Bioavailable silica and commercial products
Part I: Demystifying silica by OnlyOrnamental (super genius)
Silica also called silicon dioxide (SiO2) is one of the most abundant minerals on this earth and exists in a lot of different forms and states, from crystalline quartz, over silicates (mixes with other minerals and elements), to filigree snowflake-shaped plankton skeletons made out of amorphous silica, just to name a few. Although silicates play an important role in plant physiology and soil physico-chemistry, though mostly as source for other minerals, amorphous silica is whatâs of prime interest in terms of bioavailability, the ability of plants and animals alike to take up silicon.
The most common form of bioavailable silica is orthosilicic acid, a single silicon atom surrounded by four hydroxyl groups. In solution, the deprotonated form orthosilicate may also lose one molecule of water becoming metasilicate. These two monomeric species are thought to be the sole resorbed forms of silica but they stand in an equilibrium with small and less soluble oligomers.
Like every other acid, it forms salts of which sodium and potassium silicate are quite soluble and very alkaline; in such solutions, monomers and small linear and cyclic oligomers dominate depending on the proportion of silicon to metal. These salts are usually not found in nature but are the result of a simple, high temperature reaction between silicon dioxide and hydroxides or carbonates of alkali or alkaline earth metals.
Though, the direct synthesis product consists of hard chunks difficult to get into aqueous solution. To do so, heat and pressure are necessary, a pressure cooker running at its limit may barely suffice. The hereby obtained aqueous solution can be dried, resulting in a hydrated form of silicate which is now readily soluble. Concentrated liquids obtained in such a way are commonly called water glass and tend to harden out over time, the silicic acid molecules have the tendency to polymerise, assemble themselves back to amorphous silica (more to that later). Notably, silicic acid does also react with other elements such as polyvalent cations like Ca2+ and forms all sorts of mostly insoluble mineral silicates.
The amount of monomeric silicic acid (the main soluble form under physiological conditions) derived from amorphous silica in aqueous solution is constant at about 100-120 ppm, higher concentrations of soluble monomers but also polymers are only obtained above pH 11-12. Obviously, such a high [pH] is corrosive and anything but practical.
It may sound strange to some that the solubility of alkali silicates in water adjusted to a neutral [pH] at ambient temperature is also at said ~100 ppm which coincides with the solubility of amorphous silica.
Why is that? Silicates are the salts of silicic acid; the only thing that determines whether it is in its neutral or deprotonated form is pH (or pKa). Silicic acid is a weak acid, comparable to phenol and at a pHof ~10, half the molecules are in neutral forms (e.g. H4SiO4 or Si(OH)4) and half are deprotonated anions (e.g. HSiO3-); this pH corresponds to the pKa (point of 50% âionisationâ). Bearing a charge allows for a better interaction with water molecules and renders the molecules more water soluble. Anyway, one pH unit higher, at a pH of 11, about 90% of the molecules become deprotonated (negatively charged) and at a pH of 12, only 1% remains uncharged.
On the other hand, at for example pH 7 (three units in the other direction), 99.9% will be neutral silicic acid. This behaviour is true for all acidic and alkaline substances. This means, it doesnât matter if pure silicic acid or a silicate salts is put into solution, only the pH of the solution determines the degree of deprotonation and hence solubility. The sole difference is that in the latter case, counterions are present.
Unlike stronger acids, the week acidity (high pKa) of silicic acid makes that silicate salts, depending on the amount of alkali or alkaline earth metals added for their production, become alkaline in water mostly because the equilibrium between charged silicates and neutral silicic acid is in favour of the latter and for example potassium changes its partner and forms potassium hydroxide with water and hence is the main source of the high pH of water glasses.
As you can see, at any given pH, the solubility and concentration of free acid or its corresponding salts are the same and constant over a large pH range.
In contrary, the pH of a solution of silicic acid or silicates at a given concentration highly depends on the form used. A solution of amorphous silica is about neutral whereas silicate salts are more or less basic. The latter have no concrete molecular formula such as NaCl for table salt but are a wild mix of mostly silica polymers whose degree of deprotonation (or amount of cations) solely depends on the manufacturing process. The more alkali or alkaline earth metal has been added, the higher the final pH, and the better the solubility and stability of a concentrated solution. This is mainly important for technical applications wherein neither pH adjustments nor considerable dilutions with water are employed.
It seems so far as if thereâs not much to gain when using alkali silicates instead of plain amorphous silica such as diatomaceous earth (the aforementioned âsnowflakesâ) as plant fertiliser as the pH has to be in a physiological range at which a saturated silica solution always contains <120 mg per litre.
A question that may be asked is about the exact concentration of Si in those products: In case of diatomaceous earth and precipitated silica, the exact (empirical) molecular formula or in other words the amount of silicon is not exactly known because the inner part of the particles is silicon oxide (SiO2) containing Si-O-Si bonds whereas the surface is covered by silanol groups (SiOH). Is it safe to assume that in case of silicates such as potassium silicate the exact amount of silicon is known? No, unfortunately not. As mentioned above, potassium silicate is not like table salt and the formula K2SiO3 or alike is only theoretical. Potassium silicate is just a mix of an exact amount of potassium salt with amorphous silica such as diatomaceous earth. The given % of K2O and SiO2 on the product is just as good an approximation as talking of diatomaceous earth as SiO2 and neglecting the silanol groups.
To understand solubility but also stability of silicates better, one has to know that silicic acid tends to precipitate. The mechanism is a so called nucleophilic substitution (condensation) resulting in dehydration and polymerisation; the anionic silicate attacks neutral silicic acid with its negative charge and forms a stable bond whereupon silicic acid, now having one bond too many, âlosesâ a molecule of water to regain the tetrahedral four bond state.
This newly formed dimer is more acidic than the monomer and can now attack another monomer or dimer. When the chain has become long enough, it can âbite its own tailâ forming a closed cycle (often tetramer in aqueous solutions). As explained above, both neutral and charged forms are present in sufficient amounts even at neutral pH values (0.1% is not much but enough because the reaction is quite fast and the growing polymers become more acidic). This means that only highly basic solutions wherein most everything is charged, under which conditions water can attack a forming chain and brake it apart, and where to solubility is greatly increased are truly stable. Besides, very acidic solutions containing mostly non-charged silica are semi-stable; but more to that in the second part.
The forming insoluble polymer around neutral pH is called precipitated silica. It is an amorphous silica consisting mainly of linear chains and unbranched cycles and remains soluble once diluted below 100-120 ppm. Only over time do these chains crosslink, forming branched cycles of up to 12 silicon atoms being only truly soluble at very high pH values; such particles are small enough (~2 nm) to give a transparent solution (so called pre-sol) which easily passes through a 0.1 ”m pore sized ultra-filtration membrane. These particles grow and once reaching a size of up to 100 nm form a colloidal suspension (so called sol) which can still be brought into solution by dilution. These particles are fairly dense whereas the subsequent formation of longer, also branched chains throughout the liquid and cross-linkage of these particles, like a network of small balls, results in a highly porous continuous gel of increasingly lower solubility. This process may take minutes or years as it highly depends on concentration, pH, temperature, and nucleation particles such as rough surfaces or impurities, just to mention the most important factors. Only under extreme conditions such as volcanos and tectonic activities do organised, three-dimensional (i.e. crystalline) structures form known for example as quartz and become as good as completely insoluble (though usually, quartz forms directly from a melt or crystallises as such).
Copied (without asking, thanks nonetheless) from âThe Chemistry of Silica: Solubility, Polymerization, Colloid and Surface Properties and Biochemistry of Silicaâ by Ralph K. Iler, 1979.
The right pathway under alkaline conditions depicts also the âOstwald ripeningâ which mainly plays a role for particles smaller than 10 nm. Above, particle growth becomes slow. On the other hand, particle growth at low pH is generally slow and aggregation of the non-charged particles occurs.
Precipitated silicates, diatomaceous earth, or silica obtained from plant matter are usually amorphous silica and all result in an equal and constant amount of free and hence bioavailable silicic acid when put into water (i.e. ~100 ppm). The only thing different is the speed at which an equilibrium between solution and solid forms. In case of amorphous silica this speed determines how fast free silicic acid is liberate from solid particles and depends on the accessibility of silicon atoms for nucleophilic attack by water (i.e. hydrolysis of polysilicates, the reverse mechanism of above explained polymerisation). Accessibility is basically determined by steric hindrance, or atoms blocking the path of water.
In case of precipitated loose silica chains or diatom skeletons with their needles, holes, and sharp edges, accessibility is very good and solubilisation readily occurs. Solubilisation is a two way process, an equilibrium, also driven by the âurgeâ of silicic acid to polymerise. At the point of saturation, as many molecules go into solution as there are molecules forming solids.
Several mechanisms contribute to a phenomenon commonly called Ostwald ripening; it basically means that small particles dissolve faster whereas big particles tend to grow, the equilibrium doesnât work equally at different positions. This principle is the reason why freshly prepared silica solutions contain many small particles whereas aged solutions comprise of lesser but larger aggregates.
The lessons learned so far:
- It doesnât really matter what kind of soluble silica preparation is used, the final concentration will not be above 120 ppm at physiological pH and room temperature (400 ppm in boiling water).
- Freshly prepared silicate solutions but also suspensions of small particles with strong surface curvature like diatomaceous earth are the preferred source of bioavailable silicic acid.
âTraditionalâ commercial products usually come as precipitated, readily soluble alkali silicates containing about 20% residual water or pre-dissolved as concentrated stock solution, the latter with a somewhat limited shelf-live. As said, in case of for example potassium silicate such a solution is only acceptably stable at very high pH. It does contain up to 300 ppm orthosilicic acid, the rest are, though soluble, oligo- and polysilicates (notably, highly alkaline silicate salts are miscible with water at any ratio).
All this doesnât matter because such a solution canât be safely applied to a plants. Diluting such a solution in plain water will only reduce the pH by ~1 unit due inherent buffering capacity. The diluted solution is still aggressive and not very practical; the pH needs to be adjusted by adding an acid. As you know by now, once the pH is down to a physiological level, the silicic acid precipitates very quickly until only 100-120 ppm of orthosilicic acid remain in solution. The precipitate will re-dissolve once the soluble part is resorbed by the plants but according to the Ostwald principle with every day that passes, the big particles become bigger whereas the small ones dwindle (this effect is strongest for small non-colloidal particles).
At one point, the speed of resorption will overcome the speed of solubilisation, the free silicic acid concentrations falls below solubility levels, and the plant may not get enough. But on one hand, 100 ppm silica is already a lot and more than enough for most applications. On the other hand, adding new silica solution readily fixes the problem because toxic levels are hard to obtain. Furthermore, silica being of most interest in hydroponic gardening, many gardeners donât just dilute in plain water but solutions such as fertilisers which are either already buffered or need a pHadjustment. The additional effort is hence marginal.
This brings us to the core question: What is the big deal with stabilised silicic acid products?
First, do we really need a higher solubility than 100 ppm? I would say, no, not necessarily because plants canât safely handle more. It also seems unlikely that a âstabilisedâ form, maybe a complex or chelate, would be transported as such through the plant and if so the silicic acid is likely to lose its âactivityâ. Like so often with fertilisers and plant nutrients, more doesnât help more but rather just forms insoluble precipitates somewhere it shouldnât.
Certainly, having a physiological pH in a reasonably concentrated/diluted stock solution is an advantage. Not everyone can or will adjust the pH and diluting concentrated water glass by a factor of letâs say 10â000 is also not that convenient. But does this alone justify the likely high prices of commercial products? There has to be more to it!
A common advertisement is the higher amount of bioavailable silicic acid in stabilised solutions. As we know now, precipitated silica and diatomaceous earth, though poorly soluble, are bioavailable due the equilibrium between solid and solute; they just take more time to dissolve than alkali silicates. Furthermore, we know that, no matter what, free orthosilicic acid in solution wonât surpass 120 ppm at room temperature.
Besides, do we really know that âstabilised silicic acidâ is orthosilicic acid in solution? The answer in most cases is no, we donât.
Thatâs because for one there are no scientific publications which properly investigated the subject and for another the stability tests done for patent approvals are often based on turbidity (I admit I havenât read all of them, reading patents is very boring). Turbidity is only perceived by the naked eye or under a standard light microscope for particles larger than 0.2 micrometres (due optical resolution). This means, we do not know if the solutions really contain a higher amount of soluble silicic acid, let alone orthosilicic acid, or just smaller particles than a non-stabilised solution. As already mentioned more than once, having smaller particles is an advantage regarding shelf-life. Apropos shelf-live, the patents I found so far only claim a truly increased stability in concentrates whereas in diluted forms stability drops down to hours, days at best, which is slightly more than for non-stabilised solutions. Apart from human convenience, there is seemingly no advantage in them regarding our plants. More so, one is inclined to use plain old diatomaceous earth for the following reasons: Being pure silica, it has the highest possible silicic acid content, it is a cheap, light, and dry powder with infinitely stable making it easy to dose and compatible with all sorts of fertilisers and most additives and it doesnât require pH adjustment.
Furthermore, it exhibits several beneficial effects as soil amendment and is fine enough to be used as foliar spray which protects against insect herbivores. It may also be used in hydroponic systems, maybe wrapped in a fine mesh, to serve as long lasting depot form delivering a constant, physiological level of readily available silicic acid.
On the contrary, the proposed dilutions for commercial stabilised silicic acid are, at least from what Iâve seen so far, up to 50 ppm. It goes without saying that this is below saturation and wouldnât need stabilisation.
Is there a way to turn the tide in favour of stabilised silica solutions? A question discussed in the next partâŠ
Part II: The (ir-)rational science behind stabilised silicic acid
Silica sols with particles between 10 and 100 nm in diameter become stable for decades by the addition of a few percent alkali hydroxide to adjust the pH to 9-10. Furthermore, small quantities of added salts reduce the high viscosity and allow for very high silica concentrations (70% w/w). Certain additives render these ânanoemulsionsâ stable also under slightly acidic conditions or can be dried to a powder which readily redisperses in water.
On the other hand, freshly prepared pre-sol or even true monosilicic acid solutions are only stable under very alkaline conditions. Without stabilisation, diluting a concentrated alkali silicate solution will result in quick polymerisation likely due a drop in pH. Dilution to a concentration slightly above solubility (i.e. > 300 ppm at high alkaline pH) results in small non-colloidal particles of low nm size which remain stable for several hours. Though, at a ten times higher concentration small colloidal particles form already after an hour âagingâ and after a day or two, most monomeric silica is deposited. If done so at neutral to slightly acidic pH, this process takes seconds to minutes.
With a few exceptions, the presence of polyvalent metal cations also catalysis this reaction considerably resulting in an insoluble precipitation of metal silicates. Monovalent ions at higher concentrations may cause an effect known as salting out; the obtained precipitate remains soluble upon dilution at higher pH but may become irreversible at neutral and acidic pH. Furthermore, all kinds of ionic impurities accelerate polymerisation and gelling. Hence the recommendation to first add silicic acid when preparing a fertiliser blend or even better to avoid any additives.
At a pH around 2, orthosilicic and disilicic acid solutions at up to 1% are fairly stable but with a somewhat unpredictable shelf life of hours to days, sometimes months, and an uncontrolled dynamic between monomers and slowly forming small oligomers (particles up to 2-3 nm). For commercial purposes, this is not at all suitable and demands for two sorts of stabilisation: One that allows a more concentrated form with an increased shelf life and on in diluted and/or pH adjusted form for long lasting bioavailability or stable physico-chemical properties depending on the intended use.
Although a fair amount of todayâs knowledge on silica, dissolution and precipitation thereof as well as stabilising effects of a broad set of additives dates back sometimes to the beginning of last century, the possibly first stabilised silicic acid product on the market intended for plants, animals, and humans is BioSilÂź. It comprises of a highly concentrated and very acidic solution of choline chloride to stabilise a small percentage of in situ formed orthosilicic acid (although, it contains also an undefined mix of small non-colloidal oligomers). The stabilising effect of quaternary ammonium salts such as choline chloride on alkali silicate solutions is known for a long time and likely due a disturbed coordination hampering polymerisation. It is said that such stabilisation is not only abolished upon pH adjustment to physiological values but that choline in fact speeds up the polymerisation of such a solution most likely by inhibiting the solubilisation step in the equilibrium between dissolution â deposition/polymerisation.
Furthermore, the permanent positive charge of quaternary ammonium salts causes them to âstickâ to the surface of silica particles protecting them from interactions with other solutes; this inhibits solubilisation and precipitation alike. A second quaternary ammonium salt is carnitine phosphate used instead of choline chloride. According to patent claims, carnitine as well as phosphate increase monomer stability several fold even in diluted and pH adjusted solutions. They speak of a better chelation which is likely utter nonsense (chelates with silica exist but involve ortho -diphenols such as catechol, humic/fulvic acid and the like); earlier findings point towards a generally increased stability of larger quaternary ammonium species and carnitine is larger than choline.
Again, phosphoric acid is known to catalyse silica dissolution for over 30 years prior to the corresponding patent and phosphates and phosphonates are part of many antiscaling agents for example for dishwashers and in laundry detergents. Organic bases, some of which not as strong as tetraalkyl ammonium species, mixed with alkali metals seem to exhibit stabilising effects mostly on nm small particles rather than free orthosilicic acid though the distinction between the two is not always obvious. I found two strategies for the preparation of quaternary ammonium stabilised silica: One uses common alkali silicates and HCl resulting in NaCl (table salt) containing solutions, the other uses silicon tetrachloride or tetraethyl orthosilicate resulting in an acidic quaternary ammonium silicate solution. Such quaternary ammonium silicate salts are known for over 50 years. They are fairly stable at silica concentrations of up to 50%, are water miscible, and sometimes even soluble in water miscible solvents. Notably, they are not acidic but become very alkaline once diluted in water.
Another stabilising agent is boric acid usually combined with a high amount of a âhumectantâ (40-60%) to stabilise non-colloidal silicic acid wherein BTW no free orthosilicic acid is found but only fast dissolving nm sized pre-sol particles.
According to the corresponding patent, the advantage of such a mixture is less in its stabilisation but the yet unexplained synergism between the two micro-nutrients regarding either biological activity. Notably, the increased stability is mainly due to the polyol or polyether âhumectantâ and not boric acid. Those compounds form a protective layer around pre-sol particles much like they do in oil/water emulsions such as creams; nowadays common knowledge which predates the patent by decades. Boric acid alone does work too but is required at a too high silica/boron ratio to be of any use in or even on living organisms. Nonetheless, boric acid stabilised silica is used as industrial antiscaling agent and employed in mining and similar industrial sectors. On the other hand, these âhumectantsâ, also called osmolytes, hydrogen-bonding agents, and auxiliary solvents, such as glycerol, sugars, and other polyols, PEG and related compounds, or amino acids and derivatives, but also divalent cations (metal salts) are part of several recent patents and commercial products though their effects on stability/dynamics of non-colloidal silica solution at acidic pH is known since the 1950s.
As a side note, these agents, depending on the mode of application, can also be used as flocculants (only a very specific ratio of silica to âhumectantâ results in stabilisation, flocculation, polymerisation, and/or gelation are more likely to appear).
Commercial products are usually comprised of highly acidic solutions with shelf lives of 1-2 years. Astonishingly, calcium and other divalent metal cations donât form precipitates in highly acidic concentrated choline stabilised silicic acid solutions and also show prolonged stability upon pH adjustment. The reason therefore evades me. [I should put that somewhere else, but where?]
So far, I havenât come across any product of neutral pH. The only alkaline complex is comprised of silica, arginine, and inositol. It isnât even a liquid but a solid though readily soluble in water and nearly as alkaline as potassium silicate itself. Like most products, the resulting diluted solution is only stable for a day or two and like with the other products, the patent claim is not that original as guanidinium silicates (arginine is a guanidine derivative) have been known well before the above patent was filed.
Common methods used to determine free orthosilicic acid are under others 29Si-NMR and IR spectroscopy, turbidity determination (light scattering) and colorimetric assays with ammonium molybdate, as well as ultra-filtration and time of gelling. Not all methods work for all formulations and not all distinguish between monomers and pre-sol oligomers. As a rule of thumbs, silicic acid solutions, either freshly prepared or stabilised, comprise of monomers and small nm sized particles which form orthosilicic acid once diluted below the saturation concentration of orthosilicic acid.
One mechanism ascribed to the stabilising effect is the presence of so called deep eutectic solvents (DES). Although, these donât contain free water (although up to 20% bound water may be present depending on the solvent system); to my knowledge, most commercial products contain mainly water (at least >40%) which results in full solubilisation and dissociation of DES constituents. Hence, this hypothesis is refuted like so many other speculations mentioned in patent claims. Unfortunately, only very little is published in scientific literatureâŠ
Silicon particles have an inherent negative surface charge; the higher the degree of polymerisation, the more acidic silanol groups become. Some calculations imply that about half the surface SiOH groups of non-colloidal silica particles (< 2 nm) are ionized at physiological pH. This applies to concentrated alkali silicate preparations and contributes, together with an increased solubility, to their stability at elevated pH values. In contrast, silica particles in acidic solutions have as good as no surface charge causing the particles to aggregate easily. In contrary, polymerisation (i.e. particle growth) is at a minimum around pH 2 resulting in a temporary stability. At above ~1% silica, the tiny particles aggregate easily and form denser gels than at less âstableâ pH values. Addition of certain positively charged ions such as choline chloride and divalent metal salts or polyols and polyethers may prevent growing and cross-linking of such particles by electrostatic repulsion (the additives form a charged layer around the particles and are like miniature â+ polesâ of magnets repelling each other) or sterical hindrance (âpassiveâ shielding, acting like fenders around a boat), respectively.
Although, most patent claims for stabilised silicic acid solutions do not aim for pH 2 at which pure silica solutions would be most stable but at a pH preferably below 1. Such acidic solutions undergo slow acid catalysed polymerisation (possibly due fluoride contamination) and the particles therein are slightly positively charged; this results in less dense packed non-colloidal particles which are easier and faster to re-solubilise into orthosilicic acid once diluted and they associate better with polyols and polyethers. But under these conditions, addition of common salts do not influence (i.e. accelerate) gelling. Which leaves me wondering why and how choline might work so wellâŠ
