Super Cool

I'm going to start this experiment with a technical description of what we are doing here. I had to divine this before I could make the wildly different and terse internet "recipes" work. A few little tricks required my knowledge of what was happening as it relates to crystallization, enthalpy of fusion, and the exact reaction entailed. I had tried previously but failed. I had a good idea of what was happening, but not good enough to fully anticipate next steps along the way. It turned out to be a bit tricky, but I think my description here should be complete enough for others.

I have to admit this is the most amazing experiment that I have done, even more than I anticipated. If you want the Gee-whiz without any technical details, like an old BASIC program, GOTO EXCITING_STUFF_HERE

Otherwise... 

BORING STUFF HERE:

We are going to supercool a solution, and then apply a seed crystal, causing it to almost instantly crystallize the whole liquid into a solid and release heat upon fusing. 

First, we create a solution of sodium acetate in water. No big deal, every kid knows what happens when you mix 3 tablespoons (54g) of basic baking soda and 3 cups (750ml) of 5% acidic vinegar. You get a reaction with a lot of carbon dioxide bubbling, so you have to add the vinegar slowly.

5% Vinegar has 5g of pure acetic acid (CH3COOH) for every 100ml of solution. You can adjust the amount for your concentration. 

Baking soda is NaHCO3

At this point we have created CH3COONa (sodium acetate) dissolved in a bunch of water with a big release of carbon dioxide.   

After it is fully reacted, put the pot on the stove on low heat. You need to get the solution well above 136 degrees and leave it there, while evaporating most of the water away. Our goal is a simmering solution (less than 212F, say 190F) with exactly enough water left to form three water molecules for every sodium acetate molecule. Exactly, more or less. :)

CH3COOH + NaHCO3 + 2H2O -> CH3COONa*3H2O + CO2

Why 136 degrees? That is the "freezing temperature" below which sodium acetate trihydrate (a crystal with the three waters included) can form. In other words, this compound is normally a solid at room temperature. By keeping it hot, we can keep these crystals melted as long as we have enough water in our solution. 

Why 190F? Water boils at 212F, so our solution should boil just above that point. It is hard to see what is going on in your dish with a raging boil. Also, we will need to cool it later, so overheating seems unnecessary. OK, I warn you now, we are getting rid of about 90% of the water, and it has to be done precisely. This can take hours and frankly I don't have the patience. My shortcuts usually led to failure.

This one did work. I found that I could bring it up to a low or moderate boil periodically, but needed to check my volume and make sure my solution was still clear by turning down the heat every few minutes. Once you are below a cup of solution, you should watch more carefully and avoid the boil. Note that white crystals start forming on the sides of the pan early on. This is normal and expected. Ignore them.

After an hour and a half or so I saw an unexpected white powder forming in the bottom of the solution. I have no idea what this was. In theory, my 1 TBL for every 1 cup formula gave me about 2g too much baking soda. Did this somehow appear when there was no longer enough water to dissolve it? I think so. I had it happen twice, and I am pretty sure it was not crystallization. If it was not baking soda, it was most likely anhydrous sodium acetate which is not the exact chemical structure that we need.

The first time I made the mistake of thinking it was the desired crystallization and proceeding to drain and filter my solution for cooling. Don't do this. The second time I drained the solution and filtered the white powder out, but then rinsed my dish and put it back in there to keep simmering. This is what to do if you get the white powder formation. I suppose I could have also added some vinegar until it went away, but then I'd have even more water to evaporate away (assuming a healthy dilution).  In any case, the powder never came back or reformed. It is possible that I could have proceeded to crystallization (described below) while the powder hung around. I am not sure as I did not test this.

About two hours into the process, on a nice 190F simmer of my especially clear solution, I noticed crystals forming. The key is knowing what to look for and catching it quickly. The way they form is a small clear crust just beginning to form on the surface of the water like a scum. You really can't miss it, it is not like a precipitate down in the solution. Now when this happens, act immediately. Turn off the burner, and slowly pour the hot solution through a filter into your clean dish again. 

It is critical that the dish be cleaned well after prior uses. I used a funnel and a coffee filter. Still, it is hot and the filter drains slowly. You want to catch any impurities or crystals and keep them outside of your solution. In my case, I had a nice clear yellowish liquid, between a quarter and a third of a cup,  that had cooled to about 150F degrees (this needs to stay above 137F for now). Quickly I added plastic wrap to the top of my dish (to prevent further evaporation) and carefully placed it into my refrigerator where it would not get bumped. Now, you can use room temperature to supercool the solution, but the fridge is faster and we are 2+ hours in at this point. 

Supercooling is lowering the temperature to below the freezing level but without fusion (crystallization) occurring. Think water at less than 32F but not frozen. We do not want the phase change to solid to happen just yet. In this case, getting it to room temperature (about 60F below the point of fusion) is fine. Unlike supercooling water, we don't need to get it below 30F, and that is one reason this crystal is perfect for our experiment. Another reason is this chemical has a high latent heat of fusion, which means that it releases a lot of energy upon fusion (or takes a lot to melt). A supercooled liquid can "store" the energy required for crystallization to be released later in the form of heat. Obviously, a "frozen" phase of any substance is in a lower energy state than a liquid or gas where parts are expanding and moving about (more disorder is higher entropy).   

Another thing I had to learn is that heated water is a better solvent than unheated water. By heating our solution to 190F, we can dissolve a larger amount of chemical in our solution than at room temperature. When the solution is fully saturated then cools without crystallization, it becomes supersaturated. The liquid contains more dissolved chemicals than would have been possible without heating. This was an ah-ha moment for me, as I realized it was critical that we fully saturate the heated solution and no more. We needed crystals to form (the moment of full saturation) before stopping the dehydration.

At this point, taking time to clean the dishes while the solution supercooled in the refrigerator, I was ready for the good part. The solution still looked clear and liquid, and I was able to move it to the counter without accidentally kicking off the crystallization. Supercooled liquids are not very stable, and even a bump can start their fusion sometimes. What works best is nucleation, or the introduction of a "seed" crystal to which the remaining chemicals can readily bond. I was prepared with seed crystals scraped from the bottom of the pot after pouring out the final solution.

While I had planned to drop a seed crystal in, I wanted to measure the temperature of my supercooled solution. So I inserted the tip of a candy thermometer. Before I could read a temperature, a reaction formed around the tip of the probe and started expanding outwards as white crystal. I had provided just enough nucleation in the form of disturbance and/or the introduced probe. I quickly pulled the thermometer out to get my camera, and the reaction stopped, leaving me with  a nice seed on the tip of a thermometer probe. What luck! While I could not get a temperature of the solution because the crystallization releases energy in the form of heat, I had a nice seed on the end of a probe that could measure the heat of the reaction.  

Lights, camera, action! You can see the rest below... I should note that the heat of my fusion was 120F. It makes sense that it was not quite enough energy to melt it again. 

EXCITING_STUFF_HERE:

This experiment "supercools" a liquid. When I dip a bit of the type of crystal ("hot ice") that should form, a beautiful 3D upside-down chrysanthemum-like growth initiates and completes in less than 20 seconds, leaving nothing but a warm solid mass of crystals in my bowl. It is like a magic trick. Here it is in video:

And here is a closeup of the solid. The solution shrank a bit to fill almost exactly 1/4 cup of solid in the end:

And here are recovered crystals outside the dish:

Isn't that super cool?

Thanks for reading,

Paul   

The Candy Part of Rocket Candy

Shifting gears a bit, why did the sugar turn brown when I heated the Rocket Candy mixture? That process is called caramelization and is the same as what is used to make caramel candy, toffee, butterscotch, and that blow-torch induced crust on your creme brule. 

Most of us may know that much. Anyone who cooks at all probably is familiar with what happens when sugar is heated.

I thought I'd look up what is happening chemically. I had a vague feeling that hydrogens and oxygens were burning off and leaving a more carbon rich mixture behind, but I was not sure.

Apparently this process is quite complicated and not fully understood. A good definition would be the removal of water from sugar (yes!), proceeding to isomerization and polymerization of the sugars into various high-molecular-weight compounds. Huh? I think I was right about H and O being being off. After that it gets complicated. I believe what it is saying is that what is left of the sugar starts combining into other bigger molecular structures. These compounds have a characteristic color (brownish orange) and flavor. 

According to wikipedia, the brown colors are produced by three polymer groups: caramelans (C24H36O18), caramelens (with an 'e' not 'a' - C36H50O25), and caramelins (with an 'i' - C125H188O80). Those are some big molecules! Wikipedia also states that the favor comes from another smaller by-product chemical named diacetyl ((CH3CO)2) which is added to foods sometimes to impart a buttery flavor.

So this is too complicated for me to describe in the form of a reaction. I try to always mention that burning reactions are oversimplified. The reactions described in the Rocket Candy post and postscript are not going to be great models for carmelized sucrose. This may also explain why 35% sucrose is prescribed rather than the 26% that stoichiometry would specify. Some of that excess weight burns off in the form of water vapor. All of the carbons remain, as the fuel, so the overall reaction still works at a functional level.    

Note that the 26% unmelted sucrose reaction in the next post went just as planned. Perhaps if I had melted that mixture first there would have been a problem, at least in terms of how much product is useful for propulsion. 

Thanks for reading,

Paul

Rust-Away and Rubies

Never mind Rustoleum. I have a new product that turns pure rust (iron(iii) oxide, or the mineral hematite) into pure elemental iron almost instantly. As a bonus it creates rubies and sapphires. The main problem? It melts and welds itself to other metals.

Yes, I am finally messing with thermite, using Aluminum powder for fuel. Aluminum is what the most powerful motor ever made, the Space Shuttle Solid Rocket Booster (SRB), uses for fuel. We are in the big leagues now. 

Actually thermite is a famously exothermic reaction, creating temperatures up to 2500C. That is hot.  Twice the heat needed to melt iron. The reaction is:

Fe2O3 + 2Al -> 2Fe + Al2O3

The product aluminum oxide is the mineral corundum, which is known to us as rubies and sapphires. But we won't get any gems, just some white powder. We will also notice the pure iron that is left behind, a little more than 2/3rds the weight of the iron oxide used. This is visible as little balls of metal surrounded by the aluminum oxide. Ignore the big aluminum melt that was just the container for the reaction.

Stoiciometrically, the ratio of iron oxide powder to aluminum powder is almost exactly 3:1 by weight, and that is what is generally used. I wanted a small test so I did not blow myself up, and went with 1.5g of iron oxide and 0.5g of aluminum. 

Beside the potential to melt your spacecraft instantly, thermite is not used as a rocket propellant for a fairly obvious reason. Its reactants are solid. There is no gas released as thrust to push the rocket along. It would basically ignite, and then melt the entire spacecraft on the pad and anything in the vicinity. Not good.

So we see why rocket candy is good with its carbon dioxide, nitrogen, and water vapor gasses released. Propulsion forward is when gaseous product of a reaction inside the tank is directed to the atmosphere behind the rocket through a nozzle.

But thermite is extremely useful for things like welding railroad tracks together, underwater welding, and melting holes in stainless steel. You know, common everyday household things like that.

I should note that other metal oxides can be used with aluminum to create molten copper, etc. In this post, we are assuming iron, specifically iron(iii) oxide and not magnetite (the mineral F3O4 with iron(ii) and iron(iii) cations).   

  

Weird Science

After my rocket candy post, I continued to experiment with various oxidizers and fuels and ratios. More on that later, but I learned something more about the rocket candy reaction.

I tested 26% sucrose and 74% saltpeter in powdered form with a magnesium ribbon for ignition. One reason for the ribbon is that my blow torch would have blown the powder away. In case you don't know, magnesium burns very hot and bright with the reaction:

2Mg + O2 -> 2MgO

It is sort of the ultimate fuse, and was used to make incendiary bombs in WWII. Now that I have a roll of such ribbon, it will be my new method for safety reasons and improved odds of successful ignition.

Anyway, the reaction went off just fine with no excess fuel (in the form of smoke and carbon deposits). I think that rocketeers like the smoke; it makes a nicer display when launching the vehicle and aids faraway viewing as the rocket climbs. For my purposes, the smoke and carbon deposits are a nuisance and also waste good sugar.

Since I had no carbon left over, I had a very good look at the solid product of this reaction, shown here in a melted aluminum container. By the way, the black carbon deposits inside the "O" and above the melted product were from another test, so you can see the very clear difference of the purer burn. 

Note what a pure white the K2CO3 is. Additionally, it felt like hard smooth plastic and was a bit hard to pry up. I kept a few pieces for whatever reason, but that is where my discovery came in. After an hour of so in my home, it felt a bit wet and tacky (such that I had to wash my hands). This intrigued me since it had been hard and dry not much earlier.

I looked up potassium carbonate on wikipedia and learned a new word. The compound is deliquescent. I had no idea what that meant, and was pretty sure it doesn't mean that it tastes good on a sandwich. It means this material will take hydrogen and oxygen from the environment (as available in the form of humidity) to the point of not just hydrating itself but actually forming an aqueous solution and dissolving itself. At that point I threw it away, thinking it would start making a mess. Maybe I should have put it in a tub in the garage and watched over time. I had no idea such weird things exist!

Oh brave new world that has such compounds in't!

Thanks for reading,

Paul

 

Rocket Candy

 

After the failure of my gunpowder, I searched for use of burnt sugar for carbon and learned of rocket candy. This is what hobbyists use for some serious model rocketry. There are many variants and degrees of sophistication among enthusiasts, but I needed simple success. What I read told me (essentially) to combine 65% saltpeter with 35% powered sugar (sucrose) in a mortar and mix it into a fine powder. I should have powdered only the saltpeter first since powdered sugar is so fine already. I used 3g of the sugar and 5.5g of saltpeter (measured precisely this time). 8.5g also turned out to be a very manageable total size.

Using a so-called dry method, I put this mix in an outside grill with cover and temperature gauge. At about 220F the sugar starts to melt. I went up to around 250F and pulled it out. Note the KNO3 does not melt or react at this temperature, so we have to stay in this low temp area and not use a blow torch as I like to do.  This method can be potentially unsafe - I would not wish to spill the contents onto the grill's burners, for example. Once a reaction starts it will complete. After I had my melt out, it was simple matter to thoroughly stir the candy-like substance below: 

This material is very moldable, by the way. Here is what I recovered upon cooling and tapping out of the dish:

The bit on my stainless stirrer cooled hard quickly and so I broke it off as a perfect test size and place it in a safe area. I zapped it with my blow torch and it fizzled loudly and made prodigious smoke. Success! I have oxidized carbon from sugar in a chain reaction. I may come back to "real gunpowder," but this is essentially the same reaction. Oversimplified:

2KNO3 + CH2O -> 2KNO2 + CO2 + H2

In fact, you can add sulfur and charcoal as the reaction is very much like gunpowder. CH2O as an oversimplified model for sugar is not unlike that of charcoal. Sulfur is optional as a fuel that lowers the ignition point and can increase combustion. I don't like messing with sulfur more than I need to because it stinks and tends to create hydrogen sulfide gasses. Think of it like dirty coal or gasoline - we pay extra to keep the sulfur out of fuels.

Now I need a rocket. The whole thing took less than an hour from concept to test. No mistakes this time. :)

Thanks for reading,

Paul 

p.s. Aug 25:

I found a better equation for this reaction using the actual molecule of sucrose. It is a bit clunky, and recall that "burning" is not simply one reaction.

48KNO3 + 5C12H22O11 -> 24K2CO3 + 36CO2 + 24N2 + 55H20

With this better equation, I was able to multiply out the molecular weights and got exactly 26% sucrose and 74% saltpeter by weight. I am told the recommended formulation is a bit fuel rich to allow for more fluidity in preparation. Note the products are different. In reality there are many different reactions but this is a much better model than above.

I was also able to calculate the enthalpy of the reaction. With 48 moles of saltpeter (4.8kg) and 5 moles of sucrose (1.7kg) you could create 20.3MJ of energy. So I created enough product for 2.3kJ, roughly.

Fo Fizzle

Apologies to Snoop Dogg, my nizzle. Not much sizzle. This is a tale of fizzling out, for now, brought to you to show the frequent failure of experiments that provide numerous learnings along the way. One of the things I love about watching Nile Red is that he fails often, and experiments until he figures it out. His mistakes are not usually rookie mistakes like mine, but it makes me feel better.

I have a number of experiments planned, but the weekend hit and I don't have the chemicals and parts I need yet. What to do? Since I have been teaching myself about Enthalpy of reactions (energy produced by a reaction), I thought of making something "easy" that is exothermic enough to cause a chain reaction. Exothermic means energy is released, generally in the form of heat and/or light. I have already done this with propane combustion, but I decided to consider black powder, or the old-fashioned gunpowder that the Chinese somehow invented 2000 years ago. It never ceases to amaze me what ancient people discovered.

So, if you mix 75% saltpeter (KNO3), 15% charcoal (~C with some hydrogen and oxygen), and 10% sulfur, you just need to ignite with a spark (such as caused by flint striking steel in old flintlock guns) and boom! In my case I just want to make a small line of it like a fuse and watch it burn.

What is saltpeter? I first heard the word when watching the musical 1776 in junior high.  It is a crystal of potassium nitrate that people have made and used for many purposes over the years. Food preservation was one use. Gunpowder was obviously an important application. In any case, it can form naturally from animal dung and urine combining over time, and is sometimes found growing near stockyards. Urea, formed in the kidneys, is the source of nitrogen. Anyway, I don't have the months of time nor desire to put my dog to work. Instead, I bought Spectricide brand stump remover, the main use of KNO3 that I know of at this point.  

Sulfur is easy. I have a chemical bottle of it, but check out this rock I collected from a live volcano in Taiwan once. I will not destroy my nice mineral sample, even for science.

Now, for charcoal, I felt like I needed to make something the hard way. So I quickly shaved 20g of pine off of some excess lumber with a pocket knife and set it ablaze. While it burned I wondered about the yield and combustion formula. Note that burning is really a combination of many reactions and things like wood are not chemicals in the sense that they easily fit into equations. They can be modeled, however. What I found on the internet disturbed me as it was all over the place and equations did not balance (even detailed ones). These were my main assumptions:

• Wood can be modeled, roughly, as C10H15O7 by measure of its constituents, and charcoal roughly as C7H4O. We lose some carbon (smoky soot and carbon gasses), as well as hydrogens and oxygens in numerous other gasses.
• While I found reactions that indicated 5/6th's of my carbon would remain in the char, that seemed kind of crazy. Wikipedia's entry on charcoal told me the most modern industrial methods have a 35-40% yield. 
• If carbon is 48.5% by weight of C10H15O7, and my optimistic yield is 30%, I figured I could perhaps extract 3g from my wood.

After it cooled, I measured the product. Including tar and ash, it weighed 3g on the same scale used earlier to measure 20g of wood. Success! So I put it in a mortar and ground it down. A few pieces were too big for my liking so I picked them out. I figured I had about 2g left after losing the tar and bigger chunks. Basically, I was lazy and should have measured precisely. Instead, I added the sulfur and saltpeter to my mortar, precisely measured to 1.3g and 10g respectively. Whether I was distracted by the Astros game or just too happy with the deep fine black powder, I dunno, but I am sure it was not exactly 2g of char. To further the potential mistakes, my burning of wood probably made a very lousy charcoal. It was obviously not a hot and long fire. Did I have some carbon? Yes, but who knows how much at this point.

  

I mashed everything together in the mortar a bit and it looked and smelled like gunpowder. I was feeling good about this. I probably underestimated the danger here; 13g of gunpowder is more than I realized and it is possible that I could have created a spark or heat from the mortar and pestle grinding. Still recovering from my bismuth burns, I don't need another.

When I layed out a little line on the ground and tried to light it with a simple flame, I could not get a chain reaction. I took a blow touch out to ignite it, and ultimately got it to sizzle down the line somewhat haphazardly. Did it work at all? I think so, a little bit, but not good enough to fire any musket balls or even make fuses. If I was Chinese 2000 years ago, I would have registered no apparent usefulness.

Is my sulfur good? Clearly. Is my saltpeter good? Um, I just assume it is actually KNO3 based on lack of ingredients listed, color, and crystallization. Not good enough!

So ensued a comedy of more errors. I burned some to see if I could see a characteristic lilac flame of testing potassium. That didn't work. It was too bright out and I could not see any flame. I had the stump remover in a stainless steel cup, and crystallization looked like potassium nitrite (KNO2). But I was distracted by some tiny lumps of metal that formed as a vapor on the side of the cup. Potassium? Really? I added water and nothing happened. It was not potassium. Then I realized the dish must have had some zinc plating and I just vaporized and condensed zinc droplets. Galvanized steel is starting to cause me a lot of problems. From now on, I assume all steel is galvanized.

Later, after some cooling and washing of the dish, I looked under a microscope, I saw these lovely crystals which formed while the dish was cooling. It turns out these are KNO3 crystals, which possibly reformed from the nitrite by reacting with oxygen. But at the time I was not sure.

Instead of attempting other chemistry to confirm KNO3, I decided to punt and search the internet. The makers of Spectricide go to great lengths to not disclose the ingredients or comment on its unintended uses. I am not the only one asking, mostly for a variety of reasons including gunpowder, rocketry, and tricking drug tests. Someone out there might be thinking of using it as phosphorus-free fertilizer, I don't know. But I prevailed. The government requires a safety sheet to be posted somewhere, and I found it finally on their web site. It revealed the contents as 100% KNO3 with a bunch of warnings that I did not heed. I like the irony of the safety sheet being so hard to find.

So the problem was either my lousy charcoal or my lousy measurement throwing off the mix. I'll redo this with good charcoal at a later time, then try to measure it precisely with homemade charcoal and compare.

In any case, an oversimplified reaction would be something like this:

2KNO3 + S + 3C -> K2S + N2 + 3CO2

More likely, potassium carbonates and sulfates and soot and carbon monoxide and water vapor all form. Again, burning is not one simple equation, and charcoal is not pure carbon.

Stay tuned for the redo! At least I learned a few lessons.    

Thanks for reading,

Paul

 

 

Burn Notice

This post is really just a means to an end. There was no specific chemistry that I had in mind, but rather a step in creating a furnace for melting higher temperature metals, smelting ores, etc. I need a gas jet to heat the thing sufficiently.

I was thinking of Bernoulli's Law, the Venturi effect, and other elements of fluid dynamics mostly. But a funny thing happened on the way to this post. It occurred to me that I needed some chemistry for perfect combustion of propane:

C3H8 + 5O2 -> 4H2O + 3CO2

This tells me that I need a molar ratio of 5:1 between oxygen and propane. I don't plan on using pure oxygen, so I'll assume that it makes up about 21% of what we call air. 

If you don't get enough oxygen in the mix, you start getting soot and carbon monoxide and other emissions that are often undesirable. This is known as a reducing flame, with less energy (heat), and can be preferable when you want to avoid creating oxides of your metals or want to add some carbon. Here is a simple example where exactly one less oxygen pair generates soot in the exhaust (carbon). 

C3H8 + 4O2 -> 4H2O + 2CO2 + C

It is essentially impossible to keep one perfectly balanced reaction going. Here is another that can happen with almost but not exactly the right amount of oxygen in the mix:

6C3H8 + 29O2 -> 24H2O + 16CO2 + 2CO

Note that 29/30ths of the perfect oxygen amount creates carbon monoxide. That will happen just from minor variability in atmospheric air pressure. Yikes.

An oxidizing flame has more oxygen than needed. It is very hot as it fully combusts the fuel. But we don't really want extra oxygen unless we are trying to create oxides during smelting. If I ever use my graphite crucible, oxidation can turn it to gas. It seems like some sort of terrible disappearing magic, but the reaction is as simple as it gets.

C + O2 -> CO2

The bottom line is you generally want to design a sufficiently large air intake and then add an adjustable "choke" which reduces air intake so you can adjust the mixture you want. You may recall chokes from older cars, lawnmowers, tractors, and so forth. Or maybe that is just me. I had no idea what it did at the time. I knew to open it when the engine was flooded.

The general idea of the gas jet is to shoot a very thin and fast stream of gas out of very small hole into a larger pipe. Directly behind this, you have air vents that allow air to be pulled into the pipe according the the suction provided by the gas stream. Cars and airplanes have used similar mechanisms, but sometimes rely on "forced air" which gets scooped into the mix from the velocity of the vehicle travelling through air. My gas torch should not be moving (hopefully). In automobiles, the methods of obtaining more oxygen are numerous: turbochangers, superchargers, etc. 

At the end of the exit pipe, we have a flare that allows the mixed gas and air to expand outward and slow down so that we can light it. 

I looked at a few designs online, before deciding to use one similar to TKOR (God rest his soul). While his design seemed to meet my needs, I could not find a couple of the parts anywhere, and I do not have the tools he has to be able to drill and tap thick curved stainless steel.

(Picture: 3/8" brass flare to 1/4" MIP, 1/4" ball joint valve, 1/4" MIP to MIP, a couple of 1/2" washers that were probably not needed, a 1/4" female to female galvanized steel pipe coupling, a 1/4" brass plug, and a .6mm welding tip that I screwed into the endcap by drilling and tapping a hole into the brass.)

I am used to working around a lack of tools, and have found over the years that many applications can use a carved wooden replacement. I realize that sounds insane. Wood can do some amazing things. I even sailed a boat for years with the mast inserted into the hull by carved wood. It never failed me. In this case you may be thinking that wood will get hot and burn, but the heat is at the end where the flame begins. In the areas where gas escapes and pressure reduces, frost can actually form.

So I carved an interposer between my brass endcap and the 1.25" pipe reducer. I literally screwed it into the threading (like a reverse tap). Note that the hex-shaped endcap fits into the wood, which then has a slightly smaller  hole drilled through it for the welding tip and air to pass through.

(Picture: 6" 3/4" steel pipe with 1" reducer for flare and 1.25" reducer where my wooden venturi "screws" in. The hardware described and shown above is inserted into the wooden venturi.)

Initially I thought the welding tip was a bit of unnecessary artistic flare. Indeed, in the TKOR design he says it is optional but cool. But because I don't have an open space behind my endcap, I need the outlet positioned further into the pipe. So it is necessary with my design. This allowed me to drill 12 holes into the wood past the endcap but before the end of the welding tip. These are my air vents. My plan was to drill more holes if needed and fill some in with putty or plugs if it was too much. It just happened that my 12 holes give a good amount of air and don't need any further tuning at this time. As a future project, I may add more holes and an adjustable collar for the choke. 

Finally, I added a regulator and pressure gauge to my tank. This way I can keep the system at 30psi for safety and consistent results.

Now I was ready to light this candle! I tested it in daylight in case there were any issues. It made a nice loud jet noise, which is a good sign that my venturi was working. Then I tried to light it and it took immediately and easily, so there is sufficient oxygen and my flare works too. I tested the components and none of them were hot except near the end. Being daylight, I could not see the flame well, so that would be test #2 at night.

I should note again, always, not to use a device like this indoors. You don't want the fire risk or exhaust gasses in an unventilated area. Think of it like running a car in a garage.

Time for test #2 at dusk. Wow, what a huge (18" or so) nice blue flame. Not too oxidizing or too reducing. I'll keep the tuning as is. This flame will completely fill any kiln I build. I declare success!

(I will blog again after building my furnace. For now this jet is not much use to me. I need something to contain the flame and retain the heat.)

Thanks for reading,

Paul

 

 

 

Bismuth!!!! Pt I

OK, that title isn't very creative, but I don't feel like a Madame Curie or Steven Universe quote here. And it reminds me of A Streetcar Named Desire, for some odd reason.

What is Bismuth?? Well it is a heavy metal, much like lead, but without toxic accumulation in the body. Therefore it is recommended as an alternative to lead and buckshot that is not toxic. I don't know why you would care if shooting someone, but I digress.

Bismuth is an element, number 83, meaning it is between Lead and Polonium (nasty stuff!). Bi is 86% the weight of lead by volume. It has 83 protons by definition. It is almost always found as isotope number 209, so it has a lot of neutrons.

We are primarily interested for four reasons:

1) It is relatively cheap and easy to buy online

2) It's melting point is only 521F, just a bit above the temperature to burn books (Fahrenheit 451). We can do this without a furnace. Even a kitchen stove can do this.

3) As it cools, it can oxidize on the surface into a rainbow of colors

4) It makes super cool hopper crystals (we'll save that for pt II)

Melting: well, it was super quick and easy with a Bunsen burner or a propane torch. I used both a crucible and a stainless steel container. It doesn't matter much.

Pouring and cooling: well it poured into my little graphite mold easily and cooled a bit slowly. Bismuth has relatively low thermal conductivity. Nice little bar:

Mistake: I have to admit I accidentally spilled some on my hand, which was quite painful for some time. I am lucky to make this mistake at 521F instead of with molten copper. Here is a good reminder to wear gloves and eye protection working with high temperatures. The upside (besides the learning I won't soon forget) is that I flung the metal out of my hand instinctively (like a dead frog's leg twitching) and made a beautiful piece of artwork that will show off the rainbow oxidation.

Why does it make the colors? Like all metals, it bonds with oxygen in the air at high temperatures (some do this more easily, like in the case of iron and rust). It makes Bi2O3 (same as the mineral bismite which I have found as yellow powder coating other minerals). This oxide is used to make "Dragon's Egg" fireworks.

Tiny variations in the thickness of the oxide layer are caused by the amount of time that spot is hot enough to bond with oxygen. In my experiments, rapidly cooled bismuth looks silver still. Quickly cooled silver has a gold colored tarnish. Then comes purple, and finally blues and greens for more slowly cooling areas. 

Note that my hand splash is silver where thinnest, gold where some crystallization started, and blue/green with a bit of purple where a thick blob landed and cooled more slowly. The photographs just don't do it justice, you can get more of an effect in real life with better light and moving around. 

In the mold, the top portion exposed to air cooled more quickly than the bottom, but the bottom was not exposed to oxygen (just the graphite mold).

Why does the thickness of the oxide matter? This is physics. These are interference patterns of light. Some wavelengths (colors) reflect back off the mirror-like metal and cancel each other out. Other colors reflect back pass through. The oxide is like a filter, in a way, where the thickness and angles determine the result.  The phenomenon is known as iridescence and can be seen in many things such as soap bubbles and oils. But this one hangs around after cooling.

Thanks for reading,

Paul

p.s. It has come to my attention that some people would like to repeat these experiments without the hassle of finding the chemicals and/or having to buy too much, etc. I will try to keep an Ebay store alive (seller: cinnabarminerals) that offers low cost and related material for the purpose of enabling education. I will title offerings "Brave New Chemist Science Pack - (related Blog post name)." You can always comment and ask where to buy certain equipment too.

 

Out of Whose Hills You Can Dig Copper - Pt II

Please read Pt I if you have not already. 

We will start with a tub full of copper(ii) chloride, which has crusted over with unanticipated crystals. I added some water to remove the crystals and form an aqueous solution again. I can assume that there was not enough water present to fully dissolve all of the copper and chloride ions. This is because my acid was fairly concentrated. I added about 100ml of water. The greenish color was still fairly dark and hard to see through.

The goal here will be to turn iron into copper. I am also working up to turning lead into gold. Just joking, the goal will be to add iron and then have a reaction where the iron and copper swap places. This is known as a single replacement redox reaction. Then, I will have the pure copper that I want extracted from the malachite ore. Presently, that copper is in copper(ii) cations in the solution. We want pure elemental (native) copper with no charge as the product. The reaction is fairly simple:

Fe + CuCl2 -> Cu + FeCl2

The iron will be replaced with copper and the solution will contain iron(ii) chloride (or ferrous chloride). This is where I went wrong previously, adding iron filings. It was impossible to know when the reaction was done. I also did it all in one step so if anything did change, I could not see through the carbon dioxide gas bubbles. After some time and investigation, I realized I had an incomplete replacement reaction and my elemental metal was part copper and mostly iron. My solution was a strange mixture of iron and copper chlorides that kept changing colors and other weird things.

The answer is fairly obvious, it turns out. Sand the zinc galvanizing off of some nails and drop them in the solution. An iron nail and a copper crystal are easy to distinguish. After a short time, I could see copper growing around where the nail was being eaten away, creating cinnamon-like tubes of pure copper. Success!

I kept adding nails, confident that any unreacted iron would be obvious and I could just remove it. This way I would be sure to yield every atom of copper (more or less). I began to think about the copper(ii) chloride, and wondered if I could completely react it to a pure solution of iron(ii) chloride instead of a mixture due to limiting reagents. The answer is also simple: more iron nails!

I can do this because it so happens that Iron and excess HCl react to give me hydrogen gas and (voila!) iron(ii) chloride. So I kept throwing nails in until they were not changing. Then I knew all reactions had ceased. By the way, I had previously calculated that 50ml of 31.5% HCl was enough to react all of my copper (and then some).

Fe + 2HCl -> H2 + FeCl2

Two notes of warning: Perform the reactions outside or in vented areas. You do not want hydrogen gas in your house or near sparks. This is what happened to the Hindenburg. Also, ensure the galvanized nails are well-sanded so that zinc is not present in the reaction. Make sure the nails are iron or carbon steel and not some other metal. The amount of carbon in steel is small and won't impact the reaction.

When all reactions were complete, my solution was very pale and I had obviously reacted as much iron as possible:

While still reacting, I did play around with decanting the solution periodically to check my copper. Check this out!

I could not resist pulling a little bit of copper out and trying to melt it in a crucible with a propane torch. the melting point of copper is too high at 1983F. It only partially melted, and definitely oxidized on the outside to a black CuO (the mineral tenorite in nature). I also got the characteristic green flame of copper. You can clearly see copper crystals forming here as it heated and cooled. More on that later, they look better under a microscope.

Back to the final mess in my tub. I decanted the hydrated iron(ii) chlorate into a beaker for later use. I used a coffee filter and funnel and some distilled water to catch and clean any solid metal. The unreacted nails (iron) were simple to remove. I now have this solution for future experiments and some nice copper that I was able to crush in a mortar (like cinnamon again) and weigh.

Remember my goal was 10g or more. Well, I scored a whopping 12g! That is over a 90% yield, including the impurities of the mineral itself. There is no way my copper would be this pure or have yielded so high with smelting, although that is a cheaper way to go in mass quantities and you don't have corrosive by-products (known in the trade as pickle liquor).

Thanks for reading,

Paul

  

p.s. It has come to my attention that some people would like to repeat these experiments without the hassle of finding the chemicals and/or having to buy too much, etc. I will try to keep an Ebay store alive (seller: cinnabarminerals) that offers low cost and related material for the purpose of enabling education. I will title offerings "Brave New Chemist Science Pack - (related Blog post name)." You can always comment and ask where to buy certain equipment too.

    

Out of Whose Hills You Can Dig Copper - Pt I

The mines in Timna are reported to be King Solomon's Mines. The area is full of beautiful blue and green copper ore, which was known even to ancients as a valuable asset for making bronze and so forth. The title of this blog article is from Deuteronomy, where God promises this to the Israelites. Copper is still very valuable. Valuable enough that pennies are mostly zinc now, people steal copper wiring from walls, and vast operations extract ore from the earth. As a collector, the intense blues and greens are irresistible. And as a rockhound, they are easy to find and identify given the colors!  Malachite, the primary ore of copper, is known for taking a beautiful polish to expose its bubbly bands of light and dark greens. If you have the time, look up the Malachite Room in the Winter Palace of St. Petersburg, Russia.

Malachite is a copper carbonate. Its molecular formula is Cu2CO3(OH)2. In this experiment we will try to extract pure copper from malachite ore. I was too proud of my own found specimens to destroy them (yet), so I bought some bags of polished pebbles on Ebay. For starters, I dumped out 24g of this gem - and then crushed it with a hammer into small pieces. I didn't want to pulverize it completely as that is unnecessary and inhaling copper ore dust can be dangerous. This is what it looked like before and after:

Now, we could smelt the copper out with a furnace, like people have done for thousands of years, but I want 99.x% pure copper and a high yield using a chemical process.

Cu2CO3(OH)2 has a molecular weight of around ~225g/mol, of which ~127g is copper (or ~56% of the ore's mass). With 24g or ore being processed, the best we could hope for for perfectly pure malachite with no impurities (wrong, I can see bits of quartz and manganese oxides already!) is a yield of a little more than 13g. Even if I process it all perfectly, I'll probably lose microscopic bits in my container and filters. So my goal is a relatively reasonable 10g or more.

This process is not hard at all, but I made some mistakes the first time and failed miserably. My "copper" was mostly iron filings that I added later. The "copper" tested successfully with a propane torch - it gave off the green flame I wanted. However, something wasn't quite right and when I put a magnet near it, I knew my copper (which does look grey/black after being oxidized by flame) was almost all iron. I had to start over, so it was smart not to use up all of my ore at once.

(If you want to see some of my beautiful mistakes, look at the pictures from my Hello World post)

The first reaction we want (which created all sorts of odd internet arguments about the balanced equation, reaction type, and by-products) is as follows:

Cu2CO3(OH)2 + 4HCl -> CO2 + 3H20 + 2CuCl2

2 coppers, one carbon, 5 oxygens, 6 hydrogens, and 4 chlorines on each side

Solving for the oxidation numbers, I can calculate that the copper is 3+ on the left and reduced to 2+ on the right.

We will now make carbon dioxide gas, water, and something called Copper(ii) Chloride. This is fairly easy, dump about 50ml of 20 baume (31.5%) hydrochloric acid (HCl) into the container. Instantly it fizzes like crazy and turns to a very dark green liquid in almost no time. 

This almost black liquid is water from diluted acid, water created by the reaction, and a lot of copper(ii) chloride. The latter is some nasty stuff, and I would not get it on too much if I were you. The turquoise-colored crystals on the side of my tub above are copper(ii) chloride dihydrate. By the way, turquoise is another pretty copper mineral, a phosphate instead of carbonate. Here is another view of such residual crystals from my erstwhile experiment:

Back to the dark green liquid. Note that if I had added 100ml or so water to make it a ~10% solution of HCl, it would be a lighter color. This time I just forwent caution and skipped the water. The down side of that is that I could not see well into the dark liquid. That would be a problem later. In the mean time, I discovered a fringe benefit of keeping the water content down. Without any heat or other stimulus, I was shocked to see a dense cap of crystallization quickly form over the solution. These are hydrated copper (ii) chloride crystals that have a higher concentration of chlorine. The first photo shows the crust, and the second zooms in on a bit that I broke and flipped over to see crystals that were growing down into the solution.   

Ok, we are now ready for the second reaction that yields copper. Please read Pt II for that.

Thanks for reading,

Paul

p.s. It has come to my attention that some people would like to repeat these experiments without the hassle of finding the chemicals and/or having to buy too much, etc. I will try to keep an Ebay store alive (seller: cinnabarminerals) that offers low cost and related material for the purpose of enabling education. I will title offerings "Brave New Chemist Science Pack - (related Blog post name)." You can always comment and ask where to buy certain equipment too.

  

 

Head in the Clouds

Let's switch gears a bit to the lowest levels of the atom. We are going to observe cosmic rays (high energy particles from outer-space) as well as radiation from a mineral that I collected once.

Cosmic rays bombard us all day, we just do not see them. The come from our sun, our galaxy, and distant galaxies at the speed of light. These particular high-energy particles have a mass (unlike photons). Most are nuclei of far-away atoms, usually protons (+), and a few are electrons (-). The technical details of cosmic rays are complicated and still being researched. I won't even attempt to describe them further.

We will also observe radiation from a terrestrial mineral source. In this case it is a small rock that I found which has some very small microlite crystals. While pure microlite itself is not radioactive, it tends to contain some traces of uranium and thorium. I identified these crystals with the help of a geiger counter, which "clicks" when an energized particle is released from the mineral and makes its way into its tube.

Radiation.... In Chemistry, what do we mean by that? There are many types of radiation. We are talking about particles being ejected from an atom as it decays. Alpha particles are stopped by a sheet of paper or a few centimeters of air. They are basically two protons and two neutrons (or a helium nucleus). They have a positive charge. Beta particles penetrate paper but not sheet aluminum. They can be positive (positrons) or negative (electrons, in which case a neutron becomes a proton). Gamma particles can pass through lead. They are photons that have no mass or charge. 

We don't need to understand all of this, but you can research further. The main point is that energy particles that are not blocked by the atmosphere constantly bombard us from outer space (cosmic rays), and radioactive materials also emit energy particles (radioactive decay).

Now, we cannot see these particles, but if we create an observable cloud of frozen alcohol crystals, we can see the ionization of those crystals by particles that pass through them. Ionization is the acquisition of a new charge (positive or negative) by an atom or molecule. In this case the energy causes the frozen crystals to dissipate, breaking their chemical bonds, and leaving an observable trail that is much like the contrails of an airplane in the atmosphere.

This experiment required obtaining two tricky things: 90% or greater alcohol (this has been nearly unobtainable since the pandemic started), and dry ice which is frozen carbon dioxide (CO2). Dry ice is not that hard to find but does sublimate, so you have to buy just-in-time. We'll conduct this experiment at night so buy it that day and keep in a cooler. Note that 70% alcohol won't work well, as I learned the hard way.

What is sublimation? It is a frozen solid evaporating directly to a gas without going through a liquid phase. That is why it gives off fog that is popular at Halloween. Blowing on it or heating it in any way increases the fog. Dropping it in water causes the water to "boil" as carbon dioxide gas bubbles are released. Frankly, playing with dry ice is a good chemistry experiment in and of itself. 

The main safety issue here is to not burn yourself by touching (including eating and licking) the dry ice. Use tongs or gloves to handle. Also, the radioactive source (optional) should not be *too* radioactive. This is both for safety reasons and to prevent an "unseeable" number of particles. My sample was chosen because it is small (thus the name microlite) only mildly radioactive.

With this background, on to the experiment!

I used a spaghetti sauce jar (empty and cleaned, of course) with lid. The main thing is to use a container that can easily be seen through, or else you will not be able to see the particle activity inside. I cut a sponge to fit snuggly in the bottom of the jar (it has to stay in place when soaked and upside down). Also cut a circle of black construction paper that fits inside the bottom of the lid. 

Now, soak the sponge with 90% isopropyl alcohol and pour out any excess. Turn upside down. Put the black paper in the lid and close it up Now place the jar (still upside down) directly onto dry ice. Note you will wants a cutting board or something below the dry ice to absorb the shocking cold. You don't want to break a granite counter top. Since the ice never becomes a liquid, you don't have to worry about it melting and making a mess.

Make sure it is night and all lights are off now.  It can be day if you are in a fully dark room. The point is you need no light to see better. 

Wait about 20-30 minutes for the cloud to form inside the jar between the dry ice and the sponge. Use a flashlight to look down at the black paper. You are ready when you see (look carefully) tiny crystals of "ice" in a cloud against the black background. If you don't look at the black background, you likely won't see the tiny clear-white crystals.

Ok, now hold your view on the cloud and wait. Give or take some time, about every 30 seconds or so you will see a contrail form. It may be long and straight or short and squiggly. Those are cosmic rays penetrating the glass and ionizing the cloud in its path. Pretty cool. Sorry, I could get no picture of this.

Now, if you have a radiation source, repeat the experiment with the source on the black construction paper. Every now and them you will see little contrails that appear to be emitted from the source. Mine were short in this case.

Short fat tracks are caused by alpha particles. Long straight tracks are caused by murons. Squiggly lines are caused by electrons and positrons. If your track forks, you have witnessed your particle decay further.

Thanks for reading,

Paul

p.s. It has come to my attention that some people would like to repeat these experiments without the hassle of finding the chemicals and/or having to buy too much, etc. I will try to keep an Ebay store alive (seller: cinnabarminerals) that offers low cost and related material for the purpose of enabling education. I will title offerings "Brave New Chemist Science Pack - (related Blog post name)." You can always comment and ask where to buy certain equipment too.

    

 

A Penny for Your Thoughts - Pt II

  (Disclaimer: This post is for purposes of creating fine artwork! We are not attempting to create bullion for profit.)

In my last post (pt I), I explored getting the zinc out of a penny intact while losing the copper. If you have not read that post first, I suggest you do now for background on the penny composition, etc.

In this post we are going to do the opposite, we will remove the zinc from the middle while trying to leave the fragile copper shell relatively intact.

In full disclosure, I attempted a method that I read about on the internet to remove the copper. I discovered after the fact that this meant removing the copper to keep, not removing to expose the zinc.

This turned out to be a great lesson in the reactivity series which lists the order that elements are likely to react. Zinc is generally more reactive than copper. So in this example the solution we create will attack the zinc instead of the copper.

We will soak the pennies in diluted hydrochloric acid (bought more often as muriatic acid, usually for swimming pools). 

Now I honestly should have stopped and thought why the first step is to score the pennies around the edges. I just took at steel wirecutter and put four notches around each penny at 12, 3, 6, and 9 o'clock. The reason for this is obviously to give the solution some access to the zinc. Otherwise, the copper will just sit there protecting the pennies.

Into my beaker, I poured about 50ml of 20 baume (31.5%) HCl. Then I added roughly 100ml of water to dilute to around ~10%. Now drop the scored pennies into the diluted acid and watch the bubbles until the reaction is done.

I never got the reaction equation from the internet because I was reading about attacking copper, not zinc. So I will attempt to determine the equation here. 

We know the reactants are Zn and HCl in H2O, so it should be something like...

Zn + HCl + H2O -> a vigorous gas, a compound of Zn, and water

I am fairly certain this reaction, balanced, is:

Zn + 2HCl  -> ZnCl2 + H2 

Once again, note this reaction forms hydrogen gas most likely, so please do this in a vented area and take all safety precautions. I made a lot of gas doing this!

ZnCl2 (zinc (ii) chloride) has no natural analogue. It is supposed to be a white powder, but I made a mess of blackish stuff. This could be for any number of reasons, but most likely that I did not fully react the zinc before finishing my experiment. The black is likely remaining powdered zinc and the white powder dissolved in the solution.

Copper does not react because hydrogen is more reactive than copper, so it would stay bound to the chlorine. Zinc does because it is more reactive than hydrogen and thus swaps places in what is known as a single replacement redox reaction.

The elemental zinc with a neutral charge loses two electrons (oxidation), and the hydrogens gain one each (reduction) so that it can bond and form hydrogen gas.

We can remember this as OIL RIG: Oxidation is loss, Reduction is gain (loss/gain of electrons)

When both oxidation and reduction happen together, we call it a redox reaction.

When the reaction had slowed down, I fished out the pennies. They were very lightweight and fragile. Recall the copper shell is only .06g. Mine weigh a bit more because I did not fully oxidize the zinc.

Thanks for reading,

Paul

p.s. It has come to my attention that some people would like to repeat these experiments without the hassle of finding the chemicals and/or having to buy too much, etc. I will try to keep an Ebay store alive (seller: cinnabarminerals) that offers low cost and related material for the purpose of enabling education. I will title offerings "Brave New Chemist Science Pack - (related Blog post name)." You can always comment and ask where to buy certain equipment too.

 

 

A Penny for Your Thoughts - Pt I

 (Disclaimer: This post is for purposes of creating fine artwork! We are not attempting to create bullion for profit.)

Everyone knows the "Liberty Penny," the one that looks like it is made from copper and has Lincoln on one side. Due to the cost of copper far exceeding a one cent, in 1982, the US government modified the penny to to have a zinc core (97.5%) with copper plating (2.5%). A penny weighs 2.5g, so about 2.44g is Zn and the remaining .06g Cu. It is not an alloy (mixture), but rather one metal over the other.

So naturally, I wondered, can I get the zinc out of the middle? 

Yes, it turns out there are several ways to do this. The easiest is to simply hold the penny over a gas flame (with tongs), and when it starts to bloat, give it a little shake over a container full of water. If you do it correctly, the copper shell will be left in the tongs and the molten zinc will fall out to cool in the water. This is easy, but maybe only 99% pure or so. It is easy to accidentally break a little copper off in the process. Another problem is your zinc will be shaped like a meteorite or tektite, losing the penny shape/details. Finally, you will not get all 2.44g out with this method. Some molten zinc will undoubtedly remain attached to the copper shell.

Here is a picture of zinc obtained by this method. The flow lines where it entered the water, radiating from what is called the stagnation point, are easy to see at the bottom:

Here is what the remaining shell looks like (this one broke up):

You may ask why the zinc melts but not the copper. Easy. The melting point for zinc is 787F and for copper it is 1983F. You average gas stove, Bunsen burner, or propane torch will easily melt the zinc, but can't get hot enough for the copper.

So let's say we want a perfect zinc 2.44g penny, not a flying projectile cooled by water. We need chemistry!

We need something that will eat away at the copper, but leave the zinc alone. Sulfur is our answer. We need to turn our copper into CuS (copper (ii) monosulfide, which in nature is the highly collectible mineral covellite).

The exact reactions we will use start with preparing a solution of what is known as "lime sulfur," which is used as a pesticide:

Ca(OH)2 + S -> CaS + CaS2O3 + H2O 

(calcium hydroxide and sulfur react to make calcium sulfide, calcium thiosulfate, and water)

When we balance the equation we get...

3Ca(OH)2 + 4S -> 2CaS + CaS2O3 + 3H2O

I admit this one took me awhile: 3 calciums, 4 sulfurs, 6 oxygens, and 6 hydrogens on each side! The internet had varying equations here, but I think mine is correct.

In this early experiment, we won't calculate the exact weight of reactants needed... In a beaker, add a gram of calcium hydroxide and 2g of sulfur and some water so we can boil. The mixture will not dissolve as the sulfur floats. We need what is called a wetting agent. So add a drop or two of dishwashing soap and mix.

Incidentally, soap is a substance that makes greases, oils, and dirt dissolve in water. That is exactly why it exists, although usually most useful for washing your hands and body. In this case we are "connecting" water and sulfur into our solution.

Now, add some heat by putting this under a burner until it comes to a boil. We need energy to make the reaction occur. It will change colors from a light yellow to a dark orange as it reacts.

The gas released is *theoretically* steam from water boiling, but best to do this outside or under a vent. See below for more detail on this.

Now, add a few pennies (I used four), dated 1982 or later, and wait 30 minutes or until the pennies are fully coated with a black substance. This is copper (ii) monosulfide (or may I say lab created covellite?). The copper has been tarnished fully through. Note that excess reactants and by-products of the first reaction may remain in the solution. In my case I used too much sulfur.

 

I believe the balanced reaction should be:

2Cu + CaS2O3 -> 2CuS + CaO + O2

I may be wrong, the internet had no information on this reaction that I could find. Note we produce oxygen gas, and steam is released from our boil again. Again, do this in a ventilated area. I have made mistakes before and created hydrogen gas and carbon dioxide. Assume the worst. With water and excess sulfur, it is likely to product harmful hydrogen sulfide gas.

Remove the pennies and scrub them with a soft scrub, like ceramic cooktop cleaner. They should look a very shiny silver color and weigh 2.44g after a rinse. If you have bits of copper still attached, cook them some more in the solution and repeat. If you have black deposits remaining, soak and scrub again. It should not be too hard to clean the penny completely of its copper plating.

Warning: the black CuS and soft scrub can make a bit of a mess!

The final product(s) look like this:

You can weigh to prove the theory:

Thanks for reading, 

Paul

p.s. It has come to my attention that some people would like to repeat these experiments without the hassle of finding the chemicals and/or having to buy too much, etc. I will try to keep an Ebay store alive (seller: cinnabarminerals) that offers low cost and related material for the purpose of enabling education. I will title offerings "Brave New Chemist Science Pack - (related Blog post name)." You can always comment and ask where to buy certain equipment too.