tbatteryconstruction guide

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Making Economical,Green, High-Energy Nickel-Manganese (NiMn) Batteries

Small Scale/DIY Battery Makingthe Turquoise Battery Project

PRELIMINARYby Craig Carmichael, December 30th, 2013

TurquoiseEnergy.com

Quick Summary: These flooded cells utilize a new "moderately alkaline" chemistry, charging to about 2.6 volts and

nominally giving about 2.2 volts and high amp-hours per weight, for high energy storage (potent ially bet ter thanlithium ion) and very long or indefinite cycle life, from economical chemicals and supplies.

It should be noted that while the chemistries have been developed and tested, the author has not been able toproduce batteries that come even close to making good percentage utilization of the active chemicals - of anychemistry. These instructions must therefore be taken as a starting point for battery making, not an end point.

The metallic manganese "negatrode" (-)side is similar in reactions to familiar metals cadmium, zinc or iron,except for 3 things:1. The voltage is higher than any other, being around -1.4 volts.2. It's completely stable: in no state of charge is it an insulator or soluble. It should las t forever.3. Manganese has never been made to hold its high voltage charge before. Normally, it spontaneously discharges toMn(OH)2+ H2gas. A pH below 14, and 'trace' additives Sb2S3(1%) and ZrSiO4(3%), have been found to allow it to

hold its charge.

Zinc has been used for current collectors and conductivity enhancement powders. (Most metals would bubblehydrogen and discharge the electrode.)

The nickel plus manganese oxides "positrode" (+)side is a modified version of the typical 'nickel' oxides alkalineelectrode. The addition of manganese (~60% Ni to 40% Mn by element) causes mixed valence charging to nickelmanganate to increase amp-hours per weight and to reduce cost. (Mn is cheap - the stuff of throw-away dry cells.Too much manganese per nickel would however form soluble KMnO4 and deteriorate the e lectrode with cycling.) Thelower pH of the electrolyte raises the voltage of the reactions. It uses graphite materials for current collectors andconductivity enhancements. These materials are improved during fabrication. (All metals will oxidize in thepositrodes, so carbon materials are necessary.)

The Electrolyteis potassium chloride (KCl) salt, with the possible addition of calcium hydroxide (Ca(OH)2, lime),

which sets the pH to about 12.7. Without the lime, the pH generally drifts from 7 up to 12 or 13 over severalcharge-discharge cycles. Since lime is only slightly soluble, extra lime can be added, which will simpy sit on thebottom unless it's needed. (Lime may also be an additive layer in the positrode to improve oxygen overvoltagecharacteristics and conductivity.)

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 authorwarns the reader that his highest formal chemical education is a 74% grade in Chemistry 30 in grade 12, in 1972. Mosteverything learned about the subject since then has been self-directed web based studies and experimentation.

Note that preliminary editions are being written in bits and pieces as research proceeds, and the text may not beconsistent within itself: one statement might say "is expected to" or "should", while somewhere else, text written later maysimply 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. One piece of writing may contradict another.

READ THIS ==>THIS BOOK IS NOT COMPLETE AT THIS TIME AND IT CONTAINS MANY OUTDATEDINSTRUCTIONS. Only chapters 1 and 2, and "electrodes overview" in chapter 3, are up to date. SEE TURQUOISEENERGY NEWS #60 TO #69 FOR THE MOST UP TO DATE INSTRUCTIONS AND FORMULAS. PAY ATTENTION TO

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LATEST ELECTRODE INGREDIENTS, CURRENT COLLECTOR AND CONDUCTIVITY ENHANCEMENT MATERIALS,AND VARIOUS DETAILS.http://www.saers.com/recorder/craig/TENewsV2/index.html

TE News #60, #61: step-by-step procedure - with pictures (#60), revised but no pictures (#61)TE News #66: Breakthrough additive (zirconium silicate) addition to negative electrodeTE News #69: Latest positive electrode formula... sufficient nickel oxides in comparison to manganese oxides

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 Battery Electrochemistry Specific Electrode Substances - Nickel - Manganese - Nickel plus Manganese (Nickel Manganate?)

- Silver - Vanadium - Perchlorate - Zinc - Current Collectors

3. Battery Construction OverviewElectrodes OverviewBattery Layout(s)Chosen LayoutElectrode 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 plasticForming the square cylinderEnd 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 TestingInitial Rest PeriodInitial chargeInitial cycling

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Testing Specs

11. AppendicesA. Creating Unusual SubstancesB. Materials and Chemicals Supply SourcesC. Equipment & SuppliesD. Survey of Some Battery Electrode Materials

1. Foreward and Backward

[This section is an overview of the state of battery affairs and my own developments. If that's just a digress ion toyour prupose, skip to section 2.]

In one sense, bat teries are a well known technology, intellectual property of mankind. In another, they are almosta lost art. Factories churn out inferior lead-acid cells and small cells for portable electronic devices and cordlesstools, but the employees are just workers. While the theory of operation and the chemical reactions aren't hard toundersatnd, there are a few important details needed for successful construction that are hardly mentionedanywhere, much less all in one place. I'd say very few people know anything practical about battery design orbattery construction principles. I started this project knowing as little as anybody else, and it took six years of parttime study, experiments and frustration to finally get usable cells - and even then they're not meeting the

potentia ls of the chemistries, the percentage utilization of the active substances st ill being poor. I wish to savefuture batte ry makers from such trouble.

A great need has long existed for long lived, economical, high energy batteries for electric transport and off-gridpower. I decided to try my hand at creating some way to make some sort of batte ries at home. I soon felt sure that some better chemistries, probably much better, than existing types could be created, andpotentia lly for lead-acid or throw away dry cell prices, o r not so much more.

Batte ry research and commercialization have been sidel ined by human propensity to "go with the flow", to limitthoughts into narrow structured channels, good or (more often) inferior, and to extend that channel to the exclusionof wider possibilities, including superior ones. Thus for example, when Jungner found in about 1899 that nickel wouldn't oxidize at pH 14 and could be used as acurrent collector for alkaline batteries, subsequent research went into pH 14 alkaline chemistries with potassiumhydroxide electrolyte almost to the exclusion of all others. Later, with large, higher-energy alkaline batt eries such as nickel-metal hydride having been killed commercially by

corrupt interests in the late 1990 s via "dog in the manger" corporate patent wielding (the most important, mostcommon, and most heinous use of patents - which were created to give inventors, not corporate megaliths, rights totheir inventions), and with the single-minded recognition that lithium is the lightest atomic weight metal, mostresearch today has been working on trying to develop better lithium batteries, despite the cost and the complexproblems of making lithium work well, and despite the fact that patents on the best developments are acquired tokill each one as it emerges. This insanity won't continue for a second 100 years - the world is evolving - but in themeantime, DIY battery making of non-patented chemistries provides the rest of us a way to take matters into ourown hands.

Making 'normal' water based batteries is a rather involved but fascinating "DIY" project touching on severaldistinct specialties, and it creates a product valuable to civilization at this time. The process of learning and makingwill 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 theinstructions? Knowledge is power! I'm telling all that I can think of to say. But I'm organizing it into various sectionsso the reader can read as much or as little as desired - the basic instructions, a good theoretical overview, orcomplete detail. In other material, even the most basic information is lacking. For example, why is the positive electrode in astandard dry cell a conductive carbon rod instead of metal as in all other batteries? You'll dig long and deep and stillnot 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. A lower pH, onlycarbon or graphite works. Obviously battery makers know this (or once did), but it took me over two years ofcorroded electrodes in every test cell to figure it out for myself, because no one mentions it anywhere. (I put it onWikipedia, but it was soon erased.)

Much of the info herein has been acquired gradually, and often painfully, in my battery research over the past 4years. A tidbit of bas ic 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 ofpressure used to compact a battery electrode into a "briquette" - for one type of electrode in one experiment. WhenI started, I wasn't even aware of the vital role of compaction, and after eventually deducing it indirectly from somematerial density specs, it took a another year to figure out a simple way to get enough pressure. And it wasn't until January 2012 that I discovered the essent ial "Pourbaix diagrams" that show the relationshipbetween alkalinity, voltage and reaction products. These seem to show that a somewhat alkaline electrolyte is best



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for virtually all of the chemicals discussed. This can evidently be obtained by using salt but adding calciumhydroxide to the positrode. The slightly soluble Ca(OH)2 raises the pH to (theoretically) 12.3, an "ideal" moderatelyalkaline pH, tho still caustic enough to be somewhat hazardous. Likewise, it wasn't until February 2012 and four years of mysterious self-discharge problems that I understoodthat the wires in the negative electrode had to have as high a hydrogen overvoltage as the electrode substanceitse lf charges to. Most anything goes for iron, cadmium or hydride, but few common things work with a highervoltage chemical - zinc or manganese. For the Mn negatrode, it has to be z inc - not copper, brass, nickel, iron, ...(...or perhaps graphite, lead or bismuth might work. These I have not t ried.)

My original minimum battery goal was to copy proven and relatively economical NiMH EV battery chemistry, by thesimplest 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", Imight, in stumbling around, uncover overlooked information or ideas that could lead to a bet ter battery. That would have the additional advantage that being developed by me, freely and openly published by me, anddesignated by me as the inventor to be free technology. There would be no patent restrictions on it for vestedinterests to kill commercialization with. (Patents aside, it would be very difficult to make a decent hydride alloy athome.) I did indeed do a good bit of stumbling around in my ignorance, getting wild ideas and then see ing the flaws, andgradually learning many broad basics and fine details in no particular sequence. And I did uncover a few keyoverlooked things. I also developed some "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 havebeen rather successful: nickel-manganese batteries are in principle economical, "green", and superior to what 's onthe market today, including being quite economical and having about the highest feasible energy density,theoretically on a par or greatr than 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 isan excellent bonus.

2. Electrochemistry Overview

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% nickelhydroxide and 35% graphite powder, 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 lit erature usesthe terms "anode" and "cathode". The meaning of these terms is reversed when the battery is charging from when itis 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 confusionabout what is meant - I often get mixed up myself. As electrodes are ubiquitous to the subject and a specific one isso often referred to, herein I will call t hem "positrode" and "negatrode", which terms should be self explanatory. Ialso insist on spelling terminal wires as "leeds" to differentiate connections and wires from the metal "lead", theguy "in the lead", and at least a couple of other uses of the same four letter sequence, hoping not to "lead" anyoneastray.

Once an electrochemical design has been worked out, the physical design and construction becomes the moreimportant to making a battery that works. But the electrochemisty is the premiere part, the fascinating part, so itgets the first chapter.

I've tried to explain less common, specifically electrochemical terms herein, but the reader will understand thetext bet ter if he s till remembers his high school chemist ry. If you don't know what an "ion" or a "sulfate " are, justlook them up on W ikipedia. 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 otherreactions are measured. Whether a substance can be used inside a rechargeable cell depends on it charging belowthe voltage where gas is produced instead. Gas generation is more and more likely with increas ing voltage above 1.23 volts, but the exact voltage variesconsiderably with electrode substance and additives, temperature, and pH. Any amount over the theoretical gassinglimit, at which gas isn't generated, is called the "overvoltage". In acid, gas generation voltages shi ft 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 solut ion. However, the lead metal to sulfatereaction is just under the hydrogen generation limit at -.35 volts. On the other hand, in alkali, oxygen gas generation is encouraged and hydrogen more inhibited. The commonalkaline nickel positrode (+.49 volts) is just below the "oxygen overvoltage" at room temperature, and zinc justworks at -1.24 volts. The 0.0 volts in acid hydrogen voltage, in alkali is -.833 volts. The inverse of this voltage plusthe +.49 volts o f nickel gives us a theoretical open circuit voltage of the nickel-metal hydride alkaline battery, 1.32volts.



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Oxygen overvoltage fa lls a bit with temperature, and above 40C simple nickel electrodes won't charge properly. The electrode substance is also significant, and small amount of a high overvoltage potential substance as anadditive can increase the overvoltage so that the main substance works better, or works at higher temperatures. Toimprove 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 andfound 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 metalssuch as antimony are also used to improve lead-acid cell charge performance. In the case of manganese as a negatrode, lowering the pH even a littl e lowers the reaction voltage - from -1.57vtowards -1.18v at pH 8. That's the start, but it's not enough. Then, adding 1% antimony sulfide raises the hydrogen overvoltage above manganese's charging voltage -- as long

as the temperature is under about 20c. I discovered this worked in February 2012. Eureka! I was sure I had it! Then late r I got poor result s. In July 2013 I figured out the problem by putting the cell in the fridge. It didn't workat 25c. It worked fine at 6 or 8. I considered that cells with manganese electrodes might have to be kept in airconditioned spaces. Then I tried adding 3% zirconium silicate as well as the antimony sulfide, and this raised the stable temperatureto at least 30c. (...probably considerably higher, but that was as high as I happened to test it at.) The additives are the only reason it works at all. Without them, the overvoltage seems to be right on the edgeeven at low temperatures: the manganese may or may not charge, but it bubbles hydrogen as it does, and graduallydischarges itself to hydroxide. Thus manganese has never been used before as a negatrode. But its high reactionvoltage gives a "-Mn" battery an edge in energy density over any other. (Ni-Mn is higher voltage and far longerlast ing than Ni -Zn, making higher energy cells of about 2.2 or 2.3 nominal volts. In fact, NiMn moderately alkalinecells may last indefinitely.) There are lots of even higher voltage reactions, but it's hard to conceive of making any of them work with anyadditive. For example, aluminum to aluminum hydroxide is -2.33 volts in alkali. That would make for fantastic

energy density, but I doubt that will ever 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 sea ledbatteries are small. In addition, H2 has proven almost impossible to get rid of in sealed cells. Pressure would justbuild 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 passthrough. The oxygen migrates to the negatrode, discharges some of the substance (making heat), and preventscomplete charging of the negatrode. This gets rid of the oxygen, and prevents the negatrode from bubblinghydrogen 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 graduallydegrade substances within, turning them from active chemicals into carbonates. Various caps and valves canminimize the problems and vented cells aren't impractical, but they're second best to sealed.

To make sea led 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'vealso read that antimony is almost unique in its ability to react with small molecules - like hydrogen - and I picked itas 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 aspray 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 bl ind chemistry professor. A vented case reduces the dangers, but using potassium salt forelectrolyte makes it much safer.

I hated the thought of using potassium hydroxide or acid electrolytes. They're dangerous! I was using a saltbased electrolyte of neutral pH, potassium chloride. (KCl) It's a fast electrolyte (allowing high current flow), andless hazardous to handle than potass ium hydroxide - it's edible. However, the cells turn quite alkaline with cycling.The caustic OH- ions are much less concentrated, but it's still about pH 13.

In addition to chemistry, there were (and are) other novel improvements begging to be made. If one could find achemically inert but electrically conductive or even semiconductive binder 'glue' to hold the electrode powderstogether, it could permit higher current flow than the usual insula ting binders, and intense compacting of theelectrodes would be less critical to obtaining good current capacity... If a smal l, economical, high energy batterycould supply enough current to sta rt a car engine, that would be a marvel!

Battery Electrochemistry

First I'd like to point out a misleading quirk of terminology. Back in the beginning of understanding atomicparticles, someone decided electrons had a "negative" charge while protons were "positive". It doubtless all seemedpretty arbitrary, perhaps even using the words "positive" and "negative". Of course, t hese two words have other, wellknown meanings. Here they have been applied backwards. Consider that protons are stationary, within atoms, while free electrons move around between atoms... like banksand 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" ofthem when charged and soaks them up on discharge. This is all counterintuitive, and in some situations, ahindrance 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 negatrodeis the opposite. These confusing names indicate electrochemical reactions that involve loss and gain of electrons,



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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 allabout. (Hmm, "Reduction is gain!" -- another lovely litt le paradox of nomenclature!) A shorthand used for reductionand oxidation is "redox", and batt ery reactions are redox reactions. The electrochemical reactions at each electrode are called "half reactions", and the two half reactions of a batterymust balance each other. If the negat ive terminal supplies "x gazillion" electrons to an external circuit, the posi tiveterminal must soak up "x gazillion" electrons. And, the ions released internally by one electrode when the electronflow must complement those released by the other or be absorbed into it. After all, no atoms are being added to orremoved from a bat tery in use. The chemicals used in a battery are chosen both for complementing ions and such that the positive side is achemical 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. Usual ly 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 orhydroxide. There are exceptions, and many other possibi lities. In lead-acid batteries, the negatrode metallic lead oxidizes tolead sulfate, and the positrode lead dioxide reduces to lead sulfate, the sulfate ions being s tored as excess acid (orsodium bisulfate) in the electrolyte when the battery is charged, and absorbed as it's discharged. More examplesappear below.

Usual ly these positrode oxide forms aren't very good electrical conductors. Some oxides, like titanium andzirconium, are virtually insulators, so they can't convert easily between forms by electrical action as batteryelements. Often additives are used to improve the conductivity of the oxides. Zinc and cobalt oxides have beenused 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 orreduction, 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 energyin watt-hours is the amp-hours (the number of electrons) times the voltage (the pressure behind each electron). Thevoltage of both electrodes is subtracted for the total battery voltage, eg +.5 - -.93 = 1.43 open circuit volts for anickel-iron alkaline bat tery. The amp-hours or number of electrons isnt addit ive: it should match. The current flowstops and the cell is discharged when either electrode has been depleted to its un-energetic state and will pump nomore electrons and ions. For a given number of electrons moved per reaction, the lighter the atomic weights of the reacting elements, themore amp-hours per kilogram will be available, because there are more molecules to react in that kilogram. Oxygenand especially hydrogen are quite light, so the metal i s usual ly the dominating factor. If a heavier element i schosen, it must move more electrons per reaction, or have a higher reaction voltage, to provide equal energydensity. If the advantages are less than the added weight, as with lead, cadmium or mercury, batteries with these

heavier elements have lower energy densities. Heavier elements are also more costly. Thus my own searches (once Iunderstood this) 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 inthin film electrodes, often with non-aqueous electrolyte, and the substrates to hold all the thin films add their ownbulk and weight.

Usual ly 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 abattery 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 - hencethe 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 toqualify, 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 tablesindicated, about 1.4 volts.) But I didn't think this was the battery I was looking for. I suspected there would beunforseen complications over many cycles, and took i t no farther.

Zinc has been known as a frustrating battery negatrode element. It's energy is the highest available for pH 14alkaline cel ls and its electrical conductivity is good, and the charge and discharge products a re 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, whichusually short out dry cell batt eries, often after only 10-50 charge-discharge cycles. Cadmium, underneath zinc on theperiodic table, has the same problem to a lesser extent, and in my experience Ni-Cd dry cells rarely last anywhereclose to their supposed cycle life as cadmium crystals poke through the separator sheet and short the cell. NiZn andNiCd pocket cell bat teries fare much bette r. But it would seem that NiZn dry cells in recent years have improved, asa 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. Somehow I doubt it.

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 and with apparently ideal characteristicsfor a battery negatrode, sat enticingly on the voltage threshold between usable and not for a negatrode. After muchtime, frustration and confusion, I got it to work as explained herein.



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The electrolyte doesnt conduct electrons between the electrodes, it only conducts charged dissolved ions. It's theone place where protons are on the move. To have the oxidations and reductions take place, both ions and electronsmust 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 coursewhat the battery is for. When an external circuit is connected, the electron flow, the ion flow and the dischargereactions proceed spontaneously and simultaneously, releasing the chemically stored energy as electricity. The ionsflow 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 the battery depends partly on how fast the ions diffusethrough the electrolyte. Potassium chloride salt is supposed to be very fast. The discharging reactions release chemically stored energy electrically. The recharging reactions require electricalenergy from the external circuit - the bat tery 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 suchas water, with an acid, sa lt or alkali electrolyte diss olved in it. There is, however, one serious limitation to usingwater 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 C, 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 2.6open circuit volts of the Ni-Mn cell have raised the "practical limit" considerably.

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" voltagerating which might be expected in typical heavier use, such as "1.2 volts" for Ni-Fe, Ni-Cd, and Ni-MH, which readmore typically 1.33 to 1.43 volts with no load. Heavy loads may drop the output even more, eg to 1.0 volts. If suchloads are expected, it's usually best to add more batteries in parallel to reduce the load on each one, or to usebigger cells, which is effectively about the same thing. Ni-Mn should be rated when someone has made one thatutilizes at least 1/4 of the potential active material. (Mine so far, of any chemistry, are generally below 1/10th.)

If the positrode has lesser amp-hours capacity on discharge than the negative it is depleted first. The negatrodestill could have supplied more current and the battery is said to beposi tive limited. Vice-versa if it's the negatrodethat runs out first. It may also be positive or negative limited on charging, and not necessarily in the samedirection. It's also possible for the electrodes to be entirely off balance - one discharged and the other charged. Itcould be hard to eit her charge or discharge this cell . There are often good reasons for preferring one reactant to deplete first. For example, if there's no recombinationcatalyst 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, makinga bit of heat. Thus the cell stops charging - it just gets warm. Hydrogen doesn't readily discharge at the positrodeand the gas would accumulate until the cell bursts, so it's best to have the positrode charge first and not get anyhydrogen. With the catalyst, starting to generate both gasses at about the same time when the charge is completeshould be advantageous, since they can then start recombining to make water before the pressure of either gasbuilds up much. But that depends on having an effective catalyst and a design that utilizes it effectively.

Electrode Substances

Besides lead in lead-acid cells and lithium, there are two common positrode substances : nickel oxyhydroxide(NiOOH) and manganese dioxide (MnO2).

Manganese dioxide has been strictly the substance of one-use dry cells, so-called "carbon-zinc" but actuallymanganese-zinc, the carbon (as graphite or "carbon black") being in fact simply a conductivity improving additive. Insalty solution (pH 7) the energy is about +.5 volts, but in alkaline solution (pH 14) it's only +.15 volts.

Lower voltage cells might be okay given that manganese is cheap, except that when charged it doesn't readilystop charging at MnO2, valence 4. Unless charged very slowly with little voltage above the charged cell voltage, itcontinues to charge to potassium permanganate (KMnO4), valence 7, with the "K" from the electrolyte. And that's alittle soluble, so the manganese migrates and the electrode deteriorates. So makers of rechargeable alkaline batteries prefer virtually insoluble nickel oxyhydroxide, with +.5 volts.

In "moderately alkaline" cells, nickel oxyhydroxide's voltage, increasing as the pH drops, works great and makes2+ volt cells. But there's a third alternative: adding manganese to the nickel seems to cause "mixed valence"charging to nickel manganate, to provide the highest energy density. Others have noted the higher amp-hours whenmanganese is added to the nickel (at pH 14), but they explain it as "helping the nickel charge to a higher valence",which is probably not accurate. The proportion of manganese to nickel should be under about 40%, to avoid left-over manganese oxides and probable formation of potassium permanganate.

In alkaline negatrodes, iron, cadmium, zinc and metal hydride alloys have been used. Utilization of iron is poor,

and it very gradually forms ferric oxide or hydroxide and the electrode decreases in capacity. Ni-Fe is nevertheless along life battery chemistry. Cadmium crystals gradually grow owing to temporary solubility of cadmate (cadmiate?)ions. This tendency is however most pronounced in the element above camium - zinc. Flooded Ni-Cd cells last muchlonger than the dry cells, which often short out long before the 'advertised' time. However, being a heavier element(atomic weight 112) and lower voltage (-.82), the energy density is lower than wi th zinc (atomic weight 65.4 and -1.24v). Metal hydrides are alloys with mixed crystalline spacings that leave lots of nooks and crannies for hydrogen "-"



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ions to hide in without creating much pressure buildup. A typical alloy is Ni:La:Co in proportions 10:2:1. Amazingly,when the cell is charged, the H- ions can be packed in more densely than in liquid hydrogen itself. Since Ni(OH)2 (or'NiOHOH') charges to NiOOH using up one water molecule to release two "OH-" ions, and since the hydride storesone "H-" instead of "OH-", leaving one "HOH" - water, the water content of t he cell doesn't change between chargeand discharge as it does with the other types where the meta l converts to it s hydroxide. (eg, Cd + 2 OH- ->Cd(OH)2, effectively absorbing an H2O.) Lanthanum will spontaneously form La(OH)3 on contact with water, so the hydride gradually deteriorates andmany surface treatment tricks are used to try to slow this. As with Ni-Fe, Ni-MH cells will last a very long time -substantially longer than Ni-Cd or Ni-Zn. However, all of them have finit e lives for known reasons.

Until now zinc has been known as the most energetic negatrode element at -1.24 volts and 820 amp-hours/gram

of Zn, or 1016 watt-hours/kilogram. This i s much better than iron or cadmium and on a par with the best hydrides inalkali. However, its rather rapid deterioration owing to the temporary formation of zincate ions makes it unat tractivefor lasting batteries. Manganese , now that it's working, appears to be an ideal negatrode, and the only one that should lastindefinite ly. It has even higher energy density than zinc: theoretically 976 amp-hours per gram of Mn at a round -1.4volts or 1366 watt-hours per kilogram.

Nickel

Nickel hydroxide [Ni(OH)2] is the common positrode material used in most rechargeable alkaline batteries with

various negative e lectrode materials: Ni-Cd, Ni-Fe, Ni-MH and Ni-Zn. Dry and pure, it's a very fine, fluffy, turquoisegreen powder. The nickel will happily stay in the hydroxide form in the battery environment. It thus has no usableenergy. To convert it to a more energetic chemical, energy must be put into i t. To charge it, the nickel hydroxide is further oxidized to nickel oxyhydroxide by grabbing one electron from it. It

doesnt willingly give up the electron: the charger has to supply the energy to cause it to happen, exceeding +.52volts. This disengages a hydrogen ion (H+), which jumps over to an immediately adjacent hydroxide ion (OH-) in theelectrolyte to form water. The nickel is 'oxidized' from valence +2 to +3, by losing an electron (to the chargingcircuit) 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 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 usuallya prime requirement that the electrode doesn't dissolve. Normally if it does, the battery won't recharge. The valenceof 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 crystal line fo rms "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 Romannumerals) When it gets below 2.25 or so, the resistance rises and the user considers the battery to be "pretty muchdead". 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 thisreaction. Voltages seem to vary slightly with different electrode additives, and perhaps with temperature. A major advantage of salt y electrolyte is that the nickel reaction voltage is substantially higher,giving it morewatt-hours per kilogram than the pH 14 alkaline cell. This alone was a good reason to attempt to create workingsalt solution batteries.

The nickel oxyhydroxide is an "energized" substance: it would rather be just plain hydroxide and given a chancewill revert and give off energy in doing so. But it needs an electron and a hydrogen ion to do so. The amount ofenergy per electron is seen in the voltage. It can get the hydrogen "H+" ions from the water, leaving OH- in thewater. This is balanced with the negat ive 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)2to NiOOH, which has the same reaction voltage as the 2 to 4: +0.49 vo lts...or +0.52 depending where you read. If some of the nickel is oxidized to NiO2, valence 4, as shown

on the chart, it raises the number of electrons transferred and hence the amp-hours capacity.

Notice on the chart 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



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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... except that nickel metal won'toxidize in pH 14 solution. It does work in salt electrolyte, at a somewhat lower voltage.

A "Pourbaix" diagram shows the reaction products at different pHes and voltages.A Ni negatrode might discharge to Ni(OH)2 between -.4 and -.6 volts depending on

the pH of the electrolyte.A NiOOH positrode will discharge to Ni(OH)2 between +.49 and maybe +.9 depending on pH.

If the pH falls below about 8, the nickel can dissolve when it discharges.(I'm not aware that Ni(OH)3- is ever mentioned in literature...

but it's even less likely to form in moderately alkaline pH.I suspect this chart is very approximate in all its aspects.)

The valence of the nickel goes from II to 0, adding two electrons to each nickel atom. This charging reaction givesoff the negatively charged hydroxide ions that were bonded to the Ni(OH)2, the same as with iron, cadmium, zincand manganese, each at its own voltage. Moving two electrons instead of one, at -0.5 volts (instead of +0.5),provides 2 times the theoretical energy storage. Nickel hydroxide in moderate alkali, though the most commonpositrode material, makes a more energetic negatrode than it does a positive one! And it should have no limit to itscycle life . Notwiths tanding this, the voltage and energy of the reaction are lower than the usually used substances .Manganese has three times the energy, indefinite life, and is much cheaper.

So why is nickel [oxy]hydroxide so popular as a positrode chemistry if it has so lit tle energy beside the commonnegative ele ctrode materials? Well, it boi ls down to ... try and find something bet ter, that doesn't cost a fo rtune!

Manganese dioxide, while cheap, is only +.15 volts in alkaline solution. That means more cells to attain a givenvoltage. In salt, it's .5 volts, and it might be a more economical solution for stationary batteries, eg for off-gridhome power storage. It i s easy however, and considered dele terious, to charge it to a higher oxide form. (The zirconion 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 onlyrequired nickel in the salty battery is the actual active chemical, whereas in alkaline batteries nickel or nickel platingis used for all internal metalli c structures. (That could be changed with grafpoxy.)

But the traditiona l basic reaction doesn't reveal nickel's full potential. Nowadays, manganese is added to thepositrode as a major additive, perhaps 35 to 40% by weight ("wt%") of Mn to Ni. What this supposedly does is raisethe oxygen overvoltage, which evidently allows the nickel to charge to "alpha" nickel oxyhydroxide, wherein someportion of the nickel actually charges to NiO2, valence IV, moving two electrons ins tead o f 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 atta inable overall valence appears to be about 3.8. The actual nickel valence thus might change fromabout, say, 2.25 to 3.75 from discharged to charged, thus moving 1.5 electrons per nickel atom, twice as much aswith the old pure Ni(OH)2simple 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 asmuch work. Multiplying the theoretical value 289 AH/Kg * 1.5 = 433 AH/Kg. Naturally however, the theoretical maximum isn't



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going to be attained. (Experimentally about 350 AH/Kg has been attained, the forms being alpha hydroxide andgamma oxyhydroxide, which both occupy about the same volume of space. Although it's a higher volume form thanthe 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 sinterednickel cadmium "AA" batteries of just 0.5 AH in the 1970s.) Since the NiMH AA cells with this high energy weigh 30grams, and the nickel hydroxide (educated guess) probably weighs up to about half of it, an attainable figure in anactual 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, heavierelectrode, and anything that improves it can notably improve the entire energy density of the battery. For homebrewsalty 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 limitedby the nickel and the voltage obtained. A 1.8 nominal volts nickel-zinc/salt cell then will be somewhat under 180WH/Kg, eg maybe 120-170. For 2.1V with a manganese negatrode, if that can be made to work, 130-190 WH/Kgmight be attained. These figures seem dissappointing after reading the theoretical maximums, but they're stillbetter than commercial NiMH dry cells and as good as or better than lithium ion types. And it's not impossible thatwith 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 bearfurther experimentat ion) include: aluminum, cobalt , yttrium, ytterbium, erbium, and gadolinium. Other rare earthshydroxides such as samarium, neodymium and even lanthanum might be better, or at least fine, in salt solution. I'mnot sure why manganese is supposed to be "especially preferred" (or even why it should work well), or indeed whatthe 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 lowervoltage. (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 exactlywhat they're doing eithe r. Certainly inAlkaline Storage Batt eries(Falk and Salkind 1969), there was a lot o fspeculation about some of the main chemical reactions. And battery substance reactions in salty electrolyte arerelatively 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 toraise oxygen overvoltage in the postirode. If it is, the samarium or whatever, probably in considerably lesserquantity percentage-wise, s hould replace it entirely, providing highest energy density. (For a while I thought theKMnO4 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 completelydischarged.)

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

Neither nickel hydroxide, oxyhydroxide nor potassium permanganate i s a very good electrical conductor. Thebattery's current capacity would be extremely limited i f these were the only ingredients. Powdered graphite hasbeen 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 solidlyinto perforated metal tubes about the size of a pencil. The nickel flakes were made by electroplating alternatelayers of copper and nickel onto something, then dissolving away the copper. That costly arrangement was the besthe could come up with that worked well. He tried graphite flakes and found the performance was unpredictable - Ithink Edison didn't expect powder could be a good conductor across an electrode, but above a critical proportion itis . The sintered electrode is another good form for conductivity in alkali, the sintered nickel sponge connecting wellacross the whole e lectrode 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 metalelectrodes 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 withoutgraphite or nickel flakes, and I've been trying starting with monel alloy, which puts (25-33%) copper hydroxide insolid solution with the (67%) nickel hydroxide. (The monel I'm using also contains 2% Mn and 3% Fe, so the copperis 28%. Obviously Mn doesn't hurt, and the iron ei ther, I trust.)

On a practical note, it's worth mentioning that a nickel electrode can be discharged chemically to Ni(OH)2 byimmersing it in a small pool of hydrogen peroxide - the 3% drug store stuff is fine. It makes zillions of very tinybubbles as excess oxygen comes out. W hen 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 onedischarged for a sealed battery. For an unsealed one, charging and letting gas bubble off one electrode works.

Manganese Positrode

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 i s probably the bet ter source - it's known to be pure enough forbatteries and it's "pre-mixed" with conductive graphite powder. In the open, dioxide is the usua l sta te ofmanganese, but in the typical cell it's the charged state. An even better form for use in positrodes is as potassiumpermanganate.



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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 absorbedduring 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 a ll lit erature I've managed to find shows this reaction:

MnO2(s)+ H2O(l)+ e- MnOOH(s)+ OH

-(aq) [~+0.5 V]

Manganese Redox chart.

Manganese *can* be recharged to dioxide, and some "renewable " alkaline cells make use of this. However, just alittle extra charging voltage will cause it to charge to potassium permanganate (KMnO4), with the MnO4- ion beingsoluble. The soluble ion deteriorates the electrode, and it discharges to Mn(OH)2 when it touches the negatrode.The Mn(OH)2 gradually builds up in the separators and shorts out the cell. Thus manganese pos itrodes are knownfor short cycle li fe. The only listed permanganate that isn't soluble is silver permanganate, but some manganates (MnO4--) areinsoluble. Notably nickel manganate is probably insoluble, and we'll get to the mixed nickel-manganese positrode

further down.

Manganese Negatrode

Another manganese reaction of great interest is the one on the right end of the chart, going between valence 0and +2. If one simply uses manganese powder in water, this reaction is just high enough in voltage that it graduallybut spontaneously discharges into Mn(OH)2. This has always precluded the use of manganese as a negatrode. However, at any pH below 14 the voltage of this reaction drops a little . Combined with that, certain additives , inparticular heavier transition metals or their compounds, are known to raise the voltage at which hydrogen starts tobe generated. They have been 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 aresubstituted: eg, gallium, indium, tin, or bismuth.

I got the idea that if zinc could be improved in voltage and temperature performance with transit ion metal

additives, pe rhaps manganese - tantal izingly just out of range - could go from "doesn't work" to "works" by thesame route.

I tried antimony oxide (Sb2O3, stibia) with uncertain result s. In February 2012 I tried antimony sulfide (Sb2S3,stibnite). It worked! (BTW 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 bet ter. Whatever happens, adding 1% antimony sulfide raisedthe hydrogen overvoltage enough to allow manganese to charge and hold its charge... just enough, at lowertemperatures. Later - in summer - it didn't seem to work and a puzzl ing year went by. Finally in July 2013 I tried putting a cellthat wouldn't charge in the refrigerator, and lo and behold it worked! It only worked up to maybe 20c. (I didn'tmeasure the transition temperature.) Then I added 3% zirconium silicate (ZrSiO4, zircon) as well as the stibnite.This raised the working temperature at least as far as I tried it, which was to 29c. (Yes, I should have tried higherto verify that it would work up to at least, say, 35 or 40c. At the time, I was ecstatic to find that it worked at 29.I also didn't try any other concentration of zircon besides 3%.)

Charging Mn to metalli c state andhaving it hold its charge was doubtless a first, and is probably my biggestcontribution to battery technology. The reaction voltage of almost -1.5 volts, makes it the most energetic alkalineelectrode ever, and the reduced pH works better than pH 14 for both electrodes. With no soluble or insulating statesin the negatrode, the cells may last virtually indefinitely.



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Nickel and Manganese (Nickel Manganate?)

In late February 2012 I ran across a potentially better form of nickel for electrodes than nickel hydroxide: nickelmanganate [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 electricalresistance than most oxides. At first I thought it might make a good conductivity improving additive. It was farmore 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. Finally, I started to expectthat if one placed both nickel hydroxide and potassium permanganate oxides in the electrode, they mightspontaneously form nickel manganate via substitution reactions, or that nickel hydroxide and manganese dioxidemight well form nickel manganate when the cell is charged - or both - without any formation of slightly soluble

potassium permanganate.

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 nickelpermanganate would release 8 electrons and use up 8 OH- ions from the charging negatrode (Four become the otherfour "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 thought that the nickel would be more likely to change since its reaction voltages are lower, probablysimilarly 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'd probably get a bit more mileage out of t he nickel - because ofthe high conductivity, it might discharge down to 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. Later I considered that nickel manganate was the likely result, probably giving a combined effect and increas ing

the amp-hours, and (it would seem) the voltage.

Before considering that it might form naturally with charging the two-oxide electrode, I looked but couldn't findnickel manganate to buy. I tried making it chemically, but it was slow, messy and smelly. Then I mixed appropriateamounts of dry NiO and MnO2 powders (both from the pottery supply) in a stainless steel pot and simply heatedthem red hot with a propane torch (outdoors, with a respirator). This gave a lower resis tance product (*probably*nickel manganate) and was fast and pretty simple to do. I only got 15 grams, so I guess the torch blew 7 or 8grams of powder out of the pot. Later I made another batch and 'only' lost 25% of the mass.

But if it didn't form naturally with charging of the electrode, it would probably separate into the two oxides ondischarge anyway, negating the point of having made it. And in that case, the solubility of KMnO4 would causeproblems that aren't apparently observed except in my electrodes with too much manganese compared to nickel, the60:40 ratio having been carelessly ignored or ignored for the higher theoretical amp-hours of more manganese.

The most likely thought is that nickel manganate is the charge product. Who knows about discharge? Butwhatever exactly is happening in the mixed valence electrode, the combination of nickel and manganese oxidesseems to make a stable ele ctrode evident ly with higher amp-hours, higher conductivity and perhaps higher oxygenovervoltage than than nickel oxides alone.

Silver



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Silver oxide works well in strong alkali, but moving up to 1.5 electrons (later ones at lower voltages) doesn'tentirely make up for its higher atomic weight. It's also quite costly, and it tends to degrade separator sheets, so it'snot used a lot except for tiny cells like hearing aid batteries. Disso lving at pHes lower than 12 (per Pourbaix diagram) would make it almost impractical to use with Mnnegatrodes. It could work if the pH was fixed to about 12.7 by calcium hydroxide in the electrolyte, but overall itappears less effective than NiOOH or the combining of NiOOH and MnO2. One good point is that the internal conductivity, and hence current capacity per unit area of interface, should befabulous. Silver oxides are quite conductive... but even more especially when some of the oxide has been reduced tometallic silver - the best conductor of all the elements.

Vanadium

I 2011 I made a battery with a vanadium electrode. It was supposed to be the negatrode, but it didn't seem towork - unexpectedly, the vanadium seemed to become soluble and to migrate. (This was also the first cell I'd madewith transparent plexiglass sides, and I could see the vanadium pentoxide yellow color appearing on the otherelectrode.) I reversed the charges, and found that the cell charged to about 2.2 volts. The vanadium positrode sidewould have made up around 3/4 of that, and it seems surprising 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 thegrafpoxy yet and it deteriorated like my other cells of that period, as the "flexible graphite" backing sheet swelledand lost conductivity and good contact with the electrode and the carbon terminal post. Judging by the voltage andthe 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 oxidesappearedto me to remain in solid form, not dissolve, in the salt electrolyte. This appears to make it a goodpositrode, moving one electron per vanadium atom, hopefully with good stability from the double vanadiummolecular center. Taking the average of the acid and alkali voltages as being the approximate salt voltage, the voltage obtained inthe cell indicates the s ingle 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.

Vanadium may have the potentia l to be a good positrode in salt water electrolyte. Theoretical amp-hours worksout to be almost identical to the theoretical 289 amp-hours/Kg of beta nickel oxyhydroxide. The potential doublevalence change that is achieved by some of the nickel to alpha oxyhydroxide molecules wouldn't seem to bepossible with vanadium, but the voltage appears to be higher.

According to the electrochemical table, we mightsuppose that vanadium might also make a goodnegatrode in alkaliat the same potential as cadmium (-.82) and hydride (-.833), providing it wasn'toverdischarged, which might fo rm the higher oxides, (egV2O3) which might cause problems. But it didn't seem to work right for me - at least not insalt electrolyte. There's a unique line straight fromvalence 0 to 5 that probably messes things up.

However, the "Pourbaix" state diagram from W ikipediaseems different from the above table. It looks prettyscary. I don't see the common V2O5 form (or VO for thatmatter) anywhere in the pourbaix diagram, and thevoltages don't match the table above, nor do they appearto jibe with my experimental results.

Vanadium may deserve more research in salt electrolytefor use as a positrode (and maybe in alkali as anegatrode), but I have no present plans for doing it

myself. After I found the Pourbaix diagram, I figuredthere's probably soluble ions somewhere during charge ordischarge, which might make for limited cycle life.

Rare Earth Perchlorate?

Chlorine ion, Cl-, oxidizes to perchlorate, ClO4-, moving 8 electrons with it s very own electrochemical reactions

regardless of the metal+ion it's attached to. This idea may have some obvious problem and be ridiculous... or not.(I did warn that I'm not a chemist !)

I once tried to make a positrode of lanthanum perchlorate, La(ClO4)3, which would reduce on discharge to

lanthanum chloride: LaCl3+ 24 OH-. The lanthanum was (per 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 forthis last, an article saying chelated dissolved lanthanum behaved about the same as undissolved, tho I can'tremember where I read of it.) In addition, perchlorate is often much less soluble than chloride, as with only slightlysoluble 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



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density. Lanthanum perchlorate, potentia lly with 12 O-- ions forming 24 OH-ions on contact with water, is a superexample: 24 electrons per reaction where nickel moves one or two. That much would more than make up for theatomic weight of La(ClO4)3 being almost five times that of Ni(OH)2, and - if I'm not overlooking anything - suggestssome theoretical possibility of a much higher energy density electrode than nickel hydroxide. Unless a real chemist knows of reasons this wouldn't work, it may deserve more research. (Next time, I think I'lltry converting lanthanum hydroxide straight to perchlorate with perchloric acid instead of first to chloride withhydrochloric acid.) (Perchloric acid is called a "super acid" - yow! I must read the MSDS again before I start. andsomeone says I should get rid of it so it doesn't turn explosive.)

Zinc

Note: Zinc is superseded by manganese with 1% stibnite and 3% zircon 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 foundin 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 bothdischarge and charge. Addition of a transition metal or its oxide is used to raise the hydrogen gas generation voltage (the "hydrogenovervoltage") to improve charging characteristics. 2.5% to 4% mercury oxide is 'traditiona l' 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 thedifference is merely the water content of the battery charged versus discharged, since Zn(OH)2 = ZnO + H2O. (IIRCthe general consensus is that it's ZnO.)

The troublesome zincate ion that limitsthe life of NiZn alkaline cells is best seen inthe Zinc Pourbaix diagram. Here it isrevealed that this ion probably won't formbelow about pH 13.5, and it's the pH 14electrolyte that's the problem: it would befine at about pH 8 to 13.

Evidently, adding some manganese oxideto the zinc to lower the pH to 13, as themanganese negatrode does for itself, shouldstop zincate from forming and allow long lifezinc negatrodes.

The question then is, is there any point tomaking zinc negatrodes when manganeseones have more energy and are just ascheap? One possible reason is to get closeto a specific battery voltage. For example, ifNiMn cells are 1.7 volts, 6 volts is hard toattain:1.7 * 3 = 5.11.7 * 4 = 6.8

whereas four NiZn is closer:1.6 * 4 = 6.4or 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 appli cation however, including 12 volts, the NiMn would seem to be the winner, needing only 7 cells for11.9 volts , while z inc 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



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and zinc oxide powder both absorb CO2 out of the air andform zinc carbonate on the surface, which is passive in abattery and (I think) an insulator. The carbonate howevercan be removed by immersing the powder or the electrodein a hydroxide: KOH, NaOH or Ca(OH)+ (lime). The lime isthe best and safest one. A bit of the Ca(OH)2 will becomecarbonate (CaCO3, limestone). This should helpstrengthen 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 zincelectrodes, the finished electrodes are placed in KOH for aday, and the "carbonated" electrolyte is replaced beforecharging. But the soluble zincate ion causes zincelectrodes to degrade rapidly enough that NiZn hasn'tbeen a very popular choice, lasting as few as 10 to 50charges, fol lowed by a shorted cell being the norm in dry

cells. However, according to Wikipedia, NiZn alkaline cells with "stabilized" negatrodes have been much improved sinceY2K and are now commercially viable, attaining 400-1000 charge-discharge cycles a t 100 WH/Kg, probably at asubstantially lower cost than NiMH or lithium. When the patents run out, they might become available in vehiclebattery 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 rightunder zinc in the same column of the periodic table. They do have zinc's high conductivity. NiCd pocket cells,however, like other pocket cel l bat teries, have a good reputation for longevity. Since the atomic weight of cadmiumis 112.5 versus 65.5 for zinc, and since its voltage in alkaline solution is -.82 instead of -1.25, the energy density ofcadmium is only 38% that of zinc. Hydride is much higher even with the same voltage. (-.83) Nickel-iron is probablybetter too, even tho utilization of the iron isn't high, as it tends to agglomerate into larger particles with lesssurface area with cycling. (Additives such as cadmium help, and it was from using cadmium as an Fe additive toNiFe that NiCd was developed. I can't help but wonder if a sufficient quantity of graphite would keep the ironparticles 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 ofmaking it work. A good part of the effort of six years of bat tery R & D was t rying to come up wi th workable ways toactually make a battery, any battery, as a feasible DIY project. And even at the end of that time, my cells oftendon't have 10% of the amp hours they ought to have for the amount of active material they contain - counting"ought to have" as 50% of the theoretical. (That's for any chemistry, not just ones I've created.) If I get a chance toget back to the battery project, increas ing the performance will be the main priority. In the end, the prime cause of performance that doesn't match reasonable expectations must be active materialelectronically poorly connected or unconnected to the terminal pos t, or (less common) active material the electrolytecan't reach. Reasons for these may be many and varied.

Electrodes Overview

Everything else depends on the electrodes. Besides the chemistry, what's in an electrode? how is it made? What areits 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 every point to every other point. All the active material is electricallyconnected 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 ofan electrode that become insulated from the rest cease to function; they are "passivated" like sulfated lead-acidbattery plates gradually become. The lower the resistance within the electrode, the more current can flow with lessvoltage drop. Again, electronic conduction refers to conduction of electrons, wet or dry, not ions. Conduction only by ion flowwhen 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 whenthe electrolyte ions can interact with the active chemical. Again, any parts of an electrode where the electrolyte isblocked 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 sheetof metal is small, and all but the very surface atoms of the sheet are wasted, out of contact with the electrolyte. Avast multiplication of minute particles to make a porous substance is required in order that the battery electrodesneed not span a gymnasium to supply much current or st ore much energy. On the other hand, these many minute particles must all be in electrical contact with each other and they can't



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physically fall apart. To achieve this , they must be "glued" and compacted from loose powder into something morelike a dense piece of sandstone or brick - a porous electrode "briquette". The briquette must be well compacted sothe particles electronically connect, and yet consist of open pores so they all also contact the electrolyte. And thebinder 'glue', if used, can't interfere or coat the particles with a waterproof layer.

Since connections are generally still poor though the maze of particles over much distance (and increas ingly poorwith 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 morethan the electrode thickness away from this plate, mesh or conductive sponge that is connected straight to thebattery terminal.

Obviously there's an optimum compacting pressure to achieve the best compromise between electronicconductivity and pores for ionic conductivity. Doubtless this varies with the ingredients in the electrode mix. Anelectrode with fluffy nickel hydroxide and considerable graphite powder may have a diffe rent optimum pressure thana 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. Perhaps this may be taken as a maximum pressure requirement. (Onthe other hand, maybe a reason my cells don't perform better is insufficient compaction.) I describe some electrodecompactors 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.) Gettingthe pressure is one of the chief keys to making batteries.

Current Collector Materials & Conductivity Enhancement Powders

The material chosen for the current collector and the terminal leed is important. More particularly, the surfaceofthe material, in contact with the electrolyte, is important. Likewise, with battery active chemicals usually beingpoorly conductive, some sort of powder or powders are usua lly added to the mix to improve it. These materialsshare the requirements of the collectors.

Positrode

In the salty batte ry with neutral pH, every metal I tried for a current collector in the positrode dissolved. After 3years of frustration, I realized only conductive carbon based current collectors would work. What worked were carbonrods from dry cells, but they weren't shaped for my flat electrodes. Graphite sheets seemed to swell and losecontact with the electrode.

At one point, I created "grafpoxy", a 1 to 1 (by weight) mixture of epoxy resin and graphite powder. The resistancewas rather high, but workable. I tried various ways to make i t work. A relat ively fine meta llic screen (around 30mesh), with a terminal riveted or welded to it , was coated in grafpoxy for use as the current collector. The epoxy

protected the metal from contact with the electrolyte, and the graphite let the electrode substance electricallycontact with the mesh. The mix should have about as much graphite (by weight) as epoxy. This generally makesrather thick for painting or dipping, so some solvent is added, eg, 10% toluene, to thin it. (The solvent evaporates.)I found 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.

Then I found cheap, thin graphite shee ts with low electrical resis tance. They come in rolls, sold as gasketmaterial. Instead of trying to connect them to metal within the cell, I cut them with a long "finger" tab to stick outa slit in the top of the cell. This is bolted to a matching exterior plastic tab to support the flimsy graphite, and thebolt is the terminal. The slit is s ealed with heat glue. Other expedients such as a brushed on dopant layer of osmium in acetaldehyde, and a dousing with toluene(methyl benzene) or "Diesel Kleen" (contains t rimethyl benzene) improved surface conductivity to better connect theelectrode briquette. The methyl benzenes dissolve graphite, and when they evaporate, it can solidify into moreconductive forms such as nanotubes and lamilae. This is especially helpful for the conductivity enhancing graphite

powder: Electrodes read 30-40% lower resistance after being doused than before. The liquid can also be addedinstead of water when the electrode is mixed and compacted. Another improvement, since there's no "graphite mesh" current collector, is to punch the plate (except for thefinger) full of holes, to give more surface area and grip between the active briquette and the current collector. Adevice called a "pin frog" or "flower frog" is available at florist shops. The many parallel pins that stick straight outof the thick metal base (for sticking flower stems into) can be pressed into the graphite to make many holes atonce. Anything that helps make good and durable contact with the electrode briquette should improve performance.

Negatrode

In a high voltage negatrode likemanganese, a material with sufficienthydrogen overvoltage must be chosen.(Hydrogen voltage is -.833 volts in pH 14

alkali.) I tried many manganese and zincelectrodes with copper or nickel plated meshthat would self discharge and bubblehydrogen. I could understand this for theexperimental manganese, but zinc was aknown, working electrode chemical. It wasages before I finally realized it was the



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current collector doing the bubbling, and notthe active chemical substance itself. Zinc meta l itself, or sil ver, evidently works

well. Recent research in Iran showed that a tin-zinc mixture also corrodes. But this research showed that an alloy ofcopper, tin and zinc, "optalloy" evidently "acts as a noble meta l" 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 thesquare cylinder pocket electrode. However, these proved to cause a fair bit of self dis charge. (which was probablythe plating corroding through in the initial charging, as I think about it.) In normal use, the zinc remains in metalli c form because the reaction voltage of manganese is higher. Opposite tomost usage of zinc, the manganese forms a "sacrificial anode" (if it may be so termed) which protects the zinc.

Then, I got "zincate solution" for 'priming' aluminum, and zinc coated an aluminum rod (after shining it up with anylon scouring pad). This seemed to work... at least on a second try after the electrode was initially charged. Thesolution can be found at Caswell Plating [.com], or can probably be mixed from sodium hydroxide (caution: verycaustic! especially protect your eyes!) and zinc oxide. I found a very simple way to electroplate zinc onto metal inan on-line instructional "How to Electroplate a Penny", which I copied into an issue of Turquoise Energy News. A problem with the plating is that the initial reactions with manganese oxides used in the briquette, which are"overdischarged" for a negatrode, will oxidize the plating. A freshly plated piece works the second time, afterinitially charging the electrode, where the first plated piece on the uncharged electrode has failed and cause selfdischarge.

Up to this point, I had been unable to find pure zinc except in fat lumps, rods and bars. Then I found that thinzinc sheet metal is available in rolls at roofing or building supply stores! And it's cheap. It's placed on roofs to killmoss. The solid zinc metal is a much better choice than plated material and it can withstand the initial oxidationproblem. I shaped these like the graphite with a long "finger" tab to poke up through a slit in the top cover of the

cell. Holes can be punched in the thin zinc sheets with the same pin frog as the graphite to improve conductivity. Fine zinc powder and flakes can be used to improve the conductivity of the manganese briquette. I note that even solid zinc seems to need protection where it's in contact with electrolyte but away from themanganese: especially near the bottom of the finger. I had one finger/tab corrode right off in a month. If graphitedoesn't bubble hydrogen, it might be a bet ter choice for the current collector. Zinc powder is probably still a betterconductivity additive. Then again...

Electrode Constructions - maintaining compaction

After the compacting there's the wetted electrode in the cell, before and after charging. Electrodes want to swellwhen wetted (especially nickel hydroxide), and if they are able to do so, they lose their conductivity and becomepretty much useless. This was a major problem through most o f my battery research. There has to be virtually nospace they can expand into.

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 consis ting of thin perforated metal enclosures, "pockets", to hold the electrode briquettes, wereinvented in the 1890s. These work great and are highly conductive but with metal pouches holding the electrodematerials t hey're expensive to manufacture and heavier, with low energy densities by weight. Nevertheless nickel-iron alkaline pocket batteries were much better than lead-acid and Edison's best version wasin common use in early electric cars by 1910 or so. The pockets had to be nickel plated (at least in the positive electrode) to avoid dissolving away even in causticalkali solution. Since all common metals including nickel dissolve in salty electrolyte (anything lower than pH 14),metal pocket cells are out for Ni-Mn designs. However, perforated rigid plastic pocket electrodes may be a good choice for homemade DIY batte ries, especiallyif one can coax a 3D printer into producing ABS pieces with very, very small holes. The extra weight of the electrodeand bat tery case s tructures is compensated and more by better, higher energy chemistries. I tried making some of these. A 3D printer owned by Camosun College made some very fine-holed "porous" plastic

pieces, but mine were too coarse and allowed material leakage. As another possible process, something like epoxy or thermoplastic (ABS?) could have a fine powder added to itthat would dissolve. It would be formed into thin sheets or whole pockets, then the powder would be dissolved outto render it finely porous.

Sintered electrodes were invented in the late 1920's as a better way, but their manufacture and use only spreadgradually. The carbonyl or a powder of the metal , usually nickel and cadmium for Ni-Cd's, would be sintered (heateduntil it softens and flows a

tbatteryconstruction guide

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