How does one Produce the Most Likely Impurities in the Product from a Process Step?

In order to develop a good purity analysis for an organic substance, one needs to have some method to assess different methods. A better method separates and quantifies more impurities from the product. A better method increases the degree of separation between the closest impurity and the product without losing separation for any other impurity. A better method separates more cleanly an impurity designed to have a very minor structural difference from the product.

 

But is there a way to prepare a  resolution standard with larger amounts of the most likely potential impurities of the product? This becomes an important practical matter.

There are two types of impurities. Impurities that are derived from product degradation; that is, they come via the desired product and arise from reactions of the desired product occurring after it has been isolated and purified. A separate type of impurity is one that is formed at the time of the synthesis of the desired product and which is not completely removed by the isolation, separation, and purification processes performed before packaging the final product. Such impurities are characteristic of the process. 

It is this second type that is considered here. These impurities are produced in greater or lesser amounts by variations from the proper continuous variables controlling the process.

 

New previously unobserved impurities are usually created by changing the discontinuous variables of the process step, such as reagents, reagent purity, reagent/substrate ratio, solvent, solvent purity, substrate purity, processing chemicals…

The most significant continuous variables are time and temperature.

The one that can produce the most profound or substantial changes in chemical reactivity is temperature. The effect of time is usually more limited because if we are seeking a complete reaction it is likely that the transformation is self-terminating and will essentially stop when the correct time has expired with relatively less occurring after this required reaction time. Increasing the temperature by 10 C according to a rule of thumb should double the rate of reaction. This will also allow competing reactions that are limited under the most preferred conditions to compete and produce by-products.

Thus, an increase in the temperatures of each of the different stages of the reaction by 10C and a decrease in the time by half in each stage should produce more impurities in the final product and these impurities should reflect realistic possible impurities.   If possible the extent of disappearance of starting material should be kept about the same.

Product separated from a process step stressed in this way should show increased amounts of the most likely process impurities.

If it does not, this is important information for your determination of critical parameters that you will need to work on at some point. Tests that provide a resolution standard go hand-in-hand with examining critical parameters!

Uncommon Solvent Immiscibilities

 KiloMentor is always on the lookout for methods to separate components of a mixture by partitioning between immiscible liquids. Better known ones are methanol or acetonitrile with hydrocarbons. Many different solvent pairs may show immiscibility between ambient temperature and -20 C and this temperature range is easily accessible inside a jacketed reactor where liquid-liquid partitioning is done at scale. It is in the laboratory that this temperature range is inconvenient to achieve.

Below are listed some less-common immiscible pairs that may prove useful.

Dimethylsulfoxide - Xylene

Dimethylsulfoxide - Diethyl Ether

Dimethylformamide - Xylene

Dimethylformamide -Diisopropylether

Trichloroethylene - Xylene

Acetic acid - Hexane

Methyl t-Butyl Ether (TBME) -Sulfolane

The DMF /Diisopropyl ether immiscibility suggests that one look for an Upper Critical SolutionTemperature (UCST) between DMF and TBME at below room temperature. A small amount of water could be added to the DMF to raise the UCST.

Since DMSO and diethyl ether have immiscibility it suggests that one explore for a UCST between DMSO and TBME below ambient temperature.

What would the miscibility be between a mixture of xylene and diethyl ether with DMSO? Both xylene and diethyl ether are separately immiscible with DMSO.

To get rid of the diethyl ether suppose we try a mixture of xylene and TBME with DMSO?

How about a mixture of DMSO and trichloroethylene with xylene? Both DMSO and trichloroethylene are separately immiscible with xylene.

Dimethylsulfoxide or DMF reactions could be worked up by extraction into m-Xylene followed by azeotropically removing the xylene as an azeotrope with water after cold extracting the xylene to remove residual dipolar aprotic solvent.

Both m-xylene and isopropylbenzene form azeotropes with water that can be used to quickly remove the organic as a clean phase. Can either of these be useful for isolating organics formed in the solvents DMF or DMSO? You tell me- I’m retired; you have a lab.

Diphenylphosphine Oxide Containing Compounds: Intermediates almost guaranteed to be Crystalline

 

 Stuart Warren, in an article in Accounts of Chemical Research 11 (11) 401 (1978), wrote that almost all diphenylphosphine oxide-containing compounds are highly crystalline white solids. KiloMentor is, therefore, proposing the use of compounds containing the diphenylphosphine oxide substructure as one of the preferred intermediate types in ‘paper’ syntheses.

It is well known that the reaction of a primary alkyl halide with triphenylphosphine produces a quaternary phosphonium salt that is both an ionic salt and crystalline. Hydrolysis of such a compound in aqueous base liberates benzene and provides the phosphine oxides. These compounds in turn can be alkylated with other alkyl halides using butyllithium and TMEDA as co-solvent. [J. Chem. Soc. Perkin Trans. I, 550 (1977)] Warren predicts that these also will be highly crystalline solids.

The KiloMentor strategy for paper synthesis route design emphasizes the advantages of selecting a route that can easily be scaled up. To be preferred, intermediates need to have an increased likelihood of being easily separated and purified, preferably by acid-base extraction. This is proposed to be an overarching advantage over competing routes, whose intermediates almost always have to be purified by crystallization. The problem with these competing routes is that the crystallizability of an intermediate from a paper synthesis cannot be dependably predicted.  

Besides those intermediates, purifiable by extraction, other intermediates would also be preferred if, even when still unknown and existing only ‘on paper’, they contained a functional group that could pretty well guarantee they would be found to be crystalline. There are not many of these and they are not celebrated for this property. Usually, the ease of crystallization for a compound depends upon the entire molecular structure and cannot be predicted, but diphenylphosphine oxide appears to be one that should come with a guarantee.  

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.     
                 
Molecular Formula - C₇₆H₅₂O₄₆
Molecular Weight - 1700
Melting point - 218°C
Water solubility - 1g/ 0.35 ml

Speaking roughly, to produce a hydrotrope a chemical must dissolve in water to give a 1M solution. A 1M solution of tannic acid would contain 1700g of organic solid per liter of water. That would be 1.7 gm per milliliter. The solubility of tannic acid in water is 4.88 gm per milliliter. One could achieve a solubility of 2.87M if required in a saturated solution. Tannic acid is a material available in industrial quantities at a practical price. Sigma-Aldrich sells 500 grams for less than $100.00. Considering that only 60 g of urea are needed to produce a 1M aqueous solution that would give an effective hydrotrope and supposing that we provide three times as much tannic acid by weight, that would just be 180 g per liter that would not cost more than $75.00!

The molecule shown in the figure is only one representative (perhaps the major one) of the constituents of the organic mixture called ‘tannic acid’ but if we accept that it is typical, then each molecule can be approximated as containing about 25 phenolic groups and 10 ester linkages. The phenolic groups alone would comprise over 15 hydrogen bond acceptors and 25 hydrogen bond donors. Certainly these can be counted on to increase the solubility of many organic solutes in the tannic acid/water phase.

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.