What's Eating Gilbert Grape?

I wanted to make a simple battery to illustrate a voltaic cell. Put simply, two different metals, connected with a conductive wire, are inserted into a salt bridge (sea of electrolytes) and current (electrons) flow from one to the other. Because one metal is more electronegative that the other, and the salt bridge contains both positive and negative ions, we get a voltage between the two metals and flow a current in a loop between them and through the salt bridge.

In this case, I took a whopping five minutes to get what was most readily accessible to me and measure the voltage of our cell. I found a copper wire, a galvanized nail (meaning coated with zinc), and a grape. 

The copper has an electronegativity of 1.90  and the zinc coating the nail 1.65. Therefore zinc will be oxidized (lose electrons) and the copper will be reduced (gain electrons). Current will flow from the anode (nail) to the cathode (copper wire) through my voltmeter which will measure the electrical potential difference. inside the grape, there are several electrolytes. I believe the most common for a grape is Potassium bitartrate (or KC4H5O6). The K cation has a +1 charge and migrates to the cathode (copper). The bitartrate anion has a -1 charge and migrates to the anode (zinc). Other acids such as citric acid and other ions such as sodium exist in the grape.  

Electrical reactions are as follows:

Cathode: Cu2+ + 2e- -> Cu (standard electrode potential of .34V)

Anode: Zn -> Zn2+ + 2e- (standard electron potential of -.76V)

Together the potential difference is 1.1V if these reactions are under normal conditions (sufficient concentration and ion mobility at 25C).

My voltage with a cold grape measured .871V but was still slowly climbing. When I breathed on it it would always bump up a little, perhaps because the extra energy of heat put into the system allowed for more reaction (the grape was colder than 25C). Normal conditions for a grape may be more like .9V but I can't find a clear reference for this. I picked the one fruit nobody seems to have measured and published! 

Thanks for reading,

Paul

p.s. I redid the test with another firm grape, this time at room temperature (~25C). It measured .901V. It is possible that if the grape had ripened more, the voltage could go up or down from there. It is also possible that my Thompson seedless (sultana) grape varies a bit from other grape varieties. Winemaking is basically the chemistry of fermenting grapes, which like soap and baking is a whole other profession with thousands of years or art and science involved. For now, I am claiming .9V as the normal conditions for a grape with zinc and copper electrodes. 

   

 

Wash Your Mouth out With... Lye?!

So this is not really a 'modern' experiment of mine. A couple of years ago I got interested in making soap from lye and various oils. I used a calculator for soap based on which oils and how much of what, etc. I made quite a few batches of different types, and most are still around today, such as this one:

 

They are still around for a few reasons:

1) I made a LOT of it.

2) It is fairly dense, not being full of injected air like Dial and other commercial brands. A bar lasts a long time.

3) Nobody else trusted my "lye" ingredient, so I am mostly the only user. Never mind that Dial has lye in their process too. They list the ingredient formed with lye such as 'sodium palmate' or similar. That is essentially lye and palm oil combined. They lye! 

4) I made some pretty ones that serve as decoration in various homes of friends and family.

Anyway, I will not discuss the process or recipes in great detail. Like bread making, this is an art form unto itself. It might be months before I come out the other side. Most of my formulations work really well, so it isn't that hard to make good soap. I tried just about every oil and combination thereof. I trust my soap more than Dial for keeping the 'Rona away.

So this post will be more about the general chemistry.

We start with the ever-useful NaOH, sodium hydroxide. As we know by now, it dissolves in water easily to form Na+ and OH- ions.

But you know what doesn't mix with water? Oils. They are hydrophobic as they are not polarized. Remember water itself has H+ and OH- ions floating around, both with charges.

This creates a problem. We form oils on our body, but can't wash them off with water because they don't mix. Oils are nonpolar. We need an oil to wash away the oils, but that sounds like a contradiction in terms. It kind of is except for the process known as saponification. It bonds fats and oils and lipids with an aqueous alkali such as NaOH. Heat speeds up the reaction. The result has polar metal ions (so it can dissolve into water) and also the fatty acid that can bond with other oils and grease and dirt particles that don't want to wash off in water. At a high level, think of soap as a peacemaker between oil and water.    

This can be done with KOH, for example, although I never tried it. Actually this is known as sailor's soap, as the K+ ions work better in saturated saltwater than more Na+. Just about every form of metallic 1+  ion and oil/fat/lipid have been tried. Whale blubber used to be used for soap making. Lye was extracted from wood ash. Presumably they all work, with varying properties such as how sudsy they are. 

Now you also know why NaOH has a soapy or greasy feel to it. It is literally forming small amounts of soap on your skin as it bonds to our body oils.

Thanks for reading,

Paul

p.s. Soap can also act as a surfactant to get water and things that don't dissolve it in to mix. We did this in an earlier blog post when the sulfur did not mix with water - so we added some soap and it mixed happily. Now you know the water was attracted to a metal ion and the sulfur was bonded to fatty acids. When we heated the mixture enough for the reaction, between Ca(OH)2 and S, the soap just boiled off and/or remained as a harmless bystander in our solution for further reaction between one of the products and copper. 

Overreacting to Everything

People overreact a lot these days. I'm sure you know someone who reacts to everything with extreme positions. Well, in the world of elements, that person would be sodium. It pretty much reacts violently with anything. Sodium justice will not be denied I vaguely recall dropping sodium into water as a high school student, but all of that memory is fairly hazy. I had to relearn.

I wanted to practice a little electrolysis, where I run a current through a solution and collect the cations on the cathode (positive terminal) and collect the anions on the anode (negative terminal). I had just a little sodium hydroxide available, and a really small evaporating dish, so a small experiment seemed quite possible. In fact, it was MUCH faster and easier (in some ways) than I ever imagined.

I dumped about 10mg of sodium hydroxide pellets into the dish and fired it to a melt within seconds with a blowtorch. Then I dropped a pair of nails (zinc galvanization sanded off) connected to my car battery charger into the melt. I was careful not to electrocute myself. I had the charger set on 12V/6A at first (vs. the 6V and 2A settings). I figured that 10mg would yield at most 5.1g of sodium, so this small scale experiment was plenty.

The first thing I noticed is that the NaOH made a really nice and bright yellow flame when fired directly. Yes, this is the sodium flame test. The second thing I noticed is that it tended to freeze while I got the nails inserted and the charger plugged in. The solid was not conductive enough to melt the NaOH. I needed to put the nails in and then give it a really solid boost with the blowtorch. Pretty soon the electrolysis took over and sustained the liquid solution between the nails. Electricity does not meander around the long way, so the NaOH remained solid around the edges of the dish away from the nails and the path between them.

I wasn't sure what to expect. Sometimes things take hours, but almost immediately the current meter was showing was 2-3A, I had a nice production of gas bubbling at the anode, and I could see pellets of silver liquid metal forming and then turning whitish around the cathode. When it really got going, the little pellets did more than turn white: they started exploding (like minor fireworks) and occasionally whizzing around surface of the dish. This was pure sodium reacting to air and some water formed at the anode. I had to back away and unplug the charger. I did not wish to have any burn me or react with my skin.

Note that it was raining in 80-degree plus weather as Tropical Storm Beta moves onshore nearby here. There was decent humidity in the air, at least 1% water vapor. The odds of me collecting any nearly pure sodium were low. I had a bottle with mineral oil and tweezers handy, but I was also photographing and occasionally trying to film. I was also a bit careful about picking the sodium out of the melt for fear of it reacting with the air. I managed to get a few drops into the oil, but most of it reacted more than I wished before I got it there. Still, I did get some photographs, some grayish partially-unreacted metal in oil, and a nice video or two. The one thing I never got was a video of it popping and sizzling around - I was more concerned with stopping this than filming it. I do wish I had a video record of that though!

What are the reactions involved?

At the cathode (+) nail, 2Na+ + 2e- -> 2Na

The sodium ions are reduced.

At the anode (-) nail, 2OH- -> H2O + 2e- + .5O2

The hydroxide is oxidized.

A further "reverse" reaction occurs: The Na that comes into contact with water creates the original ions and a little hydrogen gas.

Na + H2O -> .5H2 + Na+ + OH-

Here is the electrolysis is action. the cathode is on the left side. The whitish looking blobs near the cathode are pure sodium. It looked more silver in real life. If you look close you can see bubbling everywhere, ostensibly H2 gas and O2 gas and some water vapor. I can't really explain the blackish/gray collection around the cathode other than sodium formed here, presumably much of it in microscopic form. Much of it reacted with the atmosphere to turn back into NaOH, but clearly bits of it just disappeared back into the melt and cooled off. I'm not totally sure why the anode side and much of the melt turned yellow/brown. Maybe this is what the solution looks like when deprived of some but not all of its sodium. Perhaps there is some carbonate reaction occurring as well. It seemed a little darker than I expected.

This closeup shows it better, including the cooled NaOH on the sides.

And this seemed really cool to me. The cathode removed (and still a bit hot) has a bit of silvery sodium on the bottom right, still protected by some liquid that is still reacting. Note the whitish-yellow sodium hydroxide near the top and the gray that I think are bits of sodium trapped in sodium hydroxide.

Finally, here is what I collected below. It is really hard to photograph, and this is far from pure sodium. Each little drop I tried to collect turned to NaOH, some partially and some completely. None of it is mercury-like in all of its metallic splendor. Nonetheless, the grayish ones that are not fully reacted will still make a nice little reaction with water, I am told.

By the way, what is mineral oil? It is a by-product of refining that is made of carbon and hydrogen without oxygen. It is generally inert for collecting species of reactive metals and minerals and storing hydrated things like iron and opal without them oxidizing or dehydrating.

Here is my impure metal in mineral oil:

This was a quick and fairly easy experiment, despite the risks of electrocution, sodium burns, and fire. While I could have probably done better, I think it proved that NaOH really has sodium, electrolysis really works, and sodium is extremely reactive with air, water, and just about everything but mineral oil. This makes me wonder about the sodium ions in our diet, from table salt and other sources such as baking soda. That is some reactive stuff pulsing around our bodies somewhere. No wonder we don't want too much of it.  

I assume the same experiment would work with potassium hydroxide. I'll have to look that up.

Thanks for reading,

Paul

 

     

 

  

  

Bismuth!!!! Pt III

This is the last post on Bismuth hopefully. I am not satisfied with my crystal growth, but unwilling to procure vast amounts of material and equipment to improve further.

 

These are the best crystals that I made apart from my geode method. If you look closely, you can see the stair-step pattern of hopper crystal growth in several places.

The latest geode is more interesting. This time I let my bath of molten Bi cool for 15 minutes inside a blanket of kaowool refractory insulation. Honestly, I thought it would never cool down. After plunking the above from the top of the melt, I poured the remaining liquid out and had a thick geode that still had relatively small but fairly interesting crystal growth. So as to see inside better, I used pliers to break the top edges and overburden off. This exposed an interesting simultaneous view of the silver (unoxidized) crystallized metal. Finally I sanded the top smooth. Here is what it looks like:

Here is a close up view of the floor of the geode:

I especially like the largest formation where the crystal obviously grew up and out in a spiral pattern, like a nautilus shell (which grows according to Fibonacci's mathematical sequence).

Note that even the edges are not completely parallel at 90 degree angles. The crystal lattice itself is rhombohedric and is known as pseudo-cubic because it is not exactly cubic. I wish I could explain the shape and pattern in great detail but I can't get from the unit cell of Bismuth to this in a fully logical way.

What I can explain is that the hoppered nature is clearly evident in almost every crystal. The edges grow faster than the crystal can fill in the middle. Halite (salt) is another mineral that is famous for hoppered crystals. The Bismuth cools slowly enough to form large crystals easily, but they still grow so fast (think in geological terms) that they don't have a chance to fully form inside the edges.

Fibonacci's Sequence is x(n)=x(n-1)+x(n-2). This is not a math blog but the pattern is 0,1,1,2,3,5,8,13 and so on...

Does the largest crystal obey this sequence?  No, not quite. I can see evidence of 1,1,2,3 when blown up and measured, but it falls apart before and after these four edges. That being said, I googled it and there are numerous cases of Bismuth crystals being referenced as following Fibonacci's Sequence. I think, though, that it may be an assumption to describe a given crystal. I am not sure a scientist would make this claim. But nature does have its patterns, and crystals are at least related in this sense.

The unit cell for Bismuth, the smallest pattern that replicates, has two axes the same length, and the third in-between two and three times that length. Atoms are so small that these distances are measured in Angstroms. Still, that could help explain ratios that may look similar to 1,1,2,3...

Now, there may be something else intriguing related to the unit cell. There are three angles in the symmetry of the Bismuth unit cell, between the axes whose lengths are discussed above. Two are 90 degrees and the third is 120 degrees. I definitely see angles close to if not exactly 120 angles in numerous crystals towards the center as the edges formed. I am not sure if they are related, but it seems more than a random coincidence as it is frequent and far from "cubic."

If anyone can shed more light on these patterns, I would appreciate a comment.

Thanks for reading,

Paul

     

Catalyst

I've always heard the term catalyst and wondered how they work. The most common term people know is probably catalytic converter (a thing in your car that improves emissions) which uses Platinum and/or Palladium as catalysts. People used to steal them to try to extract the precious metals. Maybe they still do.

I generally understood what a catalyst is, but wanted to experiment and see them in action. I thought I would pick the very well-known and widely understood metal oxide added to hydrogen peroxide to speed its decomposition into water and oxygen.

    2H2O2 -> 2H2O + O2

It turns out that it is very well known, especially with MnO2, but remains an active area of research for the exact series of reactions involved. The best-looking answer I found on the internet initially for MnO2 was completely wrong as it left MnO as a product. A true catalyst speeds the reaction but is unchanged; it is not modified by the completed process even thought it may be temporarily modified during the reaction. You should be able to reuse the catalyst. Luckily, I don't have to take my car down to Jiffy Lube for a refill of Platinum.

For kicks, I had some Fe2O3 (hematite) laying around, and decided to bake-off the natural decomposition of H2O2 vs. with MnO2 vs. with FeO3. I had done no research on hematite as a catalyst for the reaction - I just thought it might work. In retrospect that was probably silly as Fe2O3 is abundant and cheaper than MnO2, so I would probably have found more on iron oxides as a catalyst than manganese oxides initially if it worked just as well.

OK, so I dumped some hydrogen peroxide into two beakers. Not much happened. There was an occasional bubble. Not enough for me to have even thought a reaction was occurring, but sloooowly. Then I dumped some MnO2 into one and it went crazy making oxygen gas. The beaker also became very warm. This is clearly an exothermic reaction! Next, I dumped a bit of red iron oxide into the control beaker. I didn't observe much. There was some bubbling, but I'd have to compare to video to try to determine if it was any faster than before.   

What is going on with the MnO2 reactions?? I was a bit disappointed to find no good reference on this on the internet. The ones I did find were vague or detailed but clearly wrong. A catalyst creates a series of reactions that end up with the catalyst unchanged, the same products (water and oxygen here), but has a lower activation energy so that the reaction can be done faster with the same energy or with less energy. In my case, I was adding no energy to the system beyond room temperature (which is what the hydrogen peroxide was stored at in the first place). I expected the rate of reaction to change.

Here is a handy artwork I made to show the H2O2 reactant, it's formation enthalpy, the normal activation energy, and the products with the exothermic release of energy (delta H) calculated from the formation enthalpy of liquid water. Since the activation energy with MnO2 is known, I charted that as well and tried to approximate a couple of reactions along the way in a second line on the diagram. But later I read that there could be five reactions involved.

After trying much math and good searches, I did feel better to finally find a recent academic paper which states that the exact series of reactions is unknown and proposes five reactions as a possible solution. When you are having trouble finding a solution, it is always nice to know what you thought should be easily discoverable is actually unknown to science. I feel less stupid. One thing I so know is the "slowest" of that series of reactions must have an activation energy of 58KJ/mol to match overall lab results for the catalyst. But I'm not going to repeat the research in detail here.

So what about my iron oxide (with Fe(III) ions instead of Mn(IV) ions)? Further research uncovered that this is also being actively studying for potential performance in different environments (temperature, pH level, etc.). It does seem that the effectiveness of iron oxide is not nearly as good as that of MnO2. I noticed what appears to be some bubbling at a greater rate than before, but nothing worth writing home about.

I once visited a vast field of manganese oxides in New Mexico. I have been to mines in California and other places. The ocean floor is littered with manganese oxide nodules. I would love to have collected some of these natural minerals (pyrolusite - MnO2 - among others) and try them out. But these things are mostly just ugly black and so I didn't bring much back or keep track of it. Maybe some day I'll run to the Quick-E-Mart and buy some hydrogen peroxide if in the areas again. Like taking ore to the coal mine, or Mohammad to the mountain, it seems that taking a bottle of liquid to the ore field is the easier task. I would pour it on the black ground and it should fizz wildly. In the process I water any plants and add oxygen to the New Mexico sky.

Thanks for reading,

Paul

It's a Liquid, it's a Solid, it is Ferrofluid!

I attempted to make my own ferrofluid. I failed. First I bought some printer toner but it was not magnetic enough. So then I bought more toner, this time magnetized for check printing (MICR toner). I mixed with vegetable oil at different thicknesses. I made a magnetic suspension fluid but it did not form cool shapes quite like I expected. I'm not sure what the problem was, but I made a HUGE mess and spent a week trying to recover my beaker and magnets, etc. I decided that I would break my personal rule and buy the stuff ready-made. I will not share my formulations as they all failed to impress me and made a massive mess. If you have never dealt with printer toner, I don't recommend you start now. It is ink, and simply opening a bottle of it causes and ink cloud to form and settle on everything around it. Someone once joked that if you sneeze, you may have to buy a new house.

What is ferrofluid? Well, as NASA envisioned it, it is a liquid that can be controlled by magnetic fields. So it has properties of two phases, liquid and solid. In order to make ferrofluid, you need very small (10nm or so) particles of magnetite (magnetic iron oxide), polarized in a surfactant (like soap) and suspended in a carrier liquid. The carrier and surfactant can be one substance in the case of thin oils like mine, but NASA's formulation was hopefully more sophisticated.

Since I skimped on cheap toner, maybe my particles were too big. 10nm is really small. Again, I'm not really sure. I know from a job I had once that growing consistently small crystals on that scale is quite hard. You have to stop growth before the crystals are even visible.

Here is a picture of my new purchased fluid above a medium sized rare earth magnet:

It is a big hard to see black against black, so here is a view of the same picture where I edited the light and contrast:

Note the characteristic spikes arranged in a beautiful 360-degree flower shape. The shape corresponds perfectly with the magnetic field, as expertly drawn by me below. Note the field follows this pattern in 3D, not just to the right and left of the magnet as in my drawing.

And here is another picture with a smaller and therefore weaker magnet:

Note the field is weaker, so it can sustain fewer spikes and they are tending to bend down according to gravity and the smaller shape of the magnetic field.

You may be wondering why the spikes form so precisely. I did, so I looked it up. The particles are so small as to move about and bounce off each other randomly in the same way molecules do in liquid. This is known as Brownian motion. Thus the liquid can arrange itself perfectly with force fields according to probability's directions. As we showed in the magnetic field diagram, the spikes correspond to these magnetic field lines in 3D. However, they are limited by gravity and surface tension. The tips of the spikes are the point when the opposing forces are equal. The spikes grow longer with a stronger magnetic field. The direction of gravity can affect them. The largest and sharpest spikes are formed when surface tension of the liquid is minimized. 

If you wish to try something similar and not buy the stuff or figure out how to make it, a quick and easy thing to do is to dump some iron filings on paper or glass above a rare-earth magnet. Don't let the filings get on the magnet! The much-bigger filings align with the magnetic field in a similar way, generally showing the magnetic field shape. They cannot reposition themselves ideally because they are not small enough, not suspended in a fluid, and make physical connections with each other (like little magnets stacking themselves upon each other).

One last thing. What is magnetism? I am oversimplifying this because it is kind of complicated and honestly I don't fully understand it yet. For our purposes here, some materials have unpaired electrons. Iron is one such substance. Another is Magnetite (due to a combination of both Fe(ii) and Fe(iii) cations in its molecular makeup). Paired electrons cancel out a magnetic property known as spin. Unpaired electrons retain this property and are attracted to electromagnetic fields. While electromagnetic fields are usually the realm of physics, I guess my point here is it all comes from the chemistry of the substance.  

Thanks for reading, 

Paul

p.s. I kind of butchered some of the science for simplicity. In reality, what I am calling magnetic is paramagnetic (other than the magnets themselves). Paramagnetic material like iron filings and this ferrofluid can be attracted to magnetic fields spontaneously. Paramagnetic material has one or more unpaired electrons. Paired electrons oppose each other's "spin," and therefore magnetic properties. This means the bonding of molecules and even between molecules in crystal lattices makes a difference in whether electrons are paired, and therefore, whether they are paramagnetic. O2 is paramagnetic but N2 is not. Without getting into Molecular Orbital Theory (MOT), it is easiest to determine the elements and ions. Iron has four unpaired 3d-shell electrons (according the Hund's law). Iron +2 or (or Fe(II)) also has four. Iron +3 has 3 unpaired electrons. In a compound, note iron oxides hematite (diagmagnetic) and magnetite (strongly paramagnetic) are determined by MOT and not just their metal ions by themselves. A common question is why copper, silver, and gold are not paramagnetic. All three line up in group 11, meaning they have one unpaired electron, but the large size of the atoms means the many filled shells' diamagnetic properties overcome the attraction of the one unpaired electron. I believe MOT is beyond AP-level Chemistry.

      

 

The Black Snake

I wanted to make the Pharaoh's Serpent, but then I read how toxic it is and I had no plan to manage that properly. In researching this, I saw the much less cool Black Snake experiment as another type of intumescent reaction. An intumescent is a substance that swells with heat exposure. These can come in handy as fire retardants. Both "snake" reactions are of this type where a substance appears to grow out of a hole as it expands. I have never been to Diwali in India, but apparently this is one of their favorite firecracker types. Personally, I have no recollection of these fireworks at all (but also am not that experienced). 

The way the reaction goes, sucrose (powdered sugar)and sodium bicarbonate (baking soda) are mixed in a  4:1 ratio, added on top of sand soaked in a fuel such as lighter fluid (butane), and then the fuel is ignited. In my case, I used isopropyl alcohol as the fuel. Fairly quickly, several reactions start at once:

1) The fuel combusts, making carbon dioxide and water vapor and - most importantly - heat.

2) As the baking soda heats up, it releases carbon dioxide gas and water vapor as it decomposes to sodium carbonate. This is known as thermal decomposition, and is a key reaction in baking.  

3) The sucrose combusts much like the alcohol, again creating carbon dioxide and water vapor; however, some of the sucrose forms carbon. Remember burning creates lots of different reactions. See below.  

C12H22O11 + 12O2 -> 12CO2 + 11H2O

C12H22O11 + 11O2 -> 12C + 11H2O

So what happens is pure carbon and sodium carbonate are formed as solids, but pushed up and out by the carbon dioxide gasses. It cools as a long continual "ash" that can look like a black snake slipping crawling out of a hole.

I tried this, but it wasn't hugely impressive. I think my dish was too small and after a while the flames were choked out by the sand losing its exposure to air. It could be that I just did not use enough reactants to get a long reaction and product. I went for a relatively small effect as I didn't necessarily want a huge long pile of ash and I had limited 91% alcohol to conserve. Perhaps I ran out of that and lost the chain of exothermic reactions.

Note that this is super kid-friendly except for one big caveat: the fire. The biggest mistake made is people add alcohol to the flame and it splashes and spreads flames which have been known to badly injure people. Never do that! Of course, in general it is easy top burn yourself or something else with fire. That much is obvious. I'm not sure it is much more dangerous that barbecuing, however. The only real difference is we are adding gasses from within to expand the ash.

This is what my snake looked like when it went out after about 10 minutes.

As you can imagine it would look better if longer and thinner. I think my dish was too small. The video is not quite impressive enough to share. Some videos online are sped up. Nevertheless, you can probably play with the process a bit and get much better results.

I thought about using my new ash for gunpowder, but it contains a fair amount of sodium carbonate, I think. I'm not sure what that would do to the reaction. All of these experiments involving an oxidizer and sugar or charcoal are very similar reactions. I might be better off just burning sugar directly to get carbon. The only twist we really added here, chemically, is the decomposition of baking soda. I think it might have also slowed the burn rate a bit by displacing some oxygen with carbon dioxide.

Thanks for reading,

Paul