How to Make Economical,
Green, High-Energy Batteries


Small Scale/DIY Battery Making
the Turquoise Battery Project



PRELIMINARY EDITION 4
by Craig Carmichael, March 11th 2012

TurquoiseEnergy.com

DISCLAIMER: This information is provided freely and is in no instance or detail guaranteed as to accuracy or veracity. Any use made of the information is at the sole risk of the user. No liability will be accepted by the author. The author warns the reader that his highest formal chemical education is a 74% grade in Chemistry 30 in grade 12, in 1972.

Note that preliminary editions are being written as research proceeds, and the text may not be consistent within itself: one statement might say "is expected to" or "should", while somewhere else, text written later may simply say "this is how it works", or perhaps mentions that "it doesn't work", or simply omits further reference to an earlier idea that didn't work.

==> Check the catalog at TurquoiseEnergy.com website for planned availability of custom battery making tools and parts such as electrode compactors, plastic battery cases, current collector screens, and more.

==> Check editions of TurquoiseEnergy.com/news/ later than the date of this document for newer information and progress.

Contents

1. Foreward and Backward

2. Electrochemistry Overview
  The water-based battery cell environment
  pH: acidity and alkalinity
  Battery Electrochemistry
  Electrode Substances
 
- Nickel
  - Vanadium
  - Perchlorate
  - Manganese
  - Zinc
 
- Current Collectors

3. Battery Construction Overview
Electrodes Overview
Battery Layout(s)
Chosen Layout
Electrode Binder "glue"
Separators and Capacitors

4. Making the Case and Fittings
Case

Electrode Current Collector Grills & Terminal Leeds

5. Making Perforated Plastic Pocket Electrode Enclosures
Perforating the plastic
Forming the square cylinder
End caps
'Glue'/solvent

6. Making the Positrode

   6.a Permanganate/Nickel Hydroxide Positrode
   6.b Monel Positrode
   6.c Vanadium Pentoxide Positrode

7. Making the Negatrode
   7.a Zinc Negatrode
   7.b Manganese Negatrode

8. The Electrode Separators

9. Electrolyte and Cell Assembly

10. Charging, "Forming" and Testing

Initial Rest Period
Initial charge
Initial cycling
Testing Specs

11. Appendices
A. Creating Unusual Substances
B. Materials and Chemicals Supply Sources
C. Equipment & Supplies
D. Survey of Some Battery Electrode Materials



1. Foreward and Backward

   In one sense, batteries are a well known technology, intellectual property of mankind. In another, they are almost a lost art. Factories churn out inferior lead-acid cells and small cells for portable electronic devices and cordless tools, but the employees are just workers. While the theory of operation and the chemical reactions aren't hard to undersatnd, there are a few important details needed for successful construction that aren't mentioned anywhere in particular, certainly not all in one place, and very few people know anything practical about battery design and construction.

   A great need has long existed for long lived, economical, high energy batteries for electric transport and off-grid power. I decided to try my hand at creating some way to make some sort of batteries at home.
   I soon felt sure that some better chemistries, probably much better, than existing types could be created, and potentially for lead-acid or throw away dry cell prices, or not so much more. This book describes known and newly invented alkaline battery chemistries, and no less importantly, a design for DIY buildable batteries of any size, that I've come up with in a project spanning - as I write - over four years.

   Battery research and commercialization have been sidelined by human propensity to "go with the flow", to limit thoughts into narrow structured channels, good or (more often) inferior, and to extend that channel to the exclusion of wider possibilities, including superior ones. Thus for example, with large, higher-energy alkaline batteries having been killed commercially, and with the single-minded recognition that lithium is the lightest atomic weight metal, most research today has been working on trying to develop better lithium batteries, despite the cost and the complex problems of making lithium work well, and despite the fact that since patents on the best developments are acquired to suppress each development as it emerges to market stage, their work will dead-end the same way Ovshinsky's excellent nickel-metal hydride electric car batteries did. We trust this state of affairs won't continue for a second 100 years, but in the meantime, DIY battery making provides the rest of us a way to take matters into our own hands.

   Making 'normal' water based batteries is a rather involved but fascinating "DIY" project touching on several distinct specialties, and it creates a product truly valuable to civilization at this time. The process of learning and making will challenge and broaden your base of knowledge and abilities.

   How was I to write this? Should it be just "do this" and "do that" and you'll have a battery, should I provide a little background, or should the reader be given all the gory details, the reasons and reasoning behind the instructions? Knowledge is power! I'm telling all that I can think of to say. But I'm organizing it into various sections so the reader can read as much or as little as desired - the basic instructions, a good theoretical overview, or complete detail.
   In other material, even the most basic information is lacking for neutral pH salt solution cells. For example, why is the positive electrode in a standard dry cell a conductive carbon rod instead of metal as in all other batteries? You'll dig long and deep and not find the simple answer: that every common metal will corrode away in the positive electrode in salty electrolyte - including nickel, which sits inert in and enables all the various KOH saturated alkaline cells. Only carbon or graphite works. (Note: nickel manganate+epoxy mix might work) Obviously battery makers know this (or once did), but it took me over two years of corroded electrodes in every test cell to figure it out for myself, because no one mentions it anywhere. (I put it on Wikipedia, but it appears to have been erased.)

   Much of the info herein has been acquired gradually, and often painfully, in my battery research over the past 4 years. A tidbit of basic info is casually mentioned in one publication or another, most of which assume the reader is well versed in the battery making arts - and few people are.
   For example, it was only after 2-1/2 years that I finally saw for the first time an actual figure for the amount of pressure used to compact a battery electrode into a "briquette" - for one type of electrode in one experiment. When I started, I wasn't even aware of the vital role of compaction, and after eventually deducing it indirectly from some material density specs, it took a another year to figure out a simple way to get enough pressure.
   Likewise, it wasn't until February 2012 and four years of mysterious self-discharge problems that I understood that the wires in the negative electrode had to have as high a hydrogen overvoltage as the electrode substance itself. Anything goes for iron, cadmium or hydride, but few common things work with a higher voltage chemical - zinc or manganese. It has to be zinc or zinc alloy wire. (or silver.)

   My original minimum battery goal was to copy proven and relatively economical NiMH EV battery chemistry, by the simplest techniques I could find or work out, and thus create a "DIY" means of making batteries. But I also started to think that coming into the field as a newcomer without formal training in the field as to "that's how it is", I might, in stumbling around, uncover overlooked information or ideas that could lead to a better battery.
   That would have the additional advantage that being developed by me, freely and openly published by me, and designated by me as the inventor to be free technology, there would be no patent restrictions on it for vested interests to kill commercialization with.
(And patents aside, it probably would have been very difficult to make a decent hydride alloy.)
   I did indeed do a good bit of stumbling around in my ignorance, getting wild ideas and then seeing the flaws, and gradually learning many broad basics and fine details in no particular sequence. And I did uncover a few key overlooked things.
   I also developed useful "DIY" battery construction tools and techniques, such as a bolt-down electrode compactor, and perforating rigid plastic sheets with a heavy sewing machine to make solid "pocket" electrodes. Finally I have been rather successful: nickel-manganese and similar batteries are in principle economical, "green", and superior to what's on the market today, including being quite economical and having about the highest feasible energy density, perhaps on a par with lithium ion types. I picked the reacting substances out of a considerable number of possibilities because they seem to be the best. The fact that they are also common and relatively economical is an excellent bonus.

   Unless otherwise specified, quantities given as a percentage, eg "1% antimony sulfide", mean percent by weight ("wt%"). Sometimes this is in addition to the otherwise complete chemicals. So if an electrode has 65% nickel hydroxide and 35% graphite powder, and "1% Sb2S3 is added", the total weight is 101%.

   I'm introducing here some new terminology - more accurately, two terms and a new spelling. Most literature uses the terms "anode" and "cathode". The meaning of these terms is reversed when the battery is charging from when it is discharging, and while there is a convention that "anode" refers to the negative electrode (while it is the positive terminal of a diode or a non-rechargeable battery), this is not universally adhered to, and there is often confusion about what is meant - I often get mixed up myself. As electrodes are ubiquitous to the subject and a specific one is so often referred to, herein I will call them "positrode" and "negatrode", which terms should be self explanatory. I also insist on spelling terminal wires as "leeds" to differentiate connections and wires from the metal "lead", the guy "in the lead", and at least a couple of other uses of the same four letter sequence, hoping not to "lead" anyone astray.



2. Electrochemistry Overview

   The physical design and construction is more important to making a battery that works than the electrochemistry. But the electrochemisty is the premiere part, the fascinating part, so it gets the first chapter.

   I've tried to explain less common, specifically electrochemical terms herein, but the reader will understand the text better if he still remembers his high school chemistry. If you don't know what an "ion" or a "sulfate" are, just look them up on Wikipedia. If anyone asks, I'll try to answer things I haven't made clear.

The Water-based Battery Cell Environment

   Aqueous batteries tend to charge water into O2 (positrode) and H2 (negatrode) gasses. In acid, hydrogen generation starts to occur at 0.0 volts or anything negative: this is the reference voltage against which all other reactions are measured. Whether a substance can be used inside a rechargeable cell depends on it charging below the voltage where gas is produced instead.
   Gas generation is more and more likely with increasing voltage above 1.23 volts, but the exact voltage varies with electrode substance and additives, temperature, and pH. Any amount over the theoretical gassing limit, at which gas isn't generated, is called the "overvoltage".
   In acid, gas generation voltages shift to inhibit oxygen generation and hydrogen generation occurs more easily. Eg, a lead-acid battery allows the lead oxide to lead sulfate reaction to work at +1.7 volts. The lead dioxide would spontaneously discharge itself at that voltage in salt or alkaline solution. However, the lead metal to sulfate reaction is also just under the limit at -.35 volts.
   On the other hand, in alkali, oxygen gas generation is encouraged and hydrogen more inhibited. The common alkaline nickel positrode (+.5 volts) is just below the "oxygen overvoltage" at room temperature, and zinc just works at -1.24 volts. The 0.0 volts in acid hydrogen voltage, in alkali is -.833 volts. The inverse of this voltage plus the +.49 volts of nickel gives us a theoretical open circuit voltage of the nickel-metal hydride alkaline battery, 1.32 volts.
   Oxygen overvoltage falls a bit with temperature, and above 40ºC simple nickel electrodes won't charge properly.
   The electrode substance is also significant, and small amount of a high overvoltage potential substance as an additive can increase the overvoltage so that the main substance works better, or works at higher temperatures. To improve zinc's performance in alkaline solution (-1.24 volts), the traditional additive was 2.5-4% mercury oxide. Later, owing to mercury's toxicity, transition metals (gallium, indium, tin and bismuth) or their oxides were tried and found to work well even in amounts under .5%. In an Indian experiment with sealed Ni-Fe alkaline cells, .5% bismuth sulfide (Bi2S3) was used to reduce the hydrogen bubbling in the iron negatrode. Heavy transition metals such as antimony are also used to improve lead-acid cell charge performance.
   In the case of manganese as a negatrode, adding 1% antimony sulfide raises the hydrogen overvoltage above manganese's charging voltage. This is the only reason it works at all. Without it, the overvoltage seems to be right on the edge: the manganese may or may not charge, but it bubbles hydrogen as it does and gradually discharges itself to hydroxide, bubbling hydrogen. Thus manganese has never been used before as a negatrode. Its higher reaction voltage, made workable by the antimony sulfide, gives a "-Mn" battery an edge in energy density over any other. (Ni-Mn is higher voltage and longer lasting than Ni-Zn, making higher energy cells of about 1.7 nominal volts. In fact, NiMn alkaline cells may last indefinitely.)
    There are lots of even higher voltage reactions that it's hard to conceive of making work with any additive, such as aluminum to aluminum hydroxide at -2.3 volts in alkali. That surely will never be enticed to charge or to hold a charge in any aqueous solution.

   The gas produces pressure inside the cell, and the pressure problem increases with battery size, so sealed batteries are small. In addition, H2 has proven almost impossible to get rid of in sealed cells. Pressure would just build up until the cell burst. So sealed alkaline batteries are made with the negatrodes larger than the positrodes. The positrodes bubble oxygen first, and the cells are also made as dry cells with empty spaces that gas can pass through. The oxygen migrates to the negatrode, discharges some of the substance (making heat), and prevents complete charging of the negatrode. This gets rid of the oxygen, and prevents the negatrode from bubbling hydrogen gas, preventing mild overcharging from bursting the cell.

   Vented cells (a) dry out and need refilling, and (b) absorb carbon dioxide from the air, which may gradually degrade substances within, turning them from active chemicals into carbonates. Various caps and valves can minimize the problems and vented cells aren't impractical, but they're second best to sealed.

   To make sealed cells bigger than dry cells, some means to keep gas pressure low has to be found.
Recent work with catalysts to recombine O2 and H2 into water has been successful, but I haven't explored it at this point. I've also read that antimony is almost unique in its ability to react with small molecules - like hydrogen - and I picked it as an electrode material additive hopefully as a recombinant catalyst as well as for raising hydrogen overvoltage, but I don't know if it works, or if I've employed it well to do so. Antimony sulfide is cheap.
   I've given up on sealed cases for now. With alkaline liquid electrolyte, sealed cells are very dangerous, since a spray of postassium hydroxide out a leak can blind. "Blindness is for life"... one cell almost got me - only takes one - and I've met a blind chemistry professor. A vented case, and using potassium salt for electrolyte, reduces the dangers.

   I hated the thought of using potassium hydroxide or acid electrolytes. They're dangerous! I was using a salt based electrolyte of neutral pH, potassium chloride. (KCl) It's a fast electrolyte (allowing high current flow), and less hazardous to handle than potassium hydroxide - it's edible. However, the cells turn highly alkaline as they charge. It's less concentrated, but still pH 14.

   In addition to chemistry, there were (and are) other novel improvements begging to be made. If one could find a chemically inert but electrically conductive or even semiconductive binder 'glue' to hold the electrode powders together, it could permit higher current flow than the usual insulating binders, and intense compacting of the electrodes would be less critical to obtaining good current capacity... If a small, economical, high energy battery could supply enough current to start a car engine, that would be a marvel!
   I'm just now experimenting with nickel manganate, a highly conductive semiconductor, for the nickel electrode, but have no results or conclusions yet.

Battery Electrochemistry

   First I'd like to point out a misleading quirk of terminology. Back in the beginning of understanding atomic particles, someone decided electrons had a "negative" charge while protons were "positive". It doubtless all seemed pretty arbitrary, perhaps even using the words "positive" and "negative". Of course, these two words have other, well known meanings. But they have been applied backwards.
   Consider that protons are stationary, within atoms, while free electrons move around between atoms... like banks and money. With a surplus of electrons, paradoxically the charge is "negative", while if there is a deficit, it becomes "positive". The more money you spend, the higher your account balance; the more you earn, the higher your debt. The negatrode deposits electrons during charging and then supplies them to a load, while the positrode is "short" of them when charged and soaks them up on discharge. This is all counterintuitive, and in some situations, a hindrance to figuring out what's going on. Now back to our regularly scheduled program...

   When a positive battery electrode is charged, it is "oxidized". When it discharges, it is "reduced". The negatrode is the opposite. These confusing names indicate electrochemical reactions that involve loss and gain of electrons, which on this planet are frequently but not always related to oxygen reactions. (Remember the obnoxious "OIL RIG" - Oxidation Involves Loss, Reduction Involves Gain [of electrons].) Pushing electrons around is what batteries are all about. (Hmm, "Reduction is gain!" -- another lovely little paradox of nomenclature!) A shorthand used for reduction and oxidation is "redox", and battery reactions are redox reactions.
   The electrochemical reactions at each electrode are called "half reactions", and the two half reactions of a battery must balance each other. If the negative terminal supplies "x gazillion" electrons to an external circuit, the positive terminal must soak up "x gazillion" electrons. And, the ions released internally by one electrode must complement those released by the other or be absorbed into it. After all, no atoms are being added to or removed from a battery in use.
   The chemicals used in a battery are chosen both for complementing ions and such that the positive side is a chemical that gives energy when reducing while the negative chemical is one that gives energy when oxidizing - at least relative to each other, within the cell's closed environment.
   Usually the negatrode material reduces to the pure metal form when charged: iron, cadmium, zinc, lead, manganese, and oxidizes to an oxide or hydroxide during discharge.
   The positrode is likely to go between two oxide forms with charge and discharge, a higher and a lower oxide or hydroxide.
   There are exceptions, and many other possibilities. In lead-acid batteries, the negatrode metallic lead oxidizes to lead sulfate, and the positrode lead dioxide reduces to lead sulfate, the sulfate ions being stored as excess acid (or sodium bisulfate) in the electrolyte when the battery is charged, and absorbed as it's discharged. More examples appear below.

   Usually these positrode oxide forms aren't very good electrical conductors. Some oxides, like titanium and zirconium, are virtually insulators, so they can't convert easily between forms by electrical action as battery elements. Often additives are used to improve the conductivity of the oxides. Zinc and cobalt oxides have been used to make nickel hydroxide electrodes conductive enough to use, as have nickel powders and flakes, and powdered graphite.

   The number of amp hours depends on how many electrons the substance will release or absorb during oxidation or reduction, and the energy of each reaction is indicated by its voltage. A substance which naturally wants to oxidize (in the battery environment) will have a more negative reaction voltage than one that wants to reduce. The energy in watts-hours is the amp-hours (the number of electrons) times the voltage (the pressure behind each electron). The voltage of both electrodes is subtracted for the total battery voltage, eg +.5 - -.93 = 1.43 volts for a nickel-iron alkaline battery. The amp-hours or number of electrons isn’t additive: it should match. The current flow stops and the cell is discharged when either electrode has been depleted to its un-energetic state and will pump no more electrons and ions.
   For a given number of electrons moved per reaction, the lighter the atomic weights of the reacting elements, the more amp-hours per kilogram will be available, because there are more molecules to react in that kilogram. Oxygen and hydrogen are quite light, so the metal is usually the dominating factor. If a heavier element is chosen, it must move more electrons per reaction, or have a higher reaction voltage, to provide equal energy density. If the advantages are less than the added weight, as with lead, cadmium or mercury, batteries with these heavier elements have lower energy densities. The heavier elements are also more costly. Thus my own searches were mainly for lighter atom metals.
   Lightness of metallic substance is pursued to the ultimate in lithium battery types. But lithium has to be used in thin film electrodes, often with non-aqueous electrolyte, and the substrates to hold all the thin films add their own bulk and weight.

   Usually it is required that reaction products of both charge and discharge be solid, that is, that they don't dissolve (...or melt or turn into a gas). This greatly limits the choices. Most chlorides are soluble, so the electrodes of a battery using hydrochloric acid would dissolve and thus would be hard to recharge. The old 'standard' non-rechargeable dry cell uses ammonium chloride electrolyte, and the zinc electrode dissolves to zinc chloride in use. Most lighter elements dissolve in sulfuric acid, but lead, lead sulfate and lead dioxide are all non-soluble - hence the lead-acid battery.
   Just to prove the point, I looked for an acid that lighter metals wouldn't dissolve in. I found oxalic acid seemed to qualify, and I made a nickel-zinc test battery in oxalic acid: nickel oxide, nickel oxalate, zinc and zinc oxalate are all insoluble. Similar in concept to lead-acid, it worked and could be charged. (The voltage was lower than the tables indicated, about 1.4 volts. Acetic acid/acetates should also work.)

   Zinc has been known as a frustrating battery negatrode element. It's energy is the highest available for alkaline cells and its electrical conductivity is good, and the charge and discharge products are both solids. However, in use there is a temporary dissolved state, the zincate ion, in which form the zinc can and does gradually migrate. This causes the negatrode to gradually lose capacity, and the zinc grows dendrites, "tentacles" of zinc crystal, which usually short out dry cell batteries, often after only 10-50 charge-discharge cycles. Cadmium, underneath zinc on the periodic table, has the same problem, and Ni-Cd dry cells rarely last anywhere close to their supposed cycle life as cadmium crystals poke through the separator sheet and short the cell. NiZn and NiCd pocket cell batteries fare much better. But it would seem that NiZn dry cells in recent years have improved, as a company making AA cells (available on Amazon.com) claims 500 to 1000 charge-discharge cycles.
   It's possible that in salt electrolyte, zinc doesn't form zincate ion. Thus switching to salt might solve the problem, allowing use of this high energy density substance in long-life batteries. Or, the zirconium silicate ion blocker I paint on the electrode separator sheet may solve the problem or at least provide "500 to 1000 cycles".

   There are several choices with somewhat less energy than zinc - eg, iron, cadmium and hydride - but none with "just a little less".
   Next up, manganese at about .3 volts higher than zinc, sits on the threshold between usable and not for a negatrode. It doesn't work by itself at room temperature, bubbling hydrogen and charging at the same time, then spontaneously discharging itself too quickly to be practical. But with the right additive(s) to raise the hydrogen overvoltage, it might be made usable. It needs further research.

   The electrolyte doesn’t conduct electrons between the electrodes, it only conducts charged dissolved ions. It's the one place where protons are on the move. To have the oxidations and reductions take place, both ions and electrons must flow, as will be seen in the redox (reduction-oxidation) reactions coming up.
    A circuit connected to the battery lets the electrons flow between the electrodes - externally. This is of course what the battery is for. When an external circuit is connected, the electron flow, the ion flow and the discharge reactions proceed spontaneously and simultaneously, releasing the chemically stored energy as electricity. The ions flow mainly by diffusion through the electrolyte, spreading because like charges repel, and by attraction to the opposite electrode as they reach it. The current capacity of the battery depends partly on how fast the ions diffuse through the electrolyte. Potassium chloride salt is supposed to be very fast.
    The discharging reactions release chemically stored energy electrically. The recharging reactions require electrical energy from the external circuit - the battery charger. Charging restores the 'spent' lower energy substances to their higher energy oxidation states and valences.

There are many solutions and some solids that can pass ions, but the best - fastest - solvent is a polar liquid such as water, with an acid, salt or alkali electrolyte dissolved in it. There is, however, one serious limitation to using water as an electrolyte, as mentioned previously:

"The use of aqueous battery electrolytes theoretically limits the choice of electrode reactants to those with decomposition voltages less than that of water, 1.23 V at 25 ºC, although because of the high "overvoltage" potential normally associated with the decomposition of water, the practical limit is some 2.0 V. The liquid state offers very good contacts with the electrodes and high ionic conductivities." Lead-acid batteries are theoretically 2.05 open circuit volts, and many earlier cells were about 2 volts.

   The voltage delivered to a load circuit is somewhat lower than the open circuit voltage, depending on the internal resistances of the battery relative to the amount of current flowing. Hence batteries are given a "nominal" voltage rating which might be expected in typical heavier use, such as "1.2 volts" for Ni-Fe, Ni-Cd, and Ni-MH, which read more typically 1.33 to 1.43 volts with no load. Heavy loads may drop the output even more, eg to 1.0 volts. If such loads are expected, it's usually best to add more batteries in parallel to reduce the load on each one, or to use bigger cells, which is effectively about the same thing.

   If the positrode has lesser amp-hours capacity on discharge than the negative it is depleted first. The negatrode still could have supplied more current and the battery is said to be positive limited. Vice-versa if it's the negatrode that runs out first. It may also be positive or negative limited on charging, and not necessarily in the same direction. It's also possible for the electrodes to be entirely off balance - one discharged and the other charged. It could be hard to either charge or discharge this cell.
   There are often good reasons for preferring one reactant to deplete first. For example, if there's no recombination catalyst in a sealed dry cell, oxygen gas is much better to generate than hydrogen if the cell is overcharged. In a dry cell, it travels over form the positrode to the negatrode and there discharges an atom of metal to hydroxide, making a bit of heat. Thus the cell stops charging - it just gets warm. Hydrogen doesn't readily discharge at the positrode and the gas would accumulate until the cell bursts, so it's best to have the positrode charge first and not get any hydrogen. With the catalyst, starting to generate both gasses at about the same time when the charge is complete should be advantageous, since they can then start recombining to make water before the pressure of either gas builds up much.

Electrode Substances

   Besides lead in lead-acid cells and lithium, there are two common positrode substances: nickel and manganese. My newfound vanadium has higher voltage. It appears to work.
   Until now zinc has been the most energetic negatrode element, -1.24 volts and 820 amp-hours/gram of Zn, or 1016 watt-hours/kilogram. This is much better than iron or cadmium and on a par with typical hydrides in alkali. However, it has a temporary soluble state during discharge, and grows dendrites ("tentacles") that usually short out the cell in as few as ten recharges. (This seems to have been pretty much solved in some recent dry cells, but not for flooded cells.)
   After much trying, I've now got manganese to work. Normally it's tantalizingly borderline, charging at about the same voltage as the hydrogen and spontaneously discharging itself somewhat too fast to be practical. I found the hydrogen overvoltage can be tweaked up sufficiently by adding 1% stibnite (antimony sulfide). This makes it work. It appears to be an ideal negatrode. It's even higher energy density than zinc - virtually an amp-hour per gram of Mn at around -1.18 volts: 1150 watt-hours per kilogram on Mn.

   For the positive side, manganese dioxide has been strictly the substance of one-use dry cells, so-called "carbon-zinc" but actually manganese-zinc, the carbon (as graphite or "carbon black") being in fact simply a conductivity improving additive. But the zinc and the electrolyte are the problem with recharging the old dry cell, not the MnO2. In salty solution the energy is about +.5 volts, but in alkaline solution it's only +.15 volts, so makers of rechargeable alkaline batteries prefer nickel oxyhydroxide, with +.5 volts. However, it has high amp-hours per kilogram, and that allows it to complement more high energy negative electrode, providing higher energy cells notwithstanding somewhat lower voltages.

   Vanadium pentoxide is around +1.5 volts. I was surprised to see that this reaction actually works instead of bubbling oxygen.

   Nickel (+0.95V) works great and makes 2 volt cells, but rechargeable cells using manganese positrodes provide the highest energy density, and manganese is cheap - you can even scrounge it out of old dry cells for free.


Nickel

Note: Despite its lower reaction voltage, manganese dioxide is now preferred to nickel in any form owing to it having higher amp-hours. The reader may wish to skip reading about all the other positrode materials.

   Nickel hydroxide [Ni(OH)2] is the common positrode material used in most rechargeable alkaline batteries with various negative electrode materials: Ni-Cd, Ni-Fe, Ni-MH and Ni-Zn. Dry and pure, it's a very fine, fluffy, turquoise green powder. The nickel will happily stay in the hydroxide form in the battery environment. It thus has no usable energy. To convert it to a more energetic chemical, energy must be put into it.
    To charge it, the nickel hydroxide is further oxidized to nickel oxyhydroxide by grabbing one electron from it. It doesn’t willingly give up the electron: the charger has to supply the energy to cause it to happen, exceeding +.52 volts. This disengages a hydrogen ion (H+), which jumps over to an immediately adjacent hydroxide ion (OH-) in the electrolyte to form water. Thus the nickel is 'oxidized' from valence +2 to +3, losing an electron and a hydrogen rather than by adding oxygen. The basic half reaction is shown as:

    (beta) Ni(OH)2(s) + OH-(aq)  <==>  (beta) NiO(OH)(s) + H2O(l) + e-    [+0.49 V in alkali; +1.05 V in salt]

   (discharged <==> charged)

   Note that the "Ni" compounds are solids on both sides of the reaction -- not dissolved, liquid or gas. It is usually a prime requirement that the electrode doesn't dissolve. Normally if it does, the battery won't recharge. The valence of the nickel goes from II to III as it's charged, indicating that one electron is removed per molecule, as shown. (We'll touch on the crystalline forms "beta", "alpha" and "gamma" further on.)
   But in fact, not all of the oxyhydroxide [III] gets converted back into hydroxide [II]. When there's some of each, the nickel valence is expressed as a fraction. (which we will not attempt to describe with traditional Roman numerals) When it gets below 2.25 or so, the resistance rises and the user considers the battery to be "pretty much dead". So really, only 3/4 of an electron is moved per nickel atom, reducing the capacity below the theoretical value.

   The two voltages shown (+.49, +.52) are as listed by different sources as being the "open circuit" voltage for this reaction. Voltages seem to vary slightly with different electrode additives, and perhaps with temperature.
   A major advantage of salty electrolyte is that the nickel reaction voltage is double, about +1.05 volts, giving it double the watt-hours per kilogram of the alkaline cell. This alone was a good reason to attempt to create working salt solution batteries.

    The nickel oxyhydroxide is an "energized" substance: it would rather be just plain hydroxide and given a chance will revert and give off energy in doing so. But it needs an electron and a hydrogen ion to do so. The amount of energy per electron is seen in the voltage. It can get the hydrogen "H+" ions from the water, leaving OH- in the water. This is balanced with the negative electrode grabbing the "OH-" ions, but it will only perform this reaction when an external electrical load is connected to give it an electron.


Nickel redox chart.
Paradoxically not shown is the chief reaction of battery interest, between valences 2 and 3 in alkali (base) Ni(OH)2 to NiOOH, which has the same reaction voltage as the 2 to 4: +0.49 volts... or +0.52 depending where you read. In modern nickel formulations, some of the nickel gets oxidized to NiO2, valence 4, as shown on the chart, raising number of electrons transferred and hence the amp-hours capacity.


    Notice that nickel hydroxide can be reduced as well as oxidized, to become elemental nickel. Again, it would rather be hydroxide in the wet battery environment, and it takes energy to reduce it to elemental nickel metal. Thus, this reaction would make a "-Ni" negatrode. The reduction reaction is:

    Ni(OH)2(s) + 2 e-  ==> Ni(s) + 2 OH-(aq)    [-0.72 V]

   Again the nickel keeps a solid form, so a working Ni-Ni battery could be created. The valence of the nickel goes from II to 0, adding two electrons to each nickel atom. This charging reaction gives off negatively charged hydroxide ions that were bonded to the Ni(OH)2, the same as with iron, cadmium, zinc and manganese, each at its own voltage. (Metal alloy hydride absorbs a hydrogen ion, H+, from the water, also leaving an OH- ion.) Moving two electrons instead of one, at -0.72 volts (instead of +0.52),  2*(.72/.52) = 2.77 times the theoretical energy storage. Nickel hydroxide in alkali, though the most common positrode material, makes a much more energetic alkaline negatrode than it does a positive one! (You might need a hydrogen overvoltage raising additive to keep it from bubbling hydrogen - nickel evidently has a very low intrinsic hydrogen overvoltage, and hence nickel electrodes are often employed to generate hydrogen.)
   Notwithstanding this, the voltage and energy of the reaction are lower than the usual substances... and it's only -1/4 volt in salty solution, definitely eliminating it as a candidate.

    The theoretical energy limit of Ni(OH)2 as a "+" terminal of 289 amp hours per kilogram is presumably doubled as a "-" side to 578 AH/Kg (of Ni(OH)2), and at -.72 volts that's 417 watt-hours/Kg.

   So why is nickel [oxy]hydroxide so popular as a positrode chemistry? Well, it boils down to ... try and find something better, that doesn't cost a fortune! Silver oxide works well (eg, AgO <-> Ag2O, +.6v), but the atoms being heavier, it would have lower energy density - lower amp hours - by weight, despite the somewhat higher reaction voltages.

   Manganese dioxide, while cheap, is only +.15 volts in alkaline solution. That means more cells to attain a given voltage. In salt, it's .5 volts, and it might be a more economical solution for stationary batteries, eg for off-grid home power storage. It is easy however, and considered deleterious, to charge it to a higher oxide form. (The zircon ion and-or chelation of the Mn ions shield might alleviate this concern.)
   For transport where light weight counts, nickel's +1 volt in salt is better despite the cost. Anyway, the only required nickel in the salty battery is the actual active chemical, whereas in alkaline batteries nickel or nickel plating is used for all internal metallic structures. (That could be changed with grafpoxy.)

   But the traditional basic reaction doesn't reveal nickel's full potential. Nowadays, manganese is added to the positrode as a major additive, perhaps 35 to 40% by weight ("wt%") of Mn to Ni. What this supposedly does is raise the oxygen overvoltage, which evidently allows the nickel to charge to "alpha" nickel oxyhydroxide, wherein some portion of the nickel actually charges to NiO2, valence IV, moving two electrons instead of one. Another thought is that permanganate is a "powerful oxidizer", and it may be this that allows or causes the nickel to oxidize to a higher valence. On the other hand, the two ideas may just possibly amount to the same thing expressed differently.
   Maximum attainable overall valence appears to be about 3.8. The actual nickel valence thus might change from about, say, 2.25 to 3.75 from discharged to charged, thus moving 1.5 electrons per nickel atom, twice as much as with the old pure Ni(OH)2 simple formulation. This doesn't double energy density by weight because of the added mass of the manganese, but it does improve it, and the nickel - the costly and main ingredient - does twice as much work.
   Multiplying the theoretical value 289 AH/Kg * 1.5 = 433 AH/Kg. Naturally however, the theoretical maximum isn't going to be attained. (Experimentally about 350 AH/Kg has been attained, the forms being alpha hydroxide and gamma oxyhydroxide, which both occupy about the same volume of space. Although it's a higher volume form than the beta forms, the constancy is very desirable for long cycle life.)
   NiMH "AA" battery capacities have increased from 1.5 to 2.5 amp-hours in recent years. (This is after sintered nickel cadmium "AA" batteries of just 0.5 AH in the 1970s.) Since the NiMH AA cells with this high energy weigh 30 grams, and the nickel hydroxide (educated guess) probably weighs up to about half of it, an attainable figure in an actual battery of around 166-200 milliamp-hours/gram (= amp-hours/Kg) is suggested.
   Squeezing the most out of the nickel is important both for economy and because the nickel is the bulkier, heavier electrode, and anything that improves it can notably improve the entire energy density of the battery. For homebrew salty batteries, I'm expecting actual attainment of around 100 AH/Kg will be doing well.

   The negatrode substances being much higher energy, the energy density of the whole cell will be mostly limited by the nickel and the voltage obtained. A 1.8 nominal volts nickel-zinc/salt cell then will be somewhat under 180 WH/Kg, eg maybe 120-170. For 2.1V with a manganese negatrode, if that can be made to work, 130-190 WH/Kg might be attained. These figures seem dissappointing after reading the theoretical maximums, but they're still better than commercial NiMH dry cells and as good as or better than lithium ion types. And it's not impossible that with good design, chemicals, technique, workmanship and high compaction, even higher energy might be attained.

   Other metal oxides or hydroxides besides manganese that have been tried and appear to work (and may bear further experimentation) include: aluminum, cobalt, yttrium, ytterbium, erbium, and gadolinium. Other rare earths hydroxides such as samarium, neodymium and even lanthanum might be better, or at least fine, in salt solution. I'm not sure why manganese is supposed to be "especially preferred" (or even why it should work well), or indeed what the selection criteria are, but I've used Mn in my positrodes as well. I believe the Mn charges to higher oxides (potassium permanganate) that won't discharge until the nickel has finished discharging, and then at a lower voltage. (Manganese has so many reactions at various voltages that it's confusing to try and figure out what will actually happen in many situations, and I as far as I can see commercial battery designers often don't know exactly what they're doing either. Certainly in Alkaline Storage Batteries (Falk and Salkind 1969), there was a lot of speculation about some of the main chemical reactions. And battery substance reactions in salty electrolyte are relatively unexplored compared to alkaline.)

   It's not clear to me at the moment whether the only effect of the manganese compound is supposed to be to raise oxygen overvoltage in the postirode. If it is, the samarium or whatever, probably in considerably lesser quantity percentage-wise, should replace it entirely, providing highest energy density. (For a while I thought the KMnO4 reacted at virtually the same voltage as the nickel and would be an active chemical along with the NiOOH, but it appears it's somewhat lower and thus wouldn't start to discharge unless the nickel had completely discharged.)

   The element nickel is the biggest cost in nickel-alkaline batteries - it's not only the postrode substance, but composes over 3/4 of the hydride alloy, and the plating or substance of all the metal conductors within the cell. In the salty cell, it's just the positrode chemical, so the cell should be more economical.

   Neither nickel hydroxide, oxyhydroxide nor potassium permanganate is a very good electrical conductor. The battery's current capacity would be extremely limited if these were the only ingredients. Powdered graphite has been added for better conductivity, as in the standard and alkaline single use dry cells.
   Edison put in 80 layers per inch of alternating nickel hydroxide and ultra-thin nickel metal flakes, crammed solidly into perforated metal tubes about the size of a pencil. The nickel flakes were made by electroplating alternate layers of copper and nickel onto something, then dissolving away the copper. That costly arrangement was the best he could come up with that worked well. He tried graphite flakes and found the performance was unpredictable - I think Edison didn't expect powder could be a good conductor across an electrode, but above a critical proportion it is.
   The sintered electrode is another good form for conductivity in alkali, the sintered nickel sponge connecting well across the whole electrode for very high current capacity. NiCd cells get some of their high current ratings from this.
   But I discovered that for any salty cell battery, all metals oxidize rapidly in the salty positrode. Sintered metal electrodes are out. And graphite powder is cheap at any art supply store.
   But up to 5% cobalt hydroxide has been added to alkaline cells with good effect to improve conductivity without graphite or nickel flakes, and I've been trying starting with monel alloy, which puts (25-33%) copper hydroxide in solid solution with the (67%) nickel hydroxide. (The monel I'm using also contains 2% Mn and 3% Fe, so the copper is 28%. Obviously Mn doesn't hurt, and the iron either, I trust.)

   On a practical note, it's worth mentioning that a nickel electrode can be discharged chemically to Ni(OH)2 by immersing it in a small pool of hydrogen peroxide - the 3% drug store stuff is fine. It makes zillions of very tiny bubbles as excess oxygen comes out. When it's done, rinse out the H2O2 with clean water.
   In addition, the nickel can be charged to NiOOH using bleach, sodium hypochlorite. I haven't done this myself. 3% grocery store bleach should work fine. Again rinse out the bleach when done.
   These procedures give you a way to equalize the charge if you've ended up with one charged electrode and one discharged for a sealed battery. For an unsealed one, charging and letting gas bubble off one electrode works.

Nickel Manganate

   In late February 2012 I found a better form of nickel for electrodes than nickel hydroxide: nickel manganate [NiMn2O4], a synthesis of nickel and manganese. This little known substance (but not unknown - it's used to make thermistors) is of repute for its "spinel" crystalline structure, which gives it a much lower electrical resistance than most oxides. At first I thought it might make a good conductivity improving additive. It was far more conductive than Ni(OH)2, but nowhere near as good as graphite. Then I thought of using it in place of nickel hydroxide as the main electrode substance.

   At first I thought it might charge to nickel permanganate [Ni(MnO4)2]. Both substances have one nickel and two manganese ions. However, one has 4 oxygen ions while the other has 8. Charging nickel manganate to nickel permanganate would release 8 electrons and use up 8 OH- ions from the charging negatrode (Four become the other four "O--" ions in the permanganate, the other four become H2O). The voltage for that reaction would be around +.65 volts in alkaline solution. It would be fantastic energy density... maybe too good to be true.

   Then I realized the nickel would be more likely to change since its reaction voltages are lower, probably similarly to the reactions discussed for nickel hydroxide. A nickel valence 3 compound might be formed, for example, Ni(OH)Mn2O4, or even valence 4, eg Ni(O)Mn2O4. It'll probably get a bit more mileage out of the nickel - because of the high conductivity, it would probably discharge down to nickel valence 2.0, whereas nickel hydroxide pretty much stops supplying current when the average valence is down to 2.25 owing to increasingly poor conductivity.

   I made three small (1/4" square) cylinder electrodes from it. They worked well, and seemed to be much more conductive than simple nickel hydroxide.

The approximate formula was:

11g NiMn2O4
5 g graphite powder
.4 g Sunlight dishsoap

   I meant to put in a little neodymium oxide (maybe 1/2 a gram?), to raise the oxygen overvoltage either for better higher temperature performance or in case the higher voltage permanganate reaction applied, but I forgot.

   I couldn't find nickel manganate to buy. I tried making it chemically, but it was messy and smelly. Then I mixed appropriate amounts of dry NiO and MnO2 powders (both from the pottery supply) in a stainless steel pot and simply heated them red hot with a propane torch (outdoors, with a respirator). This gave a lower resistance product and was fast and pretty simple to do. I only got 15 grams, so I guess the torch blew 7 or 8 grams of powder out of the pot. This is what made the successful electrodes above. Later I made another batch and 'only' lost 25% of the mass.

Vanadium

   I 2011 I made a battery with a vanadium electrode. It was supposed to be the negatrode, but it didn't seem to work - unexpectedly, the vanadium seemed to become soluble and to migrate. (This was also the first cell I'd made with transparent plexiglass sides, and I could see the vanadium pentoxide yellow color appearing on the other electrode.) I reversed the charges, and found that the cell charged to about 2.2 volts. The vanadium positrode side would have made up around 3/4 of that, and it seems surprising that it didn't just bubble oxygen and spontaneously discharge itself to a lower oxide. It seemed to charge and discharge well, but at the time I hadn't made the grafpoxy yet and it deteriorated like my other cells of that period, as the graphite backing sheet swelled and lost conductivity and good contact with the electrode and the carbon terminal post. Judging by the voltage and the chart, the likely half-reaction was:

V2O5 + H2O + 2e- <==> V2O4 + 2 OH- [+1.6? V]

Unlike the case for either alkali or acid solution, and unlike it's unexpected behavior as a negatrode, the oxides appeared to me to remain in solid form, not dissolve, in the salt electrolyte. This appears to make it a good positrode, moving one electron per vanadium atom, hopefully with good stability from the double vanadium molecular center.
   Taking the average of the acid and alkali voltages as being the approximate salt voltage, the voltage obtained in the cell indicates the single valence change to V2O4 seems to apply. (average of (1.0V [acid] + 2.19V [base]) / 2 = 1.6V [salt])


The table shows that vanadium's higher oxides are "amphoteric", that is, they'll dissolve in either acid or alkali.
However, they don't seem to dissolve or break down in neutral pH salty solution even with a valence of +5.

   So vanadium seems to have the potential to be a very good positrode in salt water electrolyte. Theoretical amp-hours works out to be almost identical to the theoretical 289 amp-hours/Kg of beta nickel oxyhydroxide. The potential double valence change that is achieved by some of the nickel to alpha oxyhydroxide molecules wouldn't seem to be possible with vanadium, but the voltage is 55% higher, raising the energy density considerably.

   According to the electrochemical table, we might suppose that vanadium might also make a good negatrode in alkali at the same potential as cadmium (-.82) and hydride (-.833), providing it wasn't overdischarged, which might form the higher oxides, (eg V2O3) which might cause problems.
   I don't see why it isn't in use - the energy density should be good.

   However, the "Pourbaix" state diagram from Wikipedia probably indicates the problem I had using it as a negatrode in salt: instead of forming either VO or V(OH)2 at pH 7, vanadium instead forms a dissolved ion, VOH+, for the 'discharged' negatrode.

   I don't see the common V2O5 form (or VO for that matter) anywhere in the pourbaix diagram, and the voltages don't match the table above, nor do they appear to jibe with my experimental results.

   Vanadium probably deserves more research in salt electrolyte for use as a positrode (and maybe in alkali as a negatrode), but I have no present plans for doing it myself.
After I found the vanadium Pourbaix diagram, I figured there's probably soluble ions somewhere during charge or discharge, which might make for limited cycle life. If I'm going to stick with 'tried and true', that's nickel, and if I'm going to experiment, I'll try for a bigger prize: perchlorate or permanganate.

Perchlorate

   Chlorine ion, Cl-, oxidizes to perchlorate, ClO4-, moving 8 electrons with its very own electrochemical reactions regardless of the metal+ ion it's attached to. I once tried to make a positrode of lanthanum perchlorate, La(ClO4)3, which would reduce on discharge to lanthanum chloride, LaCl3. The lanthanum was (my intent, anyway) chelated into the substance of the electrode so that, even being in dissolved form, the heavy La+++ ions wouldn't be mobile. (There is precident for this last, the article saying chelated dissolved lanthanum behaved about the same as undissolved, tho I can't remember where I read of it.) In addition, perchlorate is often much less soluble than chloride, as with only slightly soluble potassium perchlorate versus potassium chloride salt.
   As I've said, if heavier elements were used, they'd have to move more electrons to attain the same energy density. Lanthanum perchlorate, potentially with 12 O-- ions forming 24 OH- ions on contact with water, is a super example: 24 electrons per reaction where nickel moves one or two. That much more than makes up for the atomic weight of La(ClO4)3 being almost five times that of Ni(OH)2, and suggests the theoretical possibility of a much higher energy density electrode than nickel hydroxide.
   There may be reasons I'm unaware of that this can't work. I am after all only an amateur chemist. However in the absence of any known reasons it deserves more research, and I hope to experiment more with it. (Next time, I think I'll try converting lanthanum hydroxide straight to perchlorate with perchloric acid instead of to chloride with hydrochloric acid. (That's called a "super acid" - yow! I must read the MSDS again before I start.))

Manganese

   Manganese dioxide is a dark gray, blackish powder, fairly dense. It can be scrounged from [non-alkaline] dry cells, or purchased at pottery supply stores. The dry cell is probably the better source - it's known to be pure enough for batteries and it's "pre-mixed" with conductive graphite powder. In the open, dioxide is the usual state of manganese, but in the typical cell it's the charged state. An even better form for use in positrodes is as potassium permanganate.

   Manganese can be recharged, and some "renewable" alkaline cells make use of this. Sometimes the discharge product is given as MnOOH and sometimes as Mn2O3. It matters little as both are valence three after moving one electron, the difference only affecting the amount of water released or absorbed during charge and discharge.

The literature says the discharge reaction in alkaline solution is:

    MnO2(s) + H2O(l) + e- <==> Mn2O3(s) + OH-(aq)    [+0.15 V]

In salt solution, however, the voltage is much higher, and all literature I've managed to find shows this reaction:

    MnO2(s) + H2O(l) + e- <==> MnOOH(s) + OH-(aq)    [~+0.5 V]
 

Manganese Redox chart.

   Another manganese reaction of great interest is on the right end of the chart, going between valence 0 and +2. If one simply uses manganese powder in water, this reaction is just high enough in voltage that it gradually but spontaneously discharges into Mn(OH)2. This has always precluded the use of manganese as a negatrode.
   However, additives, in particular heavier transition metals or their compounds, can raise the voltage at which hydrogen starts to generate. They are used to help zinc electrodes charge better and work at higher temperatures. Traditionally about 2.5-4% mercury oxide was used. Now smaller amounts of less toxic transition metals are substituted: eg, gallium, indium, tin, or bismuth.
   I tried antimony oxide with uncertain results. Antimony sulfide, stibnite, seems to work well. The usual ore of antimony is stibnite.) I believe the Sb2S3 converts to keresemite (Sb2S2O) or possibly to Sb2S in the cell, and it works better. Whatever happens, adding 1% antimony sulfide raises the hydrogen overvoltage enough to allow manganese to charge and hold its charge.

   I finally got an alkaline cell to charge Mn to metallic state and hold its charge in February, 2012. This is probably a first.
   As it charges in alkalinity pH 14, a bit of the Mn (starting as MnO2) becomes a soluble ion, probably Mn(OH)3-, or else MnO4--. (per the Pourbaix diagram below) When this soluble ion touches the positive electrode, it's charged to KMnO4. This is indicated by the water turning purple. The KOH electrolyte solution is normally pH 14, but this appears to be reduced to about pH 13 by the KMnO4. (How does that work? Like I said, I'm not a chemist, and I don't see the answer on Wikipedia. It may even just be bad coloring of the pH test paper.) At pH 13, the soluble ions cease to form (leaving only the desired insoluble Mn(OH)2 <-> Mn reaction), so it doesn't continue to a still lower pH.
   The open circuit voltage of the cell is 2.05 volts or so. This makes it the most energetic alkaline electrode ever, and the reduced pH works better than pH 14 for both the Mn and the NiOOH: it may well make the cells last virtually indefinitely.



Mn-Mn Cell: Highest energy density yet attained!

   For a long time I thought nickel hydroxide was a better positrode because it's twice the voltage of manganese dioxide, ~+1 in salt solution versus +.5. But with all the additives to make nickel work well, it only has around 1/2 the amp-hours per kilogram of MnO2. The total energy density of the manganese positrode is thus similar, perhaps a little higher. My thought was to go for the higher voltage of nickel anyway. Then I thought of the whole cell: double the amp-hours makes for double the amount of negatrode substance for the same amount of positrode. This drawing illustrates the effect, which is greater than it seems because the negatrode is less than 1/2 the weight and volume per amp-hour (still less per watt-hour), thus making cell B only a little larger or heavier than cell A. Cell C almost doubles again the capacity of cell B for only around 1/3 additional weight.



    Although cell A has the highest voltage, and although the positive electrode has about the same watt-hours as cell B, cell B has twice the amp-hours, providing 1.5 times the energy density (theoretically 393 WH/Kg) with only slightly more weight.
   Furthermore, manganese dioxide can discharge to two lower oxide states after discharging to MnOOH (or Mn2O3), for 1.5x or 2x the amp-hours if the equipment being powered can tolerate a lower cell voltage.
   In this case, one or two more high energy negative electrodes can be added to the battery to utilize these additional amp-hours, depending on acceptable voltage. (Cell C) The drooping voltages as the cell is 1/2 discharged and then 3/4 discharged will provide good warning that recharging is required.
   So for a while, I thought Mn-Mn would be the outstanding choice. And the main materials for Mn-Mn cells are essentially the same - and hence the same price - as those for throw-away dry cells.

   But when I made a cell, it charged right up to (potassium) manganate or permanganate, which are slightly soluble. I suspect it would have relatively short cycle life. So I went back to my recently discovered nickel manganate, which had about the same voltage, and which I expect probably isn't soluble. I suspect it probably also has similar reactions, so it probably has both the amp hours and the voltage. So so far, it's my idea of the best choice - at least this week.

Zinc

Note: Zinc is superseded by manganese with 1% stibnite added to raise its hydrogen overvoltage. Manganese is the better choice in every way.

   Zinc's reactions make it suitable only for a negatrode, but quite a high energy one. The dissolved ion form found in discharge and shown in the diagram is clarified in the Pourbaix diagram beneath it.
   The conductivity of zinc oxide or hydroxide is better than most, and cells with zinc are usually high-rate for both discharge and charge.
   Addition of a transition metal or its oxide is used to raise the hydrogen gas generation voltage (the "overvoltage") to improve charging characteristics. 2.5% to 4% mercury oxide is 'traditional' in alkaline cells. 1% antimony sulfide is better and environmentally benign.



   It was long debated whether the zinc forms Zn(OH)2 as shown or ZnO as it discharges, but as usual the difference is merely the water content of the battery charged versus discharged, since Zn(OH)2 = ZnO + H2O. (IIRC the general consensus is that it's ZnO.)


    The troublesome zincate ion that limits the life of NiZn alkaline cells is best seen in the Zinc Pourbaix diagram. Here it is revealed that this ion probably won't form below about pH 13.5, and it's the pH 14 electrolyte that's the problem: it would be fine at about pH 8 to 13.

   Evidently, adding some manganese oxide to the zinc to lower the pH to 13, as the manganese negatrode does for itself, should stop zincate from forming and allow long life zinc negatrodes.

   The question then is, is there any point to making zinc negatrodes when manganese ones have more energy and are just as cheap? One possible reason is to get close to a specific battery voltage. For example, if NiMn cells are 1.7 volts, 6 volts is hard to attain:
 1.7 * 3 = 5.1
 1.7 * 4 = 6.8
whereas four NiZn is closer:
 1.6 * 4 = 6.4
or to get very close, use 3 NiZn and one NiMH:
 1.6 * 3 + 1.2 = 6.0

   Thus it would seem that NiZn could have uses in specific situations.

For general application however, including 12 volts, the NiMn would seem to be the winner, needing only 7 cells for 11.9 volts, while zinc is way off at 11.2 or 12.8 with 7 or 8 cells. NiMH takes 10 cells.

"Active" high surface area zinc oxide (ZnOxide.org)
   An issue with zinc in salt solution is that zinc powder and zinc oxide powder both absorb CO2 out of the air and form zinc carbonate on the surface, which is passive in a battery and (I think) an insulator. The carbonate however can be removed by immersing the powder or the electrode in a hydroxide: KOH, NaOH or Ca(OH)+ (lime). The lime is the best and safest one. A bit of the Ca(OH)2 will become carbonate (CaCO3, limestone). This should help strengthen the brittle zinc electrode.
   Not only does the carbonate become zinc oxide, evidently it becomes the finest, high surface area "active" zinc oxide, ideal for a battery electrode.

   In traditional manufacture of alkaline batteries with zinc electrodes, the finished electrodes are placed in KOH for a day, and the "carbonated" electrolyte is replaced before charging. But the soluble zincate ion causes zinc electrodes to degrade rapidly enough that NiZn hasn't been a very popular choice, lasting as few as 10 to 50 charges, followed by a shorted cell being the norm in dry cells.
   However, according to Wikipedia, NiZn alkaline cells with "stabilized" negatrodes have been much improved since Y2K and are now commercially viable, attaining 400-1000 charge-discharge cycles at 100 WH/Kg, probably at a substantially lower cost than NiMH or lithium. When the patents run out, they might become available in vehicle battery sizes instead of just small dry cells.

   Cadmium also forms a soluble ion and NiCd dry cells often don't fare much better than zinc, cadmium being right under zinc in the same column of the periodic table. They do have zinc's high conductivity. NiCd pocket cells, however, like other pocket cell batteries, have a good reputation for longevity. Since the atomic weight of cadmium is 112.5 versus 65.5 for zinc, and since its voltage in alkaline solution is -.82 instead of -1.25, the energy density of cadmium is only 38% that of zinc. Hydride is much higher even with the same voltage. (-.83) Nickel-iron is probably better too, even tho utilization of the iron isn't high, as it tends to agglomerate into larger particles with less surface area with cycling. (Additives such as cadmium help, and it was from using cadmium as an Fe additive to NiFe that NiCd was developed. I can't help but wonder if a sufficient quantity of graphite would keep the iron particles from merging.) But I digress.



3. Battery Construction Overview

   For batteries, one thinks immediately of electrochemistry, but the construction of a battery is no trivial part of making it work. A good part of the effort of four years of battery R & D was trying to come up with workable ways to actually make a battery, any battery, as a feasible DIY project.

Electrodes Overview

Everything else depends on the electrodes. Besides the chemistry, what's in an electrode? how is it made? What are its properties?

   First, all points inside an electrode must be electronically connected together, that is, connected for electron flow. Ideally it is one total "short circuit" from any point to any other point. All the active material is electrically connected straight to the battery terminal. In practice there may be resistance, even considerable resistance, between points because many active materials are semiconductors, but there can't be any insulated points. Parts of an electrode that become insulated from the rest cease to function; they are "passivated" like sulfated lead-acid battery plates gradually become. The lower the resistance within the electrode, the more current can flow with less voltage drop.
   Again, electronic conduction refers to conduction of electrons, wet or dry, not ions. Conduction only by ion flow when it's wet may read "connected" on an ohm meter, but it won't work.

   Second, all active points of the electrode must be wetted by the electrolyte. The reactions only take place when the electrolyte ions can interact with the active chemical. Again, any parts of an electrode where the electrolyte is blocked are passivated and do nothing.

   These two requirements, electron flow and electrolyte penetration, are in conflict for actual physical construction. A good battery requires an immense active surface area in contact with the electrolyte. The surface area of a sheet of metal is small, and all but the very surface atoms of the sheet are wasted, out of contact with the electrolyte. A vast multiplication of minute particles to make a porous substance is required in order that the battery electrodes need not span a gymnasium to supply much current or store much energy.
   On the other hand, these many minute particles must all be in electrical contact with each other and they can't physically fall apart. To achieve this, they must be "glued" and compacted from loose powder into something more like a dense piece of sandstone or brick - a porous electrode "briquette". The briquette must be well compacted so the particles electronically connect, and yet consist of open pores so they all also contact the electrolyte. And the binder 'glue' can't interfere or coat the particles.

   Since connections are generally still poor though the maze of particles over much distance (and increasingly poor with oxidation level), some sort of continuous metal or carbon conductor spans the entire area of the electrode, the "current collector". The briquette is compacted around this for good contact throughout. None of the grains are more than the electrode thickness away from this plate, mesh or metallic sponge that is connected straight to the battery terminal.

   Obviously there's an optimum compacting pressure to achieve the best compromise between electronic conductivity and pores for ionic conductivity. Doubtless this varies with the ingredients in the electrode mix. An electrode with fluffy nickel hydroxide and considerable graphite powder may have a different optimum pressure than a dense zinc electrode with few additives.
   The only figure I've seen for compacting pressure was in one research paper where the authors mentioned an "optimum" pressure of 675 Kg/sq.cm - 9600 pounds per square inch - for an iron oxide electrode. For the chosen 1.5" x 3" electrode size, that would be 21.5 tons. I trust this may be taken as a maximum pressure requirement. I describe some electrode compactors and ways to get sufficient pressure in the appendices. (I hope to offer a good compactor press with a "steering wheel" type tightening handle - but it can be done by tightening some bolts with a wrench, too.) Getting the pressure is one of the chief keys to making batteries.

   The material chosen for the current collector and the terminal leed is important. More particularly, the surface of the material, in contact with the electrolyte, is important.

   In the salty battery with neutral pH, every metal I tried for a current collector in the positrode dissolved. To manage this (after over 3 years of frustration) I created "grafpoxy", a 1 to 1 (by weight) mixture of epoxy resin and graphite powder. A relatively fine metallic screen (around 30 mesh), with a terminal riveted or welded to it, is coated in grafpoxy for use as the current collector. The epoxy protects the metal from contact with the electrolyte, and the graphite lets the electrode substance electrically contact with the mesh.
The mix should have about as much graphite (by weight) as epoxy. This generally makes rather thick for painting or dipping, so some solvent is added, eg, 10% toluene, to thin it. (The solvent evaporates.) I find that two coats are needed, and it should be inspected in a good light. If any trace of copper color is visible, the metal will dissolve away until the cell quits working.
   The grafpoxy coating does for salty batteries what nickel plating did for alkaline batteries in 1900 - makes them practical. It replaces the carbon rods and graphite sheets I was trying previously, or graphite impregnated plastic contact sheets, none of which make very good and durable contact with the electrode briquette. (But in January 2012 discovered "Pourbaix diagrams", which show that a somewhat alkaline electrolyte is best for virtually all of the chemicals discussed. This can evidently be obtained by using salt but adding calcium hydroxide to the positrode. The slightly soluble Ca(OH)2 raises the pH to (theoretically) 12.3, an "ideal" moderately alkaline pH, tho still caustic enough to be somewhat hazardous.)

   In a higher voltage negatrode, a material with sufficient hydrogen overvoltage must be chosen. (Hydrogen voltage is -.833 volts in pH 14 alkali.) I tried many manganese and zinc electrodes with copper or nickel plated mesh that would self discharge and bubble hydrogen. I could understand this for the experimental manganese, but zinc was a known, working electrode chemical. It was ages before I finally realized it was the current collector doing the bubbling, and not the active chemical substance itself.
   Zinc metal itself, or silver, evidently works well. Recent research in Iran showed that a tin-zinc mixture also corrodes. But this research showed that an alloy of copper, tin and zinc, "optalloy" evidently "acts as a noble metal" with a high overvoltage and works well.
   I thought that the simplest thing to do would be to use a long, thin zinc plated or galvanized nail or bolt in the square cylinder pocket electrode. However, these proved to cause a fair bit of self discharge.

   Instead, I got "zincate solution" for 'priming' aluminum and zinc coated an aluminum rod (after shining it up with a nylon scouring pad). This seemed to get rid of the remaining self discharge. (As of today - 2012/03/11.) The solution can be found at Caswell Plating [.com], or can probably be mixed from sodium hydroxide (caution: very caustic! especially protect your eyes!) and zinc oxide.

   For an electrode made with a grill, the simplest thing might be to to melt some tin or tin-silver solder in a pot on the stove, and mix in some zinc, which will gradually melt as it alloys with the solder even if the temperature isn't hot enough to melt zinc by itself. Put some soldering flux on a copper grill and wire, and dip it in the pot for a moment to get a coating. The tin-zinc coating may or may not corrode away, but when it reaches where the copper is present, it should stop, providing a thin layer of the copper-zinc-tin alloy. "Optalloy" (copper:tin:zinc, 55:25:20) was specifically used in the research. "White bronze"
is also commonly over 1/2 copper with the other two metals in fairly equal proportions.

   After the compacting there's the wetted electrode in the cell, before and after charging. Electrodes want to swell when wetted (especially nickel hydroxide), and if they are able to do so, they lose their conductivity and become pretty much useless. This was a major problem through most of my battery research. The best solution appears to be the perforated rigid plastic pocket electrodes, which hold the substance in and can take the pressure.

   There are at least 3 types of electrode construction: Pocket, Sintered Plate, and Paste Electrodes. All of them use powders of the active material, usually with additives mixed in.

   Pocket electrodes consisting of thin perforated metal enclosures, "pockets", to hold the electrode briquettes, were invented in the 1890s. These work great and are highly conductive but with metal pouches holding the electrode materials they're expensive to manufacture and heavier, with low energy densities by weight. Nevertheless nickel-iron alkaline pocket batteries were better than lead-acid and Edison's best version was in common use in early electric cars by 1910 or so. They had to be nickel plated (at least in the positive electrode) to avoid dissolving away. Since all common metals including nickel dissolve in salty electrolyte, metal pocket cells would be impractical for them.
   However, having dismissed pocket electrodes for most of the duration of the battery project, I ended up adopting perforated rigid plastic pocket electrodes as the best choice for homemade DIY batteries. The extra weight of the electrode and battery case structures is compensated and more by better, higher energy chemistries.

   Sintered electrodes were invented in the late 1920's as a better way, but their manufacture and use only spread gradually. The carbonyl or a powder of the metal, usually nickel and cadmium for Ni-Cd's, would be sintered (heated until it softens and flows a bit, and the particles just barely melt together where they touch) into a porous "metal sponge" structure full of minute open cavities -- 80 - 95 % empty space. The electrode active particles would be impregnated into these spaces, then the whole thing compacted. As the conductive metal permeates every little recess of the entire electrode, these are highly conductive and have great current capacity from small cells, eg ~20 amps from ~"AA" sizes. The sintered metal also holds the compaction without an external shell. Energy density is reduced by the heavy inert sintered "sponge", which would make up a considerably greater percentage of the electrode volume after compaction than prior. The Ni-Cd sintered "AA" size cell might have up to around one amp-hour capacity. Again the sintered metals would dissolve in salty electrolyte, making this type impractical - unless perhaps a porous sponge of 'grafpoxy' could be created. I have little confidence in this idea.

   In alkaline paste electrodes, the powders are simply compacted around a nickel or nickel plated metal mesh or perforated foil "collector plate", with a binder "glue" in the mix. The compacted briquette is the finished electrode, with a nickel leed welded to an edge of the foil or mesh. Since there's not much there besides the active chemical and its additives, the highest energy density by weight is attained. The metal case of the dry cell prevents decompaction of the electrodes, which virtually fill the entire space within the cell. Generally the current capacity is lower per square centimeter of electrode than sintered types. The amazing 2.6 AH Ni-MH size "AA" (100 WH/Kg) will only put out about 7 amps. (There are 2.0 AH "high rate" NiMH's that are good for 20 amps. These are probably the sintered type.) Powder/paste is also the newest type, the easiest to make, and the most fragile electrodes to handle for insertion. Zinc electrodes are especially crumbly. But electrodes harden up in use inside the cell as pathways establish themselves.

   I found the perforated plastic pocket cells were the most reliable to make with DIY construction methods.

   Common binders for the positrode include CMC (AKA CMC gum, AKA sodium[?] carboxy methyl cellulose) in nickel electrodes.
   For the negative, PVA (poly vinyl alcohol)  PTFE (AKA teflon, AKA poly tetra fluoro ethene, AKA poly tetra fluoro ethylene, AKA (C2F2)n) suspension, with a fairly coarse particle size.
   I've variously tried in different mixes with different techniques, sometimes for different reasons: Sunlight dishsoap, fried beans (to the point of them catching fire and burning for up to 60 seconds), acetaldehyde, VeeGum (a bentonite clay mixture), and agar agar gel.
   After many experiments, I read that CMC should be "under 1%" of the electrode substance. PVA of up to 2% has been used in zinc electrodes. These figures mean I was using substantially too much of whatever I tried.
   I hear that PTFE is the best for "ordinary chemistry" alkaline electrodes, but it seems hardest to get. Also because a substance works well in alkaline cells doesn't mean it'll necessarily work well in salty cells, as is illustrated by nickel platings of electrode structures.

   An important consideration is how thick to make the electrodes. There are two considerations limiting the thickness of flat plate electrodes: electronic conductivity and ion conductivity.
   Naturally, with an electrode that has semiconductor active material connected to a collector sheet or grill, the thicker it is, the more the internal resistance from the surface layer to the collector.
   And, the thicker an electrode is, the farther removed its back recesses are from the other electrode and the farther the ions have to travel. Thus thinner electrodes may be expected to effectively have lower resistance and higher current capacity even for highly conductive electrodes. Voltage with thicker ones will drop off more at high currents when their charge is lower, as the remaining charged material at rear must come into play.
   I estimate that for flat plate electrodes in KCl, reasonable thicknesses are around 3mm for high rate, 6mm for medium rate, and 9mm for low rate batteries. For electric transport, they should probably be under 6mm unless there are quite a lot of batteries sharing the load. These are all considered very thick electrodes in most batteries. Typical alkaline cell electrodes may be 1mm or less. Lithiums generally have thin films.

   For the square cylinder pocket electrodes, the .5" square will have substantially higher resistance than the .375" (3/8") square, but both types will be substantially higher resistance than a thinner flat plate. A larger diameter of central current collector wire will reduce the distance to the particles and lower the resistance, but unless it's hollow, it'll add weight without adding storage capacity.
   If the internal resistance can be lowered, and-or if the electrolyte can penetrate better, larger diameter cylinders will perform better. Having picked this simple construction and got working electrodes but with rather low conductivity, this will be a bigger focus in future development.
   The manganese standard dry cell "+" electrode occupies almost the whole diameter of the cell - the construction does work, but is generally for low current rates.
   Of course, the taller the cylinder is, the more electrode cylinders there are in the cell, and the more cells that are in parallel, the lower the overall resistance will be and the higher the current that can be driven, but they will still charge and discharge at about the same rate.

   Another aspect to this problem is the speed of ion diffusion through the electrolyte. If an electrolyte diffuses ions twice as quickly, the electrodes may be considerably thicker and still have the same current capacity. Potassium chloride is, I believe, about the fastest electrolyte. Potassium compounds (KCl, KOH) are known to be faster than their sodium equivalents (NaCl, NaOH).


Battery Layouts

   One common type of battery construction is wrapped, spiral electrodes, common in "AAA" to "D" NiMH dry cells. A "V" of separator paper encloses one electrode.
   Another common type is "prismatic", where alternate positive and negative flat plates, separated by sheets, are connected together in parallel to each terminal. This is usually used in flooded cells such as lead-acid.
   These constructions have the advantage of providing the maximum interface area between electrodes, each plate being adjacent on both sides through separators to an opposite electrode, except of course for the two end plates, or the outside of the spiral.

   An older construction is "pocket electrodes", wherein a minutely perforated, nickel plated shell of thin metal holds the electrode chemicals compacted. (Nickel is the only metal that doesn't corrode away in the positrode in KOH or NaOH alkaline electrolyte.) Typically there's a wall every 1 to 2 cm, dividing the pocket plates up into rows or columns - they would bulge out if wider. These plates, just over a couple of millimeters thick, are then used in "prismatic" form with spacers between the electrodes. The metal pockets are thin and perforated, but they do add some weight to the battery. Nonetheless, the prototypical pocket electrode battery dating back to about 1902, nickel-iron, substantially outperforms lead-acid and lasts far longer - some decades old NiFe batteries still work today.

Chosen Battery Layout:
The Checkerboard of Perforated Plastic Square Cylinder Pocket Electrodes

First battery with two 1/2"
square cylinder electrodes
   At the start of February 2012 I suddenly conceived of a completely new construction, easier and more certain to succeed as DIY construction: the Perforated Plastic Pocket Electrode. This harkens back to the early days of batteries - but they didn't have plastic back then. A number of previous ideas came together, and a couple of new ones were soon developed, to make this work. A battery of any size could be assembled from easy to make, square cylinder plastic pocket electrodes, layed out checkerboard style. Each electrode was a separate unit, individually compacted and held that way. Initially the biggest problems were bursting of the perforated cylinders and high internal electrode resistance.

   The first plastic cylinders were 1/2" square inside. This made electrodes that were just too fat, and the conductivity was very poor. So this electrode size was changed to 1/4" square, more in keeping with the electrochemical requirements
(see "Electrodes Overview", next section), and it increased conductivity an order of magnitude. The cost was making four times as many electrodes. I may try 5/16" and see how that works, but they'll probably be pretty low rate cells. At that point, I decided a special jig to help compact the powders into the small tubes was needed to speed things up.



Right jig: channel for folding the plastic around a 1/4" steel rod, after heating it in an oven (rod shown is 5/16")
Left jig: electrode stuffing jig. Powder is dropped/brushed into slot,
1/4" rod pushes it in, and then tamps it down (not too hard).
After the top is glued on, a zinced nail is driven in.
Right electrode is the original 1/2" size, replaced by four 1/4" size.

   The sides of the first square tubes were made from .063" ABS sheet plastic. Next will be tougher .020" styrene plastic to cut waste space, weight, and the distance between electrodes. They are perforated with a heavy-duty sewing machine.
   The perforated plastic is cut into sections about 25 x 65 mm. This is heated in a kitchen oven on an unwanted cookie sheet or shallow baking pan to 350 degrees F, for about 3 minutes if the oven is preheated. (The longer it's left in, the more it shrinks.) In 3 minutes, it should be pretty much limp and can be formed into any desired shape. That shape is a four walled cylinder of 1/4" square inside dimensions, by 60mm tall, formed around a 1/4" square steel rod with a jig. A bottom and a top end cap close the ends. The top cap has a hole for the connection wire. The overlapped seam and the end caps are glued with methylene chloride, a solvent which dissolves the plastic, thus making the plastic its own glue as it evaporates.

  Filling the electrode is to be done with an "Electrode Stuffing Jig" having a slot for electrode powder mix to fall into place for a 1/4" square steel plunger rod to stuff it into the cylinder and tamp it down. Once it's filled, the end caps are glued on.
   The straight connection wire runs right through the cylinder from top to bottom, and sticks out the top far enough to poke through the top of the battery to solder to. At the roof of the battery, it's sealed with RTV cement, or epoxy.
   For the negatrode, the connection wire is a galvanized box nail, pounded into the electrode after both caps are glued on. The zinc coating on the nail has the required hydrogen overvoltage.
   For the positrode, a nail is used to make a hole, then a grafpoxy (or "nimangapoxy"?) wire (too fragile for pounding) is stuffed into the hole.

Like the carbon rod in the manganese center of a "carbon"-zinc dry cell, a single grafpoxy coated wire sticks out the center of the top of the electrode for connection, for both polarity of electrodes, of any chemistry. The electrode chemical mix is compacted by tamping it directly into the plastic shell with a hammer and a punch. (The punch has a center hole to fit around the wire.) Once the bottom is glued on, there's nowhere the chemicals can expand except by bulging the stiff plastic sides. (A bit of expansion ability is vital for most pocket electrodes, depending on the chemistry, temperature, etc.)

Separator Sheets

   Separator sheets up until the 1970s were pretty simple: any insulator to keep the two electrodes from touching. Recently there have been developments of interest.
   The first is the zircon (ZrSiO4) or zirconia (ZrO2, which becomes Zr(OH)4 when wetted) ion shield. While allowing the passage of chlorine (Cl-) or hydroxyl ions (OH-) in neutral to alkaline solution, a layer of zircon prevents migration of metallic cations. Depending on how effective this shield really is, that can mean that not all reaction products need to be solids.
   The first potential use is to prevent migration of zincate ions, which should greatly extend the life of NiZn batteries even if the shield is only partly effective.
   The second potential use, if the shield proves fully effective, is to permit new chemistries where either the charge of discharge product is partly or primarily a dissolved ion. A prime example might again be zinc, which forms soluble zinc chloride on discharge in the standard salt electrolyte dry cell. If these can be blocked from crossing the separator, the standard MnZn dry cell might become rechargeable.
   Another possible use is according to the diagram - for a single use battery where some of the products are dissolved ions.

   A very good separator sheet is Arches 90# watercolor paper, tho various papers from coffee filters to writing paper to cardboard might work. The Arches is uniformly thick and even.
   To make an ion shield on this, it is simply necessary to buy zircon powder ("Ultrox" is a pottery trade name for the purest type - use the purest), wet it, and paint it onto the paper with an artist's paintbrush. Be sure the target electrode is fully wrapped so no electrolyte can get around the shield. Zirconium oxide costs more and is less readily available.

   For the perforated pocket electrodes, separator sheets are hard to place, and a space between the electrodes can be enough. It does make it hard to try any of the fancy things!

Working Up

   The battery isn't ready to use immediately on assembly and filling. First it should be left to soak for a day or so - at least overnight. For a monel electrode, the color will change from purple water to blue-green solid as the monel is oxidized by the permanganate and swells up to fill the available space.



4. Making the Case and Fittings

   I had hoped to simply buy suitable cases of molded plastic, but I couldn't find any. Then I thought I'd look for rectangular plastic tubing of a suitable size to cut to length and glue in bottom pieces, but I couldn't find any of that, either.
   It's frustrating because round ABS and PVC plumbing pipe and fittings are everywhere. These could be heated in an oven (about 300ºF) to soften them and bent into rectangular shape, but it's hard to control the dimensions properly and consistently so everything fits well.

   I finally decided that the best flat material to make a rectangular case from seems to be acrylic plastic ("plexiglass") or lexan. I prefer to use clear stuff so the quality of the seam joins can be inspected, as well as the battery inside. The edges can be 1/2" or 5/8" thick acrylic/lexan turned sideways, so the battery is exactly 1/2" or 5/8" thick inside with no gaps around the edges and no need to sand everything flush. The edge pieces can be cut about 3/8" wide.
   Electrodes are fragile and keeping even 1.5" x 3" ones intact until they're in the cell can be challenging. But a 4mm thick electrode this size has as much active material as one 3" x 8" that's only .75mm thick, as is typical of alkaline. By the time some paper is wrapped around one and maybe things are a little off dimensionally, the case needs to be preferably about 3.2" x 1.8" on the inside so they go in readily.


Prototype battery case of acrylic plastic,
with prototype grafpoxy coated electrode collector screens.

   I tried ABS before acrylic, but in sealed cells I found it kept leaking at the seams - it seems a little too flexible, so it deforms a bit under pressure, putting high stress on the seam bonds. It would probably work well for vented cases, but at the moment I'm using the clear acrylic for one-off cases. ABS faces with acrylic edges might also work.

Face Pieces: 1/8" or 3/16" acrylic or lexan plastic (plexiglass), or 3/16" or 1/4" ABS, cut size: 4" x 2-3/8".
Left and Right Edge Pieces: 1/2" acrylic, cut size: 3/8" x 2".
Bottom piece: 1/2" acrylic, cut size 3/8" x 4".
Top Piece: Here for my prototypes I prefer to cut a custom piece of plexi with some sort of lip at both ends.


   For gluing plastic cell wall pieces together from sheets of plastic, it's about 4" x 2.4" x .8". Cut the main faces 4" x 2.4" or 100mm x 60mm. Using 1/2"/12mm thick plastic for the edges, cut them about .3" wide. Turning them sideways, the 1/2" uniform thickness goes between the wide sides as the internal width, and your cuts won't have to be perfect or sanded down to exact uniform thickness, except the bottom ends of the two side pieces should contact the edge of the bottom piece well. The chief place to watch for leaks is the bottom corners.
   I would LOVE to have injection molded cases and lids, and I hope I can make the molds some day, or perhaps two or three of them of different thicknesses. (January 11th 2012 note: I've just acquired a milling machine and ordered a kit to make it into a CNC milling machine - the vital tool for making injection molds.)

   Molded ABS (or other plastic) boxes would eliminate the seams except around the lid. A molded ABS lid with a lip made to fit over would solve that. They'd be my ideal for production - sealed or vented - when the visual inspection aspect loses its meaning.

   A possible way to do seamless rectangular boxes without injection molds would be to make a simple mold from a solid block of polyethylene, wrap polypropylene fabric ("landscaping fabric") around it, and epoxy it. When it sets, add more layers if needed to get to the desired thickness. Spray the mold block with wax to help release the finished box from around it. I would think making the block taller than required and putting in a handle hole to pull on would be the way to go.

Grafpoxy Current Collector Grills and Terminal Leeds

   The first thing needed is a compatible metal. Nickel is suitable but for some reason seems harder to get than gold. Copper seems to work, but is very hard to tack-weld. Stainless steel mesh doesn't seem to work.
   The grafpoxy is a one to one mixture of epoxy resin and graphite. It's rather thick with West System epoxy, so I'm going to try diluting it. Evidently acetone is a common solvent, also MEK, xylene... Let's see... I have toluene. I guess I'll try that.


Electrode Current Collector Grills & Terminal Leeds

  Current collectors first and foremost must not corrode away in the electrolyte during charging and discharging. Since every common metal corrodes away in salty electrolyte (very quickly in the positrode), only conductive carbon substances such as graphite can be used. This is why the standard dry cell has a conductive carbon rod for a terminal. But these rods are hard to make and brittle, and graphite isn't as conductive as one might wish. Graphite sheets or even flakes tend to degrade in the cell during charging. Fine graphite powder fares better, also as demonstrated by the standard dry cell.
   In 2011 I invented 'grafpoxy', simply a 50-50 (by weight) mixture of epoxy resin and graphite powder. The resin makes it impervious to the electrolyte, and the graphite makes it at least somewhat conductive. This could possibly be used and molded by itself, but metal is still a much better conductor. So the grafpoxy is used to coat a metal current collector and the parts of the terminal inside the cell.

   The best current collectors are those that provide the best conductance to every bit of the electrode, yet allow electrolyte to pass though so they can be in the middle of the thickness, and to which the electrode substance will remain affixed. A fine grill, eg, 20 to 40 wires per inch, is ideal. Compatible metals include (at least) copper, brass, nickel-brass (AKA "nickel-silver"), monel or other copper-nickel alloy, and nickel. Pure copper is so conductive it's hard to tack weld. My ersatz tack welder won't make any sort of join to copper.


Rivetted copper grills before grafpoxy coating,
and electrode compactor revised to work with book (or hydraulic) press.
Left electrode grill has one installed rivet,
one not yet spread with center punch and hammer,
and one empty 1/8" hole
   I got some expanded copper mesh for my prototypes at an art supply store, but it's only about 10 x 15 'wires' per inch, which is rather coarse, and I'm riveting it to copper foil for the terminal leed. I'm still looking for something better. Here I scrunched up the mesh somewhat (and then hammered it flat again) to get a few more wires per unit area. The small rivets, 'post and cap' (IIRC) rivets, are available from leather supply stores (i.e. Tandy Leatherword). I only use the posts, spreading the thin end with a center punch and then hammering it flat (see foto - the fat 'head' end of all the rivets is underneath).

   With this third batch of electrodes, I went back to using a wire for the terminal instead of the foil, and I made the slot in the compactor for wires only. The inner end is flattened and runs the length of the electrode inside the foil.
   If the electrolyte gets through the grafpoxy and inside the foil, it'll spread along the wire and it'll all dissolve from the inside. I did two coats of grafpoxy, and touched up some spots with a third.
   I used some toluene solvent to thin the grafpoxy. It seemed about right at first, but it gradually became thicker as I worked, as the solvent evaporated.
   I would very much like to automate production of the grills, as making and coating them is very tedious. With nickel or perhaps brass alloys, one could at least tack-weld the wire and mesh together and skip the foil and rivets.















5. Making Electrodes & the Positrode

Electrode Making Procedure

My general procedure for making an electrode briquette, positive or negative, is:

1. Make a grafpoxied mesh grid collector grill with a terminal leed sticking out one corner.

2. Put a 1.5 x 3" piece of thin polyethylene plastic sheet into the electrode compactor. Put in the current collector grill on top of it. (Instructions are elsewhere in this book for making compactor boxes and the grills.)

3. Have all ingredients prepared and on hand, a sub-gram weigh scale, and lightweight plastic "dishes" for weighing ingredients on the scale.

4. Measure and mix the dry ingredients. If necessary, grind them with a mortar and pestle to get a fine, uniform powder. (I ordered a glass mortar and pestle through a local drug store. A pestle ("4 oz") one size smaller than the mortar ("8 oz") was helpful.)

5. Add the liquid ingredients and mix thoroughly. The mix should seem barely damp. If there is too much liquid, it will simply ooze out the cracks instead of compacting. You should be able to tamp it down in the mortar and check electrical conductivity with an ohm meter. If it's above 100 ohms, you might want to let it dry a bit until you can tamp it down harder, or if it seems dry enough, add more graphite. If it's too dry of 'diesel kleen', the graphite won't process properly in the compactor.

6. Put about 1/2 of the mix into the compactor. Try to get it evenly spread around in the box. Put in the current collector grill on top of it, putting the terminal leed through its slot in a corner of the box. (Instructions for making the grills are .)

7. Put a 1.5" x 3" sheet of polyethylene plastic over top, then put the die in over that.

8. Put the lid on and do up the bolts (or clamp the box in a heavy book press or 20+ ton hydraulic press) to compact the electrode into a "briquette". Leave it in with the top tightened for 1/2 an hour or more. During this time, the graphite will dissolve in the diesel kleen (probably the methyl benzene is the active ingredient) and form lamillae - random linear formations that connect the active elements randomly across the electrode for the best short circuit conductivity throughout.

9. Remove the briquette from the compactor.

The next steps show what fun you miss by using perforated plastic pocket electrodes. Skip to step 13.

10. For positrodes only, paint calcium oxide on the surface. (Making calcium oxide from calcium carbonate is in the appendix.)

11. Dry it in a toaster oven outdoors (not indoors - the diesel kleen reeks) for over an hour at about 100ºC/212ºF.

12. Play a propane torch over the surface for about 5 seconds to sinter and hence harden the surface layer. The battery will last longer. If the electrode isn't wholly dry, it may suddenly pop apart from steam pressure - hence the oven step above.

13. Drip some toluene onto the briquette. This will absorb in and dissolve a bit of graphite, which will (theoretically) form into carbon nano-tubes as the toluene (or turpentine) evaporates, creating the best connections.

   Naturally, you'll save considerable effort by making multiple identical electrodes at a time. (It wouldn't hurt to have two, or even three, compactor boxes.) If some of this seems rather intricate and involved, it's because it is. It's more than has traditionally been done to create electrodes, but they should (I hope) be the best they can be and extremely long lasting.

   I caution however that some of the steps are just my own ideas which aren't all verified as to their efficacy or even their actual effect. It should be realized that I have created these things on my own on a very low budget with very minimal equipment, concurrently with other inventive projects. In no case am I certain all the ingredients are given in optimum proportions - many are just fair guesses. The amount of compaction is just whatever the box can do - it's probably below optimum pressure in most cases. Sintering with a torch to toughen the surface is a theoretical idea. The diesel kleen and toluene do appear to improve conductivity. But to verify fine scale graphite lamillae surely needs at least a microscope, and to verify carbon nanotubes would require an electron microscope. A calcium layer has also been used by others to help with toughness and oxygen overvoltage, but whether oxide (lime) is the best form I can't verify. (It's likely to turn into calcium carbonate if left exposed to the air too long. This must be what happens to bags of cement.) Torched barium carbonate (turns to oxide on heating) might be better. Thiamin to chelate the rare earth and ions is experimental, as is the specific choice of antimony sulfide, to raise hydrogen overvoltage.

Nickel Manganate Positrode

   I presently think this is the best choice. It has high voltage and I believe should have very long life, and the high amp hours of manganese.
   I won't try to give a formula this time - instead a principle. Measure out 'X' amount of NiMnO4(?), then add some graphite powder. Add 1 or 2% Sunlight dishsoap, a few % rare earth oxide or hydroxide, and a bit of water.
   Mix it well then tamp it down. Measure the resistance. If it's well tamped down and the resistance is over tens of ohms, add more graphite and try again. If it doesn't tamp down well, add more water or let some evaporate to get a better consistency.

Manganese Positrode

   The ideal powder mix for Mn positive electrodes is the positive electrode powder salvaged from throw-away dry cells.
   To this, add perhaps about 1% Sunlight dishsoap added to help permanently glue the powder together, and enough Diesel Kleen to dampen it to a dry paste. The paste becomes more moist as it is compacted into the electrode cylinder since liquid doesn't compress. If it oozes out the cracks, it has too much liquid. Best to wait for some to evaporate off. Diesel kleen evaporates more slowly than water (despite the smell). Methylbenzene (in the Diesel Kleen and in toluene) dissolves graphite. As it evaporates, the graphite forms into random lamilar nanotube structures that are more conductive, improving current capacity. Another way of doing this might be to add toluene after and let it soak in... if that doesn't hurt the plastic enclosure.
   Since the voltage is lower than nickel, there should be no need to add oxygen or chlorine overvoltage raising ingredients.
   Come to think of it, with the graphite forming conductive nanotubes, it might be worth experimenting with lesser amounts of graphite in the mix, eg by adding a percentage of pure MnO2 to it to dilute the graphite.

   Since I've already written instructions for making some nickel electrode types, I'll leave them in here. But I no longer recommend them.

Nickel Positrode

   The main dry ingredients, except for the graphite powder, each can take more than one form. See the list of supply sources to find each form of each metallic element.

   Nickel can be had as fine nickel powder, nickel oxide (NiO), nickel hydroxide (Ni(OH)2, nickel sulfate, or nickel carbonate. The recommended form is NiO, followed by Ni(OH)2. The sulfate or carbonate have to be converted to Ni(OH)2 in alkali. Unless you have a pre-existing stash, why bother? Ni(OH)2 can be converted to the charged form, NiOOH with bleach, and back again with hydrogen peroxide. (see appendix)
   In the salty battery it's not clear (to me) whether the nickel will take the typical alkaline forms above or: NiO, Ni2O3 and NiO2 for valence states 2, 3 and 4. I suspect NiO is the discharged form, and NiO2 is most likely the valence four form. For valence 3, toss a coin.
   Remember these various oxide and hydroxide forms only affect the consumption and release of water during charge and discharge. It's probably relevant in a 'dry' cell, but not really in a flooded cell.

   The manganese is best added as potassium permanganate, which is likely what it'll become when the cell is charged anyway. That way, it already has its potassium and there's no chlorine left over from charging with the potassium form the potassium chloride electrolyte. But if that's too hard to obtain (owing to bizarre substance laws or local supply) it can also be added as manganese dioxide. That will charge to KMnO4, but in the process use up some of the KCl salt electrolyte, releasing a bit of chlorine gas. (doubtless undesirable in a sealed cell.)

   Many of the rare earth oxides (REO) might work fine. I'd say the order of preference is: samarium, neodymium or lanthanum. Cerium is one that should probably be avoided, owing to a possible charging reaction from valence 3 to 4 (Ce(OH)3 <-> CeO2). Mischmetal (unseparated blend of rare earth metals) oxides should be avoided because they'd contain cerium.

Dry ingredients:

30g - NiO (or 37g Ni(OH)2)
20g - KMnO4 (= 9g MnO2)
22g - fine graphite powder
1g - Sm(OH)3 (or other rare earth oxide/hydroxide)

Liquid ingredients:

2.5g - Diesel Kleen
.5g - Lemon Fresh Sunlight dishsoap - no other dishsoap (or PVA or teflon powder)

Monel Positrode

   The monel positrode is still somewhat experimental as to ingredients and proportions. It is essentially a nickel positrode in which the nickel and manganese mix is replaced by nickel and copper oxides in solid solution. In order to attain the solid solution state, the nickel and copper are purchased as fine monel alloy powder. A typical monel alloy is about 2/3 nickel. The powder I obtained contains: Ni 67%, Cu 28%, Fe 3%, Mn 2%.
   All of these metals will oxidize/hydroxidize when the battery is charged. The copper oxide in solid solution improves the conductivity of the nickel oxide, hence less graphite is added. The manganese is at worst benign. (Perhaps one should add more.) The iron is probably benign as well.

   The only form to obtain the substance in is as metal alloy powder.

   As with the other nickel electrode formulations, a rare earth oxide/hydroxide is added to increase the oxygen overvoltage, in the same order of preference as to element. Some thiamin (bean sauce) is added to chelate the ingredients and make the electrode very long lasting.

Dry ingredients:

40g - fine monel powder
10g - fine graphite powder
1g - Sm(OH)3 (or other rare earth oxide/hydroxide)

Liquid ingredients:

2.5g - Diesel Kleen
2g - tinned bean/bean sauce
.5g - Lemon Fresh Sunlight dishsoap - no other dishsoap (or PVA or teflon powder)

   The torching step is extra important to "bake the beans" and set the thiamin.

Vanadium Positrode

   This electrode is very speculative as to ingredients and proportions. The voltage appears to be somewhat higher than for nickel.

Half reactions. (charged  <==>  discharged)
V2O5 + 2 H2O + 4e-  <==>  V2O3 + 4 OH-  [~ +1 V]

The reason for adding nickel oxide is to complement the vanadium in the structure and (perhaps) to provide a means to raise oxygen overvoltage. I'm not quite certain of the exact forms the nickel or its reactions will take. Here are the chief possibilities as I see them:

NiOOH + H2O + e-  <==>  Ni(OH)2 + OH- [~ +1 V] (same as in alkaline cells but voltage is higher.)
NiOOH + e-    <=====>  NiO + OH-         [~ +1 V]
NiO2 + 2e-     <=====>  NiO + 2 OH-      [+1.1 V ?]

It's less likely that NiO2 will be formed, and that instead the nickel will provide some oxygen overvoltage protection, to the voltage level required to form it. At least, I suspect that's how it works.

Antimony has an almost unique ability to interact with small molecules like hydrogen and oxygen. The antimony sulfide is added as a catalyst to recombine O2 and H2 gasses that get generated during charging (in spite of all measures to stop or reduce them) especially towards the end of the charge, keeping gas pressure from bursting the cell and allowing larger sealed cells. Adding this Sb2S3 to the electrode materials allows O2 & H2 to be recombined wherever they happen to meet, easing requirements for gas transport to specific locales.

The main chemicals used in the positive electrode briquette are:

- 66wt% (% by weight) Vanadium pentoxide, V2O5 (Caution: very POISONOUS if ingested)
- 33wt% Nickel oxide or hydroxide, NiO or Ni(OH)2
- 1wt% Stibnite = antimony sulfide, Sb2S3 (also very POISONOUS)

Also in the mix:

- fine graphite powder - about 50wt% as much as the total of the above ingredients
- a small wad of chopped carbon fibers, immersed in Diesel Kleen (If it looks too hairy after compacting, use less.)
- Lemon Fresh Sunlight dishsoap
- enough Diesel Kleen to dampen. Not much - too much and the substance will ooze out of the compactor when attempting to compact it, and it evaporates more slowly than water. Too little and the metallic looking surface with high conductivity won't appear. (pictures later.)

   The Sunlight solidifies during charging and discharging into a sort of binder ('glue') to hold the electrode together and keep active ingredients from migrating. This should extend cycle life 'ad infinitum'.

   The graphite powder and carbon (graphite) fibers are to improve the conductivity. The nickel and vanadium oxides are are poorly conductive, and the battery would scarcely work without some means to improve it. The Diesel Kleen disperses the graphite and helps it to form a connective network throughout the electrode and into the expanded graphite backing sheet (current collector sheet) as the electrode is compacted from loose powder into a "briquette".


- Calcium Hydroxide ("slaked lime")

The vanadium is the active ingredient, with higher amp-hours by weight and volume than nickel oxyhydroxide.
The graphite powder improves conductivity.
Cobalt improves the conductivity of nickel hydroxide positrodes, in "solid solution" with it. I'm only guessing that it will help with vanadium as well, and thinking it might fuse with it in some useful way. Perhaps it's just wishful thinking...
The dishsoap ingredients 'freeze out' and form the 'glue' to hold the electrode together.
The calcium is supposed to raise the oxygen overvoltage to reduce self discharge and improve higher temperature performance.


Chemically made Ni(OH)2 is of about 50% beta and 50% alpha forms. The alpha form is an electrical insulator or poor conductor and there shouldn't be too much of it in the battery. There is a chemical means to convert it. First, pour in bleach and leave it for a while. (10 minutes?) This oxidizes it, converting it chemically to nickel oxyhydroxide. Pour that off and rinse it thoroughly. Then pour in some hydrogen peroxide and leave it a while. Tiny bubbles come off it. This reduces it to nickel hydroxide again, but preferentially to the beta form. Again pour this out and rinse thoroughly. The beta form is also substantially denser than the alpha, so more can be packed into the battery.

I bought nickel hydroxide from Palm Inc (palminc.com). I had to get a 10 Kg bag, which cost about $400. The bag says it came from OMG Harjavalta Oy in Finland, with an e-mail address of nickel.sales@omgi.com . (It's probably only about as hazardous as any fine, dusty stuff, but check the MSDS.)

The cobalt oxide (ceramics/pottery supplies store) is prepared as well (rats, and you thought you could just dump it in!), by heating it to about 500ºF in the kitchen oven for an hour or two. This turns more of it from Co3O4 (structurally Co2O3:CoO) into the Co2O3 form. (gloves -  I think cobalt is poisonous.) This is added in the amount of about 1% to the Ni(OH)2, and then enough dishsoap (grocery) is added to make a paste. Use only the yellow "Lemon Fresh Sunlight", not Ivory, etc. Use no more than necessary: the denser the powder the better. Ni(OH)2 has a listed capacity of 289 mAH/g, and more grams of powder stuffed into the same space also conduct electrons better.

This paste is plastered in as-is. Once it is in the battery (or in the jar if left there), the dishsoap hardens, "gluing" the electrode together. Also, zinc, a desirable element for increasing the conductivity in nickel hydroxide electrodes, leaches out from the nickel-brass cell wall sheet into the paste. This is in fact a main reason for using nickel-brass instead of straight nickel plates.

According to some recent literature, it might also be beneficial to add some yttrium oxide to this mix to improve performance at higher temperatures. At the moment, the experiment seems academic for Victoria BC's climate, and too pricey, there being none at Victoria Clay Arts. (160 $/Kg from HEFA Rare Earths Sept 2008.) In a hot climate, or if the batteries prove to heat up notably when in use, it might make a notable difference.



6. Making the Negatrode

Zinc Negatrode

   The electrode briquette is prepared similarly to the positrode, see chapter 5, except for the actual ingredients in the mix. Zinc electrodes are especially crumbly and fragile.

Dry ingredients:

50g - zinc oxide powder (theoretical maximum: 35 amp-hours)
5-10g - graphite powder
.5g - antimony sulfide (stibnite) or antimony oxide (stibia) (sulfide preferred)

Liquid ingredients:

1.5g - Diesel Kleen
1g - Lemon Fresh Sunlight dishsoap - no other dishsoap (or PVA or teflon powder)

   Vary the diesel kleen as required to get the best amount of dampness. Add graphite if required to get lower resistance readings on tamped-down mix or finished electrodes. Make electrode as per instructions in chapter 5.

After putting the electrode together, it is wrapped in zircon painted separator paper (next chapter). Then it's immersed for several hours in a solution of calcium hydroxide (slaked lime). (Calcium carbonate can be purchased at a pottery supply. Making lime from calcium carbonate in a kiln and dissolving it is detailed in an appendix.) Then the Ca(OH)2 is rinsed or diluted out. The alkali removes passivating carbonate from the zinc oxide, purifying it.
   Some Ca(OH)2 absorbs the carbonate from the zinc and becomes calcium carbonate, CaCO3 - limestone. If it doesn't rinse out, this probably will, if anything, contribute a bit of strength to the fragile electrode.

Manganese Negatrode

   This is (if it works) one of my major battery making coups. No one has been able to make a manganese negatrode before, because its voltage is just a little over the hydrogen generation voltage. Thus, hydrogen bubbles up as it charges, and continues to do so until the electrode has discharged itself to Mn(OH)2 or MnO.
   Hydrogen generation voltage varies with electrode ingredients (as well as pH). Zinc is "improved" by ingredients to raise the hydrogen generation voltage. Manganese is "enabled" only by ingredients that sufficiently raise the hydrogen generation voltage to where it will charge effectively and stay charged without excessive self discharge.
   If it works, manganese makes a better electrode all round: it doesn't form a soluble ion, it's higher voltage (-1.37 vs -1.05), probably more amp-hours per weight (being a lighter element: 56 versus 65.5), and it's stronger and better consistency. It also doesn't appear, from what I can see, that it forms carbonate, so it shouldn't need the alkali treatment - the compacted briquette is ready to use.

Dry ingredients:

50g - manganese oxide powder (theoretical maximum: 35 amp-hours)
10g - graphite Powder(?)
.5g - antimony sulfide (stibnite)

Liquid ingredients:

1.5g Diesel Kleen (this is used as an additive in diesel engines, but it has what's needed!)
1g Lemon Fresh Sunlight dishsoap - no other dishsoap (or PVA or teflon powder)

   Vary the diesel kleen as required to get the best amount of dampness. Add graphite if required to get lower resistance readings on tamped-down mix or finished electrodes. Make electrode as per instructions in chapter 5.



7. Electrode Separators

The electrode separator consist of three layers:

separator paper
Paste: zirconium oxide ("zirconia", 8 parts) and ferric oxide ("rust", 1 part) powders,
    in "Sunlight".
separator paper


Sheets
Chemicals
Fabrication



8. Electrolyte and Cell Assembly
dd



9. Working it up

dd



10. Charging, "Forming" and Testing

   Once everything is put together, the cell might not work very well at first. Here are my results with an early manganese negatrode, placed in alkali with a commercial nickel electrode in a nickel-iron battery case, kept on charge for several days. At first, it seemed it had too much self discharge to really be useful, and little capacity. (I thought my 90:10 tin-silver solder coated current collector - stuck in later - must be bubbling hydrogen, and also that it wasn't making a very good connection.) It seemed like _your_favorite_expletive_here_. But it gradually improved. It needed a lot of patience. Then it started stay up above 1.8 volts longer and longer when the charge was removed, drifting down more and more slowly. Most of the improvements were after charging overnight, and I've skipped mentioning some cycles I tried too soon after the previous one. The figures for the first 6 cycles are from memory (I didn't record them at the time) and are only reliable as far as showing a general trend. The cell voltage could be seen gradually dropping, but occasionally it would jump up a bit and restart the drop from there, probably as little new connections formed within the electrode. This became more frequent with each passing cycle, and voltages stayed higher longer.

Discharge cycle. Held > 1.5 volts for ___ seconds/minutes, and 1.0 volts for ___ with 25Ω load, then after discharge continued to .9 volts, recovered to ___ volts.
1. 60 s, ?, 1.4 v
2. 90 s, ?, 1.45 v
3. 2 m, 6m, 1.5 v
4. 5 m, 10 m, 1.6 v
5. 6 m, 12 m, 1.63 v (and only 1-1/2 hours after test 5.)


   Electrodes are pretty soft when made, but they harden up with forming. They might not even need the plastic enclosures, tho I suspect they'd gradually fall apart without them. Evidently it's common to replace the electrolyte after the electrodes are formed, or even to form a bunch of electrodes in a bath before putting them into a battery cell. This seems like especially good advice for DIY with chemicals that might not be completely clean by the time they're in an electrode.



11. Appendices


"Acetal Polyester" Electrode Binder

   Once the electrode powders are mixed, they need a binder, "glue" to help hold them together and also to prevent migration of the "solid" chemicals during charge and discharge cycles.

   Various binders have been used: PTFE (teflon) powder and CMC (sodium carboxy methyl cellulose gum) among others. PTFE is the preferred choice.
   But these are insulators. A conductive, or perhaps even a semiconductive, binder would not only help hold things together, but improve the current handling capacity of the electrode.

   A great advantage of the sintered type of electrode was that the sintered metal conductor ran through the whole electrode. Thus we would find old Ni-Cd cells for portable power tools, in the "AA" size range, that would put out 25 amps. A modern high capacity Ni-MH "AA" cell may hold over twice the energy, but it won't put out 10 amps.
   A conductive binder would be a great replacement for the sintered metal "sponge".

   I decided to try polymerizing acetal ester hoping to gain this effect. This is the real "chemistry lab" part of making the battery. It's a four step process:

1. Convert potassium dichromate (AKA potassium bichromate) into potassium chlorochromate (AKA "chromic acid") with hydrochloric acid (AKA "muriatic acid").
2. Convert alcohol (AKA "ethanol", AKA "triple distilled vodka") into acetaldehyde (AKA "ethanal") with the potassium chlorochromate.
3. Convert the acetaldehyde into acetal ester (monomer solvent) with hydrochloric acid.
4. Use the acetal ester as the liquid to mix the electrode powders together with.
5. Compact the electrode.
6. Cook the electrode at 110ºc (225ºf) to polymerize the acetal ester into acetal polyester. Also freeze it. (One of these must do it...)

(You were going to mix the electrode powders with some liquid and compact the electrode anyway, so it's only four additional steps, not six.)

Making the Potassium Chlorochromate

It's hard to buy "ethanal" because it's evidently an explosion risk during shipping. To make the unavailable "ethanal" (CH3CHO) from available though pricey ethanol (CH3CH2OH) without accidentally reducing it all the way to acetic acid (vinegar, CH3CHOOH) we need another unavailable chemical: KClCrO3 (AKA KCrO3Cl).

Web info for reducing alcohols to aldehydes invariably says to use pyridinium chlorochromate, but that's pricey and also considered hazardous to ship, so it's even less available. (And don't bother to look up "pyridinium" in the periodic table of the elements - it's not there.) But the important thing is the hexavalent chromium, so we skip the imaginary pyridinium and use potassium chlorochromate. This is very poisonous and acidic, so be careful.

Nobody is likely to have potassium chlorochromate either, but this can be readily made. A ceramics supply store will have potassium dichromate (or -bichromate, K2Cr2O7, also very poisonous and acidic), bright orange pottery glaze crystals. Now, where do we get the chlorine, and how do we put it in? Turns out hydrochloric acid, HCl, has chlorine and works well. Since the dichromate is losing some oxygen and stealing the chlorine from the acid, the obvious byproduct with the hydrogen from the acid is HOH (AKA water), so according to my reckoning we're left with the desired product and water. Of course, some acid will be left over... or potassium dichromate if there wasn't enough HCl. (Actually, one could probably calculate the correct proportions. According to the hygrometer, my HCl from Rona is about 33% strength, and the K2Cr2O7 is 100%. I should probably figure out what atomic weights of the active ingredients per gram that gives.)

Pour a little HCl (RUBBER GLOVES, GOGGLES, do it in the sink; avert your nose!) into a jar and add K2Cr2O7 crystals. Heat it, stirring gently, in a shallow pot of hot water on the stove until the solids dissolve. This takes quite hot water. If the crystals don't completely dissolve, that probably means the acid is used up and there'll be some dichromate left over. Add more acid. At some point in the process, put the lid on. Then label it "KClCrO3, POISON!" and put it in a safe fridge where no one might think it's food. The KClCrO3 will precipitate out and the acidy water can be gently drained off and flushed down the sink. Label it, let it dry out, and put it in your secure chemical cupboard.

Making the Acetaldehyde

The "aldehyde" is considered a hazardous substance to ship - better to make your own.

Put 20 parts (by volume) of Alberta "triple distilled" Vodka (40% pure ethanol dissolved in 60% water by any other name) and 1 part potassium chlorochromate into a small jar and heat it in a shallow pot of hot water on a stove burner, stirring constantly. Once it’s heated, screw the lid on - acetaldehyde boils at about room temperature. (rubber gloves, goggles: KClCrO3 is acidic and poisonous... and acetaldehyde is the hangover chemical.) Shake the jar as it heats and continue until the KClCrO3 dissolves. The end product should have a strong "fruity" smell (seeing it's a major component in the smell of ripe fruit) and should be a clear liquid when everything settles. Label it carefully and put it in a fridge to cool. (Preferably not a food fridge. Label it "Aldehyde, POISON!" and take whatever precautions are necessary to ensure it’s clear that it isn’t food and to make it unavailable to young children. Again, with the chrome stuff in there, it's poisonous.)

If you could just drink the vodka and somehow extract the acetaldehyde from your body before it gave you a hangover, you wouldn’t need the nasty hexavalent chromium stuff. But the end product is unavailable. You feel lousy, and the alcohol in vodka costs more by weight than lanthanum "rare earth" metal - a sobering thought!

Making the Acetal Ester

This is made from the acetaldehyde by simply adding hydrochloric acid to it (HCl - RUBBER GLOVES, EYE PROTECTION). No heat is necessary. After it sits a while it should be a blue-greenish liquid, still with a fruity smell.



A. Materials and Chemicals Supply Sources

Obtaining Manganese: permanganate, dioxide, and metal powder.

   Potassium permanganate is used in swimming pool and other water filters to eliminate iron rust staining. For that reason, is may often be found at water treatment supply stores. [Victoria Water Treatment] Be prepared to buy at least a 5 pound, 60$ plastic jug of it. It's more than you'll need unless you're opening a battery factory, but most customers want 20 pounds or more. KMnO4 has also been used as a medical poultice for certain skin conditions. It turns water purple and stains skin and other things brown. If you can't obtain permanganate for positrodes, you can use the oxide. Oxide is the prefered initial form for a manganese negatrode.

   Battery quality "electrolytic manganese dioxide" or its (equally useful) discharged forms, Mn2O3 or MnOOH, can be salvaged for free from "carbon zinc" dry cells (use non-alkaline - alkali is corrosive to skin, and the zinc paste used can get mixed in), complete with graphite (the "carbon" of the name) added to improve its conductivity. Snip a bit of the outer zinc can of a D, F or C cell with sidecutters or something, or saw it with a hacksaw, and then take needlenose pliers, grab the flap, twist, and peel it open like a sardine can.
   For your own edification, check the resistance of the compacted powder with an ohm meter, and note its density and consistency. If your electrodes turn out as well compacted, you're doing very well.
   Remove extra stuff and bits from the manganese/graphite mix and dump the compacted powder from some cells into a jar, and fill the jar with water. (This will also show you what well compacted electrode powder is like.) In a few hours or a day, drain the water and put in fresh water. About 3 dilutions with plenty of water should pretty much eliminate the ammonium chloride electrolyte and any dissolved zinc chloride.
   Around 35% of the weight of the salvaged manganese/graphite mix is graphite; 65% is manganese. This should be taken into consideration when measuring out ingredients. The graphite is less dense, and actually accounts for about 55% of the volume of the mix.

   Manganese dioxide can also be purchased from a pottery or ceramics supply store. I'm not confident of the purity, so this is the least preferred source.

   For a negatrode charged form, manganese metal powder is available from www.micronmetals.com / Atlantic Equipment Engineers. However, if it's not well mixed with anitmony sulfate, it's likely to discharge to Mn(OH)2 or MnO anyway when it gets damp. Not only considering the expense of the metal powder but the particle size (the oxide is doubtless finer), and the likelihood of it becoming an oxide anyway, it's probably best just to use the dry cell or pottery store manganese.

This is not an exhaustive list. It's just where I got my stuff plus a few other sources I know of.
[applicable to Victoria BC]

Ceramics Supply Store [Victoria Clay Arts; gets stuff from Seattle Pottery Supply] - good source of many powdered chemicals: mostly oxides, carbonates, sulfates.
cobalt tetraoxide (a.k.a. cobalt oxide) - small qty (1/4#, 125 g)
cobalt carbonate - small qty
potassium dichromate (a.k.a. potassium bichromate) - small qty
nickel oxide

Arts Supply Store [Opus Framing]
Arches Watercolour paper 90#, electrode separator paper
expanded copper grill (used for modelling)

Plastics Supply Store [Industrial Plastics & Paints]
ABS plastic sheets, 1/4 inch thickness to make battery case from
    (Prefer white for "+" end, black for "-" end. Choice of Bk or Wh for middle parts.)
Methylene Chloride solvent to melt/glue the ABS
Methyl Ethyl Ketone (solvent): a small amount is used in the electrolyte.
    (it also seems to glue ABS - I used some by accident,
      so maybe you could skip the methylene chloride?)
Syringe
methylene chloride dispenser (has a very fine syringe-like metal tube. Careful not to get any plastic in it or it'll clog.)

 Hardware Store [Baywest Rona, Capital Iron, Canadian Tire...]
Brass bolts, 5/16" x 1.5" hex or flat head for battery terminal posts.
Brass or stainless steel nuts and washers.
Hydrochloric Acid (a.k.a. "Muriatic" acid)
Hygrometer

HEFA Rare Earths (Richmond BC Canada)
Lanthanum (2 Kg Ingots)
Zirconium Oxide (a.k.a. zirconia. powder, small qty or 1 Kg.)
    (Or try "yttria stabilized zirconia" - might give better high temperature performance.)

B. Equipment & Supplies
(This equipment is all supplied for the workshops)
Rubber Gloves (grocery etc)
Safety Goggles (hardware etc)
Syringe (Plastics Store - grind the sharp point off of it)
Hygrometer (Auto parts & supplies store)
Litmus Paper (Science/lab supplies store)
Paper Towels, cloths, dish towels (grocery)
Jars and Plastic Containers (grocery etc.)
Adhesive labels, marker pen

Sink, Stove, Fridge, Oven
sintering "furnace" (see instructions for making)
pressure cooker or autoclave (the La(OH)3 this is for making will be supplied premade)


C. Survey of  Simple Battery Electrode Materials

This charts a few possible rechargeable battery electrode materials. It includes only basic chemicals without getting into more complex organic chemistry such as the lanthanum perchlorate I've made that uses a couple of organic adjuncts for a positive electrode with better energy density.
Why do I put the words "negative" and "positive" in quotes? While it may have seemed an arbitrary thing when someone decided that an electron had a "negative" charge and a proton's was "positive", it's the electrons that move - especially in a battery - while protons essentially stay put. The more electrons you have, the more "negative" your charge. A deficit of electrons is a "positive" charge. If it worked that way with other things, it would be really simple to keep, for example, a positive bank balance: just spend more money to prevent accumulating a negative balance and stay in the positive! Hence my objection in principle to this reversed terminology: it confuses proper comprehension of elemental relationships.

Electron Receiving ("Positive") Electrode Materials

Nickel [Oxy]hydroxide

This is the "Ni" of Ni-Cd, Ni-Fe, Ni-MH, Ni-Zn, etc. The energy density isn't very good and it limits the energy density of the whole family of rechargeable alkaline batteries, but there aren't many simple things known that work very well over a lot of cycles for a positive electrode and this does.

Electron Emitting ("Negative") Electrode Materials

Nickel [Hydroxide]

Any metal hydroxide that conducts electrons, can be reduced to the metal in an alkaline environment, and with a potential of under about 2 volts, can be used as an electron emitting ("negative") electrode material, and the same stuff that oxidizes to oxyhydroxide will also reduce to the metal.

Chemical            Voltage    AH/Kg    WH/Kg    Density*
NiOOH <-> Ni(OH)2    +0.52        289        150        1.3 - 2.2

Ni(OH)2 <-> Ni         -0.72        578        416        1.3 - 2.2
Fe(OH)3 <-> Fe         -0.9
Cd(OH)2 <-> Cd
Metal <-> Metal-Hydride     -0.83        
Zn(OH)2 <-> Zn         -1.24        
Ca(OH)2.Zn(OH)2 <-> ?     -1.69

* The density of powders in pastes depends on several things, especially how hard the particles are crammed together. The higher the density, the better they conduct electrons, so higher density is best unless it doesn't provide room for any expansion that may be necessary. The most densely packed beta nickel hydroxide can cause trouble as some changes into alpha nickel hydroxide. This crystalline form is said to occupy 44% more space than the beta form. Big Edison cells have no problem with it, but a tightly packed dry cell can burst and leak.

Any metal hydroxide that conducts electrons, can be reduced to the metal in an alkaline e

Incidental Chemical Processes



Making Potassium Permanganate:
High Temperature Method

It seems it requires a license, for some reason, from "Health Canada" to purchase potassium permanganate, KMnO4. I'm not sure why this should be, as it doesn't look especially toxic according to MSDS sheets, and it's unrestricted in the USA. Instead we have to produce it ourselves, using a hazardous chemical and process.

Required are:
    1. Manganese Dioxide, MnO2
    2. Potassium Hydroxide, KOH
    3. An oven capable of 180-200ºc (360-390ºf).
        I generally suggest a toaster oven plugged in outside if available,
        though in this reaction no hazardous fumes are generated.

1. Mix a 50% w/w solution of KOH and water. (FACE SHIELD, RUBBER GLOVES, cover bare skin!) Mix it in or over a sink. Add the KOH slowly to the water, a bit at a time, as dissolving it generates considerable heat. If you add it too quickly or all at once the container will become too hot to hold. A plastic jar may melt and a glass one could perhaps crack.

2. The MnO2 powder should be in a ratio of two molar to one with the KOH (assuming pure MnO2.) (Figure out how much of each to use!) Essentially, we want to combine all the KOH and the MnO2 we put together.

3. The mixture should become K3MnO4, potassium hypomanganate.

4. Heat it in the oven at around 200ºc until all the water boils off. This turns it into K2MnO4, potassium manganate.

Now we're done. The astute reader will note this still isn't potassium PERmanganate.

Fortunately, the final part of the process is electrolysis, and for us it happens automatically: When the K2MnO4 is mixed with the electrode material, put in the battery and the battery is charged, the manganate, K2MnO4, will "automatically" become permanganate, KMnO4. (Hmm... but what happens to the other "K"?)

The existing processes for manufacturing potassium permanganate (both so-called roasting processes and liquid-phase processes) are significantly limited in their operability by problems of corrosion and wear of equipment, protracted residence time of materials being processed (slow rate of conversion), limited versatility in the use of ores which vary in their richness in MnO.sub.2 and in the level and type of impurities and, specifically in the case of liquid-phase processes, the need to operate with a large excess of potassium hydroxide (mole ratio MnO.sub.2 /KOH equal to or less than 1/10).

In general, the process of conversion of the manganese ore to potassium permanganate consists of a series of stages, comprising:

The preparation of the ore (drying and grinding).

Mixing with potassium hydroxide (KOH) in precise proportions.

The attack and disintegration of the ore through the effect of the potassium hydroxide under suitable conditions of concentration and temperature.

The oxidation of the disintegrated ore in an oxidizing atmosphere to the valency Mn6+, in the form of K2MnO4.

The dissolution in water of the potassium manganate (K.sub.2 MnO.sub.4) obtained and the separation of impurities originating from the ore.

The oxidation of manganate (Mn.sup.6+) to permanganate (Mn.sup.7+) via electrolysis.

The crystallization of the potassium permanganate obtained and the separation and drying of the crystals.






Making Potassium Permanganate:
Lower Temperature Electrolysis Method

1. Prepare an electrolytic tank. Use an anode plate of: pure nickel metal or stainless steel or iron or monel (eg); and an iron plate as a cathode.

2. Mix manganese dioxide and KOH powders. (FACE SHIELD for KOH, RUBBER GLOVES, AVOID BREATHING DUSTS) Put it in the tank. (figure out the quantities! Or do you just need "plenty" of KOH?)

3. Add HOH until the concentration OF KOH is about 15 to 25% by weight. (FACE SHIELD, RUBBER GLOVES, (figure out the quantities!)

4. Heat the tank/water to higher than 60° c, preferably about 80-90ºc. Eg, put it on a stove burner and simmer it very gently. Take all precautions with the KOH, especially FACE SHIELD! Avoid splashes and popping bubbles!

5. Turn on the power (power adapter? battery charger?). The manganese dioxide is electrolytically oxidized in the slurry of KOH using a direct current of 50 to 500 amps/m2 of anode and a current concentration of 3 to 30 amps/liter. The electrolytic conditions may be varied to suit.

6. [I presume] the KMnO4 precipitates out or rises to the surface. (A certain amount will stay dissolved.) This probably means any remaining KOH is irrelevant (except insofar as it's still dangerously corrosive). Pour it down the drain and flush some water down after it to dilute it.

"It is astonishing and unexpected that when the tetra-valent manganese oxide is electrolytically oxidized in a caustic alkali slurry having a concentration of 10 to 25% by weight at a temperature of higher than 60° C, the manganese oxide is directly oxidized into the alkali permanganate."

Tables

Energies of Common Negatrode Materials ( /Kg of metal element only)

Element
Voltage (in alkali)
Amp-Hours / Kg
Watt-Hours / Kg
Cd
0.82
477
391
H (as in MH)
0.83
500-1000
830 (max)
Fe
0.93
961
894
Zn
1.24
820
1017
Mn
1.56
976
1523

Notes:
Cd - toxic, grows crystals that short out the cells giving poor cycle life.
MH - metal hydrides store a lot of hydrogen at low pressure - the voltage is the same as for 'H' alone. The batteries work very well with long to exceptional life.
Fe - grows into larger clumps, reducing surface area, hence it has less available energy density than the figure indicates. Otherwise, the chemistry works very well with exceptional cycle life. Mixing with cadmium reduces clumping.
Zn - grows crystal "tentacles" (dendrites) that short out the cells giving poor cycle life.
Mn - any long-term issues with manganese are undetermined so far, but it looks like it might well be as close to "perfect" as it gets.

Some Overvoltage Potentials

The table below, while describing acid solution and a limited selection of materials, illustrates the differing overvoltage ("overpotential") of different substances. So while the theoretical voltage of hydrogen (the reference for all other electrochemical voltages) is 0.00 volts, it takes different voltages to actually generate it depending on the material of the electrode. Although the alkaline voltages in which zinc with mercury have been used aren't shown, one can see why mercury would be an additive to a zinc electrode, which charges to -.79 volts (in acid), to raise the hydrogen overvoltage from -.77 volts to -.85 volts. The zinc would discharge itself without such an additive.

Activation overpotential for the evolution of selected gases
on various electrode materials at 25 °C in acid solution. (from wikipedia)
Material of the electrode Hydrogen Oxygen Chlorine
Platinum (platinized) −0.07 V +0.77 V +0.08 V
Palladium −0.07 V +0.93 V
Gold −0.09 V +1.02 V
Iron −0.15 V +0.75 V
Platinum (shiny) −0.16 V +0.95 V +0.10 V
Silver −0.22 V +0.91 V
Nickel −0.28 V +0.56 V
Graphite −0.62 V +0.95 V +0.12 V
Lead −0.71 V +0.81 V
Zinc −0.77 V

Mercury −0.85 V






















MISC NOTES

Mn Negatrodes

   I looked at and became very excited by manganese as a highest energy potential negatrode, and it appeared that sealed Ni-Mn, like sealed Ni-Fe cells (first made experimentally in 2003/2004), could be maintenance free and last virtually indefinitely. Wow! The chief difference between -Fe and -Mn seems to be the metal to metal-hydroxide reaction energy, manganese being an extra 2/3 of a volt. This would make for cells of over two volts instead of 1.2.
   But no one had previously been able to use manganese in this capacity. To enable the higher voltage manganese reactions to work, I employed a previously unused 1964 discovery that organic amines and especially egg albumin, even in minute concentrations, significantly raises hydrogen overvoltage, allowing the higher voltage negatrode to charge.
   As it turned out, it might work in alkaline cells, but the voltage was still a bit too high in salt, The cells would charge to 2.3 volts, but they rapidly discharged themselves.
   It may be possible to make cells with manganese negatrodes, but in spite of eggwhite, I haven't found the technique to do it.

Diesel Kleen/Hexadecane

   The bottle labelled "Diesel Kleen" contains not only hexadecane but unspecified "petroleum distillates" (from MSDS... probably Ethylbenzene, Naphthalene, Xylene and 1,2,4-Trimethylbenzene) and "slick diesel", in unspecified proportions. Whatever it exactly is, it seems to work great!
   I believe it's forming random lamellae of graphite (graphine?) and electrode active substance(s), a random conductive network.

Lemon Fresh Sunlight Dish Soap

Contains: sodium lauryl sulfate,

Osmium-acetaldehyde separator sheet doping

   Osmium by itself is a very good metal hydride, able to store hydrogen ions (AKA protons) quite densely. It is however far too costly to consider making into an electrode. But by lightly doping the thin cellophane separator sheet with it, it divides the cell into two halves: a vanadium-metal hydride half and a metal hydride-nickel half. The protons pass easily through the sepator sheet, increasing current capacity.
   The acetaldehyde is the means of chelating, 'gluing', the osmium particles into place in the sheet.



   To prevent the osmium from oxidizing, and the cellophane from disintegrating, the thin film osmium doped cellophane is best placed between the negatrode and the paper separator sheet, contacting (if anywhere) with the negative voltage rather than the positive.




3. The Negative Electrode (OLDE)

Chemicals

The chemicals in the negative electrode are:
monel powder (Ni:Cu ~60%:40%)
graphite powder
lanthanum hydroxide powder (Neodymium oxide/hydroxide is probably better, and Samarium oxide/hydroxide is probably best.)
cobalt oxide
thiamin mononitrate (tinned bean sauce)
Sunlight yellow dishsoap
eggwhite

Monel is an alloy of nickel and copper. The fine monel powder should be purchased as such, though it can be tediously ground from solid monel with a fine sanding belt or fine (#120 or finer) grinding wheel, the dust being collected. It can be separated from the grinder grit with a supermagnet: monel is slightly magnetic. The only supplier I've found to date is www.micronmetals.com .

I bought two 1Kg lanthanum ingots at HEFA Rare Earths (Richmond BC, www.baotou-rareearth.com). To turn that into lanthanum hydroxide powder, I cut it into slices with an angle grinder cutting wheel and put it in a pressure cooker of hot water on the stove. Occasionally I opened the pressure cooker and removed and dried the white precipitate, adding water as necessary. It took weeks. An autoclave to get a hotter temperature would speed the process. Slicing the ingots made some fiery arcing tails behind the grinder: Lanthanum is what gives flint its spark, and its fiery oxidation gives a demonstration of the energy potential hidden within!

I prefer to buy some chemicals from ceramics supply stores than from chemical companies as it's local and the cost is much less. Cobalt carbonate ("cobalt blue" pottery glaze, powder) is available at ceramics supply stores but not cobalt chloride. The carbonate may be converted to chloride simply by placing it in hydrochloric acid. (Cobalt is poisonous, acid is acid: rubber gloves, goggles, at a sink!) Hydrochloric acid is available at hardware or building supply stores under the confusing code name "muriatic acid". It fizzes until the carbonate (or the acid) is used up. Refrigerate and pour the acid off the from the precipitated cobalt chloride. Rinse the sink - acid is hard on stainless steel. Use jars with plastic lids. LABEL EVERYTHING. If you don't you will soon start running across things that you can't remember what they are, or which jar is which.

After it's "dry", it's a hydrated crimson powder,  CoCl2:6H2O. Get rid of the latter (or at least four of them) by heating the powder in an oven above about 70ºC for an hour or two. It will then be a blue powder. Put a lid on it or it will soon reabsorb 4 H2O's back out of the air.

Sintering Procedure

The monel and lanthanum hydroxide are mixed in a ratio of about 3 to 1 by volume. About 1% cobalt chloride is added. To the powder is added enough "glue", tinned beans in their sauce, to make it all into a paste. This is rolled (rolling pin - grocery store etc.) into thin (1-2mm) sheets on thin pieces of aluminum sheet (eg, pieces of roofing flashing - building supplies store) and dried.

The sheets are then placed in an oven and heated until the bean sauce catches fire and burns off. Do it outside - the smoke is thick and unpleasant. The whole procedure should take less than about 90 seconds. Sometimes you don't see the blue flame, more often you do. Longer than a couple of minutes will oxidize the chemicals too much. When it cools, the crusty powder is scraped into a container. Dumping in powders with an 80cc scoop, I find I usually have around 8 sheets to burn in a "session".

The "oven" I use uses a 1500 watt electric barbecue (Value Village - doubtless available new somewhere, though I confess I've never seen another one) built up with pottery kiln bricks (Victoria Clay Arts) sliced into shapes with an abrasive cutting wheel on the radial arm saw (hardware store - kiln bricks cut like butter) to form a floor under the burner, walls, and removable roof pieces. Leave some air gaps for ventilation.

Fabrication

The powder is mixed with just enough "Lemon Fresh Sunlight" dishsoap (grocery) to make a paste. Acetaldehyde (a.k.a. ethanal, ethyl aldehyde, aldehyde, "the hangover chemical", ALcohol DEHYDrogenated) is poured in. Add a little HCl (Eye Protection, Rubber Gloves) and churn vigorously to turn the aldehyde into acetal ester. This is then dished into the battery.

Making the Acetaldehyde

Putting the electrode together is easy enough, but where do you get this "aldehyde" stuff? Turns out it's not very stable, so it doesn't keep very well and it's considered a hazardous substance to ship. (What, just because it can explode?!?) It's better to make your own. That doesn't mean it's especially easy or cheap to do so. This is the real "chemistry lab" part of making the battery.

Put 20 parts (by volume) of Alberta "triple distilled" Vodka (40% pure ethanol dissolved in 60% water by any other name) to 1 part potassium chlorochromate into a small jar and heat it in a shallow pot of water on a stove burner, stirring constantly. Once it's heated, screw the lid on - acetaldehyde boils at about room temperature. (rubber gloves, goggles: KClCrO3 is acidic and poisonous... and acetaldehyde is the hangover chemical.) Shake the jar as it heats and continue until the KClCrO3 dissolves. The end product should have a strong "fruity" smell. Label it carefully and put it in a fridge. (Preferably not a food fridge. Take whatever precautions are necessary to ensure it's clear that it isn't food and to make it unavailable to young children.)

If you could just drink the vodka and somehow extract the acetaldehyde from your body before it gave you a hangover, you wouldn't need the nasty hexavalent chromium stuff. As it is, the end product is unavailable. You feel lousy, and the alcohol in vodka costs more by weight than lanthanum "rare earth" metal - a sobering thought!

Making the Potassium Chlorochromate

Now we've got the unavailable "ethanal" but to make it we needed in its place another unavailable chemical, KClCrO3.

Web info says to use pyridinium chlorochromate, but that's pricey and also is considered hazardous to ship. But don't bother to look up "pyridinium" in the periodic table of the elements. It's not there. The important thing is the hexavalent chromium, so we skip the imaginary pyridinium and use potassium chlorochromate.

Nobody has potassium chlorochromate. (AFAIK) So again we have to make it. The ceramics supply store to the rescue! They have potassium dichromate (or -bichromate, K2Cr2O7, also very poisonous and acidic), bright orange pottery glaze crystals. (I was told that the last person to buy some was a high school student making rocket fuel a few years previously, proving that ceramics supplies are the right place to get the cool chemicals. He was planning to hit the stratosphere with his next rocket.) Now, where do we get the chlorine, and how do we put it in? Turns out hydrochloric acid, HCl, has chlorine and works well. Since the dichromate is losing some oxygen and stealing the chlorine from the acid, the obvious byproduct with the hydrogen from the acid is H2O, so according to my reckoning we're left with the desired product and water. Of course, some acid will be left over... or potassium dichromate if there wasn't enough HCl. (Actually, one could probably calculate the correct proportions. According to the hygrometer, the HCl from Rona is about 33% strength, and the K2Cr2O7 is 100%. Figure out what atomic weights of the active ingredients per gram that gives.)

It's made just about the same way as the aldehyde: Pour a little HCl (RUBBER GLOVES, GOGGLES, do it in the sink; avert your nose!) into a jar and add K2Cr2O7 crystals. Heat it, stirring gently, in a shallow pot of water on the stove until the solids dissolve. This takes quite hot water. If the crystals don't completely dissolve, that probably means the acid is used up and there'll be some dichromate left over. Add more acid. At some point in the process, put the lid on. Then label it and put it in a safe fridge. The KClCrO3 will precipitate out and the acidy water can be gently drained off and flushed down the sink. Label it, let it dry out, and put it in your secure chemical cupboard.

Making the Acetal Ester

This is made from the acetaldehyde in the presence of the negative electrode mix, which include lanthanum hydroxide. Pour some acetaldehyde into the mix. Then pour in a little HCl (RUBBER GLOVES, EYE PROTECTION). Some lanthanum chloride is formed, and this substance is called "a mild lewis acid, able to convert aldehydes to acetals in a neutral [non acidic] environment". Herein lies the unique advantage of using lanthanum over other possible metals with similar valences.

Making Monel-Lanthanum Aspic

Okay, having thrown all these ingredients into the negative electrode, there's one more thing to do: gel it. The positive electrode became a paste with the addition of the Sunlight, but that doesn't work for the negative mixture. Agar agar [a.k.a. "agar"; natural foods stores] is a white powder or flakes made from some kinds of red seaweed. It's used for a neutral microscope slide culture gel base... and for aspics and jellies that are stable at and above room temperature. This second use is of interest. The battery won't charge if the lanthanum hydroxide molecules are free to migrate. A cheesy gel holds them in place.

Take a dish. Put in ____ cc of the negative electrode mix with the acetal ester solvent in it, and add ____ grams of agar. Stir in ____ cc of water. (A lot of experimentation is due here to determine optimum proportions.) Stir or whip vigorously for a few minutes to dissolve the agar. Heat a shallow pot of water to about boiling on the stove in put the dish in it. Stir. Once it's hot, pour it into a flat tray about 6mm (1/4") deep and put it in the chemical fridge. Let it cool and gel.

Cut it into squares that fit in the battery and use a spatula to put them in. One possibility, especially for larger cell sizes, is to use trays the size of the electrodes, eg, 6" x 12", and put a cell wall sheet in the bottom. Then the gel is already on the cell wall plate, which will be easy to remove intact from the tray. Or, make the trays from the wall plates with ABS or other insulating edges, ready to pull out and stick into the battery without removing the gel at all.