All That Glitters Pt 1

 ...is gold in this case!

I have been doing this one slowly for a while since I decided to "refine" gold from some old broken electronics I was going to throw away (Compact Flash reader, etc.). Frankly, there was not much scrap and mostly only some pins were plated with gold, but it is good enough for an experiment. I am not trying to run a profitable gold reclamation enterprise. 

In this first post, I will just focus on getting the gold "foil" separated from the other electronics. Gold is typically plated onto some components such as pins and pads and fingers in connectors. These platings are very thin indeed. They exist only over the primary copper conductors so that oxidation will not interfere with the ability of a connector to make solid electrical connections over a long period of time. Yes, gold is a great conductor, but that has nothing to do with it. It is ideal for maintaining the ability to connect. 

Silver is an even better conductor, but tarnishes easily in the presence of sulfur as we read previously here. Copper forms blue-green oxides and hydroxides and so forth quite easily. Look at your pipes. These minerals do not have the same properties of conductance and malleability and ductility (yes, that is a word). They get brittle and they don't conduct electricity. The simple solution is to plate with a few atoms of gold, much like galvanizing steel with zinc.   

It does not matter what is inside wires and solder joints and underneath the gold as long as it is meets the basics needs of good-enough conductance for a decent price and environmental impact. If the outside of a wire or solder joint oxidize a bit, it really isn't terrible because the inside is still doing its job. The connector is the only place where two things made separately need to touch and conduct from the outside.

I am getting off topic, but sometimes corrosion on car battery terminals causes the vehicle to get an intermittent connection. In this case, you clean the terminals with baking soda and water, reconnect,  and you are good. But that is not practical for hundreds of microscopic connections inside a computer. 

I should mention briefly that the group 11 metals are all excellent conductors because of a quirk that fills their outermost d-shell completely with 10 electrons but leaves their outermost s-subshell only half full with one electron. That electron is freely given up, without any interference from the other 10 electrons in the d subshell, enabling many wonderful properties of these three prized metals. You might expect that the s subshell be filled and 9 electrons to inhabit the d subshell, but it doesn't. There are some theories on why this is not the case, but the best one is "because it just is."   

This topic will likely generate lots of posts, so I'll stop here for now and explain how I made gold flakes in green liquid in another post. Then we can probably continue on to refining the gold for purity (because we can!) and possibly melting it into a button. 

Thanks for reading,

Paul

Long John Silver

I wanted to explore oxidation, reduction, and electrochemistry a bit more. Luckily, I stumbled across an old tarnished silver ring that my wife needed cleaning. Silver tarnish is a coating of the mineral acanthite, a grey-black compound of Ag2S. It forms on silver over time by the following equation:

    4Ag + 2H2S + O2 -> 2Ag2S + 2H2O 

As a minor note of interest, there is more H2S in the environment (such as your home) than you might think. Eggs are a source, which is why you don't use silverware to eat eggs. In any case, this ring is over 80 years old and was quite tarnished. I didn't quite realize how tarnished it was, but more on that later. Here is a picture:

Anyway, we need to reduce the silver. Ag2S contains two silver atoms that have lost one electron each (oxidized). The silver therefore needs to be reduced, adding the electrons back so that it will let go other the annoying sulfur ions. 

    Ag2S + 2e- -> 2Ag + S2-

It just so happens that the standard reduction potential of this half-reaction is 0.69V. So we will use electrochemistry with a reduction potential above this voltage to reduce the silver. 

It just so happens that zinc has a reduction potential of 0.77V. So, in theory, I can melt enough of a penny to expose its zinc core and put that in a solution with various ions. We then place the ring in conductive contact with the zinc so the electrons can flow into the acanthite. For style points, I attached a copper wire to the half-melted penny and placed the ring onto the wire. The copper is not strong enough (0.34V) to reduce the acanthite but will conduct the electrons from the zinc to the acanthite. It would have been simpler to just melt the zinc core out and place the ring on that, but I wanted to prove the conduction theory.

The electrolyte solution I used was 100ml of water with 10g of salt (NaCl) and 5g of baking soda (NaHCO3). It really doesn't matter which ions are used in the solution. In this case, the cations are all Na+. The anions are a mix of Cl- and HCO3-. This "recipe" is common on the internet but I can't see what difference it makes personally. It should work with just one or the other salts in theory.

Here is what the setup looked like:     

After a few minutes, not much appeared to be happening. I used a voltmeter to make sure electrons were flowing into the ring. They were. After a while, I decided it was working - just really slowly. I removed the ring at this point:

Aluminum has a bit more reduction potential, 1.66V. So I put some scrap aluminum foil in the bottom of the solution instead of the zinc/copper and laid the ring onto the foil:

Again it seemed to be working but not quite fast enough for my desires. 

So I decided to just pump a massive number of electrons into the ring with an actual battery. I took a 9V and attached the anode to the aluminum foil. For the cathode, I attached it to some steel tweezers. That was an unfortunate choice of cathode, but more later on that. I just needed to conduct electricity into the solution.

As you can see in this video, the reaction is vigorous but still takes some time. The sulfur ions released kept clouding my solution with yellowish material bubbling to the top. Perhaps this was sodium sulfide, as both cations in the solution were Na+. In any case, I had to keep replacing the solution. Perhaps I didn't *have* to, but I wanted to see the progress.

In the end, I decided not to completely clean the ring. It was looking pretty good and had a nice antique-looking finish with tarnish mostly in crevices. Plus I had already proven the chemistry and would need even more solution and time and to continue. Here it is:

In retrospect, that ring was very tarnished and not suitable for a simple bath solution without external current applied. Unfortunately I didn't have anything lightly tarnished that I was willing to test. Some antiques and coins can lose value if you clean them.

I mentioned a mistake of using tweezers in the solution. While they were a cheap pair, I more or less ruined them. The sulfur ions adhered themselves to the tweezers as it gave off electrons from the iron. I ended up with what appears to be a yellow-green iron (iii) sulfide (Fe2S3) coating on my tweezers. This material was very hard and I had trouble removing it with a steel file. I consider the tweezers mostly ruined for chemical use and/or aesthetics now. I should have just used a bit of copper wire and then throw that away. 

Thanks for reading,

Paul  

 

Magically making 3g of Salt from 2g of NaOH

This is an easy-peasy little experiment to post while bigger and better things progress. In theory (which I didn't even have to verify on the internet), I can make "table salt" by adding NaOH to HCl. I'm using 31% HCl and 100% NaOH so I am going for the big reaction. 

    HCl + NaOH -> H2O + NaCl 

Basically, the ions switch partners... making simple water and salt from two extremes of the chemical world. Since 31% HCl is not fully disassociated, technically its pH is 0. Since I am using solid NaOH, its pH is 14.

This is a classic acid/base reaction where the H+ ions from the acid form with the OH- ions from the base to form water. The Na+ and Cl- ions happily bond in the aftermath. 

See what happens when I drop one of the two grams of NaOH into the fuming HCl:

It gives a serious pop (and the solution heats quite rapidly). I should note dropping even a single grain smaller than a typical grain of sand into the solution makes a noticeable impact. Immediately the NaCl forms in the solution and falls to the bottom. It cannot dissolve in this concentrated solution.

I did not want to waste all of my NaOH to neutralize all of the acid, so when the leftover solution was removed, I did test the "salt" left behind. Here is a picture:

Sure enough it tasted like table salt and did not burn my tongue. Check. When I added water it immediate dissolved into a clear solution with no further drama.

It so happens I had ~3g of NaCl. It seems kind of magic to have more white solid than what I put in, but NaCl has a molecular weight of almost 150% that of NaOH. That is because the Cl ion weighs a little more than twice as much as the hydroxide ion (35.5 vs. 17.0 g/mol). This added weight is of course offset by reducing the density of the HCl in H2O solution.

This is a super simple and surprisingly pleasing experiment. 

Thanks for reading,

Paul

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