tag:blogger.com,1999:blog-75634816952496101412024-03-18T23:38:04.792-05:00The Home ScientistDanhttp://www.blogger.com/profile/06495614313702847298noreply@blogger.comBlogger39125tag:blogger.com,1999:blog-7563481695249610141.post-42680096527797977052018-03-22T12:52:00.004-05:002021-07-30T21:36:50.663-05:00Generon Filter SupplementI recently had a tremendous opportunity to test out some cool technology on my channel: a nitrogen membrane filter by Generon. You can see that video here: <a href="https://www.youtube.com/watch?v=4-bbAwoG4Z8">Testing the Generon Nitrogen Membrane Filter</a>. You might want to watch the video first, to give context to the discussion below.<br />
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I got a lot of great comments on that video, and one in particular brought up some really great points that I wanted to take the time to address. Here's the full comment by piranha031091:<br />
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Let's address both of those points with some exciting... math!<br />
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<a name='more'></a>piranha's first point about lighting matches in an environment with a volatile fuel and enriched oxygen concentration is a good one. This can indeed be a very bad idea, for what I hope are obvious reasons. You need to carefully plan any experiment involving dangerous propositions like this. I was reasonably safe here for two reasons.<br />
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First, the fuel (ethanol) was not vaporized. My method was to soak a cotton ball (half a ball, actually) with ethanol (70% EtOH), place that in the container, and light this on fire to consume the oxygen in the contained volume. This meant that the ethanol was almost entirely liquid absorbed in the cotton, and very little had vaporized. Truly dangerous situations develop when the fuel is fully vaporized, which I believe is what piranha was referring to.<br />
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Second, I was well below the explosive limit of ethanol. Mixtures of combustible materials and fuels will only react within a well-defined concentration window. The lower end of this is the Lower Explosive Limit (LEL), which for ethanol and air is 3.3% (<a href="https://www.mathesongas.com/pdfs/products/Lower-(LEL)-&-Upper-(UEL)-Explosive-Limits-.pdf">source</a>). This means that below an ethanol vapor concentration of 3.3%, there is simply not enough fuel for an explosive situation to develop. In my experiments, I used roughly 2mL of 70% ethanol to saturate the cotton balls. The volume of the container was 630mL, so taking 3.3% of that means we need 20.8mL of <u><i>pure </i></u>ethanol <i><u>vapor</u></i> in order to be explosive. Since my fuel was a tenth of that volume, was only 70% concentration, and was not fully vaporized, I was well below the danger line here.</div>
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The situation with the enriched oxygen from the permeate stream, however, is potentially different. I could not find any data on LELs for ethanol in enriched oxygen environments. If we take Generon's expected output concentration of 40% oxygen and assume a linear relationship (which admittedly could be a very wrong assumption), that's still 1.7% or 10.7mL, far above what I used.<br />
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Side note: One other danger that he didn't mention but that I was very aware of is the container itself. Separatory funnels are not designed to withstand pressure - the wall thickness and shape are bad for pressure differences. Round bottom flasks, by contrast, are designed for pressure. They have thickened walls and the spherical shape distributes pressure evenly. The weird shape of a sep funnel could lead to stresses that could shatter the glass at high vacuum. When I was doing this experiment with the enriched oxygen from the permeate stream, I had on a face shield, heavy gloves, and a leather welding jacket (except, of course, for the one I recorded for the video. Whoops!). I believe I was reasonably safe, however, because I was only creating at most a half-atmosphere vacuum. I still wouldn't want to do this very often though, and risk repeated stresses building up to failure.<br />
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Now on to piranha's second point regarding CO<sub>2</sub> production. This is something that I had honestly not considered, and was not mentioned in any of the experimental setups I had read about. In retrospect, though, it's rather obvious. The equation for burning ethanol is<br />
<span style="color: orange;"><b>C<sub>2</sub>H<sub>5</sub>OH + 3O<sub>2</sub> == 2CO<sub>2</sub> + 3H<sub>2</sub>O</b></span><br />
As piranha said, 2 moles of carbon dioxide are produced from every 3 moles of oxygen consumed. Let's do some calculating to see if that affected my results. First, let's look at air. At 21% O2, there should be about 132mL of oxygen in the sep funnel's 630mL volume. 2/3 of that is 88.2mL, which is the amount of CO<sub>2</sub> produced. Using the Ideal Gas Law at my local temperature and pressure,<br />
<i><span style="color: lime;">PV = nRT</span></i><br />
<i><span style="color: lime;">(101 kPa)(0.0882 L) = n(8.3145 J/mol.K)(295 K)</span></i><br />
<i><u><span style="color: lime;">n = 0.0033 mol CO<sub>2</sub></span></u></i><br />
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The solubility of CO<sub>2</sub> in water at about 25 C is about 1.5 g/kg (<a href="https://www.engineeringtoolbox.com/gases-solubility-water-d_1148.html">source</a>). Using the volume of water drawn into the sep funnel in one of my air measurement runs, 154 mL,<br />
<i><span style="color: lime;">1.5 g/kg * (1kg / 1000g) * 154 mL = 0.231 g CO<sub>2</sub> soluble</span></i><br />
<i><span style="color: lime;">0.231g * (1mol / 44g) = <u>0.00525 mol CO<sub>2</sub> soluble in 154 mL water</u></span></i><br />
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Since 5 mmol is soluble in that amount of water, and we only produced 3.3 mmol, all of this gas is capable of dissolving. Thus, it should not contribute to the overall volume measurement in the end*.<br />
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Next let's consider the other end of the spectrum, the highest concentration I measured from the permeate stream. If we again assume Generon's expected 40% oxygen value, we expect there to be 252 mL in the 630 mL sep funnel volume. Following the same math as above, 2/3 of that is 168 mL which equates to 0.007 mol CO<sub>2</sub>. Using the highest water volume I measured, 296 mL, we find that 0.01 mol CO<sub>2</sub> is soluble. Thus 10 mmol is soluble and only 7 mmol was produced, so again this should all dissolve and not affect our volume measurements.<br />
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From all this we see that while a lot of carbon dioxide is produced when burning up the oxygen, it is all capable of dissolving in the amounts of water used in the experiment. There is also a lot of water produced during combustion, but that too will not contribute after it condenses (which piranha pointed out, and you can see in the video as fog on the inside of the glass).<br />
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Thanks very much to piranha031091 for taking the time to leave a detailed comment; this was a fun exercise to figure out, and I believe it helps to back up my results. Thanks to everyone else that comments as well. Your support really keeps me motivated to continue posting content!</div><div><br /></div><div>If you enjoyed reading this post and want to learn more, you might be interested in ChemTalk's article on <a href="https://chemistrytalk.org/ideal-gas-law/"><span style="color: #3d85c6;">The Ideal Gas Law</span></a>. They go into more detail on this equation, and have plenty of other articles on related topics like Charles' Law and Boyle's Law (two other very useful gas-related equations)!<br />
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*Practically speaking it may have had some effect, since it does take time for dissolution to occur. For best results, I should have waited a longer time and/or closed and shaken the sep funnel to encourage full dissolution. This also doesn't consider any CO<sub>2</sub> that may have already been dissolved, but I imagine this would be pretty minimal since it makes up only a tiny fraction of air.</div>
Danhttp://www.blogger.com/profile/06495614313702847298noreply@blogger.com2tag:blogger.com,1999:blog-7563481695249610141.post-46703192353784015872017-12-09T13:57:00.001-06:002018-03-22T10:01:30.724-05:00Boron ResearchAs mentioned in my new video, <a href="https://www.youtube.com/watch?v=4Y6H5p9wGh4">Elemental Extractions 5: Boron</a>, I've been doing a lot of research on the element and how to isolate it. My chosen reaction for that video (and my old one on boron) was the magnesiothermic reduction of boric oxide:<br />
<span style="color: orange;">B<sub>2</sub>O<sub>3</sub> + 3Mg == 2B + 3MgO</span><br />
But it turns out that this pathway only produces very impure boron, for several reasons. Read on for some of the results of my research, with links to the papers I mention. You can also see <a href="http://www.sciencemadness.org/talk/viewthread.php?tid=78684">my thread on ScienceMadness</a> for some extra discussion.<br />
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First, an overview of the overall reaction scheme. First, boric acid is dehydrated over a flame to form boric oxide:<br />
<span style="color: orange;">2H<sub>3</sub>BO<sub>3</sub> == B<sub>2</sub>O<sub>3</sub> + 3H<sub>2</sub>O</span><br />
The main reaction is a sort of Goldschmidt (thermite) between boric oxide and magnesium:<br />
<span style="color: orange;">B<sub>2</sub>O<sub>3</sub> + 3Mg == 2B + 3MgO</span><br />
However there are side reactions that produce various magnesium borides:<br />
<span style="color: orange;">Mg + 2B == MgB<sub>2</sub></span><br />
There are lots of different borides; the above is an example of one possible product.<br />
The post-reaction mass is ground up and subjected to dilute hydrochloric acid to react away the borides and other products to leave elemental boron.<br />
First, the water reacts with any leftover reactants:<br />
<span style="color: orange;">Mg (s) + 2H<sub>2</sub>O (l) == Mg(OH)<sub>2</sub> (s) + H<sub>2</sub> (g)</span><br />
<span style="color: orange;">B<sub>2</sub>O<sub>3</sub> (s) + H<sub>2</sub>O (l) == 2B(OH)<sub>2</sub> (aq)</span><br />
Then, the acid takes care of everything else:<br />
<span style="color: orange;">Mg(OH)<sub>2</sub> (s) + 2HCl (aq) == MgCl<sub>2</sub> (aq) + 2H<sub>2</sub>O (l)</span><br />
<span style="color: orange;">Mg (s) + 2HCl (aq) == MgCl<sub>2</sub> (aq) + H<sub>2</sub> (g)</span><br />
<span style="color: orange;">MgO (s) + 2HCl (aq) == MgCl<sub>2</sub> (aq) + H<sub>2</sub>O (l)</span><br />
<span style="color: orange;">Mg<sub>3</sub>B<sub>2</sub> (s) + 6HCl (aq) == B<sub>2</sub>H<sub>6</sub> (g) + 3MgCl<sub>2</sub> (aq)</span><br />
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The goal being to leave boron as the only remaining solid.<br />
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The whole focus of this new video is on purity. I wanted to make a nicer element sample than my first try. However I've been doing a lot of research recently and it appears that the magnesium reduction is really not a great path to follow if you want pure B.<br />
This paper has a good overview of this particular reaction scheme:<br />
<a href="http://www.academia.edu/4341344/Elemental_Boron_Magnesium_Boride_Synthesis_A_Review_jmsp">Elemental Boron & Magnesium Boride synthesis - Review</a> by Neale R. Neelameggham<br />
An excerpt:<br />
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<i>Moisson noted that one of the two magnesium borides formed is as resistant to water and hydrochloric acid as elemental boron causing difficulty in making pure boron. However, he was able to make 98.30% pure boron using several purification steps. The preparation of crude boron with a 90 to 95% purity practised even in the twenty-first century is called Moisson Process.</i></blockquote>
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Unfortunately that means that the acid digestion step will be <u>ineffective </u>at separating out all of the borides. I could find no other references to the Moisson Process he mentions, so I don't know what's involved with purification there.<br />
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From the Handbook of Inorganic Chemistry by Pradyot Patnaik:<br />
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<i>Boron does not react with water at ambient temperatures. The powdered amorphous form, however, reacts slowly at 100°C producing boric acid. The amorphous metal reacts slowly with dilute mineral acids at ambient temperatures; the crystalline form is inert. The former, however, reacts vigorously with concentrated nitric acid. The amorphous powder ignites in oxygen at 700°C.</i></blockquote>
Powdered amorphous B is what is produced in just about every type of synthesis, or at least the amateur-friendly ones. I considered trying other acids to get rid of borides, but the above shows that I'll be losing some B to the acid no matter which one I use, and hot sulfuric or nitric are particularly bad. So acid treatment is not ideal.<br />
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Finally I found one other reference: <a href="https://www.google.com/url?sa=t&rct=j&q=&esrc=s&source=web&cd=1&cad=rja&uact=8&ved=0ahUKEwiN2q63qO_XAhVCSiYKHS9kDjIQFggpMAA&url=http%3A%2F%2Fwww.boron-consortium.org%2Fassets%2Ffiles%2Fboronconsortium%2FSIEF%2FTechnological%2520process_Uses%2520Boron_27032013.doc&usg=AOvVaw0Q23hctBMgSYApDFTlQqHI">Uses and technological processes for member registrants of boron.doc</a><br />
This states (emphasis mine):<br />
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<i>Amorphous boron is produced in a Moisson process from boron oxide and magnesium. Both raw materials are mixed and ignited to start the strongly exothermic reaction. An excess of boron oxide is used as reaction moderator. The reaction block is cooled in a water basin and then crushed in a jaw breaker and a roller crusher. Hydrochloric acid is used to remove remaining magnesium oxide and incompletely reduced boron oxide from the raw boron. The excess of boron oxide is collected as boric acid. The boron is dried in fluidized bed driers to get boron with boron with 86%. <b>Using a higher excess of boron oxide another grade with 90% boron content is obtained.</b><br />The 86% boron grade can be cleaned further in a thermal treatment step with fluorides to convert insoluble by-products into soluble compounds which are again removed by washing with hydrochloric acid to obtain boron with 95-07% boron content.</i></blockquote>
So it's possible that using an excess of boric oxide may help with purity, although you still run into the problem of some borides not being digestible by acids.<br />
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On the topic of producing boric oxide for the reaction, it also turns out this is very difficult. Heating it all the way to B<sub>2</sub>O<sub>3</sub> is not easy, and I may in fact be producing <i>metaboric acid</i>, a polymeric species described as (HBO<sub>2</sub>)n. This would throw off the stoichiometry a bit. I did find a patent, <a href="http://www.google.com/patents/US4908196">US 4908196 A</a>, on boric oxide production that does it in stages:<br />
<blockquote class="tr_bq">
<i>More particularly, according to the present invention boric oxide is prepared by dehydrating boric acid using a method comprising:<br />(a) gradually heating the boric acid, with the absence or substantial absence of fusion phenomena, at increasing temperature to a value not exceeding about 150° C., operating at below atmospheric pressure, to eliminate water until the boric acid has been completely or substantially completely converted into metaboric acid;<br />(b) gradually heating the metaboric acid obtained in stage (a), with the absence or substantial absence of fusion phenomena, at increasing temperature to a value not exceeding about 400° C., operating at below atmospheric pressure, to eliminate water until the metaboric acid has been completely or substantially completely converted into boric oxide;</i></blockquote>
They even have a nice, detailed example section:<br />
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<i>EXAMPLE<br />Crystalline boric acid is used with a crystal size of the order of 100 microns and a purity of 99.9%. 4.5 kg of this acid are fed into a stainless steel reactor fitted with a stirrer, condenser, interspace for circulation of a diathermic fluid and a device for creating a pressure less than atmospheric.<br />While slowly agitating the solid mass, the reaction pressure is reduced to 20 mmHg and the temperature is raised to about 115° C., over a time of 30 minutes.<br />Abundant water vapour development commences at this temperature. The temperature is raised gradually over a time of about one hour to 140° C., the temperature at which water vapour development tends to substantially subside. Finally, the mass is heated to 150° C. at which heating is interrupted.<br />3150 g of orthorhombic metaboric acid are recovered from the reactor, with a yield of 98% with respect to the boric acid feed. This metaboric acid is heated rapidly while stirring to 150° C. in the apparatus used for the first stage, and the temperature is then gradually increased by 2° C. every 10 minutes to about 205° C. During this treatment there is an abundant development of water vapour, which tends to subside beyond 205° C.<br />The temperature is finally raised from 205° C. to 250° C. over a time of 30 minutes, at which heating is interrupted. 2370 g of boric oxide are recovered from the reactor with a purity of 99% and a yield of 93% with respect to the boric acid feed of the first stage.<br />The boric oxide obtained in this manner is in the form of a light, spongy friable mass which is removed without the need for mechanical action to separate it from the reactor walls.</i></blockquote>
So it might be worthwhile to try dehydrating my boric acid as I did previously, then taking the glassy product and subjecting it to higher heat in the furnace.<br />
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Finally, I found one more good book on boron: <a href="https://books.google.com/books?id=HFrsCAAAQBAJ&lpg=PA208&ots=BtiWjaV7ZT&dq=magnesium&pg=PA214#v=onepage&q=magnesium&f=false">Boron and Refractory Borides</a>, edited by V.I. Matkovich.<br />
Page 214 starts the section titled Methods of Preparation of Amorphous Boron. I can't copy the text directly, but here's a picture excerpt:<br />
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhQQ7_xj2ej11lGNHw3wvKkl4yyccH7lhxcCL8Q1Guxnkk5Y8xzC0RLN65uJXthOAsz4h-q2LvgRX21PFY4ZDh7Uj5joc4AzEZCsbQKfd1B5t93vCtLyV-WG3TTU58K-7nDsysWJMLUhX8/s1600/boron+1.jpg" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" data-original-height="740" data-original-width="580" height="320" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhQQ7_xj2ej11lGNHw3wvKkl4yyccH7lhxcCL8Q1Guxnkk5Y8xzC0RLN65uJXthOAsz4h-q2LvgRX21PFY4ZDh7Uj5joc4AzEZCsbQKfd1B5t93vCtLyV-WG3TTU58K-7nDsysWJMLUhX8/s320/boron+1.jpg" width="250" /></a></div>
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Lots of great information there to digest. The main point, I think, is that an excess of boric oxide is required. It even mentions my experience of the 'explosion' in the video: "Energetic interaction, characterized by a fast temperature increase and blaze appearance, takes place when the finely-divided reagents are heated to the reaction temperature." Neat!</div>
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Lastly, here's an excerpt about workup of the products:</div>
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEgS5PdsP2MKJl-jJ-skReLL_hTBCmlxBuEqHOtGfCX1Di3YZBdyXKm7Snj4gjUR7wQnKKj0QW2lZ6ekM16NItGbiLykBf4sHmp6MXohvxDU1QCb6pbFs4bvfvOh49TcPAQKJUgBlZB7G1E/s1600/boron+2.jpg" imageanchor="1" style="margin-left: 1em; margin-right: 1em; text-align: center;"><img border="0" data-original-height="722" data-original-width="914" height="252" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEgS5PdsP2MKJl-jJ-skReLL_hTBCmlxBuEqHOtGfCX1Di3YZBdyXKm7Snj4gjUR7wQnKKj0QW2lZ6ekM16NItGbiLykBf4sHmp6MXohvxDU1QCb6pbFs4bvfvOh49TcPAQKJUgBlZB7G1E/s320/boron+2.jpg" width="320" /></a></div>
That's it for now. All of this will inform my future boron experiments, and if I get a product of suitable purity I'll definitely make another video.Danhttp://www.blogger.com/profile/06495614313702847298noreply@blogger.com0tag:blogger.com,1999:blog-7563481695249610141.post-84415685689417407632017-02-10T17:14:00.001-06:002017-02-10T17:14:09.088-06:00Major Lab Renovation Part 3: Furniture!The renovation progresses! I received my official lab furniture and am in the process of installing it in the room. There are a few minor/major problems to resolve before I'm finished, though.<br />
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEjp523g_3WJ1BErgAIiF34ZCgiRjOl80ckaXrLXJBCt86WiJ4pDYMQD_pNxkeRhmq18H3m33H4g054MJKBFjvz9pwUo9JUMV7QI-DSciPPjVrxwsPckdTf7IS7WtXbiji_7qV8oS4s370M/s1600/20170208_155914.jpg" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="225" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEjp523g_3WJ1BErgAIiF34ZCgiRjOl80ckaXrLXJBCt86WiJ4pDYMQD_pNxkeRhmq18H3m33H4g054MJKBFjvz9pwUo9JUMV7QI-DSciPPjVrxwsPckdTf7IS7WtXbiji_7qV8oS4s370M/s400/20170208_155914.jpg" width="400" /></a></div>
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I ordered a total of four items: a nice lab table, two small "base cabinets," and a large "tall cabinet." All items were ordered from <a href="https://www.schooloutfitters.com/catalog/index">School Outfitters</a>, a supplier I found online after a long search for someone who would sell an individual some real lab furniture. The furniture is manufactured by Diversified Woodcrafts, and I have to say it's very high quality stuff.<br />
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Here are the base cabinets and table just after delivery, before assembly. (Sorry for the poor quality phone pictures.)<br />
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These items have solid black epoxy resin tops, exactly like the ones found in real school lab classrooms. In fact the vendor, as you might have guessed from the name, supplies this furniture to schools across the country. I actually have some of these exact pieces in my lab at work too! These epoxy tops are expensive, but very resistant to a wide range of chemicals as well as high temperature; perfect for lab work!<br />
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The minor problem here is that the lock on the bottom drawer of the left cabinet refuses to unlock, with any of the four sets of keys I was given. I am working with the vendor to resolve this; I think it's a defective lock. Otherwise these pieces look great and I can't wait to assemble them this weekend.<br />
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The final piece is the tall cabinet, and this thing is seriously massive.<br />
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEjCA6pA4DH0srguCXs-EP1LCvmkcI8pk9-9MgeDxp6SJ7ZGoewWOZIqILzQUrFYh1GXercPjyu2iA4bYErghhLPitkr3DJsksDsXsJpFfd1MyzqyLpOjjDGjgeYe_8Nv9CbQsKCCnh6gwg/s1600/20170208_205621.jpg" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="320" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEjCA6pA4DH0srguCXs-EP1LCvmkcI8pk9-9MgeDxp6SJ7ZGoewWOZIqILzQUrFYh1GXercPjyu2iA4bYErghhLPitkr3DJsksDsXsJpFfd1MyzqyLpOjjDGjgeYe_8Nv9CbQsKCCnh6gwg/s320/20170208_205621.jpg" width="180" /></a></div>
It measures 48" wide, 22" deep, and 84" tall. That's two feet wide and a full <b>seven feet tall</b>! I measured out the room and planned to fit it in between the two small cabinets, but here's the major problem: this monster will <u>not </u>fit through the hallway into the room. It's 5" taller than the door frame, and we can't tilt it over to fit it in because it's so wide that there isn't room in the hall to tilt it far enough. We even measured the window to try to get it through there, but it's just too big. This furniture is built incredibly solid and there's no way I can see to take it apart to put it back together in the room either.<br />
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So lesson learned: not only should you measure the space in the room for your new furniture, you also need to measure the hallway that leads to it!<br />
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I'm working with the vendor to figure out what to do. Returning it is going to be a nightmare in extra cost and effort, so currently the most promising option is putting this bad boy in the room next to the lab and ordering another small cabinet to fill the space it was supposed to fit in. We'll see though, and I'll post an update when I figure things out.Danhttp://www.blogger.com/profile/06495614313702847298noreply@blogger.com4tag:blogger.com,1999:blog-7563481695249610141.post-61587782674412899682016-12-31T22:29:00.003-06:002016-12-31T22:29:48.641-06:00Major Lab Renovation Part 2: Safety!The first pieces of my <a href="http://thehomescientist.blogspot.com/2016/12/major-lab-renovation-part-1-first-steps.html">new lab</a> have arrived: real, heavy-duty lab safety cabinets for flammables and corrosives!<br />
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEgVyroLK4s02gGS-0s0d0rtprP52f0bUJPszbS_2gStxlxiHFxKv3YQPPfNWoGxV3PqN2kbpd7Sh2XuE3o7sAnkbxmWZENmsuOiYy106FDmSNotNwroJMtcA8UjrHdjvkpIQeAKS5plb98/s1600/20161229_191012.jpg" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="225" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEgVyroLK4s02gGS-0s0d0rtprP52f0bUJPszbS_2gStxlxiHFxKv3YQPPfNWoGxV3PqN2kbpd7Sh2XuE3o7sAnkbxmWZENmsuOiYy106FDmSNotNwroJMtcA8UjrHdjvkpIQeAKS5plb98/s400/20161229_191012.jpg" width="400" /></a></div>
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Also, I'll be sealing the grout lines in the lab with this suggestively-named product:<br />
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhxUbmltRvWYhMqmEWLlyKpJ7j8liRUbz61CTnuITqJ2g1QAi0NunWcaCUxF627m_Sea7PwgnaVrIfa66Pb3w0Li9l_8lwVA2wXJksQ0EwNA2TgsNUTi0mjL_88HLQZVw0iyd76_B1RvvY/s1600/20161230_132116.jpg" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="225" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhxUbmltRvWYhMqmEWLlyKpJ7j8liRUbz61CTnuITqJ2g1QAi0NunWcaCUxF627m_Sea7PwgnaVrIfa66Pb3w0Li9l_8lwVA2wXJksQ0EwNA2TgsNUTi0mjL_88HLQZVw0iyd76_B1RvvY/s400/20161230_132116.jpg" width="400" /></a></div>
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This will help keep spills from staining (or possibly dissolving!) the grout. The tile is ceramic so it should be reasonably resistant to most things on its own.<br />
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Now to await the furniture!Danhttp://www.blogger.com/profile/06495614313702847298noreply@blogger.com2tag:blogger.com,1999:blog-7563481695249610141.post-80790594795191703252016-12-24T08:53:00.002-06:002016-12-24T08:53:47.412-06:00Major Lab Renovation Part 1: First Steps!My home lab has served me well for the past seven years on YouTube (can't believe it's been that long!), but now it's time for a major overhaul and renovation of the space. You may have already seen the <a href="https://youtu.be/WFJQvyQwj-A">announcement video</a> on my YouTube channel. Rather than post several small videos, I thought I'd document my progress here as things happen. I'll make a video to showcase the new setup once it's all in place, but that could take several months. So first, let's start with a tour of the lab space I've been using so far.
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My home chemistry lab is set up in one of the spare bedrooms of my house. With a single large window, it was probably intended to be an office space.<br />
My current table and storage cabinet is "garage work area" furniture. This is my main work area, and the cabinet holds some equipment and chemicals (mostly flammables, since this is at least a metal cabinet).<br />
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhYJE-ecPPzBzzdZfJdyhAfK-uomASEMfgTJz1KJV57sIevuKj12OBWcCDfzFkf45F6uLaytkl5YwfShmcG0MTolLYa8E-1jd9AUgg0m0yvhUZuIQpuXzD5tL3mOWmVj80B9rh6aq_RU2c/s1600/Renovation+1.png" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="179" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhYJE-ecPPzBzzdZfJdyhAfK-uomASEMfgTJz1KJV57sIevuKj12OBWcCDfzFkf45F6uLaytkl5YwfShmcG0MTolLYa8E-1jd9AUgg0m0yvhUZuIQpuXzD5tL3mOWmVj80B9rh6aq_RU2c/s320/Renovation+1.png" width="320" /></a></div>
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It looks nice, but isn't ideal. The tabletop is soft wood that dents easily, and has absorbed some spills over the years. You can also see I use up a lot of counter space! (I need to clean up more often!) The cabinet has some significant rust on the inside, probably from acidic vapors when I used to store acids in it. Some of the exterior metal is rusted as well.<br />
Above that is a shelf I installed with some extra equipment and assorted vials of small chemical samples.<br />
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On the opposite wall is my glassware storage area. These are simple wire racks from Home Depot.<br />
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEjTa0eEudC5n6g2jurIeJf534T2LrdI11ZR8-buWSdSC5sdDlM0SmRqJopofZIuOopWpvS6OuVoX4hyphenhyphenIUu1fVTu_k7Jnpddz-Y5eyeyxw8HN0ln2Y_x3HKhyphenhyphenaEluNG0rYQm-bCU_UnGv5g/s1600/Renovation+2.png" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="179" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEjTa0eEudC5n6g2jurIeJf534T2LrdI11ZR8-buWSdSC5sdDlM0SmRqJopofZIuOopWpvS6OuVoX4hyphenhyphenIUu1fVTu_k7Jnpddz-Y5eyeyxw8HN0ln2Y_x3HKhyphenhyphenaEluNG0rYQm-bCU_UnGv5g/s320/Renovation+2.png" width="320" /></a></div>
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The problem here is there just isn't enough storage space. I've got stuff strewn about all over the floor in front of the shelves.<br />
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Inside the closet is my main chemical storage area.<br />
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEjGfPy0W_WgquGHBfdJkj5WUvls3dgB6L5F54LDbWOMiAW3q32reDLwyh1FcQost8B94Wwsh5RlBOIW82lu-81jDlq0MHWJj15z2Toz58D0C0HVffpN3h-zjN8GBMSwYWqhRQub301sj6g/s1600/Renovation+3.png" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="180" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEjGfPy0W_WgquGHBfdJkj5WUvls3dgB6L5F54LDbWOMiAW3q32reDLwyh1FcQost8B94Wwsh5RlBOIW82lu-81jDlq0MHWJj15z2Toz58D0C0HVffpN3h-zjN8GBMSwYWqhRQub301sj6g/s320/Renovation+3.png" width="320" /></a></div>
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The wire rack is segregated by storage code, with general storage on the top shelf, oxidizers on the middle, and (mostly) corrosives on the bottom. To the left is my "safety corner" with a fire extinguisher and bowl of sand for fires, vinegar and baking soda for base and acid spills (respectively), and sodium thiosulfate to neutralize halogens. To the right is a Rubbermaid tub that holds all my acids. Inside is a layer of baking soda to absorb acid vapors and neutralize any leaks.<br />
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Another problem with the lab space you may have noticed is that the room is carpeted. Obviously this is not ideal for a lab; carpet can be stained easily, and tends to hold onto chemical odors. I've been using a plastic chair mat (for rolling chairs) to cover the area directly under my work space at the table, which has worked decently well.<br />
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So that's been my lab environment for many years now. But it's time to make a change. The first step: tile! Here's what the room looks like with everything removed:<br />
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhKm4kJCMXYFoSn7HZk8ZudblfU7cUCXI1X2t8EokCJQg8EGrvwY9SAWshoKedZYmuK0GEA0w9u-lqSQ8w05U52h_iXKxlCH7KEm_s896IsOSJxqlmKES1aYK8I9d-hcp60y3oBNU_PDyE/s1600/20161218_104115.jpg" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="180" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhKm4kJCMXYFoSn7HZk8ZudblfU7cUCXI1X2t8EokCJQg8EGrvwY9SAWshoKedZYmuK0GEA0w9u-lqSQ8w05U52h_iXKxlCH7KEm_s896IsOSJxqlmKES1aYK8I9d-hcp60y3oBNU_PDyE/s320/20161218_104115.jpg" width="320" /></a></div>
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And now, as of a few minutes ago, tile is in!<br />
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEjCzLgRx1Q29m35e_rfegae38O8zJ0HZTo2neRM_O32FxeoqxcbwKombm0BdhBV5SGs89UbAPmFFivqk-6qiqasb_Zwu7jm_Nv53X9FB7i_EFKcSMI_38lDt7JROtoFeK_JiDFYoV9t5-g/s1600/20161223_144517.jpg" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="180" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEjCzLgRx1Q29m35e_rfegae38O8zJ0HZTo2neRM_O32FxeoqxcbwKombm0BdhBV5SGs89UbAPmFFivqk-6qiqasb_Zwu7jm_Nv53X9FB7i_EFKcSMI_38lDt7JROtoFeK_JiDFYoV9t5-g/s320/20161223_144517.jpg" width="320" /></a></div>
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The furniture is on order, and will likely take a month or two to get in. Check back here for updates!Danhttp://www.blogger.com/profile/06495614313702847298noreply@blogger.com3tag:blogger.com,1999:blog-7563481695249610141.post-74980231547100457482014-11-26T12:47:00.001-06:002021-07-30T21:40:55.567-05:00Golden RainThis post is the video companion to <a href="https://www.youtube.com/watch?v=4xSUvSGywXo" target="_blank">this video</a>.<br />
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This experiment is hands-down the most beautiful chemistry demonstration I've ever seen. I'll be synthesizing lead(II) iodide, which has a beautiful golden yellow color. This experiment is especially striking because this bright yellow solid is produced from two water-clear solutions.<br />
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<b>Materials</b><br />
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<li><span style="color: orange;"><b>1.0 g Potassium Iodide, KI</b></span></li>
<li><span style="color: orange;"><b>0.8 g Lead Nitrate, Pb(NO<sub>3</sub>)<sub>2</sub></b></span></li>
<li><span style="color: orange;"><b>500 mL Distilled Water</b></span></li>
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<b>Reaction</b></div>
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<span style="color: lime;"><b>Pb(NO<sub>3</sub>)<sub>2</sub> (aq) + 2KI (aq) --> 2KNO<sub>3</sub> (aq) + PbI<sub>2</sub> (s)</b></span></div>
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<b><span style="color: red;">Warning!</span></b> Lead compounds, especially soluble lead compounds (like lead nitrate) are very toxic. These should only be handled by an experienced chemist with maximum safety precautions taken. I do the dangerous stuff so you don't have to!<br />
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The first step is to dissolve both reactants in distilled water - 250mL for each solution. Stoichiometrically, equal amounts of the two reactants are needed for this reaction. For safety reasons, I used an excess of potassium iodide to ensure all of the soluble lead is reacted away.<br />
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhAJgR-uvXcW691hQ8tNdxIy36o0zhMQSyi33rVaBmY4RIyl6NnGxKrnvz0NHJlZaThss8z-Uk6psdVri-g74Vm_9H7dD2JD0JG1-HA3uzfDp53Fwp3G_LSLfTcTxXvMQTl7522vkw60yM/s1600/Start.jpg" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="238" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhAJgR-uvXcW691hQ8tNdxIy36o0zhMQSyi33rVaBmY4RIyl6NnGxKrnvz0NHJlZaThss8z-Uk6psdVri-g74Vm_9H7dD2JD0JG1-HA3uzfDp53Fwp3G_LSLfTcTxXvMQTl7522vkw60yM/s1600/Start.jpg" width="400" /></a></div>
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The potassium iodide is extremely soluble, and will dissolve very quickly. Lead nitrate, on the other hand, will sometimes hydrolyze and leave behind a small amount of white precipitate (likely lead oxide or hydroxide). If this happens, a few drops of nitric acid added to this flask should bring everything into solution.</div>
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Now we simply add the two solutions together. I went a few drops at a time to start. As soon as the solutions touch, bright yellow lead iodide is produced. The crystals are so small that it looks silky. </div>
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Click the following pictures for much larger detailed views.</div>
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Lead iodide is slightly soluble at room temperature, so after the first precipitate is formed a bit of swirling dissolves it back to a clear solution. I then added the rest of the KI solution to the lead nitrate. Pouring the KI solution into the lead nitrate was actually another safety measure - I wanted to minimize handling the lead solution as much as possible.</div>
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhPD9NU1s_mwqyLWSz3mK8W_SpMIO4bFTTXmYO0ZspdkIVez40fFsrLDlUhTMhCHDrVJuwH1log1tFuKX1Fu4GSZDw0WU9Q5i7EneATI9dJQgTTVnuVL6q3I_IaTg7Igfa1Szd3OfpKzcs/s1600/lead+iodide.jpg" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="318" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhPD9NU1s_mwqyLWSz3mK8W_SpMIO4bFTTXmYO0ZspdkIVez40fFsrLDlUhTMhCHDrVJuwH1log1tFuKX1Fu4GSZDw0WU9Q5i7EneATI9dJQgTTVnuVL6q3I_IaTg7Igfa1Szd3OfpKzcs/s1600/lead+iodide.jpg" width="400" /></a></div>
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The fact that a bright yellow precipitate is produced from two clear solutions is striking enough, but the real magic of this demonstration is when it is precipitated from a hot solution. I took the flask containing the lead iodide and heated it to near boiling on a hot plate. It took a little while, but eventually all of the yellow crystals dissolved. This returned the solution to being completely water-clear. This happens because lead iodide is almost <u>ten times</u> more soluble in hot solution than in cold.</div>
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I then turned off the heat and allowed the flask to cool to room temperature on its own. This causes the lead iodide to precipitate back out of solution. This time, because of the slow cooling, it forms extremely thin hexagonal plate-like crystals that shine and glitter as they move around the solution. It's incredibly beautiful to watch - if you haven't already I highly recommend watching <a href="https://www.youtube.com/watch?v=4xSUvSGywXo" target="_blank">the video</a> to see it. This part starts at <u>5:37</u>. Here are a couple of the more beautiful photos I captured of it. Again, click the photos for a much larger view.</div>
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEgKDSxezbnzci3P-alKeQIY3hHWye6IO6MbOaRcw4d9lkk0f5QPAep0aDhY0ve5vtmmJd8DcDptMI4TYf49HT6Q6EilT_IjBRyPj8-s5LWTPBGT65C_nQocJVvKS0HEkCIp4K3G_nj73a0/s1600/rain+1.jpg" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="300" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEgKDSxezbnzci3P-alKeQIY3hHWye6IO6MbOaRcw4d9lkk0f5QPAep0aDhY0ve5vtmmJd8DcDptMI4TYf49HT6Q6EilT_IjBRyPj8-s5LWTPBGT65C_nQocJVvKS0HEkCIp4K3G_nj73a0/s1600/rain+1.jpg" width="400" /></a></div>
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After watching the crystals form and float around for a few hours, I placed the flask in the lab fridge overnight to push as much of the compound out of solution as possible. I then filtered it to isolate the lead iodide. All of the tiny crystals fused together on the filter paper to form a very beautiful sheet of golden iodide.</div>
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I very carefully collected this and placed it in a vial for display. The final yield was 0.8g PbI<sub>2</sub>.</div>
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<b><span style="color: orange;">Finally, cleanup and disposal of the waste solution is an important point.</span></b> We want to ensure that all of the soluble lead has been eliminated from everything it may have come into contact with. This can be done simply with baking soda (sodium bicarbonate). I made a hot, saturated solution of bicarbonate and added this to the waste solution that the lead iodide was filtered from. This caused the solution to go cloudy with a lead carbonate precipitate. The iodide is still marginally soluble, but the carbonate is truly insoluble. This precipitate can be filtered off and disposed of as regular garbage. The solution now contains only potassium nitrate and baking soda, and can be poured down the drain with plenty of water.</div>
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Anything that came in contact with the lead-bearing solutions should also be treated with baking soda in the same way. Any beakers, funnels, filter paper, stirring rods, etc. that may have droplets of soluble lead compounds must be cleaned up, You don't want to come back later and touch a tabletop contaminated with these toxic salts.</div><div class="separator" style="clear: both; text-align: left;"><br /></div><div class="separator" style="clear: both; text-align: left;">ChemTalk also wrote an article on this reaction, which you can read at <a href="https://chemistrytalk.org/lead-iodide-reaction/"><span style="color: #3d85c6;">Golden Rain Experiment</span></a>. They have another procedure on how to perform this reaction, plus some helpful tips and another video to watch!</div>
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Danhttp://www.blogger.com/profile/06495614313702847298noreply@blogger.com29tag:blogger.com,1999:blog-7563481695249610141.post-31451672838706290472014-03-28T17:32:00.000-05:002014-03-28T17:32:26.131-05:00Getting Started With Home ChemistryAs a part of my participation in the <a href="http://gulfcoastmakers.com/gulf-coast-makercon/" target="_blank">Gulf Coast MakerCon</a>, I wanted to write a short article to cover the basics of home chemistry, for those that want to know more. We'll go over where to start, how to find supplies, and a few excellent resources I use in my own experimental pursuits.<br />
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First and foremost...<br />
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<span style="font-family: inherit;"><b><span style="color: orange; font-size: large;">What is Home Chemistry?</span></b><br />Amateur chemistry is an amazing
exploration of the world around you through the science of chemistry. Learning
how things react together and how to separate them back into their components
is a great way to learn about the natural world. Home chemistry doesn’t
necessarily need expensive glassware and equipment – many amateur chemists are
perfectly happy working with measuring cups and mason jars. Others prefer the
accuracy and quality of real lab glass. Regardless, home chemists all share an
inquisitive spirit and an appreciation for science.</span></div>
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<span style="font-family: inherit;"><br />In
the media, chemistry has unfortunately been stereotyped as only being good for
making drugs and explosives. The truth is that there is much, much more to the
subject and far more interesting and safe things than these can be done at
home. Personally, I don't work with energetic materials in particular because I don't believe the payoff is worth the risk. Making something just to "blow something up" doesn't appeal to me at all. I feel even more strongly against any kind of drug manufacture. I much prefer reactions with beautiful colors, state changes, or elemental syntheses.</span></div>
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<b><span style="color: orange; font-size: large;">Where to Start?</span></b></div>
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<span style="font-family: inherit;">Because of the vast scope of the
science of chemistry, deciding where to start can be daunting. It is best to do
some research first and try to decide what subject interests you the most.
Organic chemistry, inorganic chemistry, aqueous reactions, electrolysis experiments, isolating elements
from their compounds, and crystal-growing are just a few areas of study. Next, I highly encourage you to find an experiment that you would like to try, and acquire the equipment and
chemicals necessary for that one experiment. By doing this over and over again, a very versatile
lab can be slowly built over time. This way, you don't spend a lot of money on things you may not use for a very long time, if ever, and avoid cluttering up your work space. This is how I built and continue to build my lab.</span></div>
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<span style="font-family: inherit;">A great place to start is copper
chemistry. There are many different and beautifully colored compounds of copper
that can be created very easily at home. Copper sulfate grows beautiful deep
blue crystals and is easily obtainable as a root killer in the plumbing section of your local hardware store. I made a very interesting and unique red copper compound <a href="http://thehomescientist.blogspot.com/2013/10/chevruels-salt.html" target="_blank">in this post</a>. You can even start from copper wire, which is hard <i>not</i> to find, if you've got the right materials to dissolve it.</span></div>
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<b><span style="color: orange; font-size: large;">Acquiring Supplies</span></b></div>
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<span style="font-family: inherit;">At first glance, this the hardest
part of practicing chemistry at home – finding the chemicals you need. However,
treasures can still be found locally and online with perseverance. Home chemists tend to become
compulsive label-readers; looking at every bottle to find out what it contains
and thinking about how it could be used. Hardware stores can be a great source.
Learning how to purify over-the-counter (OTC) materials is a very important
skill to develop. I have a <a href="https://www.youtube.com/playlist?list=PLKzSD6cYNohLTCjjbPyF9i4vop04pkQ87" target="_blank">playlist of videos</a> on this subject on my YouTube channel.</span></div>
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<span style="font-family: inherit;">There are also a surprising number
of online retailers. The big chemical companies like Alfa Aesar and Sigma Aldrich will not sell to individuals, but there are many others that cater to them
specifically. <a href="http://unitednuclear.com/" target="_blank">United Nuclear</a>, <a href="http://www.hms-beagle.com/" target="_blank">HMS Beagle</a>, bio-diesel and soap-making
warehouses, and even eBay are great places to start.</span></div>
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<span style="font-family: inherit;">Finally, be sure to consult state
and local regulations regarding possession of chemicals and chemistry equipment
and glassware. Some areas are more draconian with regards to what you legally can and cannot have. I've heard (but not verified myself) that Texas is about the worst state for chemistry in the US, going so far as to restrict the use of simple glassware!</span></div>
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<b><span style="color: orange; font-size: large;">Essential Supplies</span></b></div>
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<span style="font-family: inherit;">There are many different paths to
choose when practicing chemistry, and each has its own specialized equipment.
In general, though, this is a short list of some of the most useful tools and
chemicals you can have in your lab:</span><br />
<ul>
<li>Gram-weight scale</li>
<li>Hot plate or burner (alcohol or butane)</li>
<li>Stir plate and stir bars</li>
<li>Mortar and pestle</li>
<li>Test tubes</li>
<li>Acids (HCl, H<sub>2</sub>SO<sub>4</sub>, etc.)</li>
<li>Bases (NaOH, NH<sub>3</sub>, etc.)</li>
<li>Testing reagents (acid/base indicators, AgNO<sub>3</sub>, BaCl<sub>2</sub>, etc.)</li>
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<b><span style="color: orange; font-size: large;">Safety First!</span></b></div>
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<span style="font-family: inherit;">Home-chemistry
can be very interesting and rewarding
but this comes
with risks, both in storing
chemicals and experimenting with them.
Safety must always be the top priority in every experiment. Proper gloves,
goggles, and protective clothing are essential. Read the MSDS for each chemical
<u>before</u> you handle it. Store each chemical
properly according to category (i.e. do not put acids and bases in the same
cabinet). Though they can be dangerous, as long as chemicals are treated with
great <u>respect</u> the hobby can be enjoyed safely.
Remember to consider not only yourself, but your neighbors and the environment
too. Never pour chemicals down the drain unless they have been properly
neutralized and contain no chemicals harmful to the environment (heavy metals, carcinogens, etc.). Always keep a large amount of baking soda
and vinegar handy to neutralize acid and base spills, respectively. Think of the "worst-case scenario," and plan and prepare for it accordingly. Many landfills will gladly accept any
chemical waste as “household hazardous waste.” Finally, be sure to consult
state and local regulations for proper disposal of chemical waste products in your area.</span></div>
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<b><span style="color: orange; font-size: large;">Additional Resources</span></b></div>
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<span style="font-family: inherit;">Here are some excellent websites that
can be visited to learn much more about the subject:</span></div>
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<ul>
<li><a href="http://woelen.homescience.net/science/chem/misc/index.html" target="_blank">Science Made Alive</a></li>
<ul>
<li>Wilco Oelen hosts an amazing site with lots of great experiments. Follow the link to info on starting up your own lab.</li>
</ul>
<li><a href="http://www.sciencemadness.org/talk/" target="_blank">Science Madness Forum</a></li>
<ul>
<li>Arguably the best English-language amateur chemistry discussion board on the internet. The discussions are very advanced, but the wealth of knowledge among the members is staggering.</li>
</ul>
<li><a href="http://ehs.research.uiowa.edu/chemical-storage-nine-compatible-storage-group-system" target="_blank">Chemical Storage Guidelines</a></li>
<ul>
<li>An excellent overview of the proper way to store your chemicals.</li>
</ul>
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I hope this information has been helpful to you. If you plan on starting up your own home lab, and delving into the exciting world of home chemistry, I wish you the best of luck! Always remember to consider the safety of yourself and those around you, and enjoy the ride!</div>
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Danhttp://www.blogger.com/profile/06495614313702847298noreply@blogger.com8tag:blogger.com,1999:blog-7563481695249610141.post-62749077694497175952014-03-20T09:43:00.001-05:002014-03-20T09:43:41.800-05:00Gulf Coast Makercon 2014I'm very excited to announce that I will be participating in the <a href="http://gulfcoastmakers.com/gulf-coast-makercon/">Gulf Coast MakerCon</a> on April 5-6, in Tampa FL! This event is similar (but not affilliated with) the Maker Faire in California, and is made up of people that make or do interesting things. There's a huge variety of participants - 3D printing, robotics, custom jewelry, cosplay, tabletop gaming, and much more - so you're bound to find something that interests you! The MakerCon folks <a href="http://gulfcoastmakers.com/2014/03/10/meet-the-maker-daniel-flisek-mr-home-scientist/">featured me in one of their Meet the Maker segments</a> as well.
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I will be hosting a booth featuring my <a href="http://thehomescientist.blogspot.com/2013/02/the-element-display-introduction.html">Element Display</a>, which is now nearly fully stocked with samples! I only have about 13 left to collect. I will also have an outdoor area where I will be performing live chemistry demonstrations every few hours, including exciting experiments like thermite reactions. I haven't finalized my list of demos, but I'm leaning towards 3 per show, with the theme of wresting elements from their compounds.
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If you are in the area, I definitely recommend you try coming to this event. It promises to be very educational, exciting, and lots of fun! Come visit my booth and say hi!Danhttp://www.blogger.com/profile/06495614313702847298noreply@blogger.com5tag:blogger.com,1999:blog-7563481695249610141.post-35632260393218454952013-12-24T21:32:00.000-06:002013-12-24T21:32:05.382-06:00Oil of Wintergreen<div class="separator" style="clear: both; text-align: center;">
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">This post is the video companion to <a href="http://www.youtube.com/watch?v=VTLnNWQhSMI" target="_blank">this video.</a></span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">Today's experiment will be making a
minty-scented oil for the winter time: methyl salicylate, or oil of
wintergreen. This can be made simply from two very common and easy to obtain
chemicals, and with a very simple procedure. This is also one of my first
forays into organic chemistry!</span></div>
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<li><span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";"><b><span style="color: orange;">13.0 g aspirin, C9H8O4 (the actual tablets will weigh more!)</span></b></span></li>
<li><span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";"><b><span style="color: orange;">60 mL methanol, CH3OH</span></b></span></li>
<li><span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";"><b><span style="color: orange;">8 mL concentrated sulfuric acid, H2SO4</span></b></span></li>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";"><b><span style="color: lime;">1 mol aspirin + 2 mol methanol --> 1 mol methyl acetate + 1 mol methyl salicylate</span> </b></span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">The two reagents we will need are
ordinary aspirin and methanol. My source of methanol is HEET antifreeze, which
is almost entirely the alcohol. Make sure it is not the red version, Iso-HEET,
which is isopropanol.</span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">The first step is to crush the
aspirin tablets into a powder. I'm using 40 tablets, and each has 325mg of
aspirin. This equates to 13 grams of aspirin. The weight of the tablets
themselves will be greater than this, because each pill contains binders and
other inactive ingredients. These will not dissolve in any of our solutions,
but they do not interfere with the reaction and can be ignored. Filtration is
not necessary.</span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">The crushed aspirin tablets are
combined with 60mL of methanol in a conical flask, and swirled to dissolve as
much as possible. To this mixture, 8mL of concentrated sulfuric acid is added.
This only serves as a catalyst for the reaction, so this amount might actually
be a bit overkill.</span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">This final mixture is then placed on
a hot water bath at 60 C for at least one hour. I've done this before and the
yield is pretty small, so this time I let mine run for 4 hours.</span></div>
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<br /></div>
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhaaIYKHkV23ZsmQKB-liyjz_YSk7xETcIQxIs1T99A3gMFV4oolrZKGZk_6EA4w9EcC7kEzTW18KNPkF3MkL4TwRXGYgwzTbogYDZemBang6Ld83PnJC8vfLs2OiDSXU522Bm6FdFPU6o/s1600/2.JPG" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="184" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhaaIYKHkV23ZsmQKB-liyjz_YSk7xETcIQxIs1T99A3gMFV4oolrZKGZk_6EA4w9EcC7kEzTW18KNPkF3MkL4TwRXGYgwzTbogYDZemBang6Ld83PnJC8vfLs2OiDSXU522Bm6FdFPU6o/s320/2.JPG" width="320" /></a></div>
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<br /></div>
<div class="MsoNormal" style="line-height: normal; mso-margin-bottom-alt: auto; mso-margin-top-alt: auto;">
<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">I actually had to take the reaction
off heat after two hours to go have dinner, and then resumed heating the next
day for the remaining two hours. After this time, the temperature is raised to
about 80 C to vaporize most of the methanol. I heated mine until I reduced the
volume by about half. At this point, brown blobs of oil began to come out of
solution.</span></div>
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<br /></div>
<b><span style="background: yellow; font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";"></span></b><span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";"></span>
<div class="separator" style="clear: both; text-align: center;">
<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEjoRLzQ41i_IgNzoU1OazjrFhi-KGEMqqNyc2DuFoVlV0qlM9HAQoK2HZ21ms5sKGxqG5cX8UxL98hZVKmDiKs4UqY6vkcGkPDNsNS2j4o8caRzLny8hGnONk4xyfS_5MJhyOmTVW7AVXI/s1600/3.JPG" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="192" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEjoRLzQ41i_IgNzoU1OazjrFhi-KGEMqqNyc2DuFoVlV0qlM9HAQoK2HZ21ms5sKGxqG5cX8UxL98hZVKmDiKs4UqY6vkcGkPDNsNS2j4o8caRzLny8hGnONk4xyfS_5MJhyOmTVW7AVXI/s320/3.JPG" width="320" /></a></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";"><br /></span></div>
<div class="MsoNormal" style="line-height: normal; mso-margin-bottom-alt: auto; mso-margin-top-alt: auto;">
<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">I then took it off heating to cool.
In the meantime, I prepared the washing solutions for workup of the product.
This consists of two portions of 50 mL of distilled water, and a saturated solution
of baking soda (3.9 g in 50 mL water). These solutions were placed in the
fridge to cool down to near ice cold.</span></div>
<div class="MsoNormal" style="line-height: normal; mso-margin-bottom-alt: auto; mso-margin-top-alt: auto;">
<br /></div>
<div class="MsoNormal" style="line-height: normal; mso-margin-bottom-alt: auto; mso-margin-top-alt: auto;">
<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">For the workup, we need to separate
the oil from the rest of the mixture. This is easily done because while methyl
salicylate is very soluble in methanol, it is nearly insoluble in water.
Combining this solution with water causes all of the oil to come out of
solution, and washes away any water soluble impurities.</span></div>
<div class="MsoNormal" style="line-height: normal; mso-margin-bottom-alt: auto; mso-margin-top-alt: auto;">
<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">When cool, I transferred my reaction
mixture to a 250 mL separatory funnel. To this I added one of the 50 mL
portions of water, and shook the funnel vigorously for about 1 minute, venting often. The oil is heavier than water, and so will sit on the bottom. I drained this oil layer off, emptied the funnel of the aqueous wash water, and returned the oil to the funnel. I repeated this process a second time with the other portion of water, and then finally a third time with the baking soda solution. It is especially important to vent often when shaking this last mix - as it neutralizes any leftover acid, it produces carbon dioxide that can blow the stopper off!</span></div>
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<br /></div>
<div class="separator" style="clear: both; text-align: center;">
<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEg2JgVKaez3Ngnwz8moeo79V0dvEdIKFPPf6rIKx6ocbMSrbltvAj_lda_ZrVj1FJs-SB2uZbVdIvHkgTSttdwx3plf5_8sYxiTyB0rT_z3FBgNRB2v9voQEKQarjCKGr1k6U7UY14ZsaM/s1600/4.JPG" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="213" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEg2JgVKaez3Ngnwz8moeo79V0dvEdIKFPPf6rIKx6ocbMSrbltvAj_lda_ZrVj1FJs-SB2uZbVdIvHkgTSttdwx3plf5_8sYxiTyB0rT_z3FBgNRB2v9voQEKQarjCKGr1k6U7UY14ZsaM/s320/4.JPG" width="320" /></a></div>
<br /><div class="MsoNormal" style="line-height: normal; mso-margin-bottom-alt: auto; mso-margin-top-alt: auto;">
<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">After allowing to sit and separate for a few hours, I drained my washed product oil off and into a small vial for storage. I obtained 7.3 g of oil. Theoretical yield was 10.98 g, so this equates to a 66.5% yield. This batch of oil was very brown, but previous batches I've made were more of a tan color. This might be due to running the reaction for too long, or too hot, which could decompose some of the oil. It's also possible that interrupting the reaction and letting it sit overnight wasn't good for it either. You can see the comparison below - the oil from this run is on the left and a sample from a previous batch run for a shorter duration is on the right. Much less oil, but it looks more pure.</span></div>
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<br /></div>
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEjnXD_epIeQicBhomoO8dwblXtTX0qTKwm549taXCFhMjsyujHA-TKSkR6vMPWGd4sbI8CmIFWm9v3UpAfJocz4fncqhgk2DuaI3mTwugXNO7jfzkcDj4-ZAqqX_hqc2LrQCsKPY97Dq7w/s1600/5.JPG" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="191" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEjnXD_epIeQicBhomoO8dwblXtTX0qTKwm549taXCFhMjsyujHA-TKSkR6vMPWGd4sbI8CmIFWm9v3UpAfJocz4fncqhgk2DuaI3mTwugXNO7jfzkcDj4-ZAqqX_hqc2LrQCsKPY97Dq7w/s320/5.JPG" width="320" /></a></div>
<div class="MsoNormal" style="line-height: normal; mso-margin-bottom-alt: auto; mso-margin-top-alt: auto;">
<b><span style="background: yellow; font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";"></span></b><span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";"></span></div>
<div class="MsoNormal" style="line-height: normal; mso-margin-bottom-alt: auto; mso-margin-top-alt: auto;">
<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";"><br /></span></div>
<div class="MsoNormal" style="line-height: normal; mso-margin-bottom-alt: auto; mso-margin-top-alt: auto;">
<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">Finally, the waste rinse water can be neutralized with baking soda and is safe to pour down the drain with plenty of water.</span></div>
<div class="MsoNormal" style="line-height: normal; mso-margin-bottom-alt: auto; mso-margin-top-alt: auto;">
<br /></div>
<div class="MsoNormal" style="line-height: normal; mso-margin-bottom-alt: auto; mso-margin-top-alt: auto;">
<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">The oil looks a little gross, but smells great! It's perfect for oil burners (the kind with a little dish heated by a candle or something similar), but it should NOT be used for topical or food application. This crude oil likely has unreacted aspirin dissolved in it, and skin contact might lead to a concentrated dose of aspirin well beyond the recommended level.</span></div>
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<![endif]-->Danhttp://www.blogger.com/profile/06495614313702847298noreply@blogger.com2tag:blogger.com,1999:blog-7563481695249610141.post-37553261942429144652013-10-13T20:33:00.000-05:002014-03-26T14:45:32.448-05:00Chevreul's SaltThis post is the Video Companion to <a href="http://www.youtube.com/watch?v=CitNalVs01M" target="_blank">this video</a>.<br />
<u><b><br /></b></u>Chevreul's Salt is a little-known copper compound that is quite easy to prepare, and has a few very interesting properties. In this post, I'll go over in detail what I did.<br />
<br />
This experiment only requires two materials: copper sulfate and sodium metabisulfite. The former is sold at hardware stores as root killer for plumbing, and the latter is commonly found on eBay (since it is used by gold recovery people).<br />
<br />
<a name='more'></a>To start I made two solutions:<br />
<ul>
<li><span style="color: orange;"><b>2.4g CuSO<sub>4</sub></b><b><span class="mediumtxt"><span style="font-family: "Calibri","sans-serif"; font-size: 11pt; line-height: 115%; mso-ansi-language: EN-US; mso-ascii-theme-font: minor-latin; mso-bidi-language: AR-SA; mso-bidi-theme-font: minor-latin; mso-fareast-font-family: Calibri; mso-fareast-language: EN-US; mso-fareast-theme-font: minor-latin; mso-hansi-theme-font: minor-latin;"> · </span></span>5H<sub>2</sub>O in 25mL distilled water</b></span></li>
<span style="color: orange;"></span>
<li><span style="color: orange;"><b>2.0g Na<sub>2</sub>S<sub>2</sub>O<sub>5</sub> in 10mL distilled water</b></span></li>
</ul>
This yielded a light blue solution and a colorless one.<br />
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEh_1XOf18ioEeZvf7vN9rEO03sRVBkcmweN4Ottft6Tyf1j3SCERlQdYWZKZ4-wPT3yRxmu2t8Jj3rN7rvytwkiDm8TquDzruTVIAViYKX5hsKdPY-fyvCs_-yJbSXYnaPLMzz9iEdzjyM/s1600/1.jpg" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="288" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEh_1XOf18ioEeZvf7vN9rEO03sRVBkcmweN4Ottft6Tyf1j3SCERlQdYWZKZ4-wPT3yRxmu2t8Jj3rN7rvytwkiDm8TquDzruTVIAViYKX5hsKdPY-fyvCs_-yJbSXYnaPLMzz9iEdzjyM/s320/1.jpg" width="320" /></a></div>
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Combining the two immediately changes the color to a beautiful emerald green.<br />
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<br />
Next, the solution needs to be boiled for several minutes. The solution steadily darkens, and once it reaches boiling temperature a brick red solid precipitates. It is a very nice crystalline precipitate that easily and quickly settles to the bottom of the test tube. This is Chevreul's Salt - <span style="color: orange;"><span class="mediumtxt">Cu<sub>3</sub>(SO<sub>3</sub>)<sub>2</sub></span><span class="mediumtxt"><span style="font-family: "Calibri","sans-serif"; font-size: 11pt; line-height: 115%; mso-ansi-language: EN-US; mso-ascii-theme-font: minor-latin; mso-bidi-language: AR-SA; mso-bidi-theme-font: minor-latin; mso-fareast-font-family: Calibri; mso-fareast-language: EN-US; mso-fareast-theme-font: minor-latin; mso-hansi-theme-font: minor-latin;"> · </span>2H<sub>2</sub>O</span></span>. <br />
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEi4HbkAVRpSU4clqvMGVoBBiQytOGShFw9TFUzZ9lvJZD-LQvt6GjLu5u44v1duPHmMRMofi_xbqCJCHP2yRVstT8YS02hyphenhyphenrrp6NA2NgHaP6vGnphbY2W9pamv8BLX03hMItWsxsfJfo8A/s1600/3.JPG" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="320" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEi4HbkAVRpSU4clqvMGVoBBiQytOGShFw9TFUzZ9lvJZD-LQvt6GjLu5u44v1duPHmMRMofi_xbqCJCHP2yRVstT8YS02hyphenhyphenrrp6NA2NgHaP6vGnphbY2W9pamv8BLX03hMItWsxsfJfo8A/s320/3.JPG" width="222" /></a></div>
<br />
My yield was 1 gram of dark brick red crystalline powder.<br />
<br />
Chevreul's salt is interesting for two reasons. One is its color. Just about all copper compounds are green or blue in color, and there are very few exceptions to that. So it's neat to see copper in a color like this. Second is the fact that it is a mixed valence compound. Take a look at the formula for the salt above - you'll notice that the charge only balances if some of the copper atoms are in the +1 state and some are in the +2 state! I.E. there is a +4 charge divided over 3 copper atoms. Perhaps a more descriptive formula for this compound is <span style="color: orange;"><span style="font-family: inherit;">CuSO<sub>3</sub><span style="font-size: medium;">·</span>Cu<sub>2</sub>SO<sub>3</sub><span style="font-size: medium;">·2<span style="font-family: inherit;"></span></span>H<sub>2</sub>O</span></span>.<br />
<br />
Now for some testing of this compound.<br />
<br />
First, I wanted to show that Chevreul's Salt does indeed have both copper(I) and a copper(II) behaviors. The first test (left picture below) was to take a small spatula-full and add dilute hydrochloric acid. This precipitates copper(I) chloride as a white solid. I used dilute acid here because CuCl can redissolve in excess acid. The second test (right picture below) was to add a similar amount to concentrated ammonia solution. This dissolves the salt and changes color to a deep blue. This is the commonly-seen ammonia complex with copper(II): tetraaminecopper(II).<br />
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The next test was to distinguish this compound from the only other red copper salt I am aware of - copper(I) oxide. To do this, I took a small amount of my Chevreul's salt and added it to a test tube (left in the pictures below). To this I added concentrated hydrochloric acid, which releases sulfur dioxide gas. I drew this gas into a pipette and bubbled it through an acidified solution of potassium dichromate prepared in another test tube. (This solution was made by dissolving a tiny amount of K<sub>2</sub>Cr<sub>2</sub>O<sub>7</sub> into distilled water, and adding a few drops of dilute sulfuric acid.) The solution changed from yellow to green, indicating the presence of SO<sub>2</sub>.<br />
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhVXdPUEHGSHob2TX84FFVGuOAeM3_JuGmqwVbexwuMl5AyHoh8d0OUmqFr5Ls72D4QNadrOcl0KrPttCWd3mIN87zLp5785JJFOayPAS_83qAMrMtosxRTw0N1Z5Zdi8QrZI37MmWTmpA/s1600/10.jpg" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="285" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhVXdPUEHGSHob2TX84FFVGuOAeM3_JuGmqwVbexwuMl5AyHoh8d0OUmqFr5Ls72D4QNadrOcl0KrPttCWd3mIN87zLp5785JJFOayPAS_83qAMrMtosxRTw0N1Z5Zdi8QrZI37MmWTmpA/s400/10.jpg" width="400" /></a></div>
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This is the specific gas test for SO<sub>2</sub>, and proves that this is not Cu<sub>2</sub>O (since that has no sulfur).<br />
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Here's an awesome picture of the salt under a microscope, at 16x magnification. You can really see the crystalline nature, and some crystals are transparent! Click on the photo for a larger version.<br />
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I really enjoyed this experiment because of the range of colors it produced, and how simple it was to conduct!<br />
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Danhttp://www.blogger.com/profile/06495614313702847298noreply@blogger.com11tag:blogger.com,1999:blog-7563481695249610141.post-64754256384799696302013-09-18T21:39:00.002-05:002013-11-04T16:09:02.853-06:00Potassium Chlorate from BleachThis post serves as the Video Companion to <a href="http://www.youtube.com/watch?v=7MsHq_dUfnY">this video</a>. <br />
<br />
This experiment was also <a href="http://hackaday.com/2013/09/07/potassium-chlorate-from-household-products/">featured on Hack a Day</a>!<br />
<br />
In this simple experiment you can create potassium chlorate, a powerful oxidizer that finds use in amateur rocketry, a convenient source of oxygen, and the famous "screaming gummy bear" demo (among other things), from common household items with a minimum of effort. The tradeoff is that it's a very inefficient process and yields tend to be very low. Electrolysis is a far superior method, and is something I plan on trying out in the future.<br />
<a name='more'></a><br />
To begin, I measured out 500mL of household bleach. This is about the minimum for this method to produce reasonable (but still quite low) yields. The active ingredient in bleach is sodium hypochlorite, and the higher the concentration in yours the better. Mine was listed as 8.25%.<br />
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The good news for chemists is that when shopping for household chemicals to use in experiments, the cheaper the better! The cheapest brands generally do not include extra fancy chemicals like detergents and fragrances, which is desirable for chemistry use.<br />
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This solution was then boiled on a hotplate to get rid of most of the water, and to drive the conversion of hypochlorite to chlorate (via a disproportionation reaction, with table salt as a byproduct):<br />
<span style="color: orange;"><b>3NaClO == 2NaCl + NaClO<sub>3</sub></b></span><br />
This only occurs on boiling, so slow evaporation will not work in this experiment. Once crystals start to appear in the liquid, it is time to remove from heat and let cool.<br />
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These crystals are the sodium chloride byproduct, and need to be filtered off and can be discarded. The remaining solution contains sodium chlorate, as well as more sodium chloride.<br />
Sodium chlorate is quite soluble, so to recover it and separate it from the sodium chloride it must be converted to potassium chlorate. This can be done with another common household chemical: sodium-free salt alternative. Mine was No Salt brand, which is almost entirely KCl.<br />
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For this experiment, we need enough KCl to convert all the sodium chlorate to potassium chlorate. This happens via a simple metathesis reaction:<br />
<span style="color: orange;"><b>NaClO<sub>3</sub> + KCl == KClO<sub>3</sub> + NaCl</b></span><br />
Using 500mL of an 8.25% solution of sodium hypochlorite and the two equations above, I calculated that this reaction should only require 13.3g of KCl to go to completion. Other experimenters have recommended using a saturated solution of KCl that is equal in volume to the boiled bleach solution, so that is what I went with this time. Since my bleach ended up at about 200mL, I needed about 75g of KCl to make a saturated solution. The extra ingredients generally leave a solution of No Salt rather cloudy, so I always filter it before use.<br />
<br />
Next I combined the solutions together. This should precipitate KClO<sub>3</sub> crystals immediately, but I did not see any at all! I think the culprit was using that huge excess of KCl solution. This means that there is a lot of extra water around for the products to dissolve in, so even though potassium chlorate is much less soluble there was still enough solution for it to dissolve in. I believe that if I had used the stoichiometric amount of KCl in the minimum amount of water, 13.3g in about 39mL of water at room temperature, I think I would have seen immediate crystallization. Certainly there is room for future experimentation here.<br />
To get crystallization, I put the solution in the lab fridge over night. When I took the beaker out, there was a nice layer of white, plate-like crystals covering the bottom. I was able to recover 9 grams of product this way.<br />
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEgPUGl1Erw_KuyBvPoyykhX-bYsQViyHnknNclYiIBn2qghYhyphenhyphenMZIAaf2jIn4f0qPO8HnVl9FIhyphenhyphenoRfZGAdH_7Zu7fWLjy_aP-GxCqjkPbqjnaYtAI56ZGOrAJ_-1SLWfNUrQWwRy3-O8A/s1600/Capture3.JPG" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="275" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEgPUGl1Erw_KuyBvPoyykhX-bYsQViyHnknNclYiIBn2qghYhyphenhyphenMZIAaf2jIn4f0qPO8HnVl9FIhyphenhyphenoRfZGAdH_7Zu7fWLjy_aP-GxCqjkPbqjnaYtAI56ZGOrAJ_-1SLWfNUrQWwRy3-O8A/s320/Capture3.JPG" width="320" /></a></div>
22 grams of KClO<sub>3</sub> was the theoretical yield here, meaning I had only a 41% recovery. In an attempt to get more yield, I went ahead and boiled the remaining solution down to half its volume and placed it back in the fridge. I ended up with 23.5g of additional product, which by itself is greater than the maximum possible yield! That means that this second crop is very impure; perhaps I shouldn't have reduced the volume quite as much. These crystals looked much different too - they were much more compact, and some were needle-like rather than the nice plates from the first batch. Here is a comparison of the two, with the impure product on the right:<br />
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEjxyhNQylaQn4qrxNhfeWbHZ4IMC9u3CBuOnlmLJaS8LQYdat-QzaLHVR1vF2599bV-s1p4Ea4TVbUAZs5Z34ONPq56DsOaFZ7pk0g331JC2GyM6SF1xoNcr-LiTe7KSZjwMkBbQT1fKoI/s1600/Capture4.JPG" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="176" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEjxyhNQylaQn4qrxNhfeWbHZ4IMC9u3CBuOnlmLJaS8LQYdat-QzaLHVR1vF2599bV-s1p4Ea4TVbUAZs5Z34ONPq56DsOaFZ7pk0g331JC2GyM6SF1xoNcr-LiTe7KSZjwMkBbQT1fKoI/s320/Capture4.JPG" width="320" /></a></div>
There's still a chance to recover good product from this second batch, and it would require a recrystallization or two. I have yet to try this, but if I do I will update the post. It honestly may not be worth the effort, considering how cheap the reagents were and how simple the process is.<br />
<br />
Finally, I wanted to test the product to make sure it was in fact potassium chlorate. There is a simple and exciting test for this - combine some chlorate with half its weight of sugar, and add a drop or two of concentrated sulfuric acid. The mixture crackles and quickly bursts into beautiful lilac flames, making a lot of smoke as well. I tested 1 gram of my pure product combined with 0.5g of sugar, and two drops of acid.<br />
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEitIesS8AFVPrhBAlwyh39se5MFkpK6VIlFA4zun3watIErBrIRy3lV50VGLC2LIhFPWhtvBWMxJfbKZiAdRfooo2qzs_kEbqFi5dZoH1f-Si3wdFVpC0ZBgWMFrIx2FcvRsWPLQcz5wWU/s1600/Capture5.JPG" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="171" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEitIesS8AFVPrhBAlwyh39se5MFkpK6VIlFA4zun3watIErBrIRy3lV50VGLC2LIhFPWhtvBWMxJfbKZiAdRfooo2qzs_kEbqFi5dZoH1f-Si3wdFVpC0ZBgWMFrIx2FcvRsWPLQcz5wWU/s320/Capture5.JPG" width="320" /></a></div>
See the video link at the top to see this reaction in action. It appears near the end.Danhttp://www.blogger.com/profile/06495614313702847298noreply@blogger.com25tag:blogger.com,1999:blog-7563481695249610141.post-41194522644394097082013-08-12T12:54:00.001-05:002013-08-12T12:55:17.805-05:00'Video Companion' PostsRecently I read a comment thread elsewhere on the 'net where people were arguing the pros and cons of conveying information by video vs. by text. Some prefer video because you actually get to see what's going on, and some prefer text because they can peruse it at their own pace and it is in an easily searchable format. I think it's at least partially a case of visual learners vs. reading-writing preference learners, but I digress.
<br><br>
While I presonally prefer video format (as you may have guessed from there being much more content on my YouTube page) I can understand the points of both sides, and have decided to write blog posts here that I will call "Video Companions". These posts will include a basic write up of the experiment shown in one of my videos, to include things like procedures, formulas, necessary calculations, and any other observations I didn't point out in the video. I think these posts will help to reach a wider audience, enable people to explore the science more, and hopefully take away more from my experiments. I'll try to do this for all my new videos from here on out, and also add posts for older videos (probably progressing in order of their popularity). You can easily search for these types of posts by clicking the tag in the Categories side bar to the right.
<br><br>
So if you prefer text, rejoice and stay tuned for more! If you prefer video, feel free to stick to YouTube, but I encouorage you to visit here occasionally if you have any questions.Danhttp://www.blogger.com/profile/06495614313702847298noreply@blogger.com4tag:blogger.com,1999:blog-7563481695249610141.post-74315733718328617562013-02-03T18:04:00.002-06:002013-02-05T22:53:07.752-06:00The Element Display: LED Control Circuitry<!--[if gte mso 9]><xml>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-bidi-font-weight: bold; mso-fareast-font-family: "Times New Roman";">This post
will describe the electronics I used to drive and control the LEDs for my
element display. These were designed by my good friend and very skilled
electrical engineer Bill Porter (<a href="http://www.billporter.info/">www.billporter.info</a>)
and built by him and myself. To read about the project as a whole, see the <a href="http://thehomescientist.blogspot.com/2013/02/the-element-display-introduction.html">introductory post</a>.</span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-bidi-font-weight: bold; mso-fareast-font-family: "Times New Roman";">I’m not an
electrical engineer, but I will try to describe what I can here. Bill has a
much deeper knowledge of this sort of thing, and has posted a much more
in-depth article about how all this works on his website, which he has dubbed the <a href="http://www.billporter.info/?p=1750">Elemental Illuminator</a>.</span></div>
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEgUBnfCp4IWS5TmU5YZ2cGU9GroWc2g5T6MUgKddcJg5Hn6sWLrcfvwDo5H9Sb8A4dd7W08ocJoGWhA_fK07bdqksY6U4ImNzwvjK69pgO_6DbLiZNuToiOVYREuGOgqPGu6aybnyH-o3g/s1600/P1054354.JPG" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="300" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEgUBnfCp4IWS5TmU5YZ2cGU9GroWc2g5T6MUgKddcJg5Hn6sWLrcfvwDo5H9Sb8A4dd7W08ocJoGWhA_fK07bdqksY6U4ImNzwvjK69pgO_6DbLiZNuToiOVYREuGOgqPGu6aybnyH-o3g/s400/P1054354.JPG" width="400" /></a></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-bidi-font-weight: bold; mso-fareast-font-family: "Times New Roman";">My
original plan for the display’s illumination was to simply wire the whole thing
in a “constant on” fashion, so that I could “plug in” the display and
everything would turn on at once. Bill convinced me that a programmable display
would be much cooler, so he got to work and designed all the control circuitry
himself for my project. This consists of 9 microcontroller boards that the LEDs
wire into, which are arranged in series. At one end of the chain is a master
controller Arduino shield that Bill also designed. This is mounted to the bottom left front corner of the display - it appears on the front for easy upgrading and to facilitate Mode #4 below.</span></div>
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhDs9g0A3rPny1Ey9F0pcZDHkMF1vwQAIujIllKuRXSrg1v1xW0TWr8_TKKO7jA2RNVtGxTs-CZ_H0viKu6IlP0idyIZy3fLNQcekhzcFmBDSEIVPjRWxx8KhcGifADZX9bUp3JbajHW6M/s1600/P1124383.JPG" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="300" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhDs9g0A3rPny1Ey9F0pcZDHkMF1vwQAIujIllKuRXSrg1v1xW0TWr8_TKKO7jA2RNVtGxTs-CZ_H0viKu6IlP0idyIZy3fLNQcekhzcFmBDSEIVPjRWxx8KhcGifADZX9bUp3JbajHW6M/s400/P1124383.JPG" width="400" /></a></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">The
microcontroller boards on the back of the display can handle up to 16 LED strips, meaning theoretically I
would need 8 boards total to cover all 120 LED strips. I actually ended up
needing 9 boards, because of the backboard being broken up into sections. This
meant that all the LEDs for a circuit board had to be located on the same
section of backboard, so no wires would cross a gap. The layout is such that
there are only two electrical connections between adjacent backboard segments –
one for data and one for power. We used a phone cord for data transfer, and thick gauge speaker wire for power.</span></div>
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEiX03of1796Yc7EVnKn_-45S76ARE3KCl6xrCXbEHkxYxs_YvACVJogTrtrKyUbI73aYVMKPdE35K1ihSfDApSacXE-dlipfLoVvnaQXbpEgZ9ejDZHlNKk3RtP1NvA4QE9hOfBHPR35lk/s1600/P1054349.JPG" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="300" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEiX03of1796Yc7EVnKn_-45S76ARE3KCl6xrCXbEHkxYxs_YvACVJogTrtrKyUbI73aYVMKPdE35K1ihSfDApSacXE-dlipfLoVvnaQXbpEgZ9ejDZHlNKk3RtP1NvA4QE9hOfBHPR35lk/s400/P1054349.JPG" width="400" /></a></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">Currently there are four modes
programmed into the controller, switchable by power cycling the device. These
are most easily seen in the video posted in the<span style="mso-bidi-font-weight: bold;"> <a href="http://thehomescientist.blogspot.com/2013/01/the-element-display-introduction.html">intro post</a>, but I’ll describe them here as well. </span></span><span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">Fun fact: the master controller has
1 kB of RAM, and the current programming leaves just 5 bytes left over! </span></div>
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<span style="color: orange;"><b><span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">Mode 1: Slave</span></b></span><b style="mso-bidi-font-weight: normal;"><span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";"></span></b></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">This mode allows the master
controller to accept commands. For now, this mode does nothing. I plan on
writing a program to interface with the display in this mode that will allow me
to control the lighting through a GUI interface on a tablet PC. The overall
vision is to have a graphical periodic table displayed on the tablet, where the
user can tap a particular element, group of elements (alkali metals, noble
gases, etc.), or state of matter and have the corresponding elements light up
on the display. It could even have a “quiz mode” where a multiple choice
question will be asked, the potential answers will be illuminated on the
display, and after an answer is chosen the correct one will blink. Stay tuned!</span></div>
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<span style="color: orange;"><b style="mso-bidi-font-weight: normal;"><span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">Mode 2: “Breathing”</span></b></span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">This mode slowly fades the brightness
of every LED from zero to maximum, then back down to zero. It makes for a nice
stress test of the display, as it takes it to full power on all LEDs.</span></div>
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<span style="color: orange;"><b style="mso-bidi-font-weight: normal;"><span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">Mode 3: Twinkle</span></b></span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">This mode randomly selects an
element, a brightness value, and a time value, and fades the LED strip to the
chosen brightness over the chosen duration. This is done for 15 elements
simultaneously, to achieve a nice “starry sky” effect. The rest of the display
remains on at half brightness. This will be the default mode for display, because
the light is not too bright at any given time and the twinkle effect gives the
table a bit of motion. </span></div>
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<span style="color: orange;"><b style="mso-bidi-font-weight: normal;"><span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">Mode 4: Equalizer</span></b></span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">This is the fun one, and was a last
minute surprise addition by Bill. This converts the entire display into an
equalizer, allowing it to listen to ambient noise (via a small microphone on
the Arduino shield) and display the spectrum on the columns of the table. This
one you really have to see to appreciate, courtesy of Bill's channel:</span></div>
<br />
<br />
<iframe allowfullscreen="" frameborder="0" height="315" src="http://www.youtube.com/embed/OKnnKbogDro?list=UUna4B8jf9OgU3DLd8SCHYcA" width="560"></iframe><br />
<br />
<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">All
this was the most expensive single part of the whole project. The Arduino
master controller, the 9 microcontroller boards, and all the necessary wiring
came out to about a third of the total cost. The control this affords and the
possibilities it opens up for future software modifications, though, makes it
definitely worth the price. Check out the following links for more info on the
other portions of the project: <a href="http://thehomescientist.blogspot.com/2013/02/the-element-display-backboard.html">the backboard</a>, <span style="mso-bidi-font-weight: bold;"><a href="http://thehomescientist.blogspot.com/2013/02/the-element-display-acrylic-shelves.html">acrylic shelves</a>, and the <a href="http://thehomescientist.blogspot.com/2013/02/the-element-display-led-lighting.html">LEDs themselves</a>!</span></span>Danhttp://www.blogger.com/profile/06495614313702847298noreply@blogger.com0tag:blogger.com,1999:blog-7563481695249610141.post-23459214964743013452013-02-03T18:04:00.001-06:002013-02-03T22:30:02.253-06:00The Element Display: LED Lighting<!--[if gte mso 9]><xml>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-bidi-font-weight: bold; mso-fareast-font-family: "Times New Roman";">This post
will describe the LED lighting scheme I used to illuminate my element display,
including where I sourced the LEDs, how they are attached to the backboard, and
the table’s overall color scheme. To read about the project as a whole, see the <a href="http://thehomescientist.blogspot.com/2013/02/the-element-display-introduction.html">introductory post</a><span style="background: yellow; mso-highlight: yellow;"></span>.</span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-bidi-font-weight: bold; mso-fareast-font-family: "Times New Roman";"></span><a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEgA34bF7RLg89o4u9orLxr5GZSKYFftVnuz6ESiao0-rEhKs3FCHr_0PHKaKKV8u6MDO0vI7dCPdMOIeGjHBME2nLjIFzGFJyMSJd5AnmsisnHC5kVdOxLmzTQr-ePba__iW9yxVBuWS-U/s1600/P3254174.JPG" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="300" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEgA34bF7RLg89o4u9orLxr5GZSKYFftVnuz6ESiao0-rEhKs3FCHr_0PHKaKKV8u6MDO0vI7dCPdMOIeGjHBME2nLjIFzGFJyMSJd5AnmsisnHC5kVdOxLmzTQr-ePba__iW9yxVBuWS-U/s400/P3254174.JPG" width="400" /></a></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-bidi-font-weight: bold; mso-fareast-font-family: "Times New Roman";">I wanted
the display to be illuminated to give it a real “wow factor” and really bring
out the different sections of the periodic table. In my design, I placed the
LED lighting behind each individual shelf rather than behind where the element
actually sits. This way, the shelf itself is illuminated and transfers the
light throughout its thickness, which ended up looking really slick. The
downside to this is that the actual element samples can be hard to see when the
room lights are off, but that could be solved with some soft down lighting in
the display space.</span></div>
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<br /></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-bidi-font-weight: bold; mso-fareast-font-family: "Times New Roman";">The LED
lights were purchased as 5 meter rolls from Torchstar (<a href="http://stores.ebay.com/TorchStar">http://stores.ebay.com/TorchStar</a>),
an eBay seller with phenomenal prices. I also wanted to mention ElementalLED (<a href="http://www.elementalled.com/">www.elementalled.com</a>) here – they
provided a lot of great assistance to me in the planning stages, and I used
their LED strips to build my prototype designs. Their customer service and
depth of knowledge is really great. The pink LEDs were purchased from them (as a single strip), and
the rest from Torchstar (as rolls). No companies paid me to say these things – I am just
giving some positive feedback of the great experiences I had with each.</span></div>
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEiLyVKQt2dzhWxaBmSR0LRTn2Gpv9R_D4qQchD2-gjrbhIgAtRUbTxB4jVb5ZcTuy9Wy05gYH75ZOXlkdu9CtALzHU0HSgUXbfyzeT5Q2TvT4SoBDZ25hYOKHBM95jVr3Ae2t-cfrfRgHE/s1600/LED+Rolls.jpg" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="268" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEiLyVKQt2dzhWxaBmSR0LRTn2Gpv9R_D4qQchD2-gjrbhIgAtRUbTxB4jVb5ZcTuy9Wy05gYH75ZOXlkdu9CtALzHU0HSgUXbfyzeT5Q2TvT4SoBDZ25hYOKHBM95jVr3Ae2t-cfrfRgHE/s400/LED+Rolls.jpg" width="400" /></a></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-bidi-font-weight: bold; mso-fareast-font-family: "Times New Roman";"><br /></span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-bidi-font-weight: bold; mso-fareast-font-family: "Times New Roman";">The LED ribbons
are cuttable every 3 LEDs, or 2”. This is the perfect size to fit directly
behind each shelf, since shelves are 3” wide and the pins take up about ½” of
this space. Conveniently enough, these ribbons have an adhesive backing so they
are easy to stick to their proper places on the display. I first cut the rolls
into the individual 2” strips that I needed, and then soldered thin ribbon
cable to the contacts. Here's a video of my process - the first half is assembling the shelves (which I talk about in a separate post).</span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-bidi-font-weight: bold; mso-fareast-font-family: "Times New Roman";">You can see the solder dots on the left side of each strip in the picture below. This needed to be done 120 times – though there are only
118 elements, I have two extra shelves as placeholders for the lanthanide and
actinide series. I had to take extra care to keep their colors straight - there are no distinguishing markings at all on the strips themselves!</span></div>
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEjHX4emhTEtN5VPXaeWFPJdEVsvronqQkqZS6DQPDOwrGQ3BiQaMEh7eWrnqyaEvqsbDJ5cAfz-LeiamNsFZeL-0lWBYTYjf6ATNZTTvfu08O_MKvH2gM-075J3Ikx3453TftlVen1hQHE/s1600/PC133937.JPG" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="300" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEjHX4emhTEtN5VPXaeWFPJdEVsvronqQkqZS6DQPDOwrGQ3BiQaMEh7eWrnqyaEvqsbDJ5cAfz-LeiamNsFZeL-0lWBYTYjf6ATNZTTvfu08O_MKvH2gM-075J3Ikx3453TftlVen1hQHE/s400/PC133937.JPG" width="400" /></a></div>
<br />
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhasF4npym6gglk5DIfXytXsoM0gKpir3Iw2AAODAZ_l0vtgCXZ63OW4ttpQN8yeVnDur3CmSB9UYvuTgG2T-sGTbo6ZxIVRFB5rktNYzpgGkhKXMAAY3XHum2uclvLdkjsGqsyKmTnJNI/s1600/P1044043.JPG" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="300" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhasF4npym6gglk5DIfXytXsoM0gKpir3Iw2AAODAZ_l0vtgCXZ63OW4ttpQN8yeVnDur3CmSB9UYvuTgG2T-sGTbo6ZxIVRFB5rktNYzpgGkhKXMAAY3XHum2uclvLdkjsGqsyKmTnJNI/s400/P1044043.JPG" width="400" /></a></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-bidi-font-weight: bold; mso-fareast-font-family: "Times New Roman";">I used 6
colors because that was about the most variety I could find among single-color
LEDs of this type. I considered using RGB LEDs, but it would have increased
cost and complexity on an already very expensive and complex project. Looking
back on it and knowing more about the controllers I’m using, it sure would have
made for some amazing effects! I’ll probably go this route for Display 2.0 (if
that ever happens). The six colors I chose represent different sections of the
periodic table. There might be some differences from what you are used to,
because definitions vary and there are a few different designs of the periodic
table. As of now, there is no officially recognized, definitive table. I chose
a nice-looking candidate and based my color scheme off of that.</span></div>
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEgKxXLMG7M92F_fGWQh7TlSTFSJmd6QH6SqWUwQ9ZRQSq5WkfJFkWL48Ri7bOlSat52rxL6f4XKnwICe76AeEFlIn4EZsF5gG6HLVTKd8kVV4eRl3qWNPFxPLPXqakAVyOdfwC5BiPLXfY/s1600/periodic+table.png" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="300" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEgKxXLMG7M92F_fGWQh7TlSTFSJmd6QH6SqWUwQ9ZRQSq5WkfJFkWL48Ri7bOlSat52rxL6f4XKnwICe76AeEFlIn4EZsF5gG6HLVTKd8kVV4eRl3qWNPFxPLPXqakAVyOdfwC5BiPLXfY/s400/periodic+table.png" width="400" /></a></div>
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<br /></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-bidi-font-weight: bold; mso-fareast-font-family: "Times New Roman";">My colors
are as follows:</span></div>
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<br /></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-bidi-font-weight: bold; mso-fareast-font-family: "Times New Roman";">Orange –
alkali and alkaline earth metals</span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-bidi-font-weight: bold; mso-fareast-font-family: "Times New Roman";">Blue –
transition metals</span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-bidi-font-weight: bold; mso-fareast-font-family: "Times New Roman";">Red –
post-transition or “poor” metals</span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-bidi-font-weight: bold; mso-fareast-font-family: "Times New Roman";">Green –
nonmetals, metalloids, and halogens</span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-bidi-font-weight: bold; mso-fareast-font-family: "Times New Roman";">Pink –
noble gases</span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-bidi-font-weight: bold; mso-fareast-font-family: "Times New Roman";">White –
lanthanides, actinides, the placeholder spaces for both series, and hydrogen</span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">All 120 of these LED ribbons feed
their wires to the back of the display, where they connect with the LED control
circuitry mounted there. Each 3-LED strip requires 12V @ 20mA, so all the LEDs
at full illumination require 2.4 amps! Check out the following links for more
info on the <a href="http://thehomescientist.blogspot.com/2013/02/the-element-display-backboard.html">backboard construction</a>, <a href="http://thehomescientist.blogspot.com/2013/02/the-element-display-acrylic-shelves.html">acrylic shelves</a>,<span style="mso-bidi-font-weight: bold;"> and the <a href="http://thehomescientist.blogspot.com/2013/02/the-element-display-led-control.html">LED control circuitry</a>
portions of this project!</span></span></div>
Danhttp://www.blogger.com/profile/06495614313702847298noreply@blogger.com7tag:blogger.com,1999:blog-7563481695249610141.post-60064921898870807612013-02-03T18:04:00.000-06:002013-02-03T22:29:31.439-06:00The Element Display: Acrylic Shelves<!--[if gte mso 9]><xml>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-bidi-font-weight: bold; mso-fareast-font-family: "Times New Roman";">This post
will go into detail on how I assembled the individual acrylic shelves for my
element display, including where I ordered the plastic and the process of
attaching 240 shelf pins to the 120 shelves. To read about the project as a
whole, see the <a href="http://thehomescientist.blogspot.com/2013/02/the-element-display-introduction.html">introductory post</a>.</span></div>
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEj7amJ1O7T7qZIjXw28XY0Qg4mOjcjTOMH5Yhisctu3HJxqK9P_jalXlmO5QC4dTAM8QjBlMeGyMl-QQSpqi8MY3NBYntbyv_10x-xKHDXEj8784VMgpq-PpNC3V3NJkfuCVVzxsi6O7ko/s1600/P3234156.JPG" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="300" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEj7amJ1O7T7qZIjXw28XY0Qg4mOjcjTOMH5Yhisctu3HJxqK9P_jalXlmO5QC4dTAM8QjBlMeGyMl-QQSpqi8MY3NBYntbyv_10x-xKHDXEj8784VMgpq-PpNC3V3NJkfuCVVzxsi6O7ko/s400/P3234156.JPG" width="400" /></a></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-bidi-font-weight: bold; mso-fareast-font-family: "Times New Roman";">I designed
the element display to have individual acrylic shelves for each element on the
periodic table. The entire display had to fit in an area of my house that the
realtor called the “office nook.” The size of this space governed the
dimensions of the overall display. I won’t bore you with the details (because
it was <i style="mso-bidi-font-style: normal;">quite</i> involved!), but my
calculations led me to a per shelf size of 3” square as being the largest I
could accommodate. Each one is ¼” thick.</span></div>
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<div class="MsoNormal">
<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-bidi-font-weight: bold; mso-fareast-font-family: "Times New Roman";">After a
lot of searching around for a supplier, I settled on Ridout plastics (<a href="http://www.eplastics.com/">www.eplastics.com</a>) for their great quality
and reasonable price. I had used them to buy supplies for my prototype systems
too, so I knew they were good. Their business model is to sell plastic in large
sheets, which you pay one price for, and then you can pay extra for them to
make a number of cuts to this plastic. My sales rep worked with me to find the
smallest number of cuts for what I required. I ordered one large sheet (48" x 48" x 1/4") and had
them cut it into as many 3”x3” pieces as possible. </span><span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-bidi-font-weight: bold; mso-fareast-font-family: "Times New Roman";"> I ended up with 196 squares! </span><span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-bidi-font-weight: bold; mso-fareast-font-family: "Times New Roman";">This gave me a LOT of extra
pieces at no additional cost (since you buy the sheet and pay for once for a
range of cuts, and the extra cuts kept me within that range and thus didn’t
alter the price). The acrylic came with a protective brown paper covering for
both sides of each piece. And before you ask, no they did not pay me to say all
that – I’m just giving out some positive feedback to anyone else considering
ordering from them!</span></div>
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<div class="separator" style="clear: both; text-align: center;">
<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEgaN6cOptmlgtzV_AdczTm2ZX0-ZjCmVsI2irw5Fqnrxr5kpOnNxvgRxGUEbdCT1fz64dWhoAOSrQ9cICqKFCmQVgjE5Q9mqbkHedcwzEUCYGDsT3i9-7OJH7w_fnOzwRscgZ90FLOZLxs/s1600/P9243798.JPG" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="300" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEgaN6cOptmlgtzV_AdczTm2ZX0-ZjCmVsI2irw5Fqnrxr5kpOnNxvgRxGUEbdCT1fz64dWhoAOSrQ9cICqKFCmQVgjE5Q9mqbkHedcwzEUCYGDsT3i9-7OJH7w_fnOzwRscgZ90FLOZLxs/s400/P9243798.JPG" width="400" /></a></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-bidi-font-weight: bold; mso-fareast-font-family: "Times New Roman";">I also
went through every piece I received and sorted them based on quality of the cut
(I had enough extras where I could do this). Some cuts left crystal clear
edges, which I used preferentially over those that were a little rougher and thus
gave the edge a foggy look.</span></div>
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<br /></div>
<div class="MsoNormal">
<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-bidi-font-weight: bold; mso-fareast-font-family: "Times New Roman";">The
individual shelves were fairly straightforward to assemble. These acrylic
pieces were attached to clear plastic shelf pins, two per shelf, using Weld-On
#4 (also bought from Ridout). This is a special plastic adhesive that chemically
welds the two pieces together, affording a very strong and fast-setting bond. The
adhesive is primarily dichloromethane, which works by dissolving the plastic.
When the solvent evaporates, the plastic molecules of the two pieces intertwine
and in effect become a single piece. In fact, a minor accident showed me that
the shelf pins themselves will snap in half before the actual adhesive fails!</span></div>
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<div class="MsoNormal">
<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-bidi-font-weight: bold; mso-fareast-font-family: "Times New Roman";">For
consistency, I drilled two holes in a piece of wood at the correct distance
apart and used this template to assemble every shelf. That way I could ensure
the pins were always the same distance apart for every shelf, and that holes in
the backboard would line up with the pins (for the most part – some holes
needed to be fleshed out a bit to fit properly). While the adhesive was
setting, I used some tin bars I had lying around as weights to hold things
together. Here is a short video showing the process. The first half is the shelves, and the second involves the LEDs (which I will talk about in another post).</span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-bidi-font-weight: bold; mso-fareast-font-family: "Times New Roman";"><iframe allowfullscreen="" frameborder="0" height="315" src="http://www.youtube.com/embed/Usz1ZhX91ZM" width="560"></iframe> </span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">Once this process had been repeated
120 times (plus some spares that I also made), the shelves that the actual elements
would sit on were ready to be “plugged in” to the backboard. <span style="mso-bidi-font-weight: bold;">Though there are only 118 elements, I have
two extra shelves in the display as placeholders for the lanthanide and
actinide series. </span>Check out the following links for more info on the <a href="http://thehomescientist.blogspot.com/2013/02/the-element-display-backboard.html">backboard construction</a>, <span style="mso-bidi-font-weight: bold;"><a href="http://thehomescientist.blogspot.com/2013/02/the-element-display-led-lighting.html">LEDs</a>, and the <a href="http://thehomescientist.blogspot.com/2013/02/the-element-display-led-control.html">LED control circuitry</a> portions of
this project!</span></span></div>
Danhttp://www.blogger.com/profile/06495614313702847298noreply@blogger.com0tag:blogger.com,1999:blog-7563481695249610141.post-28744169446050572752013-02-03T18:03:00.001-06:002013-11-19T15:36:25.083-06:00The Element Display: Backboard Construction<!--[if gte mso 9]><xml>
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<span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-bidi-font-weight: bold; mso-fareast-font-family: 'Times New Roman';">This post will go into detail on how I constructed the backboard for my element display, including cutting it into sections to make it transportable, drilling the required 360 holes, and attaching the sections together in the final configuration. To read about the project as a whole, see the <b><span style="background: yellow; mso-highlight: yellow;"></span></b><a href="http://thehomescientist.blogspot.com/2013/02/the-element-display-introduction.html">introduction post</a>.</span></div>
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhZCmJo_V5IecGdKIcIBowFWLpWui7WO5FQ30O_65R7G-45bBpMHT1BzppSVkK-s38L8knwCdcwQ25yuHylDxkI3ABBrRVJBRyN2UeHz-EuTQng4jVkDITE7UjGE36SizKNTNGly6jbElw/s1600/Display1.jpg" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="300" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhZCmJo_V5IecGdKIcIBowFWLpWui7WO5FQ30O_65R7G-45bBpMHT1BzppSVkK-s38L8knwCdcwQ25yuHylDxkI3ABBrRVJBRyN2UeHz-EuTQng4jVkDITE7UjGE36SizKNTNGly6jbElw/s400/Display1.jpg" width="400" /></a></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-bidi-font-weight: bold; mso-fareast-font-family: 'Times New Roman';">The backboard is constructed of ½” thick plywood with a nice birch veneer (called “Sandeply” from Home Depot), and is split into four sections. Each split was located so as to not cut a shelf in half, and to break the board into roughly equal sections. These splits also governed the placement and wiring of the LED segments, which will be addressed in a later post. Splitting the board was done to make it easy to transport, as I mentioned above. The overall dimensions are 48” tall x 65” wide (about 4ft x 5.5ft). I first drew a grid over the whole front side, with the intersection points being where shelf holes would be drilled. Obviously not every point needed a hole, just the ones that formed the shape of the periodic table.</span></div>
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhp4O-XPvdj_xvZgJIjXvW5_1R2LjEecf3kfS7ReX45VHY0UzUCCnzSh9DbLhXHyIRkeVx8bycygelUyOQYaPiTgPCZEf1r1US2ug2XdeU9uRfF3lLeEYYbDob8e_Dfx9a1IZbPFtihCqA/s1600/PB033874.JPG" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="300" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhp4O-XPvdj_xvZgJIjXvW5_1R2LjEecf3kfS7ReX45VHY0UzUCCnzSh9DbLhXHyIRkeVx8bycygelUyOQYaPiTgPCZEf1r1US2ug2XdeU9uRfF3lLeEYYbDob8e_Dfx9a1IZbPFtihCqA/s400/PB033874.JPG" width="400" /></a></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-bidi-font-weight: bold; mso-fareast-font-family: 'Times New Roman';">I then carefully drilled out ¼” holes for all of the shelf pins for the 120 shelves needed (that’s 240 holes). After that each shelf also needed a single 1/8” hole for the LED wiring to go through, for an additional 120 holes (making 360 holes drilled total!). To guide me in drilling out these holes, I came up with a schematic detailing where the grid lines should be drawn and which intersection points actually needed holes. The colored squares represent each shelf, with the top two corners of each being where the mounting holes had to be drilled.</span></div>
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEgT_MEdXs_xusvqlgFxZKCrsJdYM6FZ-HvGXEMALdA22N1CYcfPs-DQ6iVskbEC6fOiVnAD-Gsm3MZq7oWzpSDF4RYH2bq5EhVkR6OHHxbX-JoVDM2ixq72tL6Jvz8xvA5P1AtJVwc1FZw/s1600/Schematic.jpg" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="255" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEgT_MEdXs_xusvqlgFxZKCrsJdYM6FZ-HvGXEMALdA22N1CYcfPs-DQ6iVskbEC6fOiVnAD-Gsm3MZq7oWzpSDF4RYH2bq5EhVkR6OHHxbX-JoVDM2ixq72tL6Jvz8xvA5P1AtJVwc1FZw/s400/Schematic.jpg" width="400" /></a></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-bidi-font-weight: bold; mso-fareast-font-family: 'Times New Roman';">I also used a spade bit to drill one final ½” hole in the bottom left corner of the display, to allow the master controller for the LEDs to be mounted on the front (with power and data cables passing through the hole to the back). This doesn’t appear in any of the pictures, because this one was not actually drilled until right at the end of the project.</span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-bidi-font-weight: bold; mso-fareast-font-family: 'Times New Roman';">Next the boards were painted, first with a gray primer and then two coats of flat black (both latex-based paints). Before painting, I plugged all the shelf holes with wooden dowels to prevent paint from dripping in and making the shelves not fit any more. The wiring holes I wasn’t worried about, because those didn’t require a precise size.</span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-bidi-font-weight: bold; mso-fareast-font-family: 'Times New Roman';">I also wanted a border for the whole display, to help hold the sections together and give the edges a nice finished look. For this I used a ½” aluminum channel that I cut to size.</span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-bidi-font-weight: bold; mso-fareast-font-family: 'Times New Roman';">Finally, the backboard pieces needed to be connected together with a system that makes them easy to uncouple when needed. For this I used three L-brackets per split, located at the top, middle, and bottom of the boards. The brackets are secured together with snugly-tightened nuts and bolts.</span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-bidi-font-weight: bold; mso-fareast-font-family: 'Times New Roman';">Finally, I wanted the entire display to be free-standing. To do this, I attached wooden triangles to each end of the display using screws carefully drilled into the side of each end piece. I also used L-brackets as seen above to secure them in place. Each triangle is attached with one screw and one bracket. Conveniently enough, I had some spare wood left over from the original sheet of Sandeply that I bought, so I cut this into a square and made another cut diagonally to create the two triangles. These were painted in the same way as the rest of the wood.</span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-bidi-font-weight: bold; mso-fareast-font-family: 'Times New Roman';">All this work formed a sturdy base construction for my element display, to which <a href="http://thehomescientist.blogspot.com/2013/02/the-element-display-acrylic-shelves.html">acrylic shelves</a>, <a href="http://thehomescientist.blogspot.com/2013/02/the-element-display-led-lighting.html">LEDs</a>, and all their associated <a href="http://thehomescientist.blogspot.com/2013/02/the-element-display-led-control.html">control circuitry</a> could be attached. See the relevant posts for more information!</span></div>
Danhttp://www.blogger.com/profile/06495614313702847298noreply@blogger.com0tag:blogger.com,1999:blog-7563481695249610141.post-86948211904275214442013-02-03T18:03:00.000-06:002013-02-21T19:06:58.355-06:00The Element Display: Introduction<br />
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<span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-fareast-font-family: 'Times New Roman';">For a very long time I have been interested in the hobby of element collecting. That is, collecting actual samples of elements from the Periodic Table. It's amazing to me that it is possible to possess all the building blocks for every piece of matter in the universe! This desire to start an element collection was actually what got me started as a home chemistry hobbyist in the first place. In fact, many of my experiments have the ultimate goal of isolating a pure element to add to my collection. While there are currently 118 elements discovered, it’s generally agreed that only the ones up to Uranium (92) are collectable (but it’s actually possible to get a few more than that, thanks to smoke detectors!). All elements heavier than Uranium, and a few lighter ones too, have no stable isotopes – they are radioactive and will eventually decay into lighter elements. That leaves a total of 95 samples to collect (including Am).</span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-fareast-font-family: 'Times New Roman';">I've steadily grown my collection over the years, and now it's become large enough where I wanted to build a nice display case for it. This project has taken me many months to design and construct, and has evolved through a lot of different concepts and iterations. The idea quickly grew from a simple shelving unit to one that included programmable LED lighting, shelves arranged in the shape of the Periodic Table, and the ability to be broken down easily for transport. This was a massive undertaking, and even the simplest alteration required tons of work to propagate through the entire display. It has been an incredibly rewarding and very fun experience, and has taught me a lot of handy skills like painting, soldering, and carpentry in the process. I spent three months designing it, built three prototype small-scale versions, and spent another 13 or 14 months constructing it off and on during my free time. Below is a video of the finished product and a summary of how I put it all together. Much more detail can be found in the following posts here on my blog.</span><span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-fareast-font-family: 'Times New Roman';"></span><br />
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<span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-fareast-font-family: 'Times New Roman';">The general idea was to have individual acrylic shelves for each element, in the shape of the periodic table, back-lit by LEDs, and mounted on a backboard. Here are just a few of the materials I used. This picture is a bit outdated because I replaced some things as the project developed; for example the $5 cheapo power supply that tripped </span><span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-fareast-font-family: 'Times New Roman';"><span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-fareast-font-family: 'Times New Roman';">my GFI circuit breakers every time I plugged it in,</span> was replaced with a much nicer laptop power supply.</span><br />
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<span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-fareast-font-family: 'Times New Roman';"><br /></span><span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-fareast-font-family: 'Times New Roman';">I built it as large as possible for the space I wanted to put it in, a small “office nook” in my house. This size restriction was what governed the size of each element’s shelf in the final design. I also wanted to be able to take it places to show it off at events at the local science museum, for example. To that end, I designed the display to be able to break down into sections for easy transport. Finally, I built in the capability for the display’s LED lights to be programmable and controllable from a separate device. This allows for lots of really cool effects that appear in the current version, and opens up the possibility for even greater software control in the future.</span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">Fun Facts:</span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";"><span style="mso-list: Ignore;">-<span style="font: 7.0pt "Times New Roman";">
</span></span></span><span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">Time spent designing: 3 months</span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";"><span style="mso-list: Ignore;">-<span style="font: 7.0pt "Times New Roman";">
</span></span></span><span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">Time spent building: ~1 year</span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";"><span style="mso-list: Ignore;">-<span style="font: 7.0pt "Times New Roman";">
</span></span></span><span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">Dimensions: <span style="mso-bidi-font-weight: bold;">48” tall x 65” wide (about 4ft x 5.5ft)</span></span></div>
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</span></span></span><span style="font-family: "Times New Roman","serif"; font-size: 12.0pt; mso-fareast-font-family: "Times New Roman";">Number of shelves: 120</span></div>
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<div class="MsoNormal" style="margin-left: 0.5in; mso-list: l0 level1 lfo1; text-indent: -0.25in;">
<span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-fareast-font-family: 'Times New Roman';"><span style="mso-list: Ignore;">-<span style="font: 7pt 'Times New Roman';"> </span></span></span><span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-fareast-font-family: 'Times New Roman';">Number of holes drilled: 361</span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-fareast-font-family: 'Times New Roman';"><span style="mso-list: Ignore;">-<span style="font: 7pt 'Times New Roman';"> </span></span></span><span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-fareast-font-family: 'Times New Roman';">Power drawn at maximum illumination: 12.0 V @ ~3 A (2.4A for LEDs, plus extra for the controllers)</span></div>
<div class="MsoNormal" style="margin-left: 0.5in; mso-list: l0 level1 lfo1; text-indent: -0.25in;">
<span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-fareast-font-family: 'Times New Roman';"><span style="mso-list: Ignore;">-<span style="font: 7pt 'Times New Roman';"> </span></span></span><span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-fareast-font-family: 'Times New Roman';">Total Cost: ~$1000 (not including the element samples themselves, or all the prototyping materials I went through)</span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-fareast-font-family: 'Times New Roman';"><span style="mso-list: Ignore;">-<span style="font: 7pt 'Times New Roman';"> </span></span></span><span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-fareast-font-family: 'Times New Roman';">Current Samples: 57 pure elements, 10 represented by radioactive decay products, and 2 representative compounds</span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-fareast-font-family: 'Times New Roman';">FAQs </span></div>
<div class="MsoNormal" style="margin-left: 0.5in; mso-list: l0 level1 lfo1; text-indent: -0.25in;">
<span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-fareast-font-family: 'Times New Roman';"><span style="mso-list: Ignore;">-<span style="font: 7pt 'Times New Roman';"> </span></span></span><span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-fareast-font-family: 'Times New Roman';">No, I can’t build one for you, sorry. It would be quite a lot of money to rebuild and would take many months to complete. And that’s just for the display unit – filling it with elements is another matter entirely (and one which even I haven’t completed). That being said,<span style="mso-spacerun: yes;"> </span><u>exceptionally</u> generous offers will be considered!</span></div>
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<span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-bidi-font-weight: bold; mso-fareast-font-family: 'Times New Roman';">This project can be divided into several pieces, which I will go into detail on in separate posts. These are </span><span style="font-family: "Times New Roman","serif"; font-size: 12pt; mso-bidi-font-weight: bold; mso-fareast-font-family: 'Times New Roman';"><a href="http://thehomescientist.blogspot.com/2013/02/the-element-display-backboard.html">the backboard</a>, <a href="http://thehomescientist.blogspot.com/2013/02/the-element-display-acrylic-shelves.html">acrylic shelves</a>, <a href="http://thehomescientist.blogspot.com/2013/02/the-element-display-led-lighting.html">LED lighting</a>, and the LEDs associated <a href="http://thehomescientist.blogspot.com/2013/02/the-element-display-led-control.html">control circuitry</a>. See these posts for more information!</span></div>
Danhttp://www.blogger.com/profile/06495614313702847298noreply@blogger.com4tag:blogger.com,1999:blog-7563481695249610141.post-42859336682607983812012-09-14T23:43:00.000-05:002012-10-28T21:32:46.583-05:00Titanium Thermite & Calcium SulfateTitanium metal from a thermite reaction has been one of my goals for a long time. I needed it for my element collection, and this thermite is particularly interesting because of the surprising chemicals it uses. The formulation I am using was developed by a fellow blogger and colleague on the Science Madness forum, Gert Meyers. His blog post on the subject can be found <a href="http://developing-your-web-presence.blogspot.com/2008/10/on-titanium-thermite-bronze-age-and.html" target="_blank">here</a>. The following will talk about my efforts to obtain good quality titanium metal from the thermite reaction he lays out. This is a re-post of the info I shared on Science Madness, including many pictures and 3 unlisted Youtube videos never before seen by the general public!<br />
<br />
<a name='more'></a><br />
<br />
<span class="mediumtxt">My source for calcium sulfate was DAP brand
plaster of paris, which is the hemihydrate (CaSO4 * 1/2H2O). The
ingredients listed on the back are
calcium sulfate, calcium carbonate, and silicon dioxide. Assuming the
latter two were only there in small amounts, I set off using this
plaster
straight out of the box. After heating some of it in my oven for 2 hours
at 425 F to dehydrate it, I mixed up a batch of thermite according to
blogfast’s specs:<br />
<br />
TiO2 – 30g<br />
CaSO4 – 25.5g<br />
Al – 27g<br />
CaF2 – 17.5g<br />
<br />
Here’s how it went: <a href="http://www.youtube.com/watch?v=KfqKOaZyllQ">http://www.youtube.com/watch?v=KfqKOaZyllQ</a> <br />
That video is actually my second try, which used the extra booster
formula (see the video’s description for specifics). As you can see, the
burn was
very “sputtery.” While it did produce titanium metal, the pieces were
small spheres that didn’t coalesce together like they should if the
whole thing
were molten. The slag was also very sparkly throughout, likely from tiny
pieces of Ti that were trapped before they could flow together. You can
see
in the picture at the end spheres of white titanium embedded throughout
the slag. Apparently the impurities, likely added as fire retardants or
to
improve its setting properties, were more detrimental than I thought.
Here’s the Ti I recovered from this batch:</span><br />
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhyFV-zW4afEtd1nsu5mUFugzbd4RuujmqR-3CxO_oQwWEMvu_kvm4MEzNCDcsv0cavzM8eD7Z_th9T-_fvaGXQJ-zzmgyGojQrwyFq_PX1XUd1DuvwaG1B9rbB3gBw8mwf6cggLGAdyn8/s1600/TiO2+product+1.jpg" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="240" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhyFV-zW4afEtd1nsu5mUFugzbd4RuujmqR-3CxO_oQwWEMvu_kvm4MEzNCDcsv0cavzM8eD7Z_th9T-_fvaGXQJ-zzmgyGojQrwyFq_PX1XUd1DuvwaG1B9rbB3gBw8mwf6cggLGAdyn8/s320/TiO2+product+1.jpg" width="320" /></a></div>
<br />
<span class="mediumtxt"> <span class="mediumtxt">I looked up the MSDS for this particular brand of plaster: <a href="http://www.dap.com/docs/msds/00071008_english.pdf">http://www.dap.com/docs/msds/00071008_english.pdf</a> <br />
Turns out, the impurities are pretty significant – CaCO3 at 15% – 25%, and SiO2 at 0% - 1%.<br />
<br />
So, to remedy this I tried removing some of the additives. The silica
isn’t worth the effort needed to remove it, but the CaCO3 is easy enough
to
destroy with a treatment in acid. As a bonus, by using sulfuric acid I
make more of my desired reagent!<br />
<br />
CaCO3 + H2SO4 == CaSO4 + H2O + CO2<br />
<br />
I took some of the raw plaster, submerged it in a good amount of water
(so it wouldn’t harden up on me), and added roughly 9M sulfuric acid
slowly
until the bubbling stopped. I filtered off the CaSO4, let it air dry,
and then dehydrated the powder in a fused silica crucible over my small
butane
burner for about 45 minutes. Using this in the same thermite composition
listed above, the reaction ran much smoother: <a href="http://www.youtube.com/watch?v=DSY4biEB88Y">http://www.youtube.com/watch?v=DSY4biEB88Y</a><br />
When recovering the metal afterward, the pieces were much larger and I
found 4 or 5 individual pieces about ½” in diameter. This is indicative
of a
much better reaction temperature, but there were still some problems
with this batch. There were a number of large voids in the slag, the Ti
metal was
incredibly hard to separate from the adhering slag, and the metal pieces
were much more tarnished.</span></span><br />
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhTTDmUvdRSa5rdkssLYXd4fIOuGo11vpJWvE9idZpmzJ3_Qvo4pIANWkxo35usGRkLoBzpNX345Lj7VwKM9ynLDtHVKvvNjC3mWkcf3bZjbal8-KlbXrIXFy45F6hymboO2JBXHRGtq60/s1600/TiO2+product+2.jpg" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="240" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhTTDmUvdRSa5rdkssLYXd4fIOuGo11vpJWvE9idZpmzJ3_Qvo4pIANWkxo35usGRkLoBzpNX345Lj7VwKM9ynLDtHVKvvNjC3mWkcf3bZjbal8-KlbXrIXFy45F6hymboO2JBXHRGtq60/s320/TiO2+product+2.jpg" width="320" /></a></div>
<br />
<span class="mediumtxt"><span class="mediumtxt"><span class="mediumtxt">My hypothesis is that all of these are due to
water of crystallization still in the CaSO4, i.e. I didn’t heat it up
enough. The water boiling off
would cause the voids, and I suspect it would react with the hot Ti
metal to tarnish its surface. The tarnish can be easily polished off, of
course,
but I feel that I can do better.</span></span></span><br />
<span class="mediumtxt"><span class="mediumtxt"><span class="mediumtxt"><br />
</span></span></span><br />
<span class="mediumtxt"><span class="mediumtxt"><span class="mediumtxt"></span></span></span><br />
<span class="mediumtxt"><span class="mediumtxt"><span class="mediumtxt"></span></span></span><br />
<span class="mediumtxt"><span class="mediumtxt"><span class="mediumtxt"><br />
=====================================================================<br />
</span></span></span><br />
<span class="mediumtxt"><span class="mediumtxt"><span class="mediumtxt"><br /> </span></span></span><br />
<span class="mediumtxt"><span class="mediumtxt"><span class="mediumtxt">My next attempt was to make my own CaSO4, so I know it’s pure (or at
least, I know for certain what’s in it!) I tried two methods to do this.<br />
<br />
<b>Method 1: CaCl2 + H2SO4 == CaSO4 + 2HCl</b><br />
This uses calcium chloride and sulfuric acid, both of which are fairly
OTC for me – CaCl2 is used in Damp Rid brand moisture absorbers, and my
sulfuric acid is from Liquid Fire brand drain opener. It also produces
hydrochloric acid as a byproduct, so this might be a handy way to get
much
higher purity acid than hardware store brand muriatic acid (which is
highly contaminated with, mostly, iron).<br />
<br />
<b>Method 2: MgSO4 + CaCl2 == CaSO4 + MgCl2</b><br />
I like this method a lot because it uses nothing hazardous at all, and
both reactants are very easy to source: magnesium sulfate is Epsom salt
and
calcium chloride is used in Damp Rid.</span> </span></span><br />
<br />
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhbAhPimCW7BJE_LqTlTrA6pnSM8Y3XXicdi3x8NiHjTMlhzZN730IIwDWDXd-_FcsfBygf2XOBOMmFlziAuQT3yIcCCtW5L2IGfhnYaNbBXPRRwGHlvEZkjaTCNY3NLfKk3CF1y1b2Pug/s1600/Damp+Rid+&+Epsom+Salt.jpg" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="240" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhbAhPimCW7BJE_LqTlTrA6pnSM8Y3XXicdi3x8NiHjTMlhzZN730IIwDWDXd-_FcsfBygf2XOBOMmFlziAuQT3yIcCCtW5L2IGfhnYaNbBXPRRwGHlvEZkjaTCNY3NLfKk3CF1y1b2Pug/s320/Damp+Rid+&+Epsom+Salt.jpg" width="320" /> </a></div>
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<span class="mediumtxt">I tried both methods, and they both work
exceedingly well. </span><span class="mediumtxt"><span class="mediumtxt">As long as you use dilute solutions of
everything (mine were < 5M), the precipitation
doesn’t happen immediately. That means it is much less likely to form
occlusions of impurities and you end up with a beautifully snow white
product. I
then heated both samples to 480F for two hours to dehydrate – here they
are fresh out of the oven:</span> <!--5m--><!--5m--></span></div>
<br />
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEiVWp4hMjWYGySRbHanNYrKACAg16CUeq1mUw0JJwv6iDeZWDbPIB-yvKx1nJXi8yIyuxZn5NqZBQZHvDvlclLgg6f6wxfaoxLz-IdcVQWJfUENJiZxD4zEUz-klPsJmxS61T2-fzV-ZNU/s1600/CaSO4+precipitated.jpg" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="240" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEiVWp4hMjWYGySRbHanNYrKACAg16CUeq1mUw0JJwv6iDeZWDbPIB-yvKx1nJXi8yIyuxZn5NqZBQZHvDvlclLgg6f6wxfaoxLz-IdcVQWJfUENJiZxD4zEUz-klPsJmxS61T2-fzV-ZNU/s320/CaSO4+precipitated.jpg" width="320" /></a></div>
<br />
<span class="mediumtxt">The Method 2 sulfate is on top, and Method 1 on
bottom. The latter was quite a bit more powdery and a little purer
white, which is interesting
considering it was made from acid that has a dark red color to it.<br />
I was able to use all the sulfate from Method 2 for a single 100g Ti thermite charge. The result can be seen here: <a href="http://www.youtube.com/watch?v=6TFty76NyHw">http://www.youtube.com/watch?v=6TFty76NyHw</a><br />
This one was a little slow to start, but burned fairly well once it got
going. The slag was especially nice, conforming to the bottom of the
flower
pot very nicely with only one small void – meaning it was completely
molten and (mostly) free of water. The pieces I recovered were about
halfway
between the first two reactions – larger spheres than the first run with
their somewhat shiny cast, but smaller than the lumps of the second.</span><br />
<span class="mediumtxt"> </span> <br />
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEiXahyJecNA-HBCI6prTmlnycB_ga-ZcT3-Jxh2WIUtDD_Amp8Vxzg90H0HB0mt8RihH3NZKlZ1qKSDtvNvkfR0tqLbcdrMiteRMXZHylPomnE2EzqGL9t92JpYcqoC-mZAjfY0A8Rwelc/s1600/TiO2+product+3.jpg" imageanchor="1" style="margin-left: 1em; margin-right: 1em;"><img border="0" height="320" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEiXahyJecNA-HBCI6prTmlnycB_ga-ZcT3-Jxh2WIUtDD_Amp8Vxzg90H0HB0mt8RihH3NZKlZ1qKSDtvNvkfR0tqLbcdrMiteRMXZHylPomnE2EzqGL9t92JpYcqoC-mZAjfY0A8Rwelc/s320/TiO2+product+3.jpg" width="240" /></a></div>
<br />
<span class="mediumtxt">So that’s where I’m at now. Ideally, I want to
produce large, shiny lumps of Ti for my element collection, but I think
these will do nicely for now.
If anyone else is trying this reaction, I hope my posting my efforts
here was helpful!</span> <br />
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Danhttp://www.blogger.com/profile/06495614313702847298noreply@blogger.com9tag:blogger.com,1999:blog-7563481695249610141.post-89917360917963010402012-03-09T13:23:00.002-06:002012-03-09T13:25:07.149-06:00Chemophobia: An Era of Discouraging Scientific Curiosity<span style="font-family: inherit;">In doing research for my experiments, I often run across articles in the news about police raiding suspected "home meth labs" or "amateur bomb making facilities." These are very often people's homes, which the authorities, for one reason or another, suspect are conducting illicit activity. Now I'm not against the police taking down such operations, on the contrary it's a good thing that these people get taken off the streets. However, in <em>many</em> cases they go overboard and barge in on amateur scientists like myself, who have no intention or capability of producing drugs or explosives at all, let alone in any sizable or dangerous quantity. If such search and seizures are to be conducted, those responsible need to have a <u>damn good reason</u> for doing so! It's quite clear if someone is engaging in drug production or explosives manufacturing vs someone who simply has an interest in science.</span><br />
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<span style="font-family: inherit;">As an example, the following article came to my attention today: </span><a href="http://toronto.openfile.ca/toronto/text/byron-sonne-thin-line-between-terrorist-and-gardener"><span style="font-family: inherit;">http://toronto.openfile.ca/toronto/text/byron-sonne-thin-line-between-terrorist-and-gardener</span></a><br />
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<span style="font-family: inherit;">I recommend you read the entire article, but here's the part I want to call attention to, the 'laundry list' of presumably scary chemicals with dangerous-sounding names.</span><br />
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<blockquote class="tr_bq"><div class="MsoNormal" style="margin: 0in 0in 0pt;"><span style="font-family: 'Arial','sans-serif'; font-size: 9pt;">No one disputes that Sonne had a lab in his basement, stocked with glassware and neatly labelled containers (see photos here). There was potassium permanganate, potassium nitrate, ammonium nitrate, iron oxide and zinc oxide. There was stearine, copper sulfate, urea, hydrogen peroxide and aluminum powder, as well as dextrin, sulfamic acid, hexachloroethane, charcoal, potassium silicate and sodium bicarbonate. Sonne had plastic bags full of wax shavings and PVC shavings, and a container of hexamine tablets next to his camp stove. There was acetone, methyl hydrate and hydrochloric acid in his garage. In his furnace room, he had an electrochemical setup where he seemed to be turning potassium chloride into potassium chlorate, a shiny white crystal that is, Anderson said, a well-known ingredient in improvised explosives like TATP (triacetone triperoxide) and HMTD (hexamethylene triperoxide diamlene). Most of these chemicals have multiple uses.</span></div></blockquote><br />
<span style="font-family: 'Arial','sans-serif'; font-size: 9pt;"></span><span style="font-family: inherit;">Multiple uses? That's a massive understatement. Here's what I came up with off the top of my head.</span> <br />
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<div class="MsoNormal" style="margin: 0in 0in 0pt;"><span style="font-family: 'Arial','sans-serif';"><span style="font-family: inherit;"><u>Potassium Nitrate</u>: common stump remover</span></span></div><div class="MsoNormal" style="margin: 0in 0in 0pt;"><span style="font-family: 'Arial','sans-serif';"><span style="font-family: inherit;"><u>Ammonium Nitrate & Urea</u>: fertilizers</span></span></div><div class="MsoNormal" style="margin: 0in 0in 0pt;"><span style="font-family: 'Arial','sans-serif';"><span style="font-family: inherit;"><u>Iron Oxide</u>: rust</span></span></div><div class="MsoNormal" style="margin: 0in 0in 0pt;"><span style="font-family: 'Arial','sans-serif';"><span style="font-family: inherit;"><u>Zinc Oxide</u>: used in sunscreen lotions</span></span></div><div class="MsoNormal" style="margin: 0in 0in 0pt;"><span style="font-family: 'Arial','sans-serif';"><span style="font-family: inherit;"><u>Stearin</u>: used in making tallow for candles and soaps</span></span></div><div class="MsoNormal" style="margin: 0in 0in 0pt;"><span style="font-family: inherit;"><span style="font-family: 'Arial','sans-serif';"><u>Copper Sulfate</u>: root killer for plumbing pipes, also used to make beautiful crystals, “</span><span style="font-family: 'Arial','sans-serif'; font-size: 9pt;">beakers of which were found during the search</span><span style="font-family: 'Arial','sans-serif'; mso-bidi-font-size: 9.0pt;">”</span><span style="font-family: 'Arial','sans-serif';"></span></span></div><div class="MsoNormal" style="margin: 0in 0in 0pt;"><span style="font-family: 'Arial','sans-serif';"><span style="font-family: inherit;"><u>Hydrogen Peroxide</u>: disinfectant found in nearly every household</span></span></div><div class="MsoNormal" style="margin: 0in 0in 0pt;"><span style="font-family: 'Arial','sans-serif';"><span style="font-family: inherit;"><u>Dextrin</u>: can be used as a pyrotechnic binder and fuel, aligning with his interest in rocketry</span></span></div><div class="MsoNormal" style="margin: 0in 0in 0pt;"><span style="font-family: 'Arial','sans-serif';"><span style="font-family: inherit;"><u>Wax & PVC Shavings</u>: used as binders for rocket engines, also aligning with his rocketry hobby</span></span></div><div class="MsoNormal" style="margin: 0in 0in 0pt;"><span style="font-family: 'Arial','sans-serif';"><span style="font-family: inherit;"><u>Sulfamic acid</u>: common ingredient in limescale and rust remover</span></span></div><div class="MsoNormal" style="margin: 0in 0in 0pt;"><span style="font-family: 'Arial','sans-serif';"><span style="font-family: inherit;"><u>Charcoal</u>: BBQ grills</span></span></div><div class="MsoNormal" style="margin: 0in 0in 0pt;"><span style="font-family: 'Arial','sans-serif';"><span style="font-family: inherit;"><u>Potassium Silicate</u>: a form of ‘water glass,’ which has a long list of innocuous uses (see: </span><a href="http://en.wikipedia.org/wiki/Water_glass#Uses"><span style="font-family: inherit;">http://en.wikipedia.org/wiki/Water_glass#Uses</span></a><span style="font-family: inherit;">)</span></span></div><div class="MsoNormal" style="margin: 0in 0in 0pt;"><span style="font-family: 'Arial','sans-serif';"><span style="font-family: inherit;"><u>Sodium Bicarbonate</u>: BAKING SODA (this one really makes it sound like they are trying to make him into the bad guy, by using the chemical name that not everyone will recognize)</span></span></div><div class="MsoNormal" style="margin: 0in 0in 0pt;"><span style="font-family: 'Arial','sans-serif';"><span style="font-family: inherit;"><u>Hexamine</u>: solid fuel tablets for camping stoves, found next to a camp stove (shocking)</span></span></div><div class="MsoNormal" style="margin: 0in 0in 0pt;"><span style="font-family: 'Arial','sans-serif';"><span style="font-family: inherit;"><u>Acetone</u>: paint thinner found in nearly every garage</span></span></div><div class="MsoNormal" style="margin: 0in 0in 0pt;"><span style="font-family: 'Arial','sans-serif';"><span style="font-family: inherit;"><u>Methyl Hydrate</u>: aka methanol, used as automotive antifreeze, found in his garage presumably near his car</span></span></div><div class="MsoNormal" style="margin: 0in 0in 0pt;"><span style="font-family: 'Arial','sans-serif';"><span style="font-family: inherit;"><u>Hydrochloric Acid</u>: readily available as a concrete etchant & pool chemical</span></span></div><div class="MsoNormal" style="margin: 0in 0in 0pt;"><br />
</div><div class="MsoNormal" style="margin: 0in 0in 0pt;"><span style="font-family: inherit;"><span style="font-family: 'Arial','sans-serif';">The only items I can't think of common household uses for are <span style="font-family: 'Arial','sans-serif';"><u>Potassium Permanganate</u>, <span style="font-family: 'Arial','sans-serif';"><u>Hexachloroethane</u>,<strong> </strong></span><span style="font-family: 'Arial','sans-serif';"><u>Aluminum Powder</u><span style="font-family: 'Arial','sans-serif';">, and his </span></span></span></span><span style="font-family: 'Arial','sans-serif';"><span style="font-family: 'Arial','sans-serif';"><u>Potassium Chlorate</u> electrolysis setup. The latter two have uses in pyrotechnics, so are conceiveably related to his rocketry interest. In any case, all are very useful reagents that would likely be used by someone interested in home chemistry like I am. I have 3/4 of these in my own lab, in fact, as well as the majority of the rest. I have no interest whatsoever in illicit activity, and I certainly don't have the amounts necessary to do anything major anyway (and by the looks of it, neither did he).</span></span></span></div><div class="MsoNormal" style="margin: 0in 0in 0pt;"><br />
</div><div class="MsoNormal" style="margin: 0in 0in 0pt;"><span style="font-family: inherit;"><span style="font-family: 'Arial','sans-serif';"><span style="font-family: 'Arial','sans-serif';">Take a look at the photos of his lab too, it's neater than mine! I know this isn't conclusive evidence, but generally speaking drug cooks and bombers don't keep a tidy space and rarely use real laboratory equipment and glassware.</span></span></span></div><div class="MsoNormal" style="margin: 0in 0in 0pt;"><br />
</div><div class="MsoNormal" style="margin: 0in 0in 0pt;"><span style="font-family: inherit;"><span style="font-family: 'Arial','sans-serif';"><span style="font-family: 'Arial','sans-serif';">Now obviously I don't know anything about his guilt or innocence, but it looks to me like he was just running a simple home laboratory because of a passion for science, just like I am. Now, because of this atmosphere of paranoia where anyone that does anything unique, intelligent, or otherwise out of the ordinary immediately gets the label of 'potential terrorist,' he must fight a lengthy legal battle and get lots of unwanted negative press. His neighbors will forever know him as 'that guy.'</span></span></span></div><div class="MsoNormal" style="margin: 0in 0in 0pt;"><br />
</div><div class="MsoNormal" style="margin: 0in 0in 0pt;"><span style="font-family: inherit;"><span style="font-family: 'Arial','sans-serif';"><span style="font-family: 'Arial','sans-serif';">It used to be that passers-by would see little Johnny outside building a potato cannon and think "Wow, that kid is smart! He's going places." Now, the first thing to come to mind is "Better keep an eye on him! He might be dangerous," and parents would rather their children stay inside and watch TV or play video games*. It's a sad time indeed when scientific curiosity is discouraged in such a way. </span></span></span></div><div class="MsoNormal" style="margin: 0in 0in 0pt;"><br />
</div><div class="MsoNormal" style="margin: 0in 0in 0pt;"><span style="font-family: inherit;"><span style="font-family: 'Arial','sans-serif';"><span style="font-family: 'Arial','sans-serif';">I could go on about this all day, but let's leave it at this: If you see your neighbors mixing liquids, soldering circuit boards, or building something in their garage, don't assume the worst! Go over and talk to them, find out what their interests are, and you might just learn something.</span></span></span></div><div class="MsoNormal" style="margin: 0in 0in 0pt;"><br />
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</div><div class="MsoNormal" style="margin: 0in 0in 0pt;"><span style="font-family: inherit;"><span style="font-family: 'Arial','sans-serif';"><span style="font-family: 'Arial','sans-serif';">*Full disclosure: I watch TV and play video games often myself :) I still leave time for science though!</span></span></span></div>Danhttp://www.blogger.com/profile/06495614313702847298noreply@blogger.com10tag:blogger.com,1999:blog-7563481695249610141.post-24387678862003647322011-08-10T15:38:00.003-05:002015-07-31T07:08:32.726-05:00Neodymium Metal from Hard Drive Magnets, Part I<span style="color: red;">Note: There's been a lot of work done on this subject on Science Madness, and the actual procedure I used is a bit different (specifically Step 4 & 6). Once I actually complete this experiment, I will update or replace this thread with the final method.</span><br />
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This is the first in a many-part series of posts detailing my efforts to isolate pure neodymium metal from hard drive magnets. This has been an extremely exciting project that I've really enjoyed undertaking, and the journey is just as fun as the final result. I'll be making a video on this eventually, but it's been such a long process that I thought people might be interested in reading some occasional updates on it in the meantime.<br />
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This first post will be about the theory, and will detail my reaction scheme.<br />
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Modern hard drive magnets are made of an alloy of three elements: Neodymium, Iron, and Boron (hence their other name, "NIB" magnets). This alloy approximately has the formula Nd<sub>2</sub>Fe<sub>14</sub>B. In order to get the neodymium separated, we must go through a number of different steps. You can't just melt the magnet because it's an alloy, so any physical separation techniques like that won't work. So, we need to use chemistry! The following steps are my proposed experimental outline.<br />
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1) <u>Demagnetize the magnet by heating it to it's Curie temperature, the temperature at which a magnet loses its magnetism.</u><br />
This can be done by holding the magnet in a propane torch for a few minutes. It's striking how quickly this happens - once the right temperature is reached, the magnet just falls off the tongs I used to hold it. Generally I heat it a bit more after that, because it does remain very slightly magnetic because of uneven heating. This step isn't strictly necessary, but it makes the next one a whole lot easier.<br />
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2) <u>Remove the protective outer nickel plating from the magnet, and break it into pieces.</u>These magnets are coated with a protective layer of nickel to protect it from the air, as the Fe and Nd metals would oxidize. We need to remove this coating for the next step. I've found it's easier to do this after breaking the magnet into a few pieces with pliers or a few gentle taps from a hammer. This gives you edges of the coating to grab onto.<br />
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3) <u>Dissolve the magnet pieces in sulfuric acid, and filter off the unreacted boron.</u><br />
This is where the chemistry begins. Dissolving the magnet in acid proceeds according to the following equation:<br />
Nd<sub>2</sub>Fe<sub>14</sub>B + 17H<sub>2</sub>SO<sub>4</sub> == Nd<sub>2</sub>(SO<sub>4</sub>)<sub>3</sub> + 14FeSO<sub>4</sub> + 17H<sub>2</sub> + B<br />
Nickel is inert towards sulfuric acid at normal temperatures and concentrations, so that's why it needs to be removed. From this equation, you can see that the boron is left unreacted. In this way, we can separate out one of the three components of the alloy. The Nd<sup>3+</sup> ion is lavender in color, so if your solution turns purple you're on the right track. If not, it's probably one of the older style magnets (like samarium cobalt) and doesn't contain Nd.<br />
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4) <u>Separate the neodymium sulfate from the iron sulfate.</u><br />
Here we can exploit a very interesting property of neodymium sulfate - it's inverse solubility profile. Nearly all salts dissolve better in hot solution than in cold solution (think of dissolving table salt in water: it goes much quicker if you heat it up). Neodymium sulfate, however, is the opposite: it actually dissolves better in cold solution than in hot. From Wikipedia's Solubility Table, we see that 13g of Nd<sub>2</sub>(SO<sub>4</sub>)<sub>3</sub> can dissolve in 100g of water at 0C, but only 1.2g can dissolve at 90C! Since this is the opposite behavior of the other salt in the solution, FeSO<sub>4</sub>, we can easily separate the two by heating. Neodymium sulfate precipitates out as pink or purple crystals when the solution is heated to near boiling, and the green iron sulfate remains in solution. This is a bit more complicated in practice, and I'll go into detail about that in upcoming posts.<br />
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5) <u>Convert the neodymium sulfate to neodymium fluoride.</u><br />
<span style="color: red;"><strong>CAUTION:</strong> This step is extremely dangerous, and I do not recommend you try this yourself.</span><br />
Soluble fluoride salts can be quite dangerous, and the fluoride ion is an insidious chemical. Burns from HF acid may not be immediately visible or painful, and symptoms of exposure can be delayed for up to 24 hours. I won't go into it here, but there's a wealth of information on the dangers of ammonium bifluoride online. I'll be working at maximum safety precautions for this step. Don't try this yourself at home.<br />
With that out of the way, I'll be reacting the recovered neodymium sulfate with ammonium bifluoride:<br />
Nd<sub>2</sub>(SO<sub>4</sub>)<sub>3</sub> + 3NH<sub>4</sub>HF<sub>2</sub> == 2NdF<sub>3</sub> (s) + 3NH<sub>4</sub>HSO<sub>4</sub> <br />
I'm not 100% sure this is the correct reaction, but the point is we end up with insoluble neodymium fluoride, which can be filtered off and collected. This must be done in a plastic container, because ammonium bifluoride (specifically, the hydrofluoric acid that's released) can eat through glass.<br />
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6) <u>React the neodymium fluoride with magnesium powder in a thermite-like reaction.</u><br />
This is where we finally arrive at neodymium metal:<br />
2NdF<sub>3</sub> + 3Mg == 2Nd + 3MgF<sub>2</sub><br />
This is why it was necessary to convert the sulfate to the fluoride - because elemental neodymium is such a reactive metal. Other compounds, like the oxides that are used in more traditional thermites, won't work because of neodymium's strong attachment to them. Using the fluoride here provides a high enough reaction enthalpy (i.e. heat) so that the reaction proceeds to completion and the products are in the molten state. The Nd metal will sink to the bottom of the mix, and the magnesium fluoride slag will provide a convenient crust to protect the desired product from air.<br />
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7) <u>Recover the Nd metal from the reaction products.</u><br />
If all goes well, the neodymium should have sank to the bottom of the reaction vessel and solidified in a single lump. After cooling, the whole mess of product will be placed in mineral oil and mechanically separated out. This should be done under oil because elemental neodymium will very quickly tarnish in air, similar to sodium.<br />
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So, those are the steps I'll be taking towards isolating neodymium metal for my element collection. As you can see from the process, it would be much cheaper (and faster) to simply buy some but that's not the point. The real goal here is to explore the chemistry, and try to recover something exotic out of something as ordinary as a magnet. Subsequent posts in this series will detail my efforts towards this goal, and I hope you find them as interesting as I have! Stay tuned!Danhttp://www.blogger.com/profile/06495614313702847298noreply@blogger.com11tag:blogger.com,1999:blog-7563481695249610141.post-29771407052368386422011-05-20T23:29:00.001-05:002011-05-20T23:31:41.021-05:00Copper Compound Conundrum, Part 3Copper Compound Conundrum Part 3: Testing the Precipitate - the exciting conclusion!<br />
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Now that I had determined the identity of the solution, it was time to figure out what the blue precipitate was. Remember from Part 2, this is what my mystery concoction looked like (in the beaker on the right).<br />
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I scooped some of the blue solid out of the bottom, placed it in a filter paper, and rinsed it with water to remove any of the green solution. Surprisingly, all the solids dissolved!<br />
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I ended up with this clear blue solution. My blue crystals going easily into solution like this lead me to believe that this must be the copper sulfate I was looking for. That was the theory, so time to test it! I transferred some solution into two test tubes.<br />
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<div class="separator" style="clear: both; text-align: center;"><a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEgDa257DuzqYh-OkGs67HyUHoGbTk2xuueJ8bx9yM5jcLs-OMRtBKv4UZrxCM3nPykMyxyxkbYoBolGneGz0SjA20sAlRsqDiDQlF0gcflrk_x3nChutMXRtWr9BTQ5fB5ymol9hirsMfk/s1600/Precip+test+start.jpg" imageanchor="1" style="clear: left; float: left; margin-bottom: 1em; margin-right: 1em;"><img border="0" height="240" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEgDa257DuzqYh-OkGs67HyUHoGbTk2xuueJ8bx9yM5jcLs-OMRtBKv4UZrxCM3nPykMyxyxkbYoBolGneGz0SjA20sAlRsqDiDQlF0gcflrk_x3nChutMXRtWr9BTQ5fB5ymol9hirsMfk/s320/Precip+test+start.jpg" width="320" /></a></div><br />
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I am holding a small piece of iron wire in front of the left tube - that will be test #1. If it is indeed copper sulfate, it should react with iron to displace copper metal:<br />
CuSO4 + Fe == FeSO4 + Cu<br />
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To the right test tube, I will perform an anion test, similar to what I did in Part 2. This time I am looking for the sulfate ion, so I made up a solution of barium chloride. If sulfate is present, a white precipitate of barium sulfate will form. This is a photo of the two tests taking place:<br />
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<div class="separator" style="clear: both; text-align: center;"><a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhQSxrdAiqckESbX7XqRJhkUUIwBXHsEcsokwN8qSEvKpPh1y_jnBSofBYnYSiLAtylU7AA7Qkqv1_Q_sGa_SB7SWc0Qq7zEGqvzeq2KJrUHm9bw1bQoX67q8ylM6ec6tqaCgi0QvM3Q64/s1600/Precip+test.jpg" imageanchor="1" style="clear: left; float: left; margin-bottom: 1em; margin-right: 1em;"><img border="0" height="240" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhQSxrdAiqckESbX7XqRJhkUUIwBXHsEcsokwN8qSEvKpPh1y_jnBSofBYnYSiLAtylU7AA7Qkqv1_Q_sGa_SB7SWc0Qq7zEGqvzeq2KJrUHm9bw1bQoX67q8ylM6ec6tqaCgi0QvM3Q64/s320/Precip+test.jpg" width="320" /></a></div><br />
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We can see the piece of iron on the left is visibly reddish - this is due to copper plating out of solution as iron ions enter it. On the right, we see a white precipitate that immediately forms when the barium chloride solution is added, which looks just like Part 2's test.<br />
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Testing complete:<br />
<div class="separator" style="clear: both; text-align: center;"><a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEid9izoBDh8iICfDdpmvrFavj_WllQyhfPacYybqznUUK1uRJKDcTQ3Tl65ugdtXASeN1DT3NCmU8XrAFet_20zDMnKXphNICfoxkJiyJnc2zxQrLsT63BKrzwiLkP126ufwhDXnc9nxD0/s1600/Precip+test+results.jpg" imageanchor="1" style="clear: left; float: left; margin-bottom: 1em; margin-right: 1em;"><img border="0" height="240" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEid9izoBDh8iICfDdpmvrFavj_WllQyhfPacYybqznUUK1uRJKDcTQ3Tl65ugdtXASeN1DT3NCmU8XrAFet_20zDMnKXphNICfoxkJiyJnc2zxQrLsT63BKrzwiLkP126ufwhDXnc9nxD0/s320/Precip+test+results.jpg" width="320" /></a></div><br />
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After a few minutes, quite a bit of copper has come out of solution and plated onto the wire in the left test tube. Notice also that the blue color of copper(II) ions has disappeared as iron(II) ions have replaced them. A white cloudiness of suspended barium sulfate is in the right test tube.<br />
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<div class="separator" style="clear: both; text-align: center;"><a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEjVhSUyhTdUenZvScPFzZcfNGueqrNNGHzN1CGJUgpM_KvzVMb2YidbpNYflI8r24glHV2vjoXa_xwF3F5OLlMiip5J8eUPy3M98e2JX7mFs8SshmeVeJvurA6jUd6nREiTVomfIHbFrbE/s1600/Barium+Sulfate.jpg" imageanchor="1" style="clear: left; float: left; margin-bottom: 1em; margin-right: 1em;"><img border="0" height="240" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEjVhSUyhTdUenZvScPFzZcfNGueqrNNGHzN1CGJUgpM_KvzVMb2YidbpNYflI8r24glHV2vjoXa_xwF3F5OLlMiip5J8eUPy3M98e2JX7mFs8SshmeVeJvurA6jUd6nREiTVomfIHbFrbE/s320/Barium+Sulfate.jpg" width="320" /></a></div><br />
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Here's another look at the barium sulfate precipitate. Again, the blue color of the solution has faded somewhat as a dilute solution of copper(II) chloride has been formed.<br />
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<u><b>Conclusion</b></u><br />
The blue solid is indeed copper(II) sulfate.<br />
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So we've solved it - step 4 in my initial reaction scheme produced a concentrated solution of copper(II) chloride and a precipitate of copper(II) sulfate. Apparently, copper sulfate is not soluble (or maybe slightly soluble) in a copper chloride solution.<br />
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What's strange is that the chloride appeared at all. Adding sodium hydroxide in step 3 should have neutralized any remaining hydrochloric acid and yielded NaCl, and washing the precipitate in step 4 should have carried away most or all of this salt. I would not expect that there would be enough chloride left over to form such a concentrated solution of copper chloride.<br />
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So in the end, I figured out what exactly was in my solution which was quite satisfying, but I'm still unsure of why this happened. I'll update these posts once I do some more research and determine the cause. Over this three part saga I learned a few things and had a lot of fun along the way, and I hope you did too.Danhttp://www.blogger.com/profile/06495614313702847298noreply@blogger.com5tag:blogger.com,1999:blog-7563481695249610141.post-12037156074135490372011-05-20T22:39:00.003-05:002011-05-20T23:31:05.999-05:00Copper Compound Conundrum, Part 2<span class="mediumtxt">Copper Compound Conundrum Part 2: Testing the Solution</span><br />
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<span class="mediumtxt">If you remember from my last post, my reaction scheme for this was the following:</span><br />
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<span class="mediumtxt">1) CuSO4 + 2NaHCO3 == CuCO3 + Na2SO4 + CO2 + H2O<br />
2) CuCO3 + 2HCl == CuCl2 + CO2 + H2O<br />
3) CuCl2 + 2NaOH == Cu(OH)2 + 2NaCl<br />
4) Cu(OH)2 + H2SO4 == CuSO4 + 2H2O</span><br />
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<span class="mediumtxt">After performing step 3 and filtering and drying the precipitate, I ended up with a dark green powder which was likely a mix of copper(II) hydroxide, oxide, and carbonate. All three of these compounds should react with sulfuric acid to form copper sulfate, so I figured I'd go ahead with the last step in my video's reaction scheme anyway. I should have ended up dissolving all the solids into a nice blue solution of copper sulfate like I started with. Here's what actually happened.</span><br />
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<a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhuhRiwSSkD6MEkeA-VEAoHCsQiZf6hBIeqaq01bPz893WeFs_dqNE9DtipfWlHQYw4GZxB_PhPLYj8O5sNn33KTBUGIcsQQRdjhMZN5ib0DU7kVFSGXXtm1PM2ykfiR9sOtwr8mq6gQaY/s1600/Copper+Compounds.jpg" style="clear: left; float: left; margin-bottom: 1em; margin-right: 1em;"><img border="0" height="240" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEhuhRiwSSkD6MEkeA-VEAoHCsQiZf6hBIeqaq01bPz893WeFs_dqNE9DtipfWlHQYw4GZxB_PhPLYj8O5sNn33KTBUGIcsQQRdjhMZN5ib0DU7kVFSGXXtm1PM2ykfiR9sOtwr8mq6gQaY/s320/Copper+Compounds.jpg" width="320" /></a><span class="mediumtxt"> </span><br />
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<span class="mediumtxt">The original solution is in the beaker on the right. You can see that I ended up with an emerald green solution, and a blue precipitate!</span><br />
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<span class="mediumtxt">This was very puzzling, so I decided to perform some experiments. You can see one of these in the test tube on the left. The solution looks very similar to my copper chloride step from above, so I wanted to see if this is what it was. Copper chloride solutions show different colors based on the concentration of chloride ions - concentrated solutions are green, and on dilution become blue. So, I transferred a few mL of the green solution into the test tube and diluted it with a few mL of distilled water. Sure enough, as you can see in the photo, the color changed to blue! This was a strong indication that I was correct.</span><br />
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<span class="mediumtxt">To test this further, I performed the standard test for the chloride ion. This is done by simply adding a solution of silver nitrate. If chloride is present, a white precipitate of silver chloride will form. This is exactly what happened:</span><br />
<div class="separator" style="clear: both; text-align: center;"><a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEiHo9aHcwzqmvdA8QtdNSjwielUACz2vgauaCNA2iJdcYIp5MfRfBhlxioRG9LeH7DoNMmNyNNRCxCCIcxIxtWJAOa703MZIq2Iuh6hEhUx9rP-uyfAFTqHnlovfbD0V8hcIwfkumgY7wE/s1600/Adding+AgNO3.jpg" imageanchor="1" style="clear: left; float: left; margin-bottom: 1em; margin-right: 1em;"><img border="0" height="240" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEiHo9aHcwzqmvdA8QtdNSjwielUACz2vgauaCNA2iJdcYIp5MfRfBhlxioRG9LeH7DoNMmNyNNRCxCCIcxIxtWJAOa703MZIq2Iuh6hEhUx9rP-uyfAFTqHnlovfbD0V8hcIwfkumgY7wE/s320/Adding+AgNO3.jpg" width="320" /></a></div><br />
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<span class="mediumtxt"> The first photo is while dropping in the AgNO3 solution, which immediately forms the white precipitate. The second photo is after some had settled. Leaving it standing for a while longer, more settled and the turbidity disappeared with the color remaining the same. This to me suggests the presence of copper(II) nitrate in solution, as it should be.</span><br />
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<u><span class="mediumtxt"><b>Conclusion</b></span></u><br />
<span class="mediumtxt">The green liquid of my mystery product is indeed copper(II) chloride.</span><br />
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<span class="mediumtxt"><b> </b>But how did this happen? Where did the chloride come from? There may have been some salt left over from reaction 3, but I washed the precipitate a number of times with water so that should have been mostly removed. Certainly there wasn't enough around to make such a concentrated solution of chloride. These are interesting questions that I am currently searching for an answer for. I'll update this post when I have it figured out.</span><br />
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<span class="mediumtxt">Stay tuned for Part 3, where I investigate the identity of the blue precipitate!</span><br />
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<span class="mediumtxt"><a href="http://thehomescientist.blogspot.com/2011/05/copper-compound-conundrum-part-3.html">Part 3: Testing the Precipitate</a> </span>Danhttp://www.blogger.com/profile/06495614313702847298noreply@blogger.com0tag:blogger.com,1999:blog-7563481695249610141.post-74818442284490131002011-05-20T22:26:00.004-05:002011-05-20T22:41:55.626-05:00Copper Compound Conundrum, Part 1If you've been following my blog, you've probably realized I like to write a lot. Well this one's no exception. I like to provide as much detail as possible so others can replicate my results and in reading it get the same enjoyment out of the experience as I did (hopefully). Enjoy!<br />
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In the process of filming a new video on copper compounds for my channel, I ran across a bit of a snag that ended up taking me down a different path that was both fun and instructive.<br />
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<a name='more'></a>The scheme of reactions I had planned was the following:<br />
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<span class="mediumtxt">1) CuSO4 + 2NaHCO3 == CuCO3 + Na2SO4 + CO2 + H2O<br />
2) CuCO3 + 2HCl == CuCl2 + CO2 + H2O<br />
3) CuCl2 + 2NaOH == Cu(OH)2 + 2NaCl<br />
4) Cu(OH)2 + H2SO4 == CuSO4 + 2H2O</span><br />
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<span class="mediumtxt">Each compound is a different color and exhibits different properties, and I thought it was nice to end back where I started at copper sulfate.</span><br />
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<span class="mediumtxt">The reactions proceeded as planned, with each step working out smoothly.The following are still images from the video I was recording at the time, so they might not be as nice as some of my other shots.</span><br />
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<span class="mediumtxt">Copper(II) sulfate solution: </span><br />
<div class="separator" style="clear: both; text-align: center;"><a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEgjK6AAwHAU4L5XJGuT-dzOE2vGqLk0lvlDy0nNV1UmTS5VfWVN8mAF98RN2xS3QLQGkK0H84OqYHZg1UAjh-tDlXMpw86TgDv07kwPhSISOlUu8zXpdsA7Gm0iqIOD3ChX_P9SsBtvQ9o/s1600/snapshot_001.jpg" imageanchor="1" style="clear: left; float: left; margin-bottom: 1em; margin-right: 1em;"><img border="0" height="180" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEgjK6AAwHAU4L5XJGuT-dzOE2vGqLk0lvlDy0nNV1UmTS5VfWVN8mAF98RN2xS3QLQGkK0H84OqYHZg1UAjh-tDlXMpw86TgDv07kwPhSISOlUu8zXpdsA7Gm0iqIOD3ChX_P9SsBtvQ9o/s320/snapshot_001.jpg" width="320" /></a></div><br />
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<span class="mediumtxt">Copper(II) carbonate, filtered and dried precipitate:</span><br />
<div class="separator" style="clear: both; text-align: center;"><a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEgO3-EpqvId1IaCGK_IkjHnVC4mWbpQaC-uWW4CpyDBaCku432Bj-D2TXObn-O4sikXbYP2AuZbdolsPmDKQBvFZsYCUei4G3F4HNTrnbIPv2hMxRPHEjXTo4L5G_RNIg-uxbKW0PvlpGE/s1600/snapshot_002.jpg" imageanchor="1" style="clear: left; float: left; margin-bottom: 1em; margin-right: 1em;"><img border="0" height="180" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEgO3-EpqvId1IaCGK_IkjHnVC4mWbpQaC-uWW4CpyDBaCku432Bj-D2TXObn-O4sikXbYP2AuZbdolsPmDKQBvFZsYCUei4G3F4HNTrnbIPv2hMxRPHEjXTo4L5G_RNIg-uxbKW0PvlpGE/s320/snapshot_002.jpg" width="320" /></a></div><span class="mediumtxt"> </span><br />
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<span class="mediumtxt">Copper(II) chloride solution (concentrated):</span><br />
<div class="separator" style="clear: both; text-align: center;"><a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEiqa12FkyFfK6pRXACJ2gwZJEgxvawOx-K9Svv25rKefxAtSPTKtOsmHlNjoaY6vFoGcXFP7hzJx6kZwpitREkXqueXxgIYUCsM4QsMz3DAhAX6Az63Ke1cnyk7uRoJ4j1dcON276RJoOU/s1600/snapshot_003.jpg" imageanchor="1" style="clear: left; float: left; margin-bottom: 1em; margin-right: 1em;"><img border="0" height="180" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEiqa12FkyFfK6pRXACJ2gwZJEgxvawOx-K9Svv25rKefxAtSPTKtOsmHlNjoaY6vFoGcXFP7hzJx6kZwpitREkXqueXxgIYUCsM4QsMz3DAhAX6Az63Ke1cnyk7uRoJ4j1dcON276RJoOU/s320/snapshot_003.jpg" width="320" /></a></div><span class="mediumtxt"> </span><br />
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<span class="mediumtxt">Copper(II) hydroxide thick, gooey precipitate:</span><br />
<div class="separator" style="clear: both; text-align: center;"><a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEgIKkaPTD07KHUOFLF3DHgJ-Rl8Z6AmZdaymxfV3dOFL3G8MPMmoZYN2SoCOhVYLsivHsxdAR_bFODzBLXMKIQZ0oAJINZrWsI4bav1KlbgpS65Z72PNvXbwflhr4cvU83tTDGRwQe82vo/s1600/snapshot_004.jpg" imageanchor="1" style="clear: left; float: left; margin-bottom: 1em; margin-right: 1em;"><img border="0" height="180" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEgIKkaPTD07KHUOFLF3DHgJ-Rl8Z6AmZdaymxfV3dOFL3G8MPMmoZYN2SoCOhVYLsivHsxdAR_bFODzBLXMKIQZ0oAJINZrWsI4bav1KlbgpS65Z72PNvXbwflhr4cvU83tTDGRwQe82vo/s320/snapshot_004.jpg" width="320" /></a></div><span class="mediumtxt"> </span><br />
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<span class="mediumtxt">Filtered and dried copper(II) hydroxide (?):</span><br />
<div class="separator" style="clear: both; text-align: center;"><a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEjgVnCGBPYxNnMRFgKbUWeQiu7A_Rrh_cNAaugcfRilIPzgnfhYKULNesyeJjOq99KUzBTVWpe-aldnsloK4jV-LPgmCRuLPjaR6QmqIC-IkMHfwuETtorFzC8jVVGyk6Uw6TsBQEg73BI/s1600/snapshot_005.jpg" imageanchor="1" style="clear: left; float: left; margin-bottom: 1em; margin-right: 1em;"><img border="0" height="180" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEjgVnCGBPYxNnMRFgKbUWeQiu7A_Rrh_cNAaugcfRilIPzgnfhYKULNesyeJjOq99KUzBTVWpe-aldnsloK4jV-LPgmCRuLPjaR6QmqIC-IkMHfwuETtorFzC8jVVGyk6Uw6TsBQEg73BI/s320/snapshot_005.jpg" width="320" /></a></div><span class="mediumtxt"> </span><br />
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<span class="mediumtxt">That last photo is where things went wrong. I filtered, washed, and dried my copper hydroxide and I ended up with this dark green powder instead of the turquoise blue that the hydroxide should be. I've had this problem before, and it's due to copper hydroxide's tendency to decompose and react with air readily, especially when wet. The two reactions I suspect occurred are:</span><br />
<span class="mediumtxt">Cu(OH)2 == CuO + H2O</span><br />
<span class="mediumtxt">Cu(OH)2 + CO2 == CuCO3 + H2O</span><br />
<span class="mediumtxt">Copper(II) oxide is black and copper(II) carbonate is green, yielding this dark green overall color. There were also bits of blue in there that aren't really visible - that was my surviving copper hydroxide.</span><br />
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This result was unfortunate, but I didn't think it would cause any problems. All three of these should react with sulfuric acid to form copper sulfate, as in step 4 above. However, this is where things really got strange. Stay tuned for Part 2, where I show what happened and perform some standard anion tests to test my hypotheses.<br />
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<a href="http://thehomescientist.blogspot.com/2011/05/copper-compound-conundrum-part-2.html">Part 2, Testing the Solution</a>Danhttp://www.blogger.com/profile/06495614313702847298noreply@blogger.com0tag:blogger.com,1999:blog-7563481695249610141.post-34415157145759683142011-03-28T22:40:00.002-05:002011-03-28T22:41:53.199-05:00New GlasswareI just got in a sizable order of some fancy new glassware! I got a number of specialized pieces that will allow me to do many new and more advanced experiments, and I'm very excited to try them out!<br />
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I got all this at a real bargain, from someone dismantling a start-up bio-tech business that didn't work out. The beakers, flasks, and graduated cylinders are just so I have more standard glassware available. The real exciting bits are those at the bottom. They (along with the round bottom flasks) make up an all 24/40 ground glass distillation kit. I'll be able to use this for reactions that are too corrosive for rubber stoppers or produce a lot of fumes. The ones I have in mind are producing and distilling nitric acid and elemental bromine. I'm looking forward to setting these up and starting some new and exciting chemistry. Keep an eye out for some of this glass in upcoming posts here or videos on YouTube!Danhttp://www.blogger.com/profile/06495614313702847298noreply@blogger.com1tag:blogger.com,1999:blog-7563481695249610141.post-83777252330144975452011-03-21T13:25:00.003-05:002013-08-12T13:36:03.123-05:00Potassium Iodide and RadiationThis old post serves as the Video Companion to my Potassium Iodide video, linked at the bottom to make sure you read the post first!<br><br>
This is another rather long post, but I feel it is warranted because of the importance of the content.<br />
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Because of the recent earthquake and tsunami tragedy in Japan, there's been a lot of interest recently in an otherwise ordinary chemical compound called potassium iodide. I'll get to that in a minute, but first a brief summary of the situation.<br />
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As you probably are aware, this disaster has caused one of the nuclear power plants on the island to fail and threaten to disperse radioactive materials across the country. In truth, the reactor itself shut down automatically directly after the earthquake, as planned, but the following tsunami overwhelmed the plant and destroyed its cooling systems. Even after the main fission reactions have stopped in a reactor, it remains very hot (both radioactively and thermally) for weeks or months afterward due to secondary decay processes of the products of the primary reaction. So, even though the reactor is technically "off" it still needs to be cooled for quite some time before it is safe to remove and dispose of. Since the cooling systems have failed, the fuel rods in the core threaten to overheat and melt, causing a host of problems. Nearly all of these were planned for in the construction of the plant, but there is still some risk of containment failure and dispersal of radioactive isotopes into the environment. This may have already happened to a small extent, but details are currently scarce or conflicting. It may be some time before all the facts become available. My heart goes out to those in Japan during these troubling times.<br />
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So what does this have to do with chemistry? Potassium iodide can be found as anti-radiation pills, taken to protect against exposure to radioactive iodine. Iodine-131 is a radioactive isotope of the regular iodine-127, and is one of the secondary products of the nuclear fission reaction I mentioned above. Iodine is used in the body and concentrated in the thyroid gland. When exposed to significant amounts of radioactive iodine (i.e. from nuclear fallout), it will become concentrated in the thyroid and remains there for much longer than it's half-life of 8 days. There, it can irradiate surrounding tissue and quickly lead to cancer. By taking potassium iodide in "large" amounts (large being only one 130mg tablet a day), you saturate the thyroid gland with normal iodine and this effectively prevents the radioactive variety from being accumulated.<br />
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You should remember two important things about this treatment, however. First, it only helps against radioactive iodine. This is only one of many radioactive elements present in fallout, and there isn't much you can do about the others. Second, it only protects the thyroid gland. Radiation damage can still occur in other parts of the body and from those other elements as well. That being said, the potassium iodide defense is still <i>much </i>better than nothing, but it pays to know all the facts.<br />
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Some time ago I made a video on my YouTube page on how to make this compound, as I needed it for some lab work. That has gotten a huge surge in views and comments recently, as a direct result of everything I have discussed above. Many suppliers of the tablets are sold out, and people are looking for other ways to obtain it. While this is indeed one way to make this compound, I have to stress the following: <b>I do not recommend anyone ingest potassium iodide made by the method in my video!</b> I do not have the means to guarantee the purity of my reagents or products, and I would never consider eating anything made in the lab. If you did not start with food-grade materials, you will not end up with food-grade product. For those who know a bit about chemistry, you can easily purify iodine by sublimation and recrystallization. You could also purify your potassium iodide from multiple recrystallizations from solution. <u>But</u>, even if you <i>did</i> have exceptionally pure I<sub>2</sub> and KOH, however, if you didn't combine them in <i>exactly</i> the right ratios you will have contamination in the finished product. There will also be some KIO<sub>3</sub> that you likely won't be able to remove. I can't be sure that there's any way that a home chemist can purify the potassium iodide produced this way into something safely edible.<br />
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My video on making this appears below. Enjoy it for the chemistry content, but keep in mind the warnings I pose above.<br />
<span style="font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%;"><a href="http://www.youtube.com/watch?v=SNXIY4InWFU"><span style="color: blue;">http://www.youtube.com/watch?v=SNXIY4InWFU</span></a></span><br />
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Periodic Videos also has a great video talking about all of this. They are an excellent science channel, and one of my favorites on YouTube. The Professor is an incredible person! See the video here: <a href="http://www.youtube.com/watch?v=-bcrLiATLq0">http://www.youtube.com/watch?v=-bcrLiATLq0</a><span style="font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%;"><span style="color: blue;"> </span></span>Danhttp://www.blogger.com/profile/06495614313702847298noreply@blogger.com3