What and How Can Be Obtained from Sand? The Unknown Face of Silicon
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The following article was originally published in the journal for educators Chemia w Szkole (eng. Chemistry in School) (6/2016):
An Unknown Face of Silicon
It’s fair to say that silicon Si is one of the most important elements for our civilization. We find it in most of today’s indispensable electronic devices, such as computers and mobile phones. Contemporary times are often called the “information age” and thanks to its semiconductor properties, silicon serves as the foundation for nearly all devices used to process information. These reasons alone make it worth examining this substance more closely.
Silicon is a semimetal element belonging to group 14 of the periodic table. Antoine Lavoisier took a closer look at silicon in the second half of the 18th century. In 1800, Humphry Davy mistakenly believed it was a chemical compound. Not until 24 years later did Jöns Jacob Berzelius obtain pure silicon, proving that it is indeed a chemical element rather than a compound [1] [2].
Silicon is among the most abundant elements on Earth. In fact, it’s the second most common element by weight (27%) in the Earth’s crust, right after oxygen [3].
In nature, silicon never occurs in its free state. Silicon dioxide, also called silica SiO2, is found in various polymorphic forms — one of these is quartz (Photo.1). Silica and other silicon compounds, primarily silicates and aluminosilicates, make up the vast majority of rocks that form the Earth’s crust. Along with aluminum Al, silicon is responsible for the old name “sial” referring to the Earth’s outermost layer.
In popular science and educational publications, silicon often appears almost exclusively in the form of silicates — for instance, in the experiment known as the “chemical garden” [4]. I believe that’s a big oversight, because this element forms a vast number of compounds, many of which exhibit fascinating properties.
In this article, I’d like to present a method of obtaining elemental silicon, along with several of its interesting compounds.
Right from the start, I’d like to ask anyone who decides to repeat these experiments to heed the safety warnings and proceed with caution, because any carelessness can be extremely dangerous!
What Can Be Obtained from Sand and How
To isolate free silicon, we need a suitable source of this element. In this case, that source is silica. Of course, we could use chemically pure silicon(IV) oxide, but there’s no real need. As already mentioned, silicon compounds are among the main components of the Earth’s crust — ordinary sand is almost pure silica. The best choice for our experiment is so-called quartz sand, which is typically light in color, sometimes nearly white, and contains very few impurities (Photo.2A).
Another good source of silicon(IV) oxide is diatomaceous earth, also known as diatomite (Photo.2B). It’s a sedimentary rock formed primarily from the remains of diatoms (Diatomophyceae), which are single-celled algae [5]. Diatoms incorporate silica into their cell walls (known as the theca, comprising the lid epitheca and the base hypotheca), making them rigid [6]. Diatomaceous earth is yellowish-white, very light, and highly porous.
In my experiments, I used quartz sand, regular purified sand, as well as diatomaceous earth. All these raw materials performed equally well.
To extract free silicon from its oxide, we of course need to reduce it. For that purpose, we’ll use a fairly strong reducing agent: metallic magnesium Mg [7].
The mixture we need is prepared with the following composition (by weight):
- quartz sand SiO2 – 47%
- magnesium Mg – 37%
- magnesium oxide MgO – 16%
The sand must be clean and thoroughly dried, just like the other substances. Retaining any moisture can be dangerous because it increases the risk of splattering during the reaction.
As for magnesium, you should use fine filings or, ideally, powder. Magnesium oxide MgO has a supporting function; if it’s unavailable, you can omit it but then must adjust the proportions of the other ingredients accordingly.
All substances, ground as finely as possible, are mixed thoroughly and then placed in a ceramic crucible (Photo.3). It’s not advisable to use more than a few tens of grams of this mixture.
The crucible must be placed on a nonflammable, heat-resistant surface, such as a firebrick or a thick layer of dry sand.
To initiate the reaction, you need to heat the mixture locally to a high temperature. The ideal method is to embed a strip of magnesium ribbon in the mixture. After lighting it, be sure to stand back — use personal protective equipment, especially safety goggles!
The reaction between silicon(IV) oxide and metallic magnesium is highly exothermic. It releases a large amount of heat, causing the crucible’s contents to glow white-hot (Photo.4). It’s not unusual for the crucible to crack, which is why a heat-resistant base is essential.
The reaction typically lasts only a few seconds. Then, cover the crucible and let it cool completely. Note that the reaction products occupy a larger volume than the reactants, so never fill the crucible more than two-thirds full. Otherwise, you might end up with something like what you see in Photo.5.
We can represent this reaction as follows:
As you can see, the silicon(IV) oxide in the sand is reduced by magnesium, which is oxidized in the process, yielding elemental silicon. This reaction is analogous to what happens when thermite (a mixture of iron oxides and powdered aluminum) is ignited [8].
Meanwhile, some of the newly formed silicon combines with magnesium under the reaction conditions:
Magnesium silicide Mg2Si is a dark blue solid [9].
The post-reaction mixture contains the silicon produced, but also unreacted silica and magnesium, magnesium oxide, and a small amount of magnesium silicide.
To extract the silicon from this mixture, the simplest approach is to use an acid, such as hydrochloric acid HCl. In a beaker, place 20 cm3 (0.68 fl oz) of 36% hydrochloric acid, then slowly add the powdered post-reaction mixture (Photo.6A).
As the mixture reacts with the acid, large volumes of gas are released. Interestingly, these gases ignite spontaneously on contact with air (Photo.6B). That’s because magnesium silicide reacts with the acid as follows:
The monosilane SiH4 is very reactive and ignites spontaneously in the presence of atmospheric oxygen, forming silicon dioxide and water:
The spontaneous ignition of monosilane is especially striking in the dark (Photo.7). Also note the bright streaks near the flame in this and the previous photo; they come from tiny particles of unreacted magnesium. The presence of magnesium in the post-reaction mixture can be hazardous, because its reaction with acid releases hydrogen gas H2. There is a small possibility of forming an explosive hydrogen-oxygen mixture in a 2:1 ratio (often called oxyhydrogen). Even the smallest ignition source could trigger an explosion. Naturally, with spontaneously igniting monosilane present, the risk of an explosion and scattering of glass fragments and corrosive acid is higher. Though this is rare, stringent personal protective measures are mandatory, especially around the face. To reduce the risk, you can blow a strong stream of carbon dioxide CO2 over the liquid’s surface during the reaction. This step must be carried out in a fume hood or outdoors.
During the reaction with the acid, any leftover metallic magnesium, as well as magnesium oxide and magnesium silicide, dissolves.
After a few hours, filter off the remaining solid, rinse it several times with distilled water, and dry it.
The resulting solid (Photo.8) is composed primarily of silicon Si, with a small amount of unreacted silica. The silicon here is in an amorphous form. You can attempt to recrystallize it from molten zinc Zn to obtain larger crystals [7] [10].
You could stop here and simply add silicon to your collection of elements. However, I’d like to propose continuing the synthesis to obtain a rather little-known yet, in my opinion, very intriguing silicon compound.
Wöhler Siloxene
This time, our starting material will be the finely powdered silicon from the previous experiment. We also need metallic calcium Ca.
Calcium is a silvery-white metal belonging to the alkaline earth metals. Free calcium is quite reactive; like sodium and potassium, it must be stored in oil or mineral oil. When exposed to air, it quickly forms a layer of calcium oxide CaO. Only after scraping off this layer can we see the shiny surface of the metal (Photo.9).
Because calcium oxidizes so readily, it’s easiest to store it in larger chunks, where only the surface is affected. Unfortunately, for this experiment we need the metal in the form of small filings. Luckily, calcium is a soft metal, not much harder than lead, so you can easily file it with a metal (or even wood) file. The resulting shavings are suitable for this experiment. Naturally, these should also be stored under a liquid that prevents air contact. For short-term storage, hexane C6H14 works well, as it helps rinse away residual mineral oil in which the calcium was originally kept (Photo.10).
When handling finely powdered calcium, take care not to let it contact your skin — and especially not your eyes. Calcium reacts exothermically with water, producing an alkaline solution of calcium hydroxide Ca(OH)2, so careless handling can lead to both thermal and chemical burns.
Next, prepare a mixture with the following composition:
- 2.4 g (0.085 oz) of silicon
- 1.7 g (0.06 oz) of calcium
We’ll use the finely powdered silicon from the previous experiment. A small amount of silica impurity is not a problem. Meanwhile, the powdered calcium must be carefully drained of any residual hexane and thoroughly dried before weighing. Be cautious when drying, because hexane is volatile and highly flammable. Its toxicity is relatively low, but inhaling its vapors can cause mild euphoria, drowsiness, dizziness, or nausea.
Once you’re sure the silicon and calcium mixture is completely dry and free of any hexane residue, transfer it to another ceramic crucible (Photo.11).
Cover the crucible with a ceramic lid and heat it strongly in a burner flame. Note that an alcohol burner will definitely not suffice; a gas burner is required. After the crucible glows red-hot, the reaction to form calcium silicide CaSi2) begins, according to:
Calcium silicide melts at 1033°C (1891.4°F). Once the crucible cools, you can retrieve the fairly hard lump that consists mostly of this silicide (Photo.12).
This silicide is insoluble in water but can decompose on contact with moisture (including atmospheric humidity), releasing hydrogen and converting to calcium hydroxide. It’s flammable, and in a finely divided state it can ignite spontaneously. For these reasons, it’s best to proceed to the next step immediately after synthesis.
The entire procedure involving hydrochloric acid should be carried out in a fume hood, since large amounts of irritating and toxic hydrogen chloride gas can be released. The reaction also generates hydrogen, and small amounts of monosilane may form, so the same precautions about detonating mixtures, ignition, and possible explosions apply as before.
In a beaker of at least 100 cm3 (3.38 fl oz) capacity, place 23 cm3 (0.78 fl oz) of 36% hydrochloric acid. Slowly add the crushed calcium silicide and quickly heat the solution to boiling. After a short while, add another portion of acid, this time 12 cm3 (0.41 fl oz), and bring it back to a boil. After a few minutes, add 70 cm3 (2.37 fl oz) of distilled water preheated to about 70–90°C (158–194°F) and boil again for a few minutes. Then remove the heat and let the mixture cool.
During heating, you’ll notice a large amount of precipitate forming. Filter this off and rinse it with a dilute hydrochloric acid solution. Try to discard as much of the heavier particles (e.g., residual silicon) that sink rapidly to the bottom.
The precipitate is brownish-yellow, sometimes with a greenish tint (Photo.13). This compound, with the formula Si6O3H6, is known as Wöhler siloxene, named after the 19th-century German chemist who discovered it.
The synthesis of siloxene can be summarized by the following reaction:
Unfortunately, Wöhler siloxene is not very stable — it’s best produced as needed and used within a few days. It’s somewhat more stable when kept under a weak hydrochloric acid solution rather than in a dry state. Therefore, place the purified precipitate at the bottom of a vessel and cover it with a few centimeters of dilute HCl [11].
No strong toxic properties have been confirmed for siloxene, but its exact effects on the human body have not been thoroughly studied. Hence, caution is advised.
Wöhler siloxene burns with a blue flame, yet it also has another intriguing property, which we can demonstrate very simply [12].
All you need to do is disperse a small amount of this substance in 20–50 cm3 of dilute hydrochloric acid (under 1% concentration). Siloxene is insoluble, and the precipitate settles relatively quickly, so continuous stirring is helpful (Photo.14A). Then add a few crystals or a small amount of potassium permanganate KMnO4 solution.
A very distinct red-orange chemiluminescence is visible, which may last anywhere from a few seconds to several minutes, depending on how much material you use (Photo.14B). Without stirring, the reactants sink slowly to the bottom, where they keep reacting — the glow is then weaker but lasts longer.
This reaction can also be performed in a long glass tube filled with the siloxene–acid suspension. Each crystal of permanganate dropped into the tube leaves a trail of glowing light, which looks quite impressive.
One might wonder if the violet permanganate solution alters how we perceive the color of the emitted light. You can find out easily: instead of preparing a siloxene suspension in the acidic solution, mix the acid and potassium permanganate beforehand. Then place a piece of filter paper carrying some siloxene on top of the solution (Photo.15A).
The light is clearly orange (Photo.15B), noticeably different, for example, from the chemiluminescence of singlet oxygen (Photo.16), which can be generated by reacting free chlorine Cl2 or sodium dichloroisocyanurate C3N3O3Cl2Na with an alkaline 30% hydrogen peroxide H2O2 solution [13] [14].
You can achieve another striking effect by smearing a little damp, acidified Wöhler siloxene onto filter paper and sprinkling a few crystals of potassium permanganate (VII) on the resulting patch (Photo.17).
When the lights are turned off, you’ll see something resembling a night sky filled with stars (Photo.18). On closer inspection, you can see that each crystal in contact with the chemiluminophore glows in a small surrounding zone.
These chemiluminescence experiments are both impressive and relatively safe, which makes them suitable for public demonstrations.
Explanation
The most economically important silicon compounds are silica (the main component of sand and glass), silicic acids of various structures, their salts (e.g., silicates in water glass), and chlorosilanes and alkoxysilanes, which serve as precursors for producing silica gels.
Interestingly, silicon forms chains such as Si–Si, Si–O–Si, and Si–N–Si. This feature makes it in some ways similar to carbon. The group of such silicon-based compounds is quite large, though still smaller than the analogous family of carbon compounds.
Silicon also forms a family of hydrogen compounds analogous to hydrocarbons, called silanes. These are colorless gases or liquids with a distinctly unpleasant odor. They are toxic and insoluble in water. The simplest silane is precisely the monosilane SiH4 we produced, whose structural formula is shown in Fig.1A.
The resemblance to methane CH4 (Fig.1B) is apparent. Both molecules have a tetrahedral structure with hydrogen atoms at the vertices. The Si–H bond length in monosilane is 147.98 pm, whereas the C–H bond length in methane is 108.70 pm.
Silanes form a wide range of derivatives, and their properties depend greatly on the substituents. Compared to analogous carbon compounds, silanes are often more reactive because the Si–nonmetal bond is generally more polarized than its C–nonmetal counterpart. We’ve already observed this with monosilane, which ignites spontaneously upon contact with atmospheric oxygen.
Wöhler’s siloxene also has certain parallels with carbon compounds. Specifically, it’s a cyclic substance whose molecules form the planar structure shown in Fig.2.
Depending on its degree of oxidation, Wöhler’s siloxene may appear colorless, yellow, orange, brown, or even black [15].
The mechanism of chemiluminescence here is similar to other cases: during oxidation of the substrate, an unstable chemical species is formed in an excited state. The energy of this excited state is partially emitted as light, which we see. Interestingly, it occurs under acidic conditions, whereas most chemiluminescent reactions involving organic (carbon-based) compounds take place in alkaline media.
Wöhler’s siloxene can serve as a luminescent indicator in analytical chemistry.
I believe most readers would agree that what may seem like an unremarkable substance — ordinary sand — has provided the material for many fascinating and surprising experiments.
References:
- [1] Berzelius J.J., Décomposition du fluate de silice par le potassium, Annales de Chimie et de Physique, 27, 1824, pp. 337-359 back
- [2] Davy H., Electro chemical researches, on the decomposition of the earths; with observations on the metals obtained from the alkaline earths, and on the amalgam procured from ammonia, Philosophical Transactions of the Royal Society, 98, 1808, pp. 333-370 back
- [3] Macioszczyk A., Hydrogeochemia, Wydawnictwa Geologiczne, Warszawa, 1987 back
- [4] Ples M., Chemiczny ogród (eng. Chemical Garden), w serwisie: www.weirdscience.eu, online: http://weirdscience.eu/Chemiczny%20ogr%C3%B3d.html [17.11.2016] back
- [5] Kutek J., Skały chemiczne i organogeniczne, in: Jaroszewski W., Przewodnik do ćwiczeń geologicznych, Wydawnictwa Geologiczne, 1978 back
- [6] Kawecka B., Eloranta P.V., Zarys ekologii glonów wód słodkich i środowisk lądowych, Wydawnictwo Naukowe PWN, Warszawa, 1994 back
- [7] Sękowski S., Pierwiastki w moim laboratorium, Wydawnictwa Szkolne i Pedagogiczne, Warszawa, 1989 back
- [8] Bielański A., Chemia ogólna i nieorganiczna, Wydawnictwo Naukowe PWN, Warszawa, 1981, pp. 508 back
- [9] Ehrlich P., Alkaline Earth Metals, in: Brauer G., Handbook of Preparative Inorganic Chemistry (vol. 1), Academic Press, Nowy Jork, 1963 back
- [10] Cao X.M., Ma R.N., Wu J.J., Wen M., Fan Y.Z., Du A., Influences of Si on corrosion of Fe–B alloy in liquid zinc, Corrosion Engineering, Science and Technology, 44(6), 2009, pp. 441-444 back
- [11] Kenny F., Kurtz R.B., Siloxene as Chemiluminescent Indicator in Titration, Analytical Chemistry, 22, 1950, pp. 693-697 back
- [12] Brandl H., Siloxen-Leuchtrakete, in: Wöhrle D., Tausch M.W., Stohrer W.D., Photochemie: Konzepte, Methoden, Experimente, Wiley-VCH, Weinheim, 1998, pp. 481-482 back
- [13] Ples M., Chemiluminescencja tlenu singletowego (eng. Singlet Oxygen Chemiluminescence), w serwisie: www.weirdscience.eu, online: http://weirdscience.eu/Chemiluminescencja%20tlenu%20singletowego.html [17.11.2016] back
- [14] Ples M., Światło z retorty (eng. Light from the Chemist’s Retort), Chemia w Szkole (eng. Chemistry in School), 5, 2014, Agencja AS Józef Szewczyk, pp. 33-34 back
- [15] Brandl H., Chemolumineszenz, w: Wöhrle D., Tausch M.W., Stohrer W-D., Photochemie: Konzepte, Methoden, Experimente, Wiley-VCH, Weinheim, 1998, pp. 255-258 back
All photographs and illustrations were created by the author.
Addendum
Below is a video demonstrating the silica reduction process and the pyrophoric nature of monosilane:
Marek Ples