Monday, 6 July 2020
Tannic Acid Hydrotropes
Most hydrotropes are made by dissolving organic salts at a concentration of at least 1M in water. Covalently bonded materials do exist that form hydrotropes. The best known is urea. Another inexpensive, non-ionic organic material that is highly soluble in water and that can be expected to promote the dissolution of other organic substances is tannic acid.
Monday, 29 June 2020
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.
Saturday, 27 June 2020
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.
Monday, 22 June 2020
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.
Saturday, 13 June 2020
Impinging Jet Micromixing to Solve the Problem of Small Crystal Size without Milling.
CA2044706, Crystallization Method to Improve Crystal Structure and Size expired June 14th 2011 in Canada and the family member US5314506 expired May 24th 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 6th 2003. The Canadian attempt became a dead application August 15th 2007.
Tuesday, 9 June 2020
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; TDL0; 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.,; TDLo 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.
Friday, 5 June 2020
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.
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