neither reagent or the adduct react under ice temperature conditions.
kilomentor | 15 March, 2010 05:24
In one year time, I think it is likely that I will be retiring from my present position at ratiopharm. For one thing, I will be 65 years old, and for another, ratiopharm, by that time, will have been sold off to some other organization. Most likely ratiopharm will be part of either Teva, Pfizer or Actavis. Anyway, those, Reuters reports are the front-running contenders. Since the ratiopharm organization did not itself produce any active pharmaceutical ingredients but sourced everything outside, the organization didn’t mind me writing a blog whose purpose was to mentor process development chemists anywhere in the world. I do not suppose these new people will feel the same way.
I am still in good health and cannot imagine doing anything else as enjoyable as chemistry. The action to-day is in the developing world, so I wouldn’t mind getting back into process development /research management / mentoring there with a vigorous team of young scientists determined to efficiently deliver significant projects. My wife, who is Chinese, favors China which clearly moving the fastest. As a Canadian with British roots, I may be inclined somewhat more towards India, where I have visited and appreciated what I saw. Of course there are other potential destinations as well. Any suggestions from out there in cyberspace?
Dr. Clarke SlemonThe Kilomentor
kilomentor | 21 January, 2010 07:48
Kilomentor takes the position that, in the present state of the chemical art, electronic database searching has enabled chemists of ordinary skill to design ingenious reactions schemes by little more than electronically searching for reactions to string together. It is isolation and purification procedures where there is the least technical support to support individual ingenuity. Isolation methods cannot be searched because the search terms are the solution to the problem not the starting point.
Therefore Kilomentor wants to emphasize where inexpensive transition metal complexes such as those with Chromium (III) and Cobalt (III) can simplify the work-up of chemical process steps.
Chromium (III) is the most stable and important oxidation state of the element in general and particularly in the aqueous chemistry. Advanced Inorganic Chemistry 1966 pg. 823 states, “The foremost characteristic of this state is the formation of a large number of relatively kinetically inert complexes. Ligand displacement reactions of Cr (III) complexes are only about 10 times faster than those of Co(III), with half-times in the range of several hours. It is largely because of this chemical inertness that so many complex species can be isolated as solids and that they persist for relatively long periods of time in solution, even under conditions where they are thermodynamically quite unstable.” Note that it is the kinetically inert property of chromium complexes that makes them valuable. What this is saying is that complexes that are not the thermodynamically most stable nevertheless can be isolated. This means that many more compounds are in principle accessible.
Advanced Inorganic Chemistry 1966 pg. 873 says about cobalt chemistry that “The complexes of Cobalt (III) are exceedingly numerous. Because they generally undergo ligand exchange reactions slowly, but not too slowly, they have, from the days of Werner and Jørgensen, been extensively studied and a large fraction of our knowledge of the isomerism, modes of reaction and general properties of octahedral complexes as a class is based upon studies of Co (III) complexes.” What I take this to be saying is that many different complexes of cobalt would also be readily accessible in principle.
Iron also appears to be promising in terms of offering multiple potential complexes. Iron (III) forms a large number of complexes, mostly octahedral ones, and octahedron may be considered its characteristic coordination polyhedron. The affinity of iron (III) for amine ligands is very low. No simple amine complexes exist in aqueous solution; addition of aqueous ammonia only precipitates the hydrous oxide. Chelating amines, for example, EDTA, do form some definite complexes among which is the 7 coordinate [Fe(EDTA)H2O]ion. Also, those amines such as 2,2’-dipyridyl and 1,10-phenanthroline which produce ligand fields strong enough to cause spin-pairing form fairly stable complexes, isolable in crystalline form with large anions such as perchlorate.
Transition metals now have an extensive application as catalysts in organic chemistry. Nickel, palladium and platinum complexes are today extensively used to catalyze reactions for which there is no uncatalyzed equivalent.
An extensive chemistry has also been established centering on the practical question of the recovery and recycling of the noble metal catalysts, mainly palladium and platinum, since these represent expensive inputs into a process.
From the Kilomentor perspective of using of transition metal complexes for isolations the complexes of the wide variety of less expensive chromium, cobalt and iron complexes would seem most promising.
As a first example let us consider Reinicke and Rhodalilate Salts:
The Chromium Salt NH4[Cr(NH3)2 (SCN)4] is red in color. It is soluble in ethanol or hot water and is reported to dependably yield precipitates with primary and secondary amines. The implication of many reference books seems to be that the salt does not form precipitates with t-amines, but this is untrue. According to Cotton & Wilkinson’s Advanced Inorganic Chemistry Comprehensive Text, it can be used, in general, to precipitate large cations, either organic or inorganic. It seems likely however that although, thermodynamically, precipitation of Reinecke salts may not be as selective as has been publicized, fractional precipitation based on rates of precipitation can provide purification as suggested for the closely related Rhodanilate salts (see later for rhodanilate definition).
The Reinecke and the related Rhodanilate salt possibly could be used to precipitate particular amines in the presence of others. One idea is that because the Reinicke salt is soluble in alcohol alone, a useful separation could be done on a substrate, which is sensitive to water. An example of this is that the possible difference in rates of precipitation might be useful in the case of alkylating of an amine where an excess of the starting amine could perhaps be selectively precipitated.
Amines are frequently used as reagents to neutralize acidic co-products of a reaction and thereby drive any equilibrium towards completion in a particular direction. The most frequently used amine in this regard is triethylamine. Some advantages of triethylamine are that even if it is employed in excess any unused base is
i) volatile enough to be removed by vacuum
ii) water soluble enough to be carried away in a water wash
Disadvantages are that it is volatile enough to escape from reactions that require heating and nucleophilic enough to compete in some displacements and deprotonations. Employing the Reinicke or Rhodalinate salts in a work-up of mixtures containing more complex amines may make them recoverable and recyclable and so practical as traps for acidic co-products.
Another possibility is that initial formation of an easily isolable amine salt of a complex anion X could be followed by the switch from the amine salt to the inorganic metal salt via a Rienecke or or Rhodalinate reagent which could precipitate the intermediate amine.
For example, a salt of an amine with a complex anion X might be converted into the salt of a metallic cation MX by adding that M in the form of acetate and precipitating the amine (here R3N) as the Reinicke salt precipitate.
M+ - OAc + R3NH+ X- + NH4[Cr(NH3)2 (SCN)4] going to
R3NH [Cr(NH3)2 (SCN)4] (insoluble) + NH4 OAc + M X
I do not know of any experimental examples of this, however.
Amine Recovery from lithium amide reagents
The lithium salts of many sterically hindered secondary amines, lithium diisopropylamide for example, are used for quantitative deprotonation in chemical synthesis. Because they are sterically hindered the resulting secondary amine co-products do not interfere in subsequent reactions of the carbanions they helped create. These sterically hindered secondary amines may need to be separated from desired product in the reaction work-up and if they are expensive recovered for recycling.
Reinicke Salts, Rhodanilate salts or Trisoxalatochromate salts can potentially be used to precipitate these secondary amines and remove them as filterable solids. Diisopropylamine, dicyclohexyl amine, 2,2,6,6-tetramethylpiperidine, isopropyl-cyclohexylamine, and pentamethylpiperidine need to be examined to see whether they can be quantitatively or semi-quantitatively precipitated.
According to Max Bergmann’s article in J. Biol. Chem.109, 471 (1935) proline and hydroxyproline can be precipitated from gelatine hydrolysates using Reinecke’s salt and the amines liberated by forming a complex with N,N-dimethylaniline or pyridine. This liberation shows that all amines can react.
In this Bergmann article the formation of what he regards as more selective complexing agents can be achieved replacing the ammonia ligands with other amines. Displacing the two ammonia with aniline what is called ammonium rhodanilate is formed About this Bergmann says, “ Rhodanilic acid forms rose-coloured, well crystallized salts with basic nitrogen compounds, and in particular with alkaloids and with amino-acids. Although rhodanilic acid lacks definite specificity, the various rhodanilates differ greatly in their solubilities, crystalline form, and rate of crystallization. It is therefore often possible to separate from mixtures of amines, amino acids , or peptides, single homogeneous products by fractional precipitation with rhodanilic acid. In most cases where several rhodanilates form simultaneously, a separation by fractional crystallization is often possible.”
With regard to the amount of ammonium rhodanilate in the fractional precipitation Bergmann says that “The quantity necessary was determined by examining the precipitate under the microscope in the course of successive additions.” I interpret this to mean that the precipitation was controlled by the kinetics and the fastest precipitating compound came out first followed by other compounds and the precipitate was collected in fractions that were subsequently combined on the basis of their microscopic crystal shape.
In the case of preparing proline rhodanilate, the free amino acid was simply achieved using excess pyridine.
“In order to obtain the free amino acid from proline rhodanilate, advantage was taken of the fact that pyridine rhodanilate is very difficultly soluble in water. It is therefore sufficient to suspend the solid proline rhodanilate in water and to add a little pyridine in order to precipitate almost instantly the entire rhodanilic acid as pyridine salt. On filtration a faintly colored aqueous solution of l-proline is obtained.
The by-product of such a purification is pyridine rhodanilate. It may easily be recycled and reconverted into ammonium rhodanilate with ammonia and so recovered for further use.
Thus ammonium rhodanilate can be used to precipitate a complex amine, the amine rhodaniliate can be freed from the complex with pyridine to precipitate the very poorly soluble pyridine rhodanilate and then the ammonium rhodanilate can be reformed from the pyridine salt by treatment with excess ammonia.
In a finl use, of the Reinicke salt, if mercuric acetate bound to an ion exchange resin is used as a source of mercuric ions in a reaction. Water from such a reaction, that could contain small amount s of mercury ion, can be decontaminated with Reinicke Salt which precipitates the mercury ion.
kilomentor | 28 December, 2009 19:17
Dipolar aprotic solvents such as N-methylpyrollidone, dimethyl formamide, N-methyl formamide, dimethyl acetamide and dimethyl sulfoxide are often drowned out with water and then extracted to isolate organic products. No cheap and convenient method has been worked out to separate these polar organics from the bulk of the water and return them to an anhydrous condition suitable for reuse.
Kilomentor at his present employment does not have access to a chemical laboratory where experimentation could be done so the following idea has not been tested, but on the basis of the physical properties of the chemicals it might be workable.
Diisopropyl ether (DIPE) forms an azeotrope with water that is reported to boil at 62.2 C. This is a heterogeneous azeotrope that according to the Chemical Rubber Handbook splits into a water-poor DIPE upper phase and a water rich lower phase. Addition of DIPE therefore to one of these higher boiling solvents and water, and boiling of the ternary mixture under a Dean-Stark trap with continuous return of the top DIPE phase should gradually separate a lower water rich phase which could be periodically drained away. The high boiling dipolar aprotic solvent that is being dried should theoretically be confined to the still pot at the low azeotropic boiling point. In the real laboratory situation, however, a small amount of dipolar solvent vapour entrained in the reflux stream could be all that is needed to prevent the distillate from separating into two phases in the trap and this would scupper the procedure. It is crucial for a practical process that the DIPE be recycled since the distillate is 97% DIPE and only 3% water. Recycling is essential to be able to remove a large amount of water using only a small amount of DIPE.
Other solvents that boil above 100 C that can potentially be separated from water and dried are: nitromethane, acetic acid, dioxane, ethylene diamine, sulfolane and isoamyl alcohol.
After the water has been completely removed continued distillation will separate the DIPE. Even if small amounts of DIPE might remain they are usually unreactive. If particular importance they are inert towards organometallic reagents.For safety remember that DIPE needs to be worked with under inert gas to prevent the accumulation of explosive peroxides. The solvent very readily forms peroxides.
kilomentor | 15 December, 2009 18:55
Although mixtures of carboxylic acids or mixtures of amines are each fit for separations based on extractions at controlled pHs, most functional groups are not so easy to isolate from each other. Kilomentor thinks a good deal about how to simplify separating molecules with the same functional group in slightly different structural environments. Although it has not been demonstrated in the literature yet, two different molecules each with a nitrile functionality but differing in the steric environments around them can likely be separated by selective reaction followed by a simple acid-base extraction.
Nitriles are known to react smoothly with azide in a 3+2 cycloaddition to give 1,2,3,4 tetrazoles. This is a ‘click chemistry’ reaction . Cycloadditions are typically quite sensitive to steric environment. Thus, although these reactions are generally fast, it is likely that conditions can easily optimized to get good selectivity between cyanide groups in different molecules using an insufficient amount of azide. The result will leave a nitrile in one substrate untouched and the nitrile in the other substrate converted essentially completely to tetrazole. The beauty of this is that these tetrazoles have the acidity of carboxylic acids and can be extracted into water with alkali. Thus the tetrazole derivative removes the reactive nitrile substrate from the organic phase leaving the unreactive nitrile substrate clean for a simple recovery.
Two references that I could loctate concerning the kinetics of the reaction of nitriles with azide are: Khimiya Geterotsiklicheskikh Soedinenii (1992) (9) 1214-17, which is in Russian [C.A. 11; 8945a, 8948a]; and Inorganica Chimica Acta (1985). 102(2), 157-62 that does the condensation with the nitrile coordinated in a Co(III)complex.
The synthesis of tetrazoles from nitriles and azide has been studied intensively because of its relevance to the preparation of the sartan family of drugs. Relevant patents are US5744612, US6040454, WO2005014602, WO2007054965 and CN1718574