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My adventures building a Zn-MnO2 battery

Dendrites are the main reason why you don't see any rechargeable static Zn or Li metal batteries available commercially. Anything that would solve them for real, is a very big deal since it would open up a lot of cheap chemistries (like Zn) or much more energy density (in the case of a Li battery with a Li metal anode).

Modern Li batteries are a marble of engineering, the genius of John Goodenough was discovering you could intercalate Li in graphite to obtain an anode that doesn't form any dendrites. That's why modern Li batteries can get away with 50 micrometer separators without shorts.

Hey Daniel. I first replied to your thread on Zinc/Bromine batteries some time back. I was supposed to send you some discharge curves, but life got in the way. I am still following closely and love your updates!

I experimented with concentrated ZnCl + NaBr but struggled to go much past 30 Wh/L. I was using large gaps (3 mm) between anode and cathode which prevented dendrites from shorting out the battery. I was using graphite felt dipped in a graphite/nitrocellulose mix. I believe I was getting a degree of intercalation which was exciting! Self-discharge was much lower in these water-in-salt electrolytes.

Anyway, I struggled with repeatability and the bromine always seemed to destroy my seals, and so I moved away from bromine to try Manganese. I was pleasantly surprised at deposition onto felt but ran into an early realization. Depositing manganese dioxide on the felt would introduce hydrogen ions lowering the pH. So, the more mAh you deposited, the worse your hydrogen evolution would become on the anode! I tried using acetic acid as a buffer and although I had less hydrogen issues, I still had tiny bubbles forming (my cells are flooded so I can see what's going on visually).

I'm starting to think I have impure Zinc, because all these fantastic Zn papers just gloss over hydrogen evolution, calling Zinc the "ideal anode" for aqueous batteries. All my experiments have led to hydrogen evolution to some degree, except for ones with water-in-salt and very basic electrolytes.

I'm really impressed with the stability you are achieving with your latest tests! How are you preventing the lowering of pH and hydrogen evolution?

PS:I don't think John Goodenough discovered that lithium can intercalate into graphite. I believe it was a Japanese researcher. Goodenough was doing work on cathode materials and discovered a high voltage cobalt based host.
 
Hey Daniel. I first replied to your thread on Zinc/Bromine batteries some time back. I was supposed to send you some discharge curves, but life got in the way. I am still following closely and love your updates!

I experimented with concentrated ZnCl + NaBr but struggled to go much past 30 Wh/L. I was using large gaps (3 mm) between anode and cathode which prevented dendrites from shorting out the battery. I was using graphite felt dipped in a graphite/nitrocellulose mix. I believe I was getting a degree of intercalation which was exciting! Self-discharge was much lower in these water-in-salt electrolytes.

Anyway, I struggled with repeatability and the bromine always seemed to destroy my seals, and so I moved away from bromine to try Manganese. I was pleasantly surprised at deposition onto felt but ran into an early realization. Depositing manganese dioxide on the felt would introduce hydrogen ions lowering the pH. So, the more mAh you deposited, the worse your hydrogen evolution would become on the anode! I tried using acetic acid as a buffer and although I had less hydrogen issues, I still had tiny bubbles forming (my cells are flooded so I can see what's going on visually).

I'm starting to think I have impure Zinc, because all these fantastic Zn papers just gloss over hydrogen evolution, calling Zinc the "ideal anode" for aqueous batteries. All my experiments have led to hydrogen evolution to some degree, except for ones with water-in-salt and very basic electrolytes.

I'm really impressed with the stability you are achieving with your latest tests! How are you preventing the lowering of pH and hydrogen evolution?

PS:I don't think John Goodenough discovered that lithium can intercalate into graphite. I believe it was a Japanese researcher. Goodenough was doing work on cathode materials and discovered a high voltage cobalt based host.

Thanks for commenting!

It is interesting to hear about your experiments with water-in-salt electrolytes for Zn-Br. I never tried ZnCl2 + NaBr, since I moved away from this chemistry a while ago. It makes sense for ZnCl2+NaBr to have less self-discharge and less hydrogen evolution. The water-in-salt electrolytes do have lower conductivity due to their much higher viscosity values, so it makes sense your energy density was never much higher than 30 Wh/L. With a 3mm gap the energy efficiency of the batteries was probably also on the lower side (50-60% at best would be my guess).

About hydrogen evolution. It is the elephant in the room for sure. While few papers address it, we know that this is one of the main reasons why rechargeable static Zn batteries have never been successfully commercialized.

The papers with Zn-MnO2 batteries that are mildly acidic probably just have enough acid in the electrolyte to avoid these reactions killing batteries within the cycling experiments (enough acid to just produce some H2 and still function). However, you cannot just keep Zn in an aqueous acid for too long without it deteriorating heavily and generating hydrogen, my guess is if these batteries were assembled and then cycled after two weeks, they would not work well. This might be the reason why none of these papers have ever addressed the point of cycling batteries a longer time after assembly.

I also would be shocked if I'm not getting a lot of hydrogen evolution in the anode, given that my average charging voltage is quite high (>2V) and there isn't anything there that would suppress HER, it would be quite impressive if somehow these reactions were suppressed. My concentrations are not high enough to call this water-in-salt, as it's ZnCl2 4m + MnCl2 2m, so water activity should be high enough for hydrogen evolution to happen readily at these values. My CE is around 95-98%, I bet most of the electrons that are not recovered are going into these reactions.

Note that I am running things at really low capacities (~3-5 Wh/L), just to see if I can get long term cycling stability. I bet when I cycle to higher values I will get a lot of additional issues with hydrogen evolution (as I did with Zn-Br batteries as well).

Also, thanks for the correction about Goodenough, I have corrected my post to reflect that.
 
I finished preparation of my MnCl2 solution from MnSO4+CaCl2. The precipitate was hard to filter - as CaSO4 has quite fine particle size - but in the end I was able to obtain the solution as planned. The density of the solution was 1.428 g/cm3 (measured with a pycnometer) which gave me a final molal concentration of 4.8m. This should allow me to prepare electrolytes of varying concentration with much lower sulfate content.
 
Exp30, still running strong at 700 cycles. So far, this shows that placing the Swagelok cell vertically - with the anode at the bottom - leads to much lower dendrite formation. I will run this cell to 1000 cycles, then I will increase charging to 1.290mAh (1mAh/cm2) and see if we see dendrites show up or not. Let's see if we can make it to 1000 ?

1639072294236.png
 
Anode at the bottom? Eh? Dendrites fighting against Gravity? Wonder how it goes if you put on the anode on the top? Or is it just when it's sideways as gravity is pulling on the separator?
 
Anode at the bottom? Eh? Dendrites fighting against Gravity? Wonder how it goes if you put on the anode on the top? Or is it just when it's sideways as gravity is pulling on the separator?

Exp 29 put the anode on top, dendrites formed just as fast. I don't think this has anything to do with "dendrites fighting gravity" - the force of gravity is known to have very weak effects on dendrite formation - but just that the cathode carbon cloth doesn't let electrolyte flow past it very well (as it wicks quiet well) so putting the anode at the bottom allows the excess solution to flow out much more easily, as electrolyte can flow past a metal plate without problems.
 
I ran Exp30b for 32 cycles, no problems happened. I am now going to double the discharge capacity again, to 2.580 mAh, and see what happens. This will be Exp30c. Note all of these Exp30 runs are using the same battery - I haven't taken it apart - so I am basically now pushing it to failure, to see to what extent I have actually reduced dendrite formation and other issues.

1639176458007.png
 
That test didn't last for long, cell died pretty quickly (12 cycles). I am now doing a test (Exp31), with the newly prepared MnCl2 solution, to create a 2m MnCl2 + 4m ZnCl2. As expected, this time the electrolyte had zero precipitates upon preparation, due to the complete absence of sulfates. I have built a battery using this electrolyte, a Spectracarb carbon paper cathode, zinc anode and 2 layers of fiberglass as separators.

I am cycling to 0.5mAh and discharging to 0V at 10mA/cm2. Given the much lower volume, this cell will have at least 3x the energy density at the same capacity as the 6 separator layer cell.
 
I received my new carbon paper electrode materials so I decided to continue my tests with them.

I started testing the P50 carbon material, which is working fine.

I wanted to go to higher capacity values with smaller separator length, but noticed I was getting a lot of dendrites really quickly with the MnCl2+ZnCl2 solutions, without being able to get any higher capacity with high coulombic efficiency. When using these solutions, all I could get was around 0.5mAh, if I tried charging any higher, the CE would drop dramatically and dendrites would form very quickly. I didn't add any of these tests to the experiment log as I rarely got more than 5 cycles at these capacities.

Since charging to higher capacities involved going to voltages above the 2V mark for a significant period of time, I believe the higher charging generated significant amounts of Cl2, which caused very high losses of CE.

For this reason, I decided to go back to MnSO4 + ZnSO4 electrolytes. I prepared a 1.5m MnSO4 + 1.5m ZnSO4 electrolyte, saturated with K2SO4 using 15% vinegar/80% distilled water as the solvent. I then ran tests using 3 fiberglass separator layers, a Zn anode and a P50 carbon cathode. All charging and discharging being done at 5mA/cm2.

I noticed something interesting. When you start charging these batteries, you initially get to very high voltages quite quickly (getting above 2.2V before getting to 0.5mAh) but as you go to higher voltages your charging voltage drops on each cycle. If you for example charge to 1mAh and cycle there a few times, but then only charge/discharge the battery to a lower threshold, say 0.5mAh, you get much better CE values.

This is because there are mainly two sets of reactions that happen in this battery. Initially you have Mn2+ ions in solution and these are oxidized upon charging to form MnO2 in the cathode. When you have MnO2 formed, you can then store H+ and Zn2+ ions inside the MnO2 structure - provided it's the correct type of MnO2 - but only when you already have formed the MnO2 in the cathode. This is why many papers, especially those using sulfates, will have some sort of preconditioning process before the batteries are cycled.

These are the half reactions that happen upon initial charging of the battery:

Mn2+ + 2H2O ⇌ MnO2(s) + 4 H+ + 2 e− -1.224V
Zn2++ 2 e− ⇌ Zn(s) -0.7618V

These can only happen at around 1.98V, which is the thermodynamic potential of the full reaction at 25C. Most likely there is some over potential due to internal resistance, reason why we only reach the charging plateau of this reaction at around 2.1-2.2V. At least in the batteries you've seen I've built.

In the 1K cycle experiment we saw how initially the potential was very high and with time it decreased to where the average charging potential was much lower than the >1.98V needed to create MnO2. This is because, with time, we shift from a reaction where we create MnO2, to a reaction where we store Zn2+ and H+ ions inside MnO2 structures.

2021-12-16_10-03-43.png
In a Cl rich electrolyte, we can likely get high CE at low capacities from the start, because the MnO2 reaction is more reversible than in a chloride-free environment. We shift to ion storage slowly with time, but we seemingly cannot do this at anything but really low capacity values.

In a sulfate-only environment, the CE is lower if you do the same thing, but if you generate ample MnO2 by having a few cycles to much higher capacity than you intend to use - losing electrons intentionally to a less reversible MnO2 formation - you can then use the created MnO2 for ion storage.

After cycling the battery 5 times to 2.6mAh, I was then able to get really high CE values to 1mAh. Something that I just cannot do in chloride containing electrolytes. Exp31a shows 54 charge/discharge cycles to 1mAh, after this initial priming process:

1639667530008.png

I then decided to hold the battery at 2.2V for 1 hour, to see if I could create more MnO2 and charge the battery to 1.5mAh. I could then get a discharge capacity of 1.4mAh, with a CE of 94% and an EE of 59% on the first cycle. Given this geometry and this capacity, this battery would be at 16Wh/L. We can likely take it higher with the proper preconditioning, given these observations.

This likely shows that the easiest path for these batteries to give decent capacity is to use a sulfate-rich electrolyte and perform some pre-conditioning process so that we can generate all the MnO2 we need, then cycle the battery through an intercalation process, avoiding reaching potentials significantly above 2V for any long period of time.

The trick is that this preconditioning needs to be carried out quite slowly, if you try to do it fast, then the strong overpotential heavily stimulates dendrite formation and you short the battery quite fast. Saturating the electrolyte with K2SO4 seems to increase the conductivity and ion migration issues enough to avoid this issue, at least in the short term. You also need to make sure not to go to high current densities, so I was careful to stay at or below 5mA/cm2.

I will let you know how cycling at 1.5mAh goes. I will also do some further preconditioning after that and try to go to even higher capacity.
 
I ran this experiment for 43 cycles (Exp31b, results below), charging to 1.5mAh at 5mA/cm2. The CE remains high, EE remains high, capacity was 1.38mAh on average. Energy density at this capacity, given total device volume,, is 16Wh/L.

Exp31b.png

I have now done another single conditioning cycle, charging to 2.2V at 5mAh/cm, then setting the battery to 2.2V until the current dropped to 1mAh/cm2, then fully discharging. I am now cycling by charging to 2.5mAh at 5mA/cm2 and discharging to 0V.
 
So, I read an article that came out last year that went deep into the characterization of these Mn/Zn systems and the effect that mildly acidic conditions and pH buffers have on the formation of MnO2 oxides and their charging/discharging behavior (https://hal.archives-ouvertes.fr/hal-02566126/document).

Given this information, it seemed clear that I needed to use an acetate buffer instead of just acetic acid. Note that I used acetic acid because this is allegedly what was used in the first paper I followed about acetic acid used in Zn/MnO2 batteries. I therefore decided to prepare a buffer and test a battery with it in Exp32.

I prepared an electrolyte with 1m ZnSO4 + 0.5m MnSO4 in a 0.8M pH 5 buffer that I created beforehand by neutralizing vinegar with sodium hydroxide till I got a pH reading of 5 (using a freshly calibrated Apera PH60 meter).

The buffer makes a huge difference. Here are 4 cycles, running to only 0.5mAh at 5mA/cm2 but discharging only to 1V instead of 0V.

1639757283113.png

I am getting a CE > 92% by discharging to only 1V plus, the energy efficiency of the battery has now increased to >74%. Not only this, but my average discharge voltage has now increased to ~1.47V. The acetate buffer has made the chemistry way more reversible. Capacity and efficiency values were also stable for these few cycles.

I am now going to try higher capacity, see what we can get here ?
 
So far I've been able to cycle the battery 9 times at 1mAh. I discharged to 0.8V. The discharge capacity is up to 0.90mAh, with a discharge potential of 1.42V this gives me an energy density of around 13Wh/L. The CE, EE and capacity are all stable so far. The EE>70% and the CE>90%.

1639773754646.png
 
It is worth mentioning that many papers that work on Zn/MnO2 chemistry will do constant potential, instead of constant current charging. This means that they will hold the device at a given potential until a given mAh is accumulated and then discharge at constant current to a given threshold potential.

I have now implemented such charging within the python program I use to test the batteries, so I'll be able to cycle batteries this way and see what difference it makes. This mode of charging is also easier in practice, since just putting a battery across a fixed potential difference is easier than trying to control the exact current flowing through it.
 
I decided to try higher energy density and go with a much thinner separator. The fiberglass is too porous, so I still get shorts from stray carbon fibers from the carbon fiber papers going through. For this reason, I changed to 1 layer of W42 filter paper separator, which measures around 200 micrometers. This is 1/3rd the separator thickness previously used. I also changed the current density to 10mA/cm2. The electrolyte is 1m ZnSO4 + 1m MnSO4 in a pH5 0.8M sodium acetate buffer. This run is Exp33.

I am charging to 1mAh and discharging to 0.8V. The average discharge voltage is close to 1.5V and capacity is at close to 0.9mAh. So far the configuration is stable after 7 cycles. CE is > 89% with an EE > 76%.

1639932770015.png
Energy density under these conditions is 28Wh/L. If dendrites don't pose an issue, then this is likely to be a usable configuration in practical setups.
 
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Very impressive.

Would a switch from distilled water to deionized water help make sure the chemistry is what you expect?
 

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