Diisopropyl Ether (DIPE) Solvent Can be Safely Used in Industry

 

Diisopropyl ether also trivially called isopropyl ether (analogous with ethyl ether) is an important anti-knock additive for gasoline. It is an important coproduct in the preparation of isopropanol by the hydration of propylene. As a result, it is reasonably priced.

In the Research Laboratory

In the laboratory setting, diisopropyl ether must be treated with great caution because, more than almost any prospective solvent, it readily forms explosive peroxides when exposed to atmospheric oxygen. Bottles of old solvent that are left in a laboratory or storeroom slowly evaporate through inadequately seals and the peroxides concentrate. Sometimes the peroxides even crystallize. Such residues or concentrates are extremely dangerous. If one of these concentrates is discovered, it must be handled by trained personnel with special safety equipment.

The consequence is this useful solvent does not get incorporated into scaled-up processes. This is unfortunate because at scale the dangers of the solvent are drastically mitigated. 

The Difference In the Plant At-Scale

In the plant, all process operations are executed under an inert atmosphere. This is part of standard operating procedures (SOPs). Vessels are closed. Transfers are made by piping liquids, solutions, or slurries. There is no pouring through the air! The possibility of exposure to oxygen in the air is remote. 

In addition, in the laboratory the formation of peroxides in diisopropyl ether is made more likely because exposure to light is increased and light can catalyze peroxide formation. In the plant light is blocked by working in drums, closed metal reactors, piping, and pumps. Reactions and processing involving DIPE occur either in subdued lighting or in the dark. There is no photocatalysis possible.

Finally, at scale, batch sheets require that all chemical inputs be tested to be sure they meet their specifications and one of the requirements for DIPE use is that it passes its requirement with regard to peroxide impurities. So unlike the situation in a laboratory where an old bottle of solvent might be used in an experiment, all the inputs for working in the kilo lab or pilot plant are rigorously tested. Furthermore, the capacity for the analytical testing laboratory to do retesting for peroxides during processing is also available.

So as we can show, unlike other materials, the higher danger point using diisopropyl ether occurs in the research laboratory during process research and development. Yes- special precautions need to be implemented -in the laboratory!

These laboratory dangers can be stymied a number of ways:

• Store in the dark 
• Keep bottle sealed
• Stabilize with butylated hydroxytoluene (BHT) or NaOH  
• Remove peroxides by acidic iron(II) sulfate wash
• Pass through alumina (does not destroy the peroxides; merely traps them)
• A more drastic method that also removes water/oxygen is to distill from sodium/benzophenone

But Why Bother Taking Any Risk?

DIPE readily separates from water-free sulfolane.

 

DIPE won’t separate from totally anhydrous DMF, but adding  a little water gives two layers.

 

DIPE does give phase separation from anhydrous DMSO. So you can do a reaction in dry DMSO and repeatedly extract the product into DIPE. 

A biphasic/phase transfer catalyzed reaction can be conducted using the DIPE/DMSO system. 

Diisopropyl ether (DIPE) is a clear liquid that is immiscible with water. It smells like decomposing green tea. MP: -60 °C; BP: 69 °C; Density: 0.725 g/mL . It has a reputation as a go-to solvent for recrystallizations that have failed with other solvents.

In addition to what has been established for sure, DIPE is promising in other ways. Reactions performed in dipolar aprotic solvents such as N-methylpyrollidone, dimethylformamide, N-methylformamide, dimethylacetamide and dimethylsulfoxide 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 the dipolar aprotic to an anhydrous condition suitable for reuse.

On the basis of the physical properties of the chemicals, the following might be workable but KiloMentor has seen no experiment to substantiate it. 

Diisopropyl ether (DIPE) forms an azeotrope with water that is reported to boil at 62.2 C. This is a heteroazeotrope.  The designation means that this azeotrope’s vapor is in equilibrium with two immiscible liquid phases. According to the Chemical Rubber Handbook, DIPE and water form an azeotrope that on condensation splits into a water-poor DIPE-rich upper phase and a water-rich lower phase. Thus, addition of DIPE to a mixture of 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 could be expected to gradually separate a lower water-rich phase which could be periodically drained away. The high boiling solvent that is being dried would theoretically be retained throughout in the still pot.

In the real laboratory situation, however, a small amount of the high boiling solvent as vapor entrained in the reflux stream that one is trying to free from water could be all that is needed to prevent the distillate from separating into two phases in the trap and this would scupper the procedure so this concept would need to be tested. Nevertheless, if it works and your facility has unused distillation capacity, solvent recovery could be profitably practiced.

 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. 

Before recovering the DIPE by distillation in the plant it should be tested for peroxides and washed with aq. acidic iron (II) sulfate if the peroxide test is positive.

Other solvents that boil above 100 C that can potentially be separated from water and dried using DIPE are nitromethane, acetic acid, dioxane, ethylenediamine, sulfolane, and isoamyl alcohol.

After the water has been completely removed continued distillation will drive over the DIPE itself. Even if small amounts of DIPE remained in a recovered dipolar aprotic solvent it is usually unreactive. Of particular importance… it is inert towards organometallic reagents.

Extractive Crystallization-The Use of Methyl or Ethyl Salicylate to Crystallize Solutes Insoluble in both Hydrocarbon Liquids and Water

The following is a research idea. As far as KiloMentor is aware It has not been demonstrated. Before any trials, an up-to-date literature search is recommended.

A very large number of organic compounds are essentially insoluble in both pure hydrocarbon solvents and in water. As such they would be expected to also be insoluble in a two-phase mixture of water and hydrocarbon. Solutes that belong to this large group would be candidates for a crystallization/recrystallization procedure that, as far as I know, has not been tried to date.

Methyl and ethyl salicylates are very low melting solids and high boiling liquids: methyl salicylate mp -8.6 C; Bp 220-224 C; ethyl salicylate mp 1 C; Bp 232.5 C.

They share another common property. When stirred with aqueous alkali they are hydrolyzed to 2-hydroxybenzoate salts. What is less commonly recognized is that the presence of a separate hydrocarbon phase would not effectively inhibit this hydrolysis. The reason: the free phenolic substituent essentially drags the ester into the aqueous phase where the base attacks the ester functionality irreversibly after which it no longer has any affinity for the hydrocarbon layer. 

Unlike hydrocarbons or water, these compounds will be good solvents for a wide variety of other organic materials. They can interact using Van der Waal dispersion forces, dipole-dipole interactions, and hydrogen bonding using both the phenolic hydrogen bond donor and the ester carbonyl hydrogen bond acceptor. Furthermore, these compounds are not particularly expensive and are readily available at an industrial scale. They have been demonstrated to be safe. Methyl and ethyl salicylates are flavoring and perfume chemicals. 

 
Suppose we choose to dissolve a solute of interest in a combination of a highly apolar poor solvent, like the hydrocarbon heptane for example, and as solubilizing agent methyl or ethyl salicylate. Such a combination will have the property of having a boiling point at least as high as the hydrocarbon used but will have the enhanced dissolving power provided by the additive. When the mixture is all a single solution it is cooled to ambient temperature and an immiscible aqueous solution of base is added. Even with only very weak stirring, hydrolysis in the two-phase medium will result in the salicylate being taken into the aqueous phase. Now with its solubilizer degraded neither hydrocarbon nor aqueous phases will appreciably solubilize the substrate so it is likely to slowly crystallize out.

The mechanism by which the methyl or ethyl salicylate gets hydrolyzed and retained in the aqueous phase as carboxylate salt is the same extractive hydrolysis that was featured in another KiloMentor blog.

Wide Range of Acceptable Solutes

Even solutes containing functional groups sensitive to aqueous alkali can be expected to safely undergo this treatment. Molecules without active hydrogens (such as phenols, carboxylic acids) will not be extracted out of the hydrocarbon phase and so will be protected from significant alkaline hydrolysis.

Hypothetical Examples

An example of a preparation that might be improved using this methodology can be found in Organic Synthesis Col. Vol. 1 pg. 60 Anthrone Synthesis. The anthrone is finally crystallized from 3:1 benzene and petroleum ether. It is reported that about 12 g of the 3:1 mixture is required for each gram of anthrone.  The yield percent recovery is 62/82.5. An effort is made in this preparation to recycle mother liquors and this reuses about 2/3 of the liquid. The anthrone is much more soluble in benzene than in the petroleum ether antisolvent.  It would be interesting to see how the purification would proceed with heptanes as antisolvent and one of these hydroxyl benzoate esters as the solvent with dissolution at the reflux temperature of heptanes.

Another opportunity to use this technology seems to be presented by the bromination of anthracene to 9, 10-dibromo anthracene. A process is described in Organic Synthesis Col. Vol. 1 pg. 207. This procedure uses carbon tetrachloride as solvent. This would be unacceptable in scale-up to-day since carbon tetrachloride is a recognized carcinogen.  It might work to brominates anthracene with bromine in heptanes. The dibrominated product is likely to be poorly soluble in heptanes and anthracene itself would only be somewhat better. In the heated reaction mixture, the anthracene would probably dissolve enough to allow the reaction to proceed. At the end, the crude dibromoanthracene would be precipitating. To recrystallize and recover the solvent one of our salicylates could be added with heating to get a solution; the combination then could be filtered hot; dilute aqueous alkaline added to hydrolyze and extract the hydroxyl benzoate ester. Since the 9, 10-dibromoanthracene would then be insoluble both in the aqueous and the hydrocarbon phases it should crystallize.

Other Possible applications

Other compounds from Organic Synthesis that could benefit from purification from a two-phase mixture of aq. alkali and high boiling hydrocarbon solvent: desoxybenzoin pg. 156: desyl chloride pg. 159; dibenzalacetone pg. 167; ethyl 2,3-dibromo-3-phenylpropionate pg. 270; m-nitroacetophenone pg. 434; Organic Synthesis Col. Vol. III acenaphthenequinone pg. 1; acenaphthenol-7 pg. 3.   

Phase-Transfer Chemistry

Introduction

A shortcoming of the KiloMentor blog is that its author retired in 2011. Thus, although a fair account of developments up until that time can be realistically expected anyone wanting a completely up to the present assessment needs to supplement my examination.

Simple reaction conditions and inexpensive reagents become more consequential the larger the scale at which a chemical reaction is practiced. Consequently, phase-transfer methods have become major transformers of process chemistry.

The three initial, principal reviews on Phase-transfer Chemistry were:

 G.W. Gokel and W.P. Weber, J. Chem.. Ed., 35, 350 (19780);

W.E. Keller, Compendium of phase-transfer Reactions and Related Synthetic Methods, Fluka, 1979;

C. M. Starks and C. Liotta, Phase-transfer Catalysis, Academic Press, New York, 1978.

Also, volumes 8, 9,10, 11, 12, 13, 15, and 18 of Fieser & Fieser’s Reagents for Organic Synthesis have entries under the specific heading of Phase-transfer catalysis. In these, select significant examples of applications are listed. KiloMentor has not examined F & F volumes beyond number 18.

The Phase-transfer Concept and its Range

Phase-transfer reactions, as the name conveys, involve the interplay of two phases during the course of a reaction. These two phases can be two partly immiscible liquids or one liquid phase and a solid.

The methodology applies the finding that ion pairs made up of a large organic cation, when present in organic solvents, exhibit reduced solvation of the anion partner. Consequently, the more ‘naked’ anion is more reactive. Particularly practically, when quartenary ammonium cations are paired with hydroxide to extract it from a concentrated aqueous alkaline solution into an organic solvent, the hydroxide is drawn into the organic solvent with many fewer solvating water molecules and so is called a ‘naked' hydroxide ion. This so-called ‘naked’ hydroxide has a much higher apparent base strength and as such deprotonates substances with pKa less than 37. 

According to Halpern et al. [Hydroxide Ion Initiated Reactions Under Phase-transfer Catalysis Conditions: Mechanism and Implications, Angew. Chem. Int. Ed. Engl. 25 (1986) 960-970 pg, 963 ]  “The empirical upper limit found for substrate acidity in anionic reactions occurring via the interfacial mechanism lies at about pKa=23.”


The most frequently used and most promising catalysts are: Methyl trialkyl(C₈-C₁₀)ammonium chloride, whose trivial names are Adogen 464 and Aliquat 336; Tetrabutylammonium hydrogen sulfate; Benzyltriethylammonium chloride; or Cetyltrimethylammonium bromide. Other interesting catalysts are Trident-1, Polysorbate 80, Polyethylene glycol 400, Cyclophosphazenic polypodands (made from Brig 30; JOC 59, 5059 (1994)).  

Anything that increases the effective cation radius and so moves it away from participating in a tight-ion pair to a loose or solvent separated ion pair is likely to increase the nucleophilicity of any associated anion. That is why, even in inorganic chemistry, as the size of the cation increases, as in Li<Na<K<Rb<Cs, the reactivity of any associated anion increases. It is for the same reason that crown ethers and crepitates are sometimes effective.

The earliest phase-transfer applications used an aqueous solution containing dissolved inorganic reactant as one phase; the reaction substrate was in a  second immiscible organic solvent phase. The function of the catalyst was to provide a mechanism for the water-soluble reagent and the organic-soluble substrate to meet with sufficient frequency to provide a practical reaction rate.

Later, the water phase was found to be sometimes unnecessary and sometimes even detrimental. That is, the phase-transfer catalyst could sometimes make possible reaction between a solid phase of essentially insoluble neat reagent and a substrate dissolved in an organic phase. It has been hypothesized that this most often was successful when undetected small amounts of water were adsorbed onto the bulk solid. Where it succeeded this provided an even simpler process. For example, nominally anhydrous potassium or sodium carbonate could be used as the base for the generation of carbanions in a solid-liquid two-phase system using tetraalkylammonium salts or crown ethers as catalysts. Probably the carbanions are generated on the surface of the carbonate and migrate as ion pairs into the organic medium.  [M. Fedorynski, K. Wojciechowski. Z. Matacz, and M. Makosza J. Org. 43, 4682 (1978).] in Fieser & Fieser Vol. 8 pg. 356-361.

Crown Ethers

“Crown ethers have commonly been used as catalysts for reactions between a solid-liquid interface, and quaternary ammonium or phosphonium salts have been used only as catalysts for reactions in two-phase, liquid-liquid reactions.” But crown ethers are expensive, toxic, and very often difficult to recover.  “Several laboratories have reported that less expensive catalysts can satisfactorily replace crown ethers for solid-liquid reactions. Thus, dichlorocarbene can be generated from chloroform and solid sodium hydroxide under catalysis with benzyltriethylammonium chloride in yields comparable to those of the classical Makosza method. { s. Julia and A. Ginebrada, Synthesis, 682 (1977)}. See F & F Vol. 8 pg. 390.

Cheaper than crown ethers is the Trident-1 ligand and even simple polyethylene glycol monoethers. Both act in what can be visualized as wrapping an alkali cation in its tentacles and folds which increase the effective size of the cation and thereby make the associated anion more reactive.

The attraction of phase-transfer catalysis for work at scale is the general simplification of conditions and the use of inexpensive quartenary ammonium (hereafter called quat) catalysts that allow the employment of aqueous sodium hydroxide as a base in organic synthesis instead of the classical, more sensitive, more dangerous, and more expensive alkali metal alkoxides, amides, and hydrides.

The favorable price and availability of the quat ions,  phosphonium ions, the Trident-1 ligand, and cyclophosphazenic polypodands render these catalysts of choice on-scale.

Detailed Mechanisms

There are two competing mechanisms that can be contemplated for the sub-set of phase-transfer catalyzed reactions that require a deprotonation as a step: the interfacial mechanism and the extraction mechanism. The interfacial mechanism does not require deprotonation of the weakly acidic substrate by a quat hydroxide in the bulk organic phase. Rather, the alkali hydroxide deprotonates the substrate at the interface between the two layers. Then the quat exchanges with the alkali cation and carries the reactive deprotonated organic ion pair into the bulk of the organic solvent where it reacts. The presence of the quat cation allows a higher concentration of deprotonated substrate anion in the bulk organic solvent.

 In contrast, in the extraction mechanism, a quat with many large hydrophobic groups forms an ion pair with hydroxide in the bulk aqueous phase and carries it into the bulk organic medium where both deprotonation of substrate and subsequent reaction occurs.

For synthetic organic chemists who are interested in quickly finding good reaction conditions, the importance of the existence of two mechanisms is that each predicts a different preferred structure for the most effective quat catalyst. When one cannot infer the mechanism, each of the two different types of catalyst structure ought to be investigated.

Ion-Pair Extraction (F & F Vol. 11)

Using a full equimolar amount of either a large hydrophobic cation or hydrophobic anion can lead to complete extraction of the complimentary ion into an appropriately selected non-polar organic solvent. Organic substances that have higher carbon-acid acidities such as those with two or more electron-withdrawing functional groups attached to the same carbon can, for example, be extracted as stoichiometric stable anions. Sometimes this is the preferred method for their reaction- even better than phase-transfer catalyzed anion formation and reaction.

Ion pair extraction is more frequently used as a means of separating acids or bases that are very similar in acidity/basicity but markedly different in hydrophobicity. 

Catalyzed Enantioselective Alkylation (F & F Vol. 12 pg.379-380)

Phase-transfer Quat catalysts that are themselves chiral can induce chirality in the alkylation products to promote catalytically. Examples of this are referenced in Fieser & Fieser vol. 12 pg. 379-380 and Vol. 13 pg. 239. I would expect there have been many more examples in the period since 2010.

Catalyst Removal after Reaction Completion

The removal of the catalyst from the organic phase after reaction completion is a frequent problem with strongly hydrophobic phase-transfer quats. If not effectively removed it will often contaminate an organic-soluble product after it is isolated. KiloMentor speculates that the use of a quat that contains two quarternary groups linked together might provide a means to remove that quat from the final product since it is known that dications can be precipitated as insoluble salts with pamoate dianion. Note: this has not been demonstrated.

Quats are also known whose structure contains a degradable link between the hydrophilic head and the hydrophobic tail so the catalyst can be destroyed.

The Role of Water

The higher the salt concentration in any aqueous layer used in a phase-transfer reaction, the less water is available to be extracted as part of the solvation sphere of the anion that is transported by the catalyst into the organic layer. The less the solvation sphere the more reactive the anion whether as a nucleophile or as a base. the reason: high salt concentrations reduce water activity.

Phase-transfer Catalyst Poisoning

A reaction that is being catalyzed by a phase-transfer catalyst often stops before completion. The most common cause is that a more hydrophobic anion has accumulated in the reaction medium and it is pairing with the quat holding it permanently in the organic phase so that it cannot cycle back and forth between aqueous and organic as is required for catalysis. Nucleophilic substitution reactions are particularly prone to this difficulty because the good classical leaving groups are large soft Lewis bases like bromide or iodide that can pair with and immobilize a quat.

Impinging Jet Micromixing to Solve the Problem of Small Crystal Size without Milling.

CA2044706, Crystallization Method to Improve Crystal Structure and Size expired June 14^th 2011 in Canada and the family member US5314506 expired May 24^th 2011 in the United States. The invention addresses the general problem, how to obtain a reproducible micronization of a pharmaceutical compound without milling.

Crystallization is a process step that has for a very long time has only been scaled up empirically.  

One standard crystallization procedure contacts a supersaturated solution of the substrate with an appropriate anti-solvent in a stirred vessel. The anti-solvent initiates primary nucleation as it mixes into the supersaturated solution of active and these seeds then grow. The process can be modified by using preformed seed crystals and/or further aging of the solid, once formed, which digests the crystals to change their initial sizes and/or polymorphic forms. In order to get the smaller crystals, preferred for their greater bioavailability, the saturated solution needs to be added into the anti-solvent in order to get very rapid formation of many tiny seeds. Using this reverse addition methodology a concentration gradient cannot be avoided in a large reactor because the introduction of feed solution into the anti-solvent in the stirred vessel does not afford a thorough mixing of the two fluids prior to the initiation of crystallization.  The presence of these concentration gradients and heterogeneous fluid environment both interferes with optimal crystal structure creation and allows greater entrainment of impurities. On scale even the fastest bulk mixing cannot smooth out the microenvironments in which the seeds form. Furthermore, in a large bulk reactor the number of seeds present at the beginning of the nucleation process is very different from the seeds present in the bulk when the last of the supersaturated solution enters the tank. On scale stirring cannot handle the micromixing requirement.

Another standard crystallization procedure cools a solution of the desired product in order to bring the solution to its supersaturation point, but cooling in batch processing is a slow process that becomes even slower as the batch size increases. Although the solvent gradient is solved, there is a thermal gradient and in any case the crystals are larger with the slower process. The characterisrtics of size, purity, and stability are difficult to control.

The technology taught in CA2044706 pumps both solution and anti-solvent as two impinging jets of fluid that because of their small volumes and high velocities create almost instantly a region of high intensity micromixing where they collide. Once the fast crystallization has occurred, the mixture of solution and anti-solvent can be accumulated and filtered when all the material has been processed or it can be collected after any other appropriate time.

This impinging jet technology removes the problem of scale. Larger scale just translates into a longer period pumping the same streams together. The heterogeneous slurry in which the seed crystals form becomes a function of the pumping rates, the concentration of solute in solvent and anti-solvent,  and the radii of the columnar jets of colliding fluids.  All the parameters come within engineering control.  Because of this, the surface area, crystallinity, stability and purity can be optimized. Because a milled quality material is available directly, a step is saved and the noise, dust, yield loss, equipment cost and worker exposure hazard of milling are by-passed.

The entry of this manufacturing technology into the public domain in 2011 was a significant development. 

CA2349136 is still more interesting. This is the same inventive idea but incorporates a reaction involving elements from each of the two impinging liquid jets. For example this could be used to form a salt from the free base form of a pharmaceutical in one stream and a solution of an appropriate acid in the second stream. Thus we can intuit the formation and instantaneous crystallization of a desired salt and its crystallization into micro crystals.  The US equivalent US6558435B2 expired May 6^th 2003. The Canadian attempt became a dead application August 15^th 2007.

Dichloroacetic Acid: A Solvent with Unique Dissolving Power

Dichloroacetic acid is a liquid at ambient temperature.  The melting point of the higher melting of its two forms is +9.7 ºC. The bp is 193-194 ºC, Its density is 1.563. Neat acid is 12.12 molar. In the presence of an immiscible liquid phase ( such as heptane, cyclohexane etc.) it will be the lower phase in the reactor. It is miscible with water and organic solvents (ethanol, diethyl ether, methylene chloride) but only slightly miscible with carbon tetrachloride. Although I can find no data, I would be surprised if it is miscible with trichloroethylene or tetrachloroethylene. Dichloroacetic acid has a  pKa of 1.48, meaning that one-half of the molecules are ionized at a pH of 1.48. The strongest acid that can be present in this solvent is protonated dichloroacetic acid. The strongest base that can be present is dichloroacetate. 

 Anhydrous Acidification without Excess Local Low pH

Because it is a fluid it could be added into refluxing organic solvent, using a continuous dilution head, so that a local excess of strong acid will never be experienced and the mixture will remain completely anhydrous.

High Dissolving Power

The potential dissolution power of the acid is impressive. An 80:20 v/v dichloroacetic acid/methylene chloride mixture can dissolve many otherwise poorly soluble materials including polymers.

Dichloroacetic acid may be an alternative for reactions that are presently considered for neat sulphuric acid, hydroiodic acid, hydroxymethanesulfonic acid, methanesulfonic acid, polyphosphoric acid, trifluoromethanesulfonic acid, or trifluoroacetic acid. Dichloroacetic acid is much cheaper than trifluoroacetic acid although for some applications the greater volatility of the latter will still make it more advantageous. 

Dichloroacetates are pharmaceutically acceptable salts. Dichloroacetic acid, therefore, is acceptable at low levels as an impurity in pharmaceutical products and so is attractive as a reagent in drug synthesis. 

The Safety & Toxicology data are:

Acid: LD₅₀ rat oral, 2.82 g/kg; rabbit skin, 510 μl/kg; TDL_0; rat oral, 3.195 g/kg/90 days; mouse oral, 7.1g/kg/10 weeks.

Sodium salt: LD₅₀: rat oral 5.281 g/kg; mouse, 4.845 g/kg.; i.p.,; TDL_o rat oral, 2.45 g/kg./7 weeks, 30.425 g/kg/12 weeks; dog oral, 4.55 g/kg/13 weeks.

Because dichloroacetic acid is a liquid, during the crystallization of dichloroacetate salts, it can be present in excess to contribute to their insolubility by the common ion effect.

The many distinctive properties of this acid should lead to a wider application.

Crystallization and Recrystallization from Polar Water-Miscible Organic Solvents

Dissolution for Recrystallization by Adding Small Amounts of Water to Increase Solubility in a Water-Miscible Organic Solvent

It is well known that small amounts of water have a marked influence on the solubility of many solutes when mixed into less polar organic liquids. An example is the difference in dissolving power between well dried acetonitrile and acetonitrile that has picked up just the moisture it gets standing near a steam bath for a few minutes. Years ago, Dr. Renuka Misra, an extraordinary natural products chemist then working at the University of Toronto, demonstrated this to me when I was having trouble doing a particular recrystallization.

 Removing Water from a Solvent Mixture with a Miscible Polar Solvent to Causing Crystallization by Distilling an Azeotropic Composition

Another way to access the solubilizing power of water is with water azeotropes. Co-solvents that form azeotropic mixtures with water can be used to recrystallize polar materials by first dissolving them in the co-solvent assisted with some water and then distilling the azeotropic composition to remove the water as its azeotrope leaving the solute in the less polar, pure organic solvent from which it can crystallize or precipitate, often in excellent yield.

Steam Distillation- A Guide to where it might be used and where it should not be used in Kilo Scale Process Development

Perspective

The KiloMentor blog has set as a goal to provide free chemical process development information to anyone with web access, anywhere in the world.

The KiloMentor slant on organic synthesis is that excellence in designing separations and purifications that work at scale is really what characterizes the ingenious process chemist. Why is this? There are electronic databases for searching structures and substructures, and for searching reactions, but the process chemist must depend on his/her own understanding and imagination when it comes to designing rugged elegant isolations. This is particularly important because it is the separation, not the reaction that takes most of the processing time.

Once a More Respected Method

Even technologies that in most situations have compelling disadvantages have their use in special circumstances. For example, a first edition laboratory manual, Laboratory Technique in Organic Chemistry, written by Avery Adrian Morton, McGraw-Hill Book Company, Inc. 1938 has an entire chapter devoted to steam distillation. That suggests that methods, which were useful when conditions in chemical sciences were more rudimentary, have a power and ruggedness that might usefully be rejuvenated. Thinking about the reasons that steam distillation is not favored, particularly at scale, finds part of the problem with engineering.

Engineering Disadvantages

First, the large reactor would have to be fitted with a large steam line for superheated steam to deliver the volumes of live steam needed for a high distillation rate. 

Second, modern batch-processing condensers are designed for efficient condensation with very small gaps between the condensing plates to recover even low boiling solvents like methylene chloride. The high distillation rates of water and volatile organics (not to mention solids) that could pass over in a steam distillation would quickly flood such condensers and cause a large pressure drop.

Third, if the distillate, now purified turned out to be a solid, which it often does, the condenser would plug. In the laboratory, we can use a different configuration of condensers. With laboratory steam distillation it is normal to have two different kinds of condensers in series. The first condenser has plenty of space in the vapor path where solid can gather before it is swept into the receiver. The second condenser often follows the receiver and traps out efficiently the remaining steam. Diagrams of laboratory set-ups for steam distillation of both liquids and solids can be found online or by consulting the index of the popular chemical synthesis reference text, Fieser & Fieser, Vol. 1 or Organic Synthesis.

Fourth, in a steam distillation set-up, supplemental heating is normally provided to the still pot to prevent condensed steam from accumulating, making heat transfer increasingly difficult. To this must be added the corresponding problem of the tremendous cooling burden on the condensers. The standard multi-purpose reactor and condenser train are not well suited for steam distillation.

Chemical Processing Difficulties

There are chemical processing disadvantages as well. One must deal with a very large volume of condensate containing relatively little product. In steam distillation, the volume at the point of maximum volume limits the number of kilograms per reactor/liter that can be processed. In the early steps of a multi-step process, this will probably constitute a bottleneck because early steps must be repeated a greater number of times even in the best case of batch processing. A steam distillation in one of the early steps of a route almost certainly would seriously limit the throughput.

On the other hand, in the later steps, the need for throughput is much less. In fact, because the product is by now so very expensive relative to the other reaction inputs, a company may not even want to commit a large kilogram charge into a single batch, and so the high point of maximum volume in steam distillation isolation may be of little importance. Note, however, that if the sequence is many steps the accumulating molecular weight of the product may have made the product non-volatile. Steam distillation however could still be used to remove a high boiling reaction solvent. 

How does Steam Distillation Compare with Fractional Distillation?

In regular fractional distillation, the fractionation occurs because a column mimics a series of simple distillations in which the distillate from the nth distillation becomes the pot charge for the n+1 th distillation. Since the distillate is always richer in the most volatile component, if sufficient mimics of a simple distillation ( a theoretical plate) are combined in series the more volatile component is eventually obtained pure. If a fractional distillation column is heated too strongly we say the column floods and separation is lost because there is no longer a vapor phase in equilibrium with a liquid phase. Steam distillation is just co-distillation with water under flooding conditions, where there is insufficient vaporization to balance the condensation.

The components come over in the ratio of their vapor pressures as they would in one single simple distillation.

The Benefit of Steam Distillation

For compounds that are too large and too high boiling for simple distillation and that either degrade or are at risk of degrading at their own distillation boiling point even under high vacuum, co-distillation with a large quantity of lower boiling fluid is the only way to vaporize and then recondense them. Steam distillation is just a special case of co-distillation where the cheap low-boiling fluid is water. The other physical requirement for successful steam distillation is that the compound to be distilled must be at least poorly soluble and preferably essentially insoluble in cold water. This preference arises from the need to recover the volatile substance from a massive amount of water co-distillate. Fortunately, most organic target products are poorly water-soluble. It should be noted that all that is actually mandatory is that the product be practically extractable from a large amount of water. 

Another traditional use of steam distillation is to remove an otherwise troublesome high-boiling solvent from a reaction mixture so the reaction products can be taken up cleanly in a more manageable, lower-boiling solvent for further processing, most often recrystallization. For example, both nitrobenzene and 1,1,2,2-tetrachloroethane are useful Friedel-Crafts solvents but are infrequently used for crystallization. They are high boiling so both are routinely removed by steam distillation. Other solvents that can be removed by steam distillation are benzyl alcohol, dimethylformamide, methylene bromide, 1,1,2,2-tetrachloroethane, cumene, anisole, cyclohexanone, bromobenzene, collidine, p-cymene, 1,2-dichlorobenzene, aniline, iodobenzene, o-cresol, benzonitrile, nitrobenzene, and quinoline. Together with all the solvents boiling above 100 C that have azeotropes boiling below 100 C, the number of practically useful solvents is increased significantly. This is important because changing solvent is the most frequently effective method of improving overall reaction selectivity.

In some previous blogs, KiloMentor discussed methods to make solvent switches on scale. The transition from a high-boiling, water-immiscible solvent to a lower boiling, water-immiscible can quite generally be cleanly achieved by distilling the high boiling organic with steam and then extracting the non-volatile mixture of product into the lower boiling, water-immiscible organic. This has the advantage compared to azeotropic distillation that the two organics are never mixed together at any point, so recovery and recycling both are possible. That is, there are no intermediate fractions of mixed organic solvents. In this way a solvent switch that in the laboratory, can be done by evaporation of the first solvent solution to dryness can be replaced by (i) concentrating the first solution as much as possible using regular distillation (ii) a short steam distillation to remove the final amount of this first solvent, then (iii) addition of the water-immiscible second solvent to the steam distillation pot residue and (iv) liquid-liquid extraction combining the organic extracts and finally (v) drying the second solvent solution. Such a method could, for example, replace chlorobenzene with methylene chloride or xylene with pentane. The only limitation is that none of the solutes should decompose, polymerize, oxidize, or hydrolyze.

Other Steam Distillation Applications

Another situation where steam distillation can overcome a difficulty arises with a reaction that upon quenching produces a gel that can neither be filtered nor submitted to extraction. Such a difficulty can arise in Friedel-Crafts reactions when aluminum chloride hydrolyzes to silica gel or in lithium aluminum hydride reactions. Steam distillation can get rid of the organic solvent that is gelling with the inorganic material.

Steam distillations are not fractional distillations. Substances that are volatile and form a phase of their own distill over in proportion to their vapor pressures. The only separation is between substances with detectable vapor pressure and substances with no detectable vapor pressure. But steam distillation can be combined with reactive distillation. If two volatile constituents are treated with just sufficient reagent to interact with one of them preferentially and that interaction makes the interacting component non-volatile, the remaining free component can be steam distilled out of the reaction mixture.

KiloMentor Stresses the Importance of the Integrity of the Reactor at Scale

Laboratory equipment costs just a minuscule fraction of that of process equipment. For that reason, scientists can perform a reaction that requires strong aqueous alkali in a glass round-bottomed flask even though one knows that at the end of the reaction the flask will be opaque and etched by the dissolution of a portion of the glass itself. One the other hand, precautions must be taken that a large scale reactor, that is expected to have a long useful life should nor e partially dissolved or pitted or weakened by any reactor contents. A process development chemist must never put a large scale reactor at risk. Consideration should be paid early on that reaction conditions are not incompatible with the materials of construction. Engineers are particularly knowledgable in this area and can provide an early warning that particular conditions must be vetted. This is normally done in the laboratory by placing weighed tiles of reactor surface material into the laboratory reactor throughout the process step of concern and at its conclusion, these tiles are fished out and carefully reweighed. Any experimentally significant difference between before and after weighings is suggestive that teethe reaction conditions are eroding the reactor surface material, 

At the same time, the experiment will detect any unexpected effect of the reactor’s material on the course of the process's reaction.

Loss of the surface of the reactor can also be caused by abrasion. The surface is simply rubbed off and probably remains as fine insoluble particles inside the reactor. Very little can be done about this except togged away from the abrasive reagent. Sometimes this problem can be solved by packing the abrasive agent tightly into a special column-shaped reactor tube and rapidly circulating the reaction mixture solution through the column past the insoluble abrasive agent.

Loss of the reactor surface may simply be caused by excessive pH and this cane controlled by an adjustment in the reactor material itself.

Another cause is the use of or the creation of a very strong chelating agent which simply rips metal ions out of the reactor surface. I have encountered such a situation. I was able to overcome the corrosion simply by adding a stoichiometric quantity of an inorganic iron salt into the reactor with the rest of the reagents. As the chelator formed it complexed the iron cations and left the reactor alone!

Identifying Chemical Process Stopping Points for Working in the Kilolab or Pilot Plant

It is not as if there is no planning in the laboratory. If a synthetic lab procedure is so long that the reaction and workup cannot be completed in a single day, chemists can use their experience to extrapolate from similar procedures and guess at what points manipulations can be stopped and under what conditions intermediate solutions or crude solids can be stored without damage. Occasionally there are misjudgments and surprises and a product will be prepared in lower than expected yield or poorer purity. But then even in the worst situation what is lost is no more than a couple of man-days of labor and the price of the starting materials consumed. Also, in the laboratory, because the capacities of refrigerators, freezers, and evaporators are so much greater than the quantities of material being transformed, there are do-able fixes at almost any stage for the situation where a stoppage is forced.

There is no room for such risk-taking on-scale. For advanced intermediates that are themselves the product of a series of sequential steps, one misstep can be economically disastrous. The more points in the process that have been verified as safe-to-stop, by actual test results, the more confidently the process team can be. Moreover, to be a safe stopping point it must be proven safe not just for the quantity and quality of the product but also for the protection of the processing equipment.

As a general rule once a reaction has been initiated, the kinetics must be allowed to run undisturbed to the proper end-point according to the batch sheet. The dynamic transformations cannot be expected to respond to any speed up or slow down without some quantity or quality deviation. After the endpoint condition has been reached and the reaction quenched then the mixture is likely more stable and various stopping points during the work-up can be tested by holding portions of a process mixture for given periods under controlled conditions and examining the mixtures and isolating the product to see whether an unacceptable deviation has occurred or not.

It is more difficult to demonstrate a good stopping point where the mixture in the process equipment is heterogeneous. The difficulty is taking a representative sample for analysis out of a heterogeneous mixture to show that no change affecting quantity, quality, or the protection of the reactor has occurred.  Since one cannot easily take a precise fraction of a heterogeneous mixture, working up that fraction after a pause will not accurately tell you whether the yield would have been different.

Finally, to fairly test the stability of an aliquot at a proposed stopping point the aliquot must be left in contact with a sample of the reactor material. In my experience, this is rarely ever done. At the very least it should be kept in mind where an aliquot might be corrosive to the reactor material.

Use of free radical inhibitors or antioxidants to increase the overall yield of organic synthesis steps

The use of radical inhibitors or antioxidants to improve yields does not appear to have many precedents in organic synthesis. A keyword search in 2011 provided only two references- both related to the stabilization of m-chloroperbenzoic acid towards thermal degradation during the epoxidation of resistant olefins.

Y. Kishi, M. Aratani, H. Tanino, T. Fukuyama and T. Goto, J.C.S. Chem. Comm. 1972  64 and 

D.M. Tal, Steroids (1989),  54(1), 113-22.

The best inhibitor found by Kishi for stabilizing m-chloroperbenzoic acid was 4,4’-thiobis-(6-t-butyl-3-methyl-phenol) that allowed 100% of an m-chloroperenzoic acid charge to be retained after 3 hours heating at 90 C in ethylene dichloride. Octene-1, dodecene-1 and methyl methacrylate were quantitatively epoxidized using such stabilized oxidant.

Synthetic chemists apparently assume that free radical reactions do not occur unless free radical initiators are present in the reaction mixture or unless the reaction mixture is irradiated. It might seem they think it can’t happen unless they are intending it to happen. Obviously, this is not true! Free radical reactions can take place not just during the contemplated reaction phase but during the work-up of the reaction when we might think that all the reacting is stopped. Actually, the opportunity is greater in the work-up phase; this phase usually takes more time, particularly when the process is being scaled up.

Are free-radical reactions inhibited by particular pH ranges of the solvent medium? No, they are not. The most frequent type of free radical reaction is oxidation and only the relative amounts of different species that can be oxidized are affected by pH not particularly the oxidation rates.

Oxidation often produces coloured products when it can introduce new unsaturation into molecules. The presence of unexpected colour in a reaction is suggestive of unanticipated oxidation. I recall that in the preparation of some aniline compounds the procedure teaches the addition of hydrogen sulfide to the aqueous phase during isolation to prevent colour development from exposure to air during workup and crystallization. The usual response to a colored product is to use charcoal in the recrystallization rather than trying to prevent colored by-products in the first place.

If you are performing a distillation and the contents of the still pot are darkening why wouldn't you add an antioxidant? Answer- I've never thought of it.