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