how to make economical batery

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How to Make Economical,

Green, H igh-Energy Batter ies

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", orperhaps mentions that "it doesn't work", or simply omits further reference to an earlier

idea that didn't work.

==>Check the catalog atTurquoiseEnergy.comwebsite for planned availability of

custom battery making tools and parts such as electrode compactors, plastic battery

cases, current collector screens, and more.

==>Check editions ofTurquoiseEnergy.com/news/later than the date of this

document for newer information and progress.

Contents

1. Foreward and Backward

2. Electrochemistry OverviewThe water-based battery cell environment
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pH: acidity and alkalinity

Battery Electrochemistry

Electrode Substances

-Nickel

- Vanadium

- Perchlorate- Manganese

- Zinc

- Current Collectors

3. Battery Construction OverviewElectrodes Overview

Battery Layout(s)

Chosen Layout

Electrode Binder "glue"

Separators and Capacitors

4. Making the Case and FittingsCase

Electrode Current Collector Grills & Terminal Leeds

5. Making Perforated Plastic Pocket Electrode EnclosuresPerforating the plastic

Forming the square cylinderEnd caps

'Glue'/solvent

6. Making the Positrode6.a Permanganate/Nickel Hydroxide Positrode

6.b Monel Positrode

6.c Vanadium Pentoxide Positrode

7. Making the Negatrode

7.a Zinc Negatrode7.b Manganese Negatrode

8. The Electrode Separators

9. Electrolyte and Cell Assembly


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10. Charging, "Forming" and TestingInitial Rest Period

Initial charge

Initial cycling

Testing Specs

11. AppendicesA. Creating Unusual Substances

B. Materials and Chemicals Supply Sources

C. Equipment & Supplies

D. Survey of Some Battery Electrode Materials

1. Foreward and BackwardIn 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 forportable electronic devices and cordless tools, but the employeesare just workers. While the theory of operation and the chemicalreactions aren't hard to undersatnd, there are a few importantdetails needed for successful construction that aren't mentionedanywhere in particular, certainly not all in one place, and very fewpeople know anything practical about battery design andconstruction.

A great need has long existed for long lived, economical, highenergy batteries for electric transport and off-grid power. I decidedto try my hand at creating some way to make some sort of batteriesat home.

I soon felt sure that some better chemistries, probably muchbetter, than existing types could be created, and potentially for

lead-acid or throw away dry cell prices, or not so much more. Thisbook describes known and newly invented alkaline batterychemistries, and no less importantly, a design for DIY buildablebatteries of any size, that I've come up with in a project spanning -as I write - over four years.


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Battery research and commercialization have been sidelined byhuman propensity to "go with the flow", to limit thoughts intonarrow structured channels, good or (more often) inferior, and toextend that channel to the exclusion of wider possibilities, including

superior ones. Thus for example, with large, higher-energy alkalinebatteries 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 betterlithium batteries, despite the cost and the complex problems ofmaking lithium work well, and despite the fact that since patents onthe best developments are acquired to suppress each developmentas it emerges to market stage, their work will dead-end the sameway Ovshinsky's excellent nickel-metal hydride electric car batteriesdid. We trust this state of affairs won't continue for a second 100

years, but in the meantime, DIY battery making provides the rest ofus a way to take matters into our own hands.

Making 'normal' water based batteries is a rather involved butfascinating "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 yourbase 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, orshould the reader be given all the gory details, the reasons andreasoning behind the instructions? Knowledge is power! I'm tellingall that I can think of to say. But I'm organizing it into varioussections so the reader can read as much or as little as desired - thebasic instructions, a good theoretical overview, or complete detail.

In other material, even the most basic information is lacking forneutral pH salt solution cells. For example, why is the positive

electrode in a standard dry cell a conductive carbon rod instead ofmetal as in all other batteries? You'll dig long and deep and not findthe simple answer: that every common metal will corrode away inthe positive electrode in salty electrolyte - including nickel, whichsits inert in and enables all the various KOH saturated alkaline cells.Only carbon or graphite works. (Note: nickel manganate+epoxy mixmight work) Obviously battery makers know this (or once did), but


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it took me over two years of corroded electrodes in every test cell tofigure it out for myself, because no one mentions it anywhere. (Iput it on Wikipedia, but it appears to have been erased.)

Much of the info herein has been acquired gradually, and oftenpainfully, in my battery research over the past 4 years. A tidbit ofbasic info is casually mentioned in one publication or another, mostof which assume the reader is well versed in the battery makingarts - and few people are.

For example, it was only after 2-1/2 years that I finally saw forthe first time an actual figure for the amount of pressure used tocompact a battery electrode into a "briquette" - for one type ofelectrode in one experiment. When I started, I wasn't even aware ofthe vital role of compaction, and after eventually deducing it

indirectly from some material density specs, it took a another yearto figure out a simple way to get enough pressure.

Likewise, it wasn't until February 2012 and four years ofmysterious self-discharge problems that I understood that the wiresin the negative electrode had to have as high a hydrogenovervoltage as the electrode substance itself. Anything goes foriron, cadmium or hydride, but few common things work with ahigher voltage chemical - zinc or manganese. It has to be zinc orzinc alloy wire. (or silver.)

My original minimum battery goal was to copy proven andrelatively economical NiMH EV battery chemistry, by the simplesttechniques I could find or work out, and thus create a "DIY" meansof making batteries. But I also started to think that coming into thefield as a newcomer without formal training in the field as to "that'show it is", I might, in stumbling around, uncover overlookedinformation 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 asthe inventor to be free technology, there would be no patentrestrictions on it for vested interests to kill commercialization with.(And patents aside, it probably would have been very difficult tomake 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


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many broad basics and fine details in no particular sequence. And Idid uncover a few key overlooked things.

I also developed useful "DIY" battery construction tools andtechniques, such as a bolt-down electrode compactor, and

perforating rigid plastic sheets with a heavy sewing machine tomake solid "pocket" electrodes. Finally I have been rathersuccessful: nickel-manganese and similar batteries are in principleeconomical, "green", and superior to what's on the market today,including being quite economical and having about the highestfeasible energy density, perhaps on a par with lithium ion types. Ipicked the reacting substances out of a considerable number ofpossibilities because they seem to be the best. The fact that theyare 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% graphitepowder, and "1% Sb2S3is 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 whilethere is a convention that "anode" refers to the negative electrode(while it is the positive terminal of a diode or a non-rechargeablebattery), this is not universally adhered to, and there is oftenconfusion about what is meant - I often get mixed up myself. Aselectrodes are ubiquitous to the subject and a specific one is sooften referred to, herein I will call them "positrode" and"negatrode", which terms should be self explanatory. I also insist onspelling terminal wires as "leeds" to differentiate connections and

wires from the metal "lead", the guy "in the lead", and at least acouple of other uses of the same four letter sequence, hoping not to"lead" anyone astray.

2. Electrochemistry Overview


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The physical design and construction is more important to making

a battery that works than the electrochemistry. But theelectrochemisty is the premiere part, the fascinating part, so it gets

the first chapter.

I've tried to explain less common, specifically electrochemicalterms herein, but the reader will understand the text better if hestill remembers his high school chemistry. If you don't know whatan "ion" or a "sulfate" are, just look them up on Wikipedia. Ifanyone 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 at0.0 volts or anything negative: this is the reference voltage againstwhich all other reactions are measured. Whether a substance can beused inside a rechargeable cell depends on it charging below thevoltage where gas is produced instead.

Gas generation is more and more likely with increasing voltageabove 1.23 volts, but the exact voltage varies with electrodesubstance and additives, temperature, and pH. Any amount over

the theoretical gassing limit, at which gas isn't generated, is calledthe "overvoltage".

In acid, gas generation voltages shift to inhibit oxygen generationand hydrogen generation occurs more easily. Eg, a lead-acid batteryallows the lead oxide to lead sulfate reaction to work at +1.7 volts.The lead dioxide would spontaneously discharge itself at thatvoltage in salt or alkaline solution. However, the lead metal tosulfate reaction is also just under the limit at -.35 volts.

On the other hand, in alkali, oxygen gas generation is encouragedand hydrogen more inhibited. The common alkaline nickel positrode(+.5 volts) is just below the "oxygen overvoltage" at roomtemperature, and zinc just works at -1.24 volts. The 0.0 volts inacid hydrogen voltage, in alkali is -.833 volts. The inverse of thisvoltage plus the +.49 volts of nickel gives us a theoretical opencircuit voltage of the nickel-metal hydride alkaline battery, 1.32volts.


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Oxygen overvoltage falls a bit with temperature, and above 40Csimple nickel electrodes won't charge properly.

The electrode substance is also significant, and small amount of ahigh overvoltage potential substance as an additive can increase the

overvoltage so that the main substance works better, or works athigher temperatures. To improve zinc's performance in alkalinesolution (-1.24 volts), the traditional additive was 2.5-4% mercuryoxide. Later, owing to mercury's toxicity, transition metals (gallium,indium, tin and bismuth) or their oxides were tried and found towork well even in amounts under .5%. In an Indian experiment withsealed Ni-Fe alkaline cells, .5% bismuth sulfide (Bi2S3) was used toreduce the hydrogen bubbling in the iron negatrode. Heavytransition metals such as antimony are also used to improve lead-acid cell charge performance.

In the case of manganese as a negatrode, adding 1% antimonysulfide raises the hydrogen overvoltage above manganese'scharging voltage. This is the only reason it works at all. Without it,the overvoltage seems to be right on the edge: the manganese mayor may not charge, but it bubbles hydrogen as it does and graduallydischarges itself to hydroxide, bubbling hydrogen. Thus manganesehas never been used before as a negatrode. Its higher reactionvoltage, 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 ofabout 1.7 nominal volts. In fact, NiMn alkaline cells may lastindefinitely.)

There are lots of even higher voltage reactions that it's hard toconceive of making work with any additive, such as aluminum toaluminum hydroxide at -2.3 volts in alkali. That surely will never beenticed 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 sealedcells. Pressure would just build up until the cell burst. So sealedalkaline batteries are made with the negatrodes larger than thepositrodes. The positrodes bubble oxygen first, and the cells arealso made as dry cells with empty spaces that gas can passthrough. The oxygen migrates to the negatrode, discharges some of


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the substance (making heat), and prevents complete charging ofthe negatrode. This gets rid of the oxygen, and prevents thenegatrode from bubbling hydrogen gas, preventing mildovercharging from bursting the cell.

Vented cells (a) dry out and need refilling, and (b) absorb carbondioxide from the air, which may gradually degrade substanceswithin, turning them from active chemicals into carbonates. Variouscaps and valves can minimize the problems and vented cells aren'timpractical, but they're second best to sealed.

To make sealed cells bigger than dry cells, some means to keepgas pressure low has to be found. Recent work with catalysts torecombine O2 and H2 into water has been successful, but I haven't

explored it at this point. I've also read that antimony is almostunique in its ability to react with small molecules - like hydrogen -and I picked it as an electrode material additive hopefully as arecombinant 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 liquidelectrolyte, sealed cells are very dangerous, since a spray ofpostassium hydroxide out a leak can blind. "Blindness is for life"...

one cell almost got me - only takes one - and I've met a blindchemistry professor. A vented case, and using potassium salt forelectrolyte, reduces the dangers.

I hated the thought of using potassium hydroxide or acidelectrolytes. They're dangerous! I was using a salt based electrolyteof neutral pH, potassium chloride. (KCl) It's a fast electrolyte(allowing high current flow), and less hazardous to handle thanpotassium 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 novelimprovements begging to be made. If one could find a chemicallyinert but electrically conductive or even semiconductive binder 'glue'to hold the electrode powders together, it could permit highercurrent flow than the usual insulating binders, and intense


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compacting of the electrodes would be less critical to obtaining goodcurrent capacity... If a small, economical, high energy battery couldsupply enough current to start a car engine, that would be amarvel!

I'm just now experimenting with nickel manganate, a highlyconductive semiconductor, for the nickel electrode, but have noresults or conclusions yet.

Battery Electrochemistry

First I'd like to point out a misleading quirk of terminology. Backin the beginning of understanding atomic particles, someonedecided 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 twowords have other, well known meanings. But they have beenapplied backwards.

Consider that protons are stationary, within atoms, while freeelectrons 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 moneyyou spend, the higher your account balance; the more you earn, thehigher your debt. The negatrode deposits electrons during charging

and then supplies them to a load, while the positrode is "short" ofthem when charged and soaks them up on discharge. This is allcounterintuitive, and in some situations, a hindrance to figuring outwhat'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 thatinvolve loss and gain of electrons, which on this planet are

frequently but not always related to oxygen reactions. (Rememberthe obnoxious "OIL RIG" - Oxidation Involves Loss, ReductionInvolves Gain [of electrons].) Pushing electrons around is whatbatteries are all about. (Hmm, "Reduction is gain!" -- another lovelylittle paradox of nomenclature!) A shorthand used for reduction andoxidation is "redox", and battery reactions are redox reactions.

The electrochemical reactions at each electrode are called "half


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reactions", and the two half reactions of a battery must balanceeach other. If the negative terminal supplies "x gazillion" electronsto an external circuit, the positive terminal must soak up "xgazillion" electrons. And, the ions released internally by one

electrode must complement those released by the other or beabsorbed into it. After all, no atoms are being added to or removedfrom a battery in use.

The chemicals used in a battery are chosen both forcomplementing ions and such that the positive side is a chemicalthat gives energy when reducing while the negative chemical is onethat 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 formwhen 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, andthe positrode lead dioxide reduces to lead sulfate, the sulfate ionsbeing stored as excess acid (or sodium bisulfate) in the electrolytewhen the battery is charged, and absorbed as it's discharged. Moreexamples appear below.

Usually these positrode oxide forms aren't very good electricalconductors. Some oxides, like titanium and zirconium, are virtuallyinsulators, so they can't convert easily between forms by electricalaction as battery elements. Often additives are used to improve theconductivity of the oxides. Zinc and cobalt oxides have been used tomake nickel hydroxide electrodes conductive enough to use, ashave nickel powders and flakes, and powdered graphite.

The number of amp hours depends on how many electrons thesubstance will release or absorb during oxidation or reduction, andthe energy of each reaction is indicated by its voltage. A substancewhich naturally wants to oxidize (in the battery environment) willhave a more negative reaction voltage than one that wants toreduce. The energy in watts-hours is the amp-hours (the number ofelectrons) times the voltage (the pressure behind each electron).


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The voltage of both electrodes is subtracted for the total batteryvoltage, eg +.5 - -.93 = 1.43 volts for a nickel-iron alkaline battery.The amp-hours or number of electrons isnt additive: it shouldmatch. The current flow stops and the cell is discharged when either

electrode has been depleted to its un-energetic state and will pumpno 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-hoursper kilogram will be available, because there are more molecules toreact in that kilogram. Oxygen and hydrogen are quite light, so themetal is usually the dominating factor. If a heavier element ischosen, it must move more electrons per reaction, or have a higherreaction voltage, to provide equal energy density. If the advantagesare 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 searcheswere mainly for lighter atom metals.

Lightness of metallic substance is pursued to the ultimate inlithium battery types. But lithium has to be used in thin filmelectrodes, often with non-aqueous electrolyte, and the substratesto 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 orturn into a gas). This greatly limits the choices. Most chlorides aresoluble, so the electrodes of a battery using hydrochloric acid woulddissolve and thus would be hard to recharge. The old 'standard'non-rechargeable dry cell uses ammonium chloride electrolyte, andthe zinc electrode dissolves to zinc chloride in use. Most lighterelements dissolve in sulfuric acid, but lead, lead sulfate and leaddioxide 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 Imade a nickel-zinc test battery in oxalic acid: nickel oxide, nickeloxalate, zinc and zinc oxalate are all insoluble. Similar in concept tolead-acid, it worked and could be charged. (The voltage was lowerthan the tables indicated, about 1.4 volts. Acetic acid/acetatesshould also work.)


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Zinc has been known as a frustrating battery negatrode element.It's energy is the highest available for alkaline cells and its electricalconductivity is good, and the charge and discharge products areboth solids. However, in use there is a temporary dissolved state,

the zincate ion, in which form the zinc can and does graduallymigrate. This causes the negatrode to gradually lose capacity, andthe zinc grows dendrites, "tentacles" of zinc crystal, which usuallyshort out dry cell batteries, often after only 10-50 charge-dischargecycles. Cadmium, underneath zinc on the periodic table, has thesame problem, and Ni-Cd dry cells rarely last anywhere close totheir supposed cycle life as cadmium crystals poke through theseparator sheet and short the cell. NiZn and NiCd pocket cellbatteries fare much better. But it would seem that NiZn dry cells inrecent years have improved, as a company making AA cells

(available on Amazon.com) claims 500 to 1000 charge-dischargecycles.

It's possible that in salt electrolyte, zinc doesn't form zincate ion.Thus switching to salt might solve the problem, allowing use of thishigh energy density substance in long-life batteries. Or, thezirconium silicate ion blocker I paint on the electrode separatorsheet may solve the problem or at least provide "500 to 1000cycles".

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 thethreshold between usable and not for a negatrode. It doesn't workby itself at room temperature, bubbling hydrogen and charging atthe same time, then spontaneously discharging itself too quickly tobe practical. But with the right additive(s) to raise the hydrogenovervoltage, it might be made usable. It needs further research.

The electrolyte doesnt conduct electrons between the electrodes,it only conducts charged dissolved ions. It's the one place whereprotons are on the move. To have the oxidations and reductionstake place, both ions and electrons must flow, as will be seen in theredox (reduction-oxidation) reactions coming up.

A circuit connected to the battery lets the electrons flow betweenthe electrodes - externally. This is of course what the battery is for.


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When an external circuit is connected, the electron flow, the ion flowand the discharge reactions proceed spontaneously andsimultaneously, releasing the chemically stored energy aselectricity. The ions flow mainly by diffusion through the electrolyte,

spreading because like charges repel, and by attraction to theopposite electrode as they reach it. The current capacity of thebattery depends partly on how fast the ions diffuse through theelectrolyte. Potassium chloride salt is supposed to be very fast.

The discharging reactions release chemically stored energyelectrically. The recharging reactions require electrical energy fromthe external circuit - the battery charger. Charging restores the'spent' lower energy substances to their higher energy oxidationstates and valences.

There are many solutions and some solids that can pass ions, butthe best - fastest - solvent is a polar liquid such as water, with anacid, salt or alkali electrolyte dissolved in it. There is, however, oneserious limitation to using water as an electrolyte, as mentionedpreviously:

"The use of aqueous battery electrolytes theoretically limits thechoice of electrode reactants to those with decomposition voltagesless than that of water, 1.23 V at 25 C, although because of the

high "overvoltage" potential normally associated with thedecomposition of water, the practical limit is some 2.0 V. The liquidstate offers very good contacts with the electrodes and high ionicconductivities." Lead-acid batteries are theoretically 2.05 opencircuit volts, and many earlier cells were about 2 volts.

The voltage delivered to a load circuit is somewhat lower than theopen circuit voltage, depending on the internal resistances of thebattery relative to the amount of current flowing. Hence batteries

are given a "nominal" voltage rating which might be expected intypical 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. Heavyloads may drop the output even more, eg to 1.0 volts. If such loadsare expected, it's usually best to add more batteries in parallel toreduce the load on each one, or to use bigger cells, which iseffectively about the same thing.


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If the positrode has lesser amp-hours capacity on discharge than

the negative it is depleted first. The negatrode still could havesupplied more current and the battery is said to bepositive limited.

Vice-versa if it's the negatrode that runs out first. It may also bepositive or negative limited on charging, and not necessarily in thesame direction. It's also possible for the electrodes to be entirely offbalance - one discharged and the other charged. It could be hard toeither charge or discharge this cell.

There are often good reasons for preferring one reactant todeplete first. For example, if there's no recombination catalyst in asealed dry cell, oxygen gas is much better to generate thanhydrogen if the cell is overcharged. In a dry cell, it travels over formthe positrode to the negatrode and there discharges an atom of

metal to hydroxide, making a bit of heat. Thus the cell stopscharging - it just gets warm. Hydrogen doesn't readily discharge atthe positrode and the gas would accumulate until the cell bursts, soit's best to have the positrode charge first and not get anyhydrogen. With the catalyst, starting to generate both gasses atabout the same time when the charge is complete should beadvantageous, since they can then start recombining to make waterbefore the pressure of either gas builds up much.

Electrode Substances

Besides lead in lead-acid cells and lithium, there are two commonpositrode substances: nickel and manganese. My newfoundvanadium 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 apar 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. (Thisseems 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'stantalizingly borderline, charging at about the same voltage as thehydrogen and spontaneously discharging itself somewhat too fast to


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be practical. I found the hydrogen overvoltage can be tweaked upsufficiently by adding 1% stibnite (antimony sulfide). This makes itwork. It appears to be an ideal negatrode. It's even higher energydensity 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 thesubstance of one-use dry cells, so-called "carbon-zinc" but actuallymanganese-zinc, the carbon (as graphite or "carbon black") being infact simply a conductivity improving additive. But the zinc and theelectrolyte are the problem with recharging the old dry cell, not theMnO2. In salty solution the energy is about +.5 volts, but in alkalinesolution it's only +.15 volts, so makers of rechargeable alkalinebatteries prefer nickel oxyhydroxide, with +.5 volts. However, it has

high amp-hours per kilogram, and that allows it to complementmore high energy negative electrode, providing higher energy cellsnotwithstanding somewhat lower voltages.

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

Nickel (+0.95V) works great and makes 2 volt cells, butrechargeable cells using manganese positrodes provide the highest

energy density, and manganese is cheap - you can even scrounge itout of old dry cells for free.

Nickel

Note: Despite its lower reaction voltage, manganese dioxide is nowpreferred 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 usedin most rechargeable alkaline batteries with various negativeelectrode materials: Ni-Cd, Ni-Fe, Ni-MH and Ni-Zn. Dry and pure,it's a very fine, fluffy, turquoise green powder. The nickel willhappily stay in the hydroxide form in the battery environment. It


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thus has no usable energy. To convert it to a more energeticchemical, energy must be put into it.

To charge it, the nickel hydroxide is further oxidized to nickeloxyhydroxide by grabbing one electron from it. It doesnt willingly

give up the electron: the charger has to supply the energy to causeit 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 ratherthan 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 thereaction -- not dissolved, liquid or gas. It is usually a primerequirement that the electrode doesn't dissolve. Normally if it does,the battery won't recharge. The valence of the nickel goes from II toIII as it's charged, indicating that one electron is removed permolecule, 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 backinto hydroxide [II]. When there's some of each, the nickel valence isexpressed as a fraction. (which we will not attempt to describe withtraditional Roman numerals) When it gets below 2.25 or so, theresistance rises and the user considers the battery to be "prettymuch dead". So really, only 3/4 of an electron is moved per nickelatom, 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, andperhaps with temperature.

A major advantage of salty electrolyte is that the nickel reactionvoltage is double, about +1.05 volts,giving it double the watt-hours per kilogram of the alkaline cell. This alone was a good reasonto attempt to create working salt solution batteries.


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The nickel oxyhydroxide is an "energized" substance: it would

rather be just plain hydroxide and given a chance will revert andgive 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 thevoltage. It can get the hydrogen "H+" ions from the water, leavingOH- in the water. This is balanced with the negative electrodegrabbing the "OH-" ions, but it will only perform this reaction whenan 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)2to NiOOH, whichhas the same reaction voltage as the 2 to 4: +0.49 volts... or +0.52depending 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 inthe wet battery environment, and it takes energy to reduce it toelemental 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 batterycould be created. The valence of the nickel goes from II to 0, addingtwo electrons to each nickel atom. This charging reaction gives off


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negatively charged hydroxide ions that were bonded to the Ni(OH)2,the same as with iron, cadmium, zinc and manganese, each at itsown voltage. (Metal alloy hydride absorbs a hydrogen ion, H+, fromthe 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 timesthe theoretical energy storage. Nickel hydroxide in alkali, thoughthe most common positrode material, makes a much moreenergetic alkaline negatrode than it does a positive one! (You mightneed a hydrogen overvoltage raising additive to keep it frombubbling hydrogen - nickel evidently has a very low intrinsichydrogen overvoltage, and hence nickel electrodes are oftenemployed to generate hydrogen.)

Notwithstanding this, the voltage and energy of the reaction arelower 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)2as a "+" terminal of 289amp hours per kilogram is presumably doubled as a "-" side to 578AH/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 positrodechemistry? 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 lowerenergy density - lower amp hours - by weight, despite thesomewhat higher reaction voltages.

Manganese dioxide, while cheap, is only +.15 volts in alkalinesolution. That means more cells to attain a given voltage. In salt,it's .5 volts, and it might be a more economical solution forstationary batteries, eg for off-grid home power storage. It is easyhowever, and considered deleterious, to charge it to a higher oxide

form. (The zircon ion and-or chelation of the Mn ions shield mightalleviate 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 saltybattery is the actual active chemical, whereas in alkaline batteriesnickel or nickel plating is used for all internal metallic structures.(That could be changed with grafpoxy.)


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But the traditional basic reaction doesn't reveal nickel's full

potential. Nowadays, manganese is added to the positrode as amajor additive, perhaps 35 to 40% by weight ("wt%") of Mn to Ni.

What this supposedly does is raise the oxygen overvoltage, whichevidently allows the nickel to charge to "alpha" nickel oxyhydroxide,wherein some portion of the nickel actually charges to NiO2, valenceIV, moving two electrons instead of one. Another thought is thatpermanganate is a "powerful oxidizer", and it may be this thatallows or causes the nickel to oxidize to a higher valence. On theother hand, the two ideas may just possibly amount to the samething expressed differently.

Maximum attainable overall valence appears to be about 3.8. Theactual nickel valence thus might change from about, say, 2.25 to

3.75 from discharged to charged, thus moving 1.5 electrons pernickel atom, twice as much as with the old pure Ni(OH)2simpleformulation. This doesn't double energy density by weight becauseof the added mass of the manganese, but it does improve it, andthe nickel - the costly and main ingredient - does twice as muchwork.

Multiplying the theoretical value 289 AH/Kg * 1.5 = 433 AH/Kg.Naturally however, the theoretical maximum isn't going to beattained. (Experimentally about 350 AH/Kg has been attained, the

forms being alpha hydroxide and gamma oxyhydroxide, which bothoccupy about the same volume of space. Although it's a highervolume form than the beta forms, the constancy is very desirablefor 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 withthis 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 energydensity of the battery. For homebrew salty batteries, I'm expectingactual attainment of around 100 AH/Kg will be doing well.


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The negatrode substances being much higher energy, the energy

density of the whole cell will be mostly limited by the nickel and thevoltage 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 witha manganese negatrode, if that can be made to work, 130-190WH/Kg might be attained. These figures seem dissappointing afterreading the theoretical maximums, but they're still better thancommercial NiMH dry cells and as good as or better than lithium iontypes. And it's not impossible that with good design, chemicals,technique, workmanship and high compaction, even higher energymight be attained.

Other metal oxides or hydroxides besides manganese that have

been tried and appear to work (and may bear furtherexperimentation) include: aluminum, cobalt, yttrium, ytterbium,erbium, and gadolinium. Other rare earths hydroxides such assamarium, neodymium and even lanthanum might be better, or atleast fine, in salt solution. I'm not sure why manganese is supposedto be "especially preferred" (or even why it should work well), orindeed what the selection criteria are, but I've used Mn in mypositrodes 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 hasso many reactions at various voltages that it's confusing to try andfigure out what will actually happen in many situations, and I as faras I can see commercial battery designers often don't know exactlywhat they're doing either. Certainly inAlkaline Storage Batteries(Falk and Salkind 1969), there was a lot of speculation about someof the main chemical reactions. And battery substance reactions insalty electrolyte are relatively unexplored compared to alkaline.)

It's not clear to me at the moment whether the only effect of themanganese compound is supposed to be to raise oxygenovervoltage in the postirode. If it is, the samarium or whatever,probably in considerably lesser quantity percentage-wise, shouldreplace it entirely, providing highest energy density. (For a while Ithought the KMnO4 reacted at virtually the same voltage as thenickel and would be an active chemical along with the NiOOH, but it


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appears it's somewhat lower and thus wouldn't start to dischargeunless 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 thehydride alloy, and the plating or substance of all the metalconductors within the cell. In the salty cell, it's just the positrodechemical, so the cell should be more economical.

Neither nickel hydroxide, oxyhydroxide nor potassiumpermanganate is a very good electrical conductor. The battery'scurrent capacity would be extremely limited if these were the onlyingredients. Powdered graphite has been added for betterconductivity, as in the standard and alkaline single use dry cells.

Edison put in 80 layers per inch of alternating nickel hydroxideand ultra-thin nickel metal flakes, crammed solidly into perforatedmetal tubes about the size of a pencil. The nickel flakes were madeby electroplating alternate layers of copper and nickel ontosomething, then dissolving away the copper. That costlyarrangement was the best he could come up with that worked well.He tried graphite flakes and found the performance wasunpredictable - I think Edison didn't expect powder could be a goodconductor across an electrode, but above a critical proportion it is.

The sintered electrode is another good form for conductivity inalkali, the sintered nickel sponge connecting well across the wholeelectrode for very high current capacity. NiCd cells get some of theirhigh current ratings from this.

But I discovered that for any salty cell battery, all metals oxidizerapidly in the salty positrode. Sintered metal electrodes are out. Andgraphite powder is cheap at any art supply store.

But up to 5% cobalt hydroxide has been added to alkaline cellswith 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%) nickelhydroxide. (The monel I'm using also contains 2% Mn and 3% Fe,so the copper is 28%. Obviously Mn doesn't hurt, and the ironeither, I trust.)

On a practical note, it's worth mentioning that a nickel electrode


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can be discharged chemically to Ni(OH)2 by immersing it in a smallpool of hydrogen peroxide - the 3% drug store stuff is fine. Itmakes 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 storebleach should work fine. Again rinse out the bleach when done.

These procedures give you a way to equalize the charge if you'veended up with one charged electrode and one discharged for asealed battery. For an unsealed one, charging and letting gasbubble off one electrode works.

Nickel Manganate

In late February 2012 I found a better form of nickel forelectrodes than nickel hydroxide: nickel manganate [NiMn2O4], asynthesis of nickel and manganese. This little known substance (butnot unknown - it's used to make thermistors) is of repute for its"spinel" crystalline structure, which gives it a much lower electricalresistance than most oxides. At first I thought it might make a goodconductivity improving additive. It was far more conductive thanNi(OH)2, but nowhere near as good as graphite. Then I thought ofusing 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 manganeseions. However, one has 4 oxygen ions while the other has 8.Charging nickel manganate to nickel permanganate would release 8electrons and use up 8 OH- ions from the charging negatrode (Fourbecome the other four "O--" ions in the permanganate, the otherfour become H2O). The voltage for that reaction would be around

+.65 volts in alkaline solution. It would be fantastic energydensity... maybe too good to be true.

Then I realized the nickel would be more likely to change since itsreaction voltages are lower, probably similarly to the reactionsdiscussed for nickel hydroxide. A nickel valence 3 compound mightbe formed, for example, Ni(OH)Mn2O4, or even valence 4, eg


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Ni(O)Mn2O4. It'll probably get a bit more mileage out of the nickel -because of the high conductivity, it would probably discharge downto nickel valence 2.0, whereas nickel hydroxide pretty much stopssupplying 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. Theyworked well, and seemed to be much more conductive than simplenickel hydroxide.

The approximate formula was:

11g NiMn2O45 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 temperatureperformance or in case the higher voltage permanganate reactionapplied, but I forgot.

I couldn't find nickel manganate to buy. I tried making itchemically, but it was messy and smelly. Then I mixed appropriate

amounts of dry NiO and MnO2 powders (both from the potterysupply) in a stainless steel pot and simply heated them red hot witha propane torch (outdoors, with a respirator). This gave a lowerresistance product and was fast and pretty simple to do. I only got15 grams, so I guess the torch blew 7 or 8 grams of powder out ofthe pot. This is what made the successful electrodes above. Later Imade another batch and 'only' lost 25% of the mass.

Vanadium

I 2011 I made a battery with a vanadium electrode. It wassupposed to be the negatrode, but it didn't seem to work -unexpectedly, the vanadium seemed to become soluble and tomigrate. (This was also the first cell I'd made with transparentplexiglass sides, and I could see the vanadium pentoxide yellowcolor appearing on the other electrode.) I reversed the charges, and


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found that the cell charged to about 2.2 volts. The vanadiumpositrode side would have made up around 3/4 of that, and it seemssurprising that it didn't just bubble oxygen and spontaneouslydischarge itself to a lower oxide. It seemed to charge and discharge

well, but at the time I hadn't made the grafpoxy yet and itdeteriorated like my other cells of that period, as the graphitebacking sheet swelled and lost conductivity and good contact withthe electrode and the carbon terminal post. Judging by the voltageand 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'sunexpected behavior as a negatrode, the oxides appearedto me to

remain in solid form, not dissolve, in the salt electrolyte. Thisappears to make it a good positrode, moving one electron pervanadium atom, hopefully with good stability from the doublevanadium molecular center.

Taking the average of the acid and alkali voltages as being theapproximate salt voltage, the voltage obtained in the cell indicatesthe single valence change to V2O4 seems to apply. (average of(1.0V [acid] + 2.19V [base]) / 2 = 1.6V [salt])


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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.


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So vanadium seems to have the potential to be a very good

positrode in salt water electrolyte. Theoretical amp-hours works outto be almost identical to the theoretical 289 amp-hours/Kg of beta

nickel oxyhydroxide. The potential double valence change that isachieved by some of the nickel to alpha oxyhydroxide moleculeswouldn't seem to be possible with vanadium, but the voltage is 55%higher, raising the energy density considerably.

According totheelectrochemical table, wemight suppose

that vanadiummight alsomake a goodnegatrode inalkaliat thesame potentialas cadmium (-.82) andhydride (-

.833),providing itwasn'toverdischarged, which mightform thehigher 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 probablyindicates the problem I had using it as a negatrode in salt: insteadof forming either VO or V(OH)2 at pH 7, vanadium instead forms adissolved ion, VOH+, for the 'discharged' negatrode.


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I don't see the common V2O5 form (or VO for that matter)

anywhere in the pourbaix diagram, and the voltages don't matchthe table above, nor do they appearto jibe with my experimental

results.

Vanadium probably deserves more research in salt electrolyte foruse as a positrode (and maybe in alkali as a negatrode), but I haveno present plans for doing it myself. After I found the vanadiumPourbaix diagram, I figured there's probably soluble ionssomewhere during charge or discharge, which might make forlimited cycle life. If I'm going to stick with 'tried and true', that'snickel, 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 ofthe metal+ion it's attached to. I once tried to make a positrode oflanthanum perchlorate, La(ClO4)3, which would reduce on dischargeto 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 dissolvedlanthanum behaved about the same as undissolved, tho I can'tremember where I read of it.) In addition, perchlorate is often muchless soluble than chloride, as with only slightly soluble potassiumperchlorate versus potassium chloride salt.

As I've said, if heavier elements were used, they'd have to movemore electrons to attain the same energy density. Lanthanumperchlorate, potentially with 12 O--ions forming 24 OH-ions on

contact with water, is a super example: 24 electrons per reactionwhere nickel moves one or two. That much more than makes up forthe atomic weight of La(ClO4)3 being almost five times that ofNi(OH)2, and suggests the theoretical possibility of a much higherenergy density electrode than nickel hydroxide.

There may be reasons I'm unaware of that this can't work. I amafter all only an amateur chemist. However in the absence of any


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known reasons it deserves more research, and I hope to experimentmore with it. (Next time, I think I'll try converting lanthanumhydroxide straight to perchlorate with perchloric acid instead of tochloride 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 atpottery supply stores. The dry cell is probably the better source -it's known to be pure enough for batteries and it's "pre-mixed" withconductive graphite powder. In the open, dioxide is the usual stateof 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 cellsmake use of this. Sometimes the discharge product is given asMnOOH and sometimes as Mn2O3. It matters little as both arevalence three after moving one electron, the difference onlyaffecting the amount of water released or absorbed during chargeand 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 allliterature I've managed to find shows this reaction:

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


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Manganese Redox chart.

Another manganese reaction of great interest is on the right endof the chart, going between valence 0 and +2. If one simply usesmanganese powder in water, this reaction is just high enough involtage that it gradually but spontaneously discharges intoMn(OH)2. This has always precluded the use of manganese as a

negatrode.However, additives, in particular heavier transition metals or theircompounds, can raise the voltage at which hydrogen starts togenerate. They are used to help zinc electrodes charge better andwork at higher temperatures. Traditionally about 2.5-4% mercuryoxide was used. Now smaller amounts of less toxic transition metalsare 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 allowmanganese to charge and hold its charge.

I finally got an alkaline cell to charge Mn to metallic state andhold its charge in February, 2012. This is probably a first.


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As it charges in alkalinity pH 14, a bit of the Mn (starting asMnO2) becomes a soluble ion, probably Mn(OH)3-, or else MnO4--.(per the Pourbaix diagram below) When this soluble ion touches thepositive electrode, it's charged to KMnO4. This is indicated by the

water turning purple. The KOH electrolyte solution is normally pH14, 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 seethe answer on Wikipedia. It may even just be bad coloring of the pHtest paper.) At pH 13, the soluble ions cease to form (leaving onlythe desired insoluble Mn(OH)2 Mn reaction), so it doesn'tcontinue to a still lower pH.

The open circuit voltage of the cell is 2.05 volts or so. This makesit the most energetic alkaline electrode ever, and the reduced pHworks better than pH 14 for both the Mn and the NiOOH: it may well

make the cells last virtually indefinitely.


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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 saltsolution versus +.5. But with all the additives to make nickel workwell, 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 voltageof nickel anyway. Then I thought of the wholecell: double the amp-hours makes for double the amount of negatrode substance for thesame amount of positrode. This drawing illustrates the effect, whichis greater than it seems because the negatrode is less than 1/2 theweight 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 almostdoubles again the capacity of cell B for only around 1/3 additionalweight.

Although cell A has the highest voltage, and although the positiveelectrode has about the same watt-hours as cell B, cell B has twicethe 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 2xthe amp-hours if the equipment being powered can tolerate a lowercell voltage.

In this case, one or two more high energy negative electrodes canbe added to the battery to utilize these additional amp-hours,depending on acceptable voltage. (Cell C) The drooping voltages asthe 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


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would have relatively short cycle life. So I went back to my recentlydiscovered nickel manganate, which had about the same voltage,and which I expect probably isn't soluble. I suspect it probably alsohas 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 thisweek.

Zinc

Note: Zinc is superseded by manganese with 1% stibnite added toraise its hydrogen overvoltage. Manganese is the better choice inevery way.

Zinc's reactions make it suitable only for a negatrode, but quite a

high energy one. The dissolved ion form found in discharge andshown 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 andcharge.

Addition of a transition metal or its oxide is used to raise thehydrogen gas generation voltage (the "overvoltage") to improvecharging 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 orZnO as it discharges, but as usual the difference is merely the watercontent of the battery charged versus discharged, since Zn(OH)2 =


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NiZn alkaline cells is best seen in the Zinc Pourbaix diagram. Here itis 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 atabout pH 8 to 13.

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

The question then is, is there any point to making zinc negatrodeswhen manganese ones have more energy and are just as cheap?One possible reason is to get close to a specific battery voltage. Forexample, if NiMn cells are 1.7 volts, 6 volts is hard to attain:1.7 * 3 = 5.1

1.7 * 4 = 6.8whereas 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 specificsituations.

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

"Active" high surfacearea zinc oxide

(ZnOxide.org)

An issue withzinc in saltsolution is that

zinc powderand zinc oxidepowder bothabsorb CO2 outof the air andform zinc


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carbonate on the surface, which is passive in a battery and (I think)an insulator. The carbonate however can be removed by immersingthe 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 strengthenthe brittle zinc electrode.Not only does the carbonate become zinc oxide, evidently it

becomes the finest, high surface area "active" zinc oxide, ideal for abattery electrode.

In traditional manufacture of alkaline batteries with zincelectrodes, the finished electrodes are placed in KOH for a day, andthe "carbonated" electrolyte is replaced before charging. But thesoluble zincate ion causes zinc electrodes to degrade rapidly enough

that NiZn hasn't been a very popular choice, lasting as few as 10 to50 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 andare now commercially viable, attaining 400-1000 charge-dischargecycles at 100 WH/Kg, probably at a substantially lower cost thanNiMH or lithium. When the patents run out, they might becomeavailable in vehicle battery sizes instead of just small dry cells.

Cadmium also forms a soluble ion and NiCd dry cells often don'tfare much better than zinc, cadmium being right under zinc in thesame column of the periodic table. They do have zinc's highconductivity. NiCd pocket cells, however, like other pocket cellbatteries, have a good reputation for longevity. Since the atomicweight of cadmium is 112.5 versus 65.5 for zinc, and since itsvoltage in alkaline solution is -.82 instead of -1.25, the energydensity of cadmium is only 38% that of zinc. Hydride is much highereven with the same voltage. (-.83) Nickel-iron is probably better

too, even tho utilization of the iron isn't high, as it tends toagglomerate into larger particles with less surface area with cycling.(Additives such as cadmium help, and it was from using cadmium asan Fe additive to NiFe that NiCd was developed. I can't help butwonder if a sufficient quantity of graphite would keep the ironparticles from merging.) But I digress.


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3. Battery Construction Overview

For batteries, one thinks immediately of electrochemistry, but theconstruction of a battery is no trivial part of making it work. A goodpart of the effort of four years of battery R & D was trying to comeup with workable ways to actually make a battery, any battery, as afeasible 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 electronicallyconnected together, that is, connected for electron flow. Ideally it isone total "short circuit" from any point to any other point. All theactive material is electrically connected straight to the batteryterminal. In practice there may be resistance, even considerableresistance, between points because many active materials aresemiconductors, but there can't be any insulated points. Parts of anelectrode that become insulated from the rest cease to function;they are "passivated" like sulfated lead-acid battery plates graduallybecome. The lower the resistance within the electrode, the morecurrent can flow with less voltage drop.

Again, electronic conduction refers to conduction of electrons, wetor 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 theelectrolyte. The reactions only take place when the electrolyte ions

can interact with the active chemical. Again, any parts of anelectrode where the electrolyte is blocked are passivated and donothing.

These two requirements, electron flow and electrolytepenetration, are in conflict for actual physical construction. A goodbattery requires an immense active surface area in contact with the


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electrolyte. The surface area of a sheet of metal is small, and all butthe very surface atoms of the sheet are wasted, out of contact withthe electrolyte. A vast multiplication of minute particles to make aporous substance is required in order that the battery electrodes

need not span a gymnasium to supply much current or store muchenergy.On the other hand, these many minute particles must all be in

electrical contact with each other and they can't physically fallapart. To achieve this, they must be "glued" and compacted fromloose powder into something more like a dense piece of sandstoneor brick - a porous electrode "briquette". The briquette must be wellcompacted so the particles electronically connect, and yet consist ofopen 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 ofparticles over much distance (and increasingly poor with oxidationlevel), some sort of continuous metal or carbon conductor spans theentire area of the electrode, the "current collector". The briquette iscompacted around this for good contact throughout. None of thegrains are more than the electrode thickness away from this plate,mesh or metallic sponge that is connected straight to the batteryterminal.

Obviously there's an optimum compacting pressure to achieve thebest compromise between electronic conductivity and pores for ionicconductivity. Doubtless this varies with the ingredients in theelectrode mix. An electrode with fluffy nickel hydroxide andconsiderable graphite powder may have a different optimumpressure than a dense zinc electrode with few additives.

The only figure I've seen for compacting pressure was in oneresearch paper where the authors mentioned an "optimum"

pressure of 675 Kg/sq.cm - 9600 pounds per square inch - for aniron oxide electrode. For the chosen 1.5" x 3" electrode size, thatwould be 21.5 tons. I trust this may be taken as a maximumpressure requirement. I describe some electrode compactors andways to get sufficient pressure in the appendices. (I hope to offer agood compactor press with a "steering wheel" type tighteninghandle - but it can be done by tightening some bolts with a wrench,


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too.) Getting the pressure is one of the chief keys to makingbatteries.

The material chosen for the current collector and the terminal leed

is important. More particularly, the surfaceof the material, incontact with the electrolyte, is important.

In the salty battery with neutral pH, every metal I tried for acurrent collector in the positrode dissolved. To manage this (afterover 3 years of frustration) I created "grafpoxy", a 1 to 1 (byweight) mixture of epoxy resin and graphite powder. A relativelyfine metallic screen (around 30 mesh), with a terminal riveted orwelded 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 withthe 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 solventevaporates.) I find that two coats are needed, and it should beinspected in a good light. If any trace of copper color is visible, themetal will dissolve away until the cell quits working.

The grafpoxy coating does for salty batteries what nickel platingdid for alkaline batteries in 1900 - makes them practical. It replaces

the carbon rods and graphite sheets I was trying previously, orgraphite impregnated plastic contact sheets, none of which makevery good and durable contact with the electrode briquette. (But inJanuary 2012 discovered "Pourbaix diagrams", which show that asomewhat alkaline electrolyte is best for virtually all of thechemicals discussed. This can evidently be obtained by using saltbut adding calcium hydroxide to the positrode. The slightly solubleCa(OH)2 raises the pH to (theoretically) 12.3, an "ideal" moderatelyalkaline pH, tho still caustic enough to be somewhat hazardous.)

Inahighervolt


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age negatrode, a material with sufficient hydrogen overvoltage mustbe chosen. (Hydrogen voltage is -.833 volts in pH 14 alkali.) I triedmany manganese and zinc electrodes with copper or nickel platedmesh that would self discharge and bubble hydrogen. I could

understand this for the experimental manganese, but zinc was aknown, working electrode chemical. It was ages before I finallyrealized it was the current collector doing the bubbling, and not theactive chemical substance itself.

Zinc metal itself, or silver, evidently works well. Recent researchin Iran showed that a tin-zinc mixture also corrodes. But thisresearch showed that an alloy of copper, tin and zinc, "optalloy"evidently "acts as a noble metal" with a high overvoltage and workswell.

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 cylinderpocket electrode. However, these proved to cause a fair bit of selfdischarge.

Instead, I got "zincate solution" for 'priming' aluminum and zinccoated an aluminum rod (after shining it up with a nylon scouringpad). This seemed to get rid of the remaining self discharge. (As oftoday - 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 toto melt some tin or tin-silver solder in a pot on the stove, and mixin some zinc, which will gradually melt as it alloys with the soldereven if the temperature isn't hot enough to melt zinc by itself. Putsome soldering flux on a copper grill and wire, and dip it in the potfor a moment to get a coating. The tin-zinc coating may or may notcorrode 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 theresearch. "White bronze" is also commonly over 1/2 copper with theother 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


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(especially nickel hydroxide), and if they are able to do so, they losetheir conductivity and become pretty much useless. This was amajor problem through most of my battery research. The bestsolution 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 theactive material, usually with additives mixed in.

Pocket electrodes consisting of thin perforated metal enclosures,"pockets", to hold the electrode briquettes, were invented in the1890s. These work great and are highly conductive but with metalpouches 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 thanlead-acid and Edison's best version was in common use in earlyelectric cars by 1910 or so. They had to be nickel plated (at least inthe positive electrode) to avoid dissolving away. Since all commonmetals including nickel dissolve in salty electrolyte, metal pocketcells would be impractical for them.

However, having dismissed pocket electrodes for most of theduration of the battery project, I ended up adopting perforated rigid

plastic pocket electrodes as the best choice for homemade DIYbatteries. The extra weight of the electrode and battery casestructures is compensated and more by better, higher energychemistries.

Sintered electrodes were invented in the late 1920's as a betterway, but their manufacture and use only spread gradually. Thecarbonyl or a powder of the metal, usually nickel and cadmium forNi-Cd's, would be sintered (heated until it softens and flows a bit,

and the particles just barely melt together where they touch) into aporous "metal sponge" structure full of minute open cavities -- 80 -95 % empty space. The electrode active particles would beimpregnated into these spaces, then the whole thing compacted. Asthe conductive metal permeates every little recess of the entireelectrode, these are highly conductive and have great currentcapacity from small cells, eg ~20 amps from ~"AA" sizes. The


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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 theelectrode 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, makingthis 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 compactedaround a nickel or nickel plated metal mesh or perforated foil"collector plate", with a binder "glue" in the mix. The compactedbriquette is the finished electrode, with a nickel leed welded to anedge of the foil or mesh. Since there's not much there besides the

active chemical and its additives, the highest energy density byweight is attained. The metal case of the dry cell preventsdecompaction of the electrodes, which virtually fill the entire spacewithin the cell. Generally the current capacity is lower per squarecentimeter of electrode than sintered types. The amazing 2.6 AH Ni-MH size "AA" (100 WH/Kg) will only put out about 7 amps. (Thereare 2.0 AH "high rate" NiMH's that are good for 20 amps. These areprobably 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 electrodesharden up in use inside the cell as pathways establish themselves.

I found the perforated plastic pocket cells were the most reliableto 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 (tothe point of them catching fire and burning for up to 60 seconds),acetaldehyde, VeeGum (a bentonite clay mixture), and agar agargel.


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After many experiments, I read that CMC should be "under 1%"of the electrode substance. PVA of up to 2% has been used in zincelectrodes. These figures mean I was using substantially too muchof whatever I tried.

I hear that PTFE is the best for "ordinary chemistry" alkalineelectrodes, but it seems hardest to get. Also because a substanceworks well in alkaline cells doesn't mean it'll necessarily work well insalty 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 plateelectrodes: electronic conductivity and ion conductivity.

Naturally, with an electrode that has semiconductor activematerial 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 haveto travel. Thus thinner electrodes may be expected to effectivelyhave lower resistance and higher current capacity even for highlyconductive electrodes. Voltage with thicker ones will drop off moreat high currents when their charge is lower, as the remainingcharged 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 shouldprobably be under 6mm unless there are quite a lot of batteriessharing the load. These are all considered very thick electrodes inmost 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 havesubstantially higher resistance than the .375" (3/8") square, but

both types will be substantially higher resistance than a thinner flatplate. A larger diameter of central current collector wire will reducethe distance to the particles and lower the resistance, but unless it'shollow, it'll add weight without adding storage capacity.

If the internal resistance can be lowered, and-or if the electrolytecan penetrate better, larger diameter cylinders will perform better.Having picked this simple construction and got working electrodes


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but with rather low conductivity, this will be a bigger focus in futuredevelopment.

The manganese standard dry cell "+" electrode occupies almostthe 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 cylindersthere are in the cell, and the more cells that are in parallel, thelower the overall resistance will be and the higher the current thatcan be driven, but they will still charge and discharge at about thesame rate.

Another aspect to this problem is the speed of ion diffusionthrough the electrolyte. If an electrolyte diffuses ions twice asquickly, the electrodes may be considerably thicker and still have

the same current capacity. Potassium chloride is, I believe, aboutthe fastest electrolyte. Potassium compounds (KCl, KOH) are knownto be faster than their sodium equivalents (NaCl, NaOH).

Battery Layouts

One common type of battery construction is wrapped, spiralelectrodes, 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 inparallel to each terminal. This is usually used in flooded cells suchas lead-acid.

These constructions have the advantage of providing themaximum interface area between electrodes, each plate beingadjacent 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 minutelyperforated, nickel plated shell of thin metal holds the electrodechemicals compacted. (Nickel is the only metal that doesn't corrodeaway in the positrode in KOH or NaOH alkaline electrolyte.)Typically there's a wall every 1 to 2 cm, dividing the pocket platesup into rows or columns - they would bulge out if wider. These


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plates, just over a couple of millimeters thick, are then used in"prismatic" form with spacers between the electrodes. The metalpockets are thin and perforated, but they do add some weight tothe battery. Nonetheless, the prototypical pocket electrode battery

dating back to about 1902, nickel-iron, substantially outperformslead-acid and lasts far longer - some decades old NiFe batteries stillwork today.

Chosen Battery Layout:The Checkerboard of Perforated Plastic Square CylinderPocket Electrodes

First battery with two 1/2"square cylinder electrodes

AtthestartofFebruary2012Isudd

enlyconceivedof acompletelynewconst

ruction,easier andmorecertain to succeed as DIY

construction: the Perforated Plastic Pocket Electrode. This


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harkens back to the early days of batteries - but they didn't haveplastic back then. A number of previous ideas came together, and acouple of new ones were soon developed, to make this work. Abattery 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 heldthat way. Initially the biggest problems were bursting of theperforated cylinders and high internal electrode resistance.

The first plastic cylinders were 1/2" square inside. This madeelectrodes that were just too fat, and the conductivity was verypoor. So this electrode size was changed to 1/4" square, more inkeeping with the electrochemical requirements (see "ElectrodesOverview", next section), and it increased conductivity an order of

magnitude. The cost was making four times as many electrodes. Imay try 5/16" and see how that works, but they'll probably bepretty low rate cells. At that point, I decided a special jig to helpcompact the powders into the small tubes was needed to speedthings 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 toohard).

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" ABSsheet plastic. Next will be tougher .020" styrene plastic to cut wastespace, weight, and the distance between electrodes. They areperforated with a heavy-duty sewing machine.

The perforated plastic is cut into sections about 25 x 65 mm. Thisis heated in a kitchen oven on an unwanted cookie sheet or shallowbaking pan to 350 degrees F, for about 3 minutes if the oven ispreheated. (The longer it's left in, the more it shrinks.) In 3minutes, it should be pretty much limp and can be formed into anydesired shape. That shape is a four walled cylinder of 1/4" squareinside dimensions, by 60mm tall, formed around a 1/4" square steel


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rod with a jig. A bottom and a top end cap close the ends. The topcap has a hole for the connection wire. The overlapped seam andthe end caps are glued with methylene chloride, a solvent whichdissolves 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 itdown. Once it's filled, the end caps are glued on.

The straight connection wire runs right through the cylinder fromtop to bottom, and sticks out the top far enough to poke throughthe top of the battery to solder to. At the roof of the battery, it'ssealed 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 zinccoating 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 intothe hole.

Like the carbon rod in the manganese center of a "carbon"-zinc drycell, a single grafpoxy coated wire sticks out the center of th

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