New Alternative Electrolytes for Flooded LeadAcid Batteries.
Lets make them last longer, be safer and be
able to be reused second hand from a
depreciated state.
Sylph Dominic Hawkins – April 2016
version 1.4
This work is licensed under the Creative Commons Attribution-ShareAlike 4.0
International License. To view a copy of this license, visit
http://creativecommons.org/licenses/by-sa/4.0/
Let's get straight to the point, my interest in battery technology has always been to find,
develop, expand on and invent rechargeable batteries that are:
-
Non-Toxic to Human Health
Environmentally Sustainable to be built
Made with environmentally safe materials
Be simple to understand for easy access for all humans world wide
Be potentially source-able from waste streams in 1st, 2nd and 3rd world countries
Be totally financially accessible to all people across all income streams worldwide
While these goals are a great forward vision to continue heading towards, and I believe
that rechargeable batteries that provide all of these benefits can easily exist, I have
come across and experimented with some changes and tweaks to current existing
rechargeable battery chemistry that most of the world uses everyday, which could make
our use of them even more efficient, reliable, environmentally sound and lesstoxic, which, in the short and long term, is at least heading in the right direction.
Therefore this PDF is going to focus on the Wet-Cell Lead-Acid battery and its
chemistry.
Let me say at the outset that I am not particularly trained in Chemistry at all.
High-school HSC Chemistry years 11 + 12 was the last time I formally studied and was
trained in chemistry, but I have always been fascinated by it, and have continued
reading / learning / experimenting / observing / discussing chemistry ever since.
Everything that is in this PDF has come about from my curiosity with my work installing
and building off-grid solar power systems over the years, and working with different
battery types, and my musings, thoughts, ponderings and bursts of inspiration along the
way.
Anyone who puts their attention to this subject would surely have worked all of this out
just like I have, and while there is a fair amount of small scale experimenting and
articles than can be found searching on the internet about things like “Epsom salts and
lead-acid batteries”, “Sodium Sulphate and rejuvenating lead acid batteries”, etc..,
I have not been able to find any larger scale university / lab / Corporate funded research
paper (surprise surprise) looking in depth at what I have found, nor have I found any
databases or tables of experimented findings, calculations, chemistry equations, etc..,
just mostly a large number of armchair keyboard warriors who are happy to say that
things can't work, without actually exploring them themselves, and presenting the
chemistry behind it.
I also found lots of articles and home research done using the sulphate compounds as
“additives” in tiny quantities with the sulphuric acid, and I was theorising using only or
at least mostly the new sulphate compounds dissolved in distilled water, and to really
get at the base of the problem with using sulphuric acid in varying concentrations with
the lead battery environment.
For this reason, I got inspired to write this PDF, run workshops for people, put together
my theoretical chemical equations and create a basic table of my findings, so that other
people might be able to use it as a starting point in their own experimenting, and it
might serve as a reference material for home DIY battery experimenting. :)
------------------------------------------I have great thanks and appreciation for all the people who I have picked the brain of in
helping me refine my equations and understanding more chemical theory (thanks
Chris), assist me with suppliers (You rock Jim), safety notes, answer my questions,
provide me with huge amounts of free heavily sulphated deep-cycle batteries to run my
tests on (Thanks so much Richard and Kaggie), and generally be interested (or at least
keep eye contact with me) when making the mistake of asking me about what I'm up
to. :)
------------------------------------------This PDF is broken into two parts, aimed at different people.
Part 1 gives you a very quick rundown on the lead-acid battery chemistry and a simple
step-by-step instruction on how to change the usual Sulphuric Acid liquid electrolyte of a
standard flooded-lead-acid wet-cell battery to a liquid electrolyte solution of a more
environmentally friendly, non-toxic to humans, low impact chemical that anyone can
source easily in most countries world wide.
Part 2 of this PDF is where I will outline the specifics of my experiments and findings,
the chemical equations and tables of experimented outcomes, tables of potential
chemical alternatives and some comments. :)
In this section I also provide voltage charts to represent SOC% for the new alternative
sulphate solutions that I propose and compare them to traditional Sulphuric Acid based
wet-cell Lead-Acid batteries, Solubility and molarity charts and notes (to assist those
who want to try this for them selves in creating specific chemical compounds for their
batteries to work as best as they can), and a general summary and personal opinion on
the options and features of the different sulphate's, and how best we might utilise them
for different purposes.
Index:
Introduction and Outline
Index
1-2
3
Part 1 (Basic Into and Walkthroughs)
- Basic Chemistry Info For Lead Acid Batteries
- Restoring old Sulphated Batteries – A Step-by-step Walkthrough
- Make an Alternative Sulphate electrolyte for brand new/clean Lead-Acid
Batteries - A step-by-step Walkthrough (Sodium Sulphate/Bisulphate)
Part 2
4-7
8-12
13-14
(A deeper look at the Chemistry)
Alternative Liquid Electrolytes
15
Magnesium Sulphate / Bisulphate Theory
16
(First Charge Process Chemical Equations for conversion to Bisulphate)
17
(Ongoing Charge Process Chemical Equations using Bisulphate)
19
Thoughts on Other Alternative Sulphate Compound Options
20
A Deeper look at Solubility and Molarity
21
SO4 Stripping Potential of our Alternative Sulphate Compounds
22-24
The Role that Hydrogen (H+) is Playing
25
Bisulphate H+ Availability Chart & Sulphate/Bisulphate Comparison Chart
26
PDF Wrap up Summary + Recommendations
27-28
My Experiment Notes and Further Comments
29-32
Chart – Voltage of Compounds in Batteries representing SOC%
33
7x cell Series Arrangement for Alternative Sulphate Electrolyte Chart
34
End notes and Dedication
35
Part 1---Basic Chemistry Intro for Lead-Acid---A Basic introduction to Lead-Acid Batteries and a simplified outline of the alternative
liquid electrolytes that I have been experimenting with.
For more complex understandings of the chemistry, half equations and observations, see
Part 2.
Also note that this PDF focuses on Sulphation, tackles an outline of how it is created,
observations of it, and its chemical reversal. It is not going to focus on the other
potential reasons for lead-acid battery failure, such as Shedding, Corrosion, Shorting,
mechanical failure, etc..
- The Lead-Acid wet-cell battery is actually a very simple chemical redox (reduction
and oxidation) reaction.
The standard chemistry for Lead-Acid wet-cell batteries, using Sulphuric Acid, is:
CHARGED
Pb (Lead) + Pb02 (Lead Dioxide) + 2(H2SO4) (Sulphuric Acid)
<=>
DISCHARGED
(2)PbSO4 (Lead Sulphate) + (2)H2O (water) + 2volts eIn this full equation, when the traditional lead-acid battery is fully charged, the +
electrode (which I will call the Positrode4, as suggested by Craig Carmichael from
Turquoise Energy in the US) is a solid of Lead-Dioxide, which is a dark red/chocolate
brown colour.
The – electrode (negatrode4), is a solid paste of pure Lead, which is a dull grey pastemetal colour.
The liquid will be approximately half Sulphuric Acid. The pH of the solution will be low
(<1 pH) (very acidic), and the cell will output approx. 2volts.
Standard Molarity (Concentration) of the Sulphuric Acid is 4.5M / L (calculated to equal
roughly 441.27g/L of H2SO4)
As the battery discharges, the O2 (oxygen) from the Lead-Dioxide (PbO2) joins with the
free H+ (Hydrogen) from the Sulphuric Acid, to form H2O (water).
SO42- (sulphate ions, or “sulphate potential”) joins with the Pb (Lead metal), to form an
insoluble compound called Lead Sulphate (PbSO4) which builds up as a crystal layer on
both lead-plates. This Lead Sulphate layer is commonly seen as white/grey crystal build
up.
* Here-in lies the major problem with this version of the Lead-Acid battery. *
As the battery discharges, both + and – electrodes become plated with this LeadSulphate (PbSO4), and Lead-Sulphate is one of the most insoluble sulphate compounds.
This means that the Lead-Sulphate will not readily dissolve back into the solution of its'
own accord, to become ionically available, join with free Hydrogen in the solution, and
then be able to turn back into Sulphuric Acid (H2SO4) and Pb / PbO2 (lead and leaddioxide). Although the + positrode more easily releases the SO42- during charging as it
undergoes Oxidation and the O2 takes the place of the SO42- on the Pb (lead), the –
negatrode, undergoing Reduction during the charge cycle, does not as readily release
the SO42- from the PbSO42- layer. For this reason it is fairly common to see old lead-acid
batteries with heavily sulphated – negatrodes (negative plates), and relatively clean +
positrodes.
Over time the – negatrode surface area becomes completely covered in the PbSO4
crystals (a white/grey crystal), as well as a more internal version of the PbSO4 crystal
within the pores of the Lead (a black, non-light reflecting crystal).
This reduces the available plate area, which limits the ion flow and subsequent electron
production, as well as limiting the Hydrogen production on the negatrode -, and I
explore this idea more in detail later on in Part 2.
This is one of the major reasons for Lead-acid batteries having such a shallow
depth of total discharge available before you begin to damage the battery and
massively reduce it's total number of possible life-cycles. (That is, the total amount of
power you can draw out of the battery, measured in Amps, before you must stop or risk
damaging the battery).
The standard maximum available depth of discharge (called DOD) for a wet-cell leadacid battery is 30% of total capacity (rated in Amp hours (Ah) ),
and standard maximum DOD for so-called marine or deep-cycle lead acid batteries is
approx. 50%.
This means, quite literally, that if you have a 12v 100Ah lead acid battery (this is a fairly
common size / capacity of battery for 4WD's and cars wanting a second backup battery,
camping, battery for solar power storage, etc..)
you can only draw 50% of its total capacity, that is, in this example, 50Amp's, out of the
battery before you must stop, or risk battery damage. (Risk heavy plate sulphation and
possibly, to a much lesser extent, shedding of the active material from the positrode +)
For a standard car battery, the DOD is only 30%, so for a 12v 50Ah (amp hour) battery,
you can only draw 30%, or 15Amps, total or risk damage to the battery when using
Sulphuric Acid.
In reality, once you draw beyond 50% DOD from a standard lead-acid battery, it's cell
voltage drops so much, that it falls below the required voltage to drive the starter motor,
or other 12v based appliances, especially devices that draw high amps.
It is for this reason that we need to hugely over-size battery storage for off-grid houses,
to compensate for the fact that we can only draw out, lets say, 30% of the total capacity
from the battery bank. I am also ignoring the reality that in an off-grid house example,
we also need to plan for, usually, a minimum of 3 days of backup battery draw (called
autonomy) to compensate for many days of no sun or cloudy weather during winter, as
this PDF is not specifically dealing with off-grid solar setups, I will just use our simple
example as it could be applied to any lead-acid battery that we assume will be fully
charged the following day/use cycle.
---------------------------------------------------Moving onwards with the main purpose of this PDF,
As we attempt to draw more power out of the battery, it discharges further, which puts
more lead-sulphate onto the + and – plates, which will not totally dissolve back into the
solution at the next charge, and the sulphation process begins.
Ironically, when a standard 12v lead-acid battery has its' voltage drop below 12v, this is
usually a sign that more than 50% of its capacity has been drained (since a normal 12v
lead acid battery will actually have a voltage of around 12.7v at full 100% State Of
Charge), and it will begin to “sulphate” more rapidly, which is a term that describes
when the white/grey coloured lead-sulphate crystals begin to build up on the surface of
the metallic lead + and – plates. Research data and test results that I have read on
lead-sulphate crystal formation actually suggests that the PbSO4 crystal begins to
“grow” or be precipitated onto the lead electrodes as soon as the total charge voltage
drops from its float voltage of 13.3v (2.21v per cell)
The more this happens, the less “charge” the lead-acid battery will be able to hold,
relative to it's rated capacity, and eventually, it will not be able to sustain the amount of
charge it needs to keep it's voltage above 11v, in order to run things like starter-motors
in cars, etc.., this is because the Amp-hour capacity of the battery is directly
proportionate to the surface area of the lead-plates coming into contact with
the electrolyte.
When some of that surface area is converted into PbSO4 crystal, it stops that portion of
the plate from interacting with the electrolyte, and the surface area of the battery has
been reduced, thus, the amp-hour capacity has been reduced also.
Ironically, the lead-sulphate compound is relatively non-conductive, and it actually stops
the liquid electrolyte (the H2SO4 sulphuric acid) from making contact with the metallic
lead underneath the layer, effectively stoping the reaction, but also, conveniently,
insulating the underlying Lead from further conversion into PbSO4, so
underneath that white/grey/black crystalline coating should be plenty of pure
pasted Lead, just waiting to be put back to work in our little
oxidation/reduction reaction as a battery.
Just for ease of reference, these are the standard average voltages that I have come
across in regard to State of Charge (SOC %) of lead-acid batteries using sulphuric acid
electrolyte. This chart is based on original 12v Flooded Wet cell lead-acid batteries found
here: (http://modernsurvivalblog.com/alternative-energy/battery-state-of-charge-chart)
and cross referenced with my own data from my own battery experiments, and I found
this chart to be surprisingly accurate, and so have just used the original 12v figures
from the chart, and divided them by 6 in order to get the “per cell” voltages for the
various SOC %.
SOC % Per Cell (2v X6 cells (Standard
nominal)
12v Batt)
100
2.12v
12.73v
90
2.10v
12.62
80
2.08v
12.5
70
2.06v
12.37
60
2.04v
12.24
50
2.01v
12.1
40
1.99v
11.96
30
1.96v
11.81
20
1.94v
11.66
10
1.91v
11.51
Notice that there is an approximate 200mV (thats 0.2v) difference between 100%
charged and 10% charged (so almost drained “flat” to a 90% DOD)
I have found in my experimenting that this figure remains consistent for the different
sulphate electrolyte chemistries that I am testing, even if the cell resting voltage is less
than when using Sulphuric Acid, say for example, how in Sodium Sulphate / Bisulphate
based electrolytes, the lead electrodes exhibit a 1.9v potential across the terminals,
instead of the usual 2.1v (roughly), and the figures are similar to the chart above with
the new relative voltage, so a fully charged NaSO4 based lead-acid battery sits around
1.9v, and it could be said to be fully discharged at around 1.6/1.7v (about 0.2v/0.3v
difference). I will discuss this further later on in regards to compensating for this and
sizing battery systems to work with the different overall voltage.
---------------------------------Now, onto some fun news! :)
The Lead-Acid battery is simply a mix of some very simple chemicals, performing some
very simple, observable and reversible reactions.
Just like in all chemistry, there are rules that determine what possible chemical
compounds may be formed under certain conditions between different elements.
The PbSO4 (Lead Sulphate) that becomes hardened on the surface of the electrodes is
still just a chemical compound, and can be reversed chemically using the right methods.
So, this PDF will present a viable alternative liquid electrolyte which Wet-Cell
Lead-Acid batteries can be filled with, and I will outline one variation of this
alternative liquid in order to de-sulphate old batteries and return the plates to
near-new condition, as well as outline a different variation of this alternative liquid
which, when used in brand new Lead-Acid batteries, should reduce or potentially even
stop the PbSO4 sulphation problem from occurring at all over the entire life of the
battery.
(I say potentially, because currently what I have are theories of the chemistry and
multiple short-term experiment case studies, including on-site case studies of batteries
in solar-powered off-grid homes currently being used daily by families, as well as my
own personal sheds of lead-acid batteries in varying states of de-sulphation and
operation, but no super long term observations (like, lets say, over 20 years :).
As of March 2016, the longest going lead-acid wet cell battery that I have converted to
use MgSO4, (Magnesium Sulphate / Bisulphate) is 1.5 years. It is performing really well,
but I personally would like to see test samples of the different chemistries over a 5, 10,
15, 20+ year study. I also have on-site case studies of off-grid battery systems, being
used in normal daily operation, at 2 locations around Victoria, Australia,
taking battery and cell voltages, shunt-based amp-outputs and performance under load
data for me to add to my study.
I am currently running hands-on workshops taking groups of people through the process
of converting lead-acid batteries to use these alternative sulphate liquids, with the aim
of spreading the information as well as building up more on-site case studies of use to
collaborate my data. If you are interested in running one of these workshops, usually for
people who are building off-grid houses or restoring old battery sets, and would like to
help add data to this study, feel free to get in contact with me using my email address
at the end of this PDF.
Restoring old, Sulphated Batteries –
Step by Step Walkthrough
To get a full background of the chemistry and my theory of how this all works, see Part
2, but for ease sake for people wanting to just get started with converting their
sulphated lead-acid batteries, here is a step-by-step run down of the process of
emptying second-hand batteries and refilling them with a simple, currently tested to be
working, alternative liquid electrolyte. (For this simple example we will make up a
Sodium Bisulphate electrolyte, but many others can be used, see Part 2 for details).
For people looking for the walkthrough for preparing the Bisulphate compounds to use in
brand-new Lead Acid batteries, it follows after this section.
-------------------------------------------------------------------------------------------A list
–
–
–
–
–
–
–
–
–
–
–
–
of required ingredients + tools is:
Sodium Sulphate (powder)
Distilled water
Large tub with sealable lid to house all the old battery acid
Large tub to mix up the new alternative electrolyte
Some kind of pouring device to pour the new electrolyte back into the cells
Funnel's to help pour the new electrolyte back into the cells
Some source of charging current (solar panels without a controller,
or a non-smart battery charger)
Goggles / lab glasses
Chemical resistant gloves
Face mask / Vapour mask
Wear OLD clothes than you are happy to get holes in
A friend to help you lift heavy batteries and assist in case of medical needs,
splashes, first aid, etc..
Note: Sulphuric Acid concentrate from used old lead acid batteries is still very
acidic and highly dangerous. Take all steps to minimise exposure, and ensure
that absolutely NO acid can possibly enter the eyes or get around goggles in
case of splashing as Sulphuric Acid splashes can cause permanent Blindness.
Sulphuric Acid is a skin irritant and will cause chemical burns on contact.
The fumes of sulphuric acid can cause respiratory problems if breathed in.
Before starting to work with sulphuric acid, even if just for the small period of time that
we will be converting the batteries, it is important to download and read, and keep a
copy of the MSDS (Material Safety Data Sheet) with you during the conversion.
A copy of an MSDS for H2SO4 can be found here:
(http://avogadro.chem.iastate.edu/MSDS/H2SO4.htm)
Step 1 - Make up a mix of your new electrolyte solution.
For this example, let's use Sodium Sulphate (Na2SO4) and Distilled water (or Deionised
water). We will make up a 1.37M concentration solution of the electrolyte. (This is the
maximum we can achieve based on the solubility of Na2SO4 in water at room
temperature of 20*c, see more on this in my solubility v's molarity chart in Part 2)
To do this, refer to the periodic table of the elements to calculate the molar mass of
each element in the chemical to get a total molar mass (weight) of sodium sulphate that
is required to be added to the distilled water to arrive at a concentration of 1.37M (M
means mol/L).
For Na2SO4, it has a molar mass of 142.04g/mol, so this means that a 1M concentration
will be to dissolve 142.04g per Litre of H2O, and thusly, if we want a 1.37M
concentration, we will need: 142.04g x 1.37-mol = 194.6g / L H2O
Here is where we need to know the total number of Litres of needed electrolyte for our
battery.
Check the battery label / google the model number / or just measure out the old battery
acid as you remove it. An easier alternative is also to empty the battery of all old acid,
then get a 1L measuring jug and pour tap water back into the cells to measure the total
amount needed per cell. Then, after getting this figure, dump out the tap water into the
same tub as the old acid and empty the battery once again.
Let's say that for our example we need a total of 5 Litres electrolyte to refill the battery
for our solar battery cell (remember that the electrolyte needs to well cover above the
top of the lead plates inside). (Car batteries will take much less electrolyte)
Now we know we need 5 Litres total, and we want a 1.37M concentration,
so finally, in order to find out exactly how much Sodium Sulphate we need to mix into
our 5L Distilled Water, lets multiply 5L x 194.6g = 973g Sodium Sulphate powder
Step 2 – Make sure you are wearing chemical resistant gloves, goggles and
facemask, as well as old clothes that can get destroyed, overalls, etc... Then,
open up each of the second-hand cells, assuming they are wet-cell flooded lead-acid
type, and if they are VRLA (valve regulated lead acid) and have no obvious caps, then
you will need to drill a hole in the top.
Step 3 – Once opened, each cell will need to be fully emptied out into a tub. This tub
will be full of Battery Acid (sulphuric acid), no doubt with flakes of lead floating in it from
the battery sludge. Make sure you use a tub that is deep enough to contain any splashes
from the dumping of the battery acid, and have a friend there to help you, 2 people
holding a battery upside down and dumping out the liquid in it is much safer and more
controlled than one individual struggling by themselves. Lead acid batteries are very
heavy. No need to wash out the battery cell unless you are trying to remove the lead
sludge from the bottom of the battery (unless you can get in around the sides of the
battery, don't bother trying)
Step 4 – Using plastic funnels and a pouring bucket / jug, carefully pour the new
electrolyte back into the battery cells, filling each cell one at a time. Ensure that the
liquid is well above the lead plates, but not overflowing out of the battery. Replace the
battery caps, or if you drilled into a VRLA battery, fill the hole with rubber/cork/silicone.
Step 5 – Wash down all the jugs / tubs / utensils used, with water
Final Step 6 – Next is to put the battery cells back on charge. Depending on the depth
of sulphation, this step will vary, but generally, the first thing to try is to charge up the
battery as you normally would with your usual battery charger or solar setup.
Sometimes the battery has such a high internal resistance due to the sulphation that it
will need to trickle charge slowly over many days to allow the current (amps) to start
flowing into the battery. (This process will also be electrolysing the water into Hydrogen
and Oxygen, forming Hydrogen at the negative – electrode, and Oxygen at the positive
+ electrode. The Hydrogen (H+) formed will act with the Sodium Sulphate (Na2SO4) and
the extra Sulphate (SO42-) trapped as Lead Sulphate on the Lead plates, this will form a
new compound called Sodium Bisulphate (Or also known as Sodium Hydrogen
Sulphate), with the chemical formula of NaHSO4.
You will be able to see the lead-sulphate coming off firstly from the negatrode, and
slowly cleaning out from there throughout the battery plates over numerous charge +
discharge cycles. This phenomena is fun to observe in clear acrylic case batteries, where
you can clearly see the lead of the negatrode become clean matte grey-coloured metal
again, looking almost brand new.
----------------------------------------------------------------------------------------------Note: This process takes time to work, on average, a few weeks to a month of regular
charge / discharge cycles to fully strip the sulphate off of the plates.
Depending on the “depth of sulphation”, that is, how much lead-sulphate there is on the
plates, this use of Sodium Sulphate, as it converts to Sodium Bisulphate, may not be
enough to fully strip the sulphate away from the lead, and this process will need to be
repeated, but not more than 3 times. This is all to do with the total solubility of our
alternative metal sulphates (so, in this example, Sodium Sulphate), and the maximum
possible molar concentrations.
I go into this in more detail in Part 2, and offer a database chart of the different
sulphates, and the theoretical maximum amounts of SO42- that the compounds can strip
from the lead.
For usual light to mild-sulphation, the Sodium Sulphate will suffice.
For more heavy sulphation, this process will not be able to convert enough of the
sulphate to the bisulphate to fully strip clean the plates. See my notes and database in
Part 2 about this for further info and deeper understanding of the chemical process as I
theorise it to be.
So long as the battery will accept some charge and come up to float voltage (so 2.1v for
a single cell, or 12.5v for a whole 12v battery) you will notice a constant but slow
increasing in the depth of capacity and “omph” of the battery over many weeks as the
Sodium ions slowly do their work of stripping the sulphate ions away from the Lead
plates, and forming Sodium Bisulphate in the process.
If after a few days the battery will still not accept charge, the battery will need to have
some funky voltage tricks applied to it to kick it into action.
First is to try using a higher voltage than the battery. Usually, for a 2v lead-acid cell,
hitting it with 12v straight at about 1 or 2amps will do the trick, (so for a 12v battery,
hit it with a 24v source load at about 1 or 2amps). Do this for a few hours at a time,
watching the temperature of the cell (it will heat up considerably). If the battery
becomes hot to touch, or, if you are measuring it, rises to over 50*c, then you need to
stop the charge input, let the battery cool, and try again when it has cooled down.
-----------------What if after a few weeks, it still hasn't helped?--------------------* If after a day of charging at this higher voltage rate, it still will not begin to accept
charge, then an alternative option (which I am still experimenting with, so take it with
some caution) is to totally drain the battery (attach a low power drain such as a resistor
or a low power light bulb (car filament bulb) directly across the + and – terminals.
Once it is completely drained (the voltage reads about 0v on your voltmeter),
then reverse the polarity of the charging cables. That is, from your charger or solar
system, attach the + to the blue/black/negative terminal on the battery, and attach the
– to the red/positive terminal on the battery.
This will cause the sulphate crystals on the – Negatrode, and some of the Oxide layer on
the Positrode + plates to return to the liquid solution, and slowly an oxide layer will form
on the other plate (what was marked as the – negative plate, which will now become
the + positive plate, and will turn from a dull grey colour to a dull red/chocolate brown
colour).
From there, you would then discharge the cell again with a controlled load (like a
resistor or an incandescent light bulb), and once it has discharged to flat again, swap
the polarity of the charging cables once more, and charge it up the “correct” way round.
The idea of this is to cause the sulphate to come away from the lead on the negative
electrode more easily by giving the electrode a positive + potential, and getting Oxygen
from the water to take the place of the sulphate on that terminal.
Then, ideally, after discharging once more, and then by charging the battery up again
the “correct” way around, the Oxygen from the Lead Dioxide (PbO2) will be reduced
from the negative electrode, and now it will be pure lead plate again, since the sulphate
got removed by the oxygen under the reverse charge.
This can bring back heavily sulphated cells from the dead.
(I may comment more on this in my personal comments and experiments at the end of
Part 2)
Remember that this chemical process will need time to work.
Expect to see the battery slowly increase in depth of capacity over many weeks
as it goes through full charge + discharge cycles, especially given that,
according to the chemical equations, the majority of the action of the Sodium
ions becoming free and stripping Sulphate ions away from the Lead, bonding
with hydrogen in the water, most probably formed at the negatrode –, and
forming Sodium Bisulphate (basically just Sodium Sulphate with a Hydrogen
atom bonded to it, after losing one of it's Sodium ions) happens during both
DISCHARGE and CHARGE processes.
Ideally, the presence of the Sodium in the electrolyte will protect the lead plates from
sulphation during discharge, and allow you to deeply discharge the cell to aid in and
speed up this “stripping” process, which you could not normally do in a Sulphuric Acid
based electrolyte, since deeply discharging the cell, would cause the lead plates to
sulphate very effectively. :)
The process of de-sulphating old used lead-acid batteries using the sulphate solutions
can be a tricky one.
On one hand, we are introducing a new sulphate-based electrolyte into the battery to
allow the redox reaction to occur as normal, but to keep the lead-plates clean during the
discharge process, so as to stop sulphation from occurring, and hopefully extend the life
of the battery, keeping as much of the lead-plate surface area as possible being able to
interact with the electrolyte.
On the other hand, the old, sulphated battery has a proportion of sulphate on the lead
plates in the form of Lead-Sulphate, so the total amount of “stripping” power that our
alternative solutions will be able to do will be based on the relative amounts of
concentration of both “how much sulphate is on the lead plates” and “how much
total sulphate could the new electrolyte strip during the formation of the Bisulphate”. I explore this in great detail in part 2, however for now, a general rule can
be to look at the total possible sulphation of the battery.
* If the battery is near new, or has been on float charge for its' life, or has a small
degree of sulphation on the plates, but not substantially, then I would expect a 1.37M
solution of our Sodium Sulphate electrolyte that we made above, to be capable of fully
stripping the sulphate from the plates, converting it into Sodium Bisulphate, and then
arriving at a point of equilibrium regarding the amount of Sodium ions in the solution
compared to the total Sulphate ions in solution, and hopefully ensuring that in future
discharges of the battery there is enough Sodium ions in solution ready to pull the
Sulphate away from the lead during discharge, if it even forms at all.
* If the battery is very heavily sulphated, and you can clearly see that the majority of
the plates (especially the negatrode) is covered in lead-sulphate crystals, then we need
to do some more playing around.
This is because the concentration of SO42- ions (Sulphate ions) will be substantially high.
Assuming, for over-estimation sake, that an average lead-acid battery uses a 5M
concentration of Sulphuric Acid, then we can calculate that 1M of SO4 = 96.06g/L,
so 5M = 96.06g x 5 = 480.3g/L.
This means that, theoretically, there will be up to 480.3g/L of SO4 in the battery.
If even half of this amount (e.g. 2.5M) joins with the Lead to form Lead-Sulphate, then
we could possibly have up to 240.15g/L SO4 to try and strip away from the plates, as
well as lock away in the solution as the Bisulphate.
I go into this in more detail in Part 2, but for an example, @ temperatures of 20*c,
1.37M of Sodium Sulphate, as it converts to Sodium Bisulphate in a sulphated battery,
will be able to strip up to 1.37M extra of SO4 that is on the lead plates.
You can clearly see that in this scenario, we do not have enough stripping potential from
the Sodium at a 1.37M concentration, to strip and bind with all of the potential
2.5M SO4 that could be on the plates.
If we have 2.5M (240.15g/L SO4) on the lead-plates, and our solution can strip a max. of
1.37M (131.6g/L of SO4), then we potentially still have 108.55g/L of SO4 on the Leadplates.
This accounts for some of the observations people make, using varying
concentrations of sulphate solutions, on batteries with varied levels of sulphation, and
have varied results. One simple solution would be to dump out and replace the sodium
sulphate solution each time it has fully done its' job of stripping away the SO4, and we
could assume that after 3 electrolyte changes each time we allow the electrolyte to strip
away its' max. amount of SO4, then we could possibly remove all the lead-sulphate. I
will go into more detail on this in Part 2.
Make an Alternative Sulphate Electrolyte for
brand new / clean Lead-Acid Batteries –
Step by Step Walkthrough
If you are looking at using the alternative sulphate solutions as an electrolyte to fill
brand new Lead-Acid batteries, or if you have gotten your hands on a battery that has
no sulphation on the lead-plates, then our electrolyte needs to be made in a slightly
different way, because we need to create the fully balanced Bisulphate solution first,
since the battery will not be able to provide any extra SO4 to help the creation process.
To get the battery to operate correctly, we need to have a solution that provides H+ ions
which are available to strip the O2 oxygen from the Positrode and release electrons in the
process as the Lead on the positrode + is reduced from Pb4+ to Pb2+. We also want to
ensure that there is a balance of the free metal ions that we are introducing, (for
example, Na+ Sodium in our Sodium Sulphate electrolyte) with the amount of SO4 in the
electrolyte, so that ideally, the Sodium will always strip the SO4 away from the lead
every time during discharge (to find out why SO4 might still be bonding with the Lead
into PbSO4 even when using a new sulphate solution, and the step-by-step process of
how the Na+ works during discharge, see my theoretical equations in Part 2).
To make a Bisulphate solution, we can simply look at the chemistry:
Na2SO4 + H2SO4 = 2NaHSO4
In this example, the mixing is simple, as it is a 1:1 ratio.
So, if we make a 1Litre, 1.37M solution of Sodium Sulphate (remember, this is the
maximum concentration we can get from the Sodium Sulphate based on its' solubility in
water at 20*c room temp), and a 1Litre, 1.37M solution of Sulphuric Acid, then we can
simply mix them together to form a 2Litre, 1.37M solution of Sodium Bisulphate. (Since
molar concentrations are based on the atomic weight of the individual elements in a
compound, and through this reaction we have not changed the total number of elements
we are dealing with, only their combinations, the molarity will be unchanged, however,
the solubility of Sodium Bisulphate I greater than Sodium Sulphate, more on that later)
We get a compound that has an equilibrium of Sodium ions to SO4 ions, and will provide
the H+ to do the reducing of the Oxygen on the Positrode +.
We could mix the solutions in this way:
Making a 1.37M solution of Sodium Sulphate would be to dissolve 194.6g of Sodium
Sulphate into 1-Litre of Distilled/Deionised water,
and we will mix this with a 1.37M solution of Sulphuric Acid, which we can make by
dissolving 134.34g of H2SO4 in 1-Litre of Distilled/Deionised water.
Obviously, this will give us 2 Litres of electrolyte, and depending on the size of your
batteries, you will need to scale up the amounts of reagents to make however many
Litres you need, but still at a 1.37M concentration for our use of Sodium
Sulphate/Bisulphate.
Now that we have obtained our starting liquid electrolyte in the concentration and
volume that we need to completely fill our new/clean batteries, we can follow the Steps
2-6 from above (listed under the previous section, “Restoring old Sulphated Batteries –
Step by Step Walkthrough”) to empty out the new batteries (if they are already filled
with acid, otherwise skip this step) and refill them with our new 1.37M Sodium
Bisulphate liquid compound.
The final step is to charge the batteries in the correct polarity.
A clean/brand new lead acid battery, filled with this alternative sulphate solution, should
begin working straight away using normal chargers / solar systems with normal solar
controllers and not require any messing around to get the battery to work as required.
Note: While running the clean/new lead-acid battery with this Sodium Bisulphate
alternative electrolyte, it should ensure that no lead-sulphate ever forms onto the
lead plates, even during deep discharge.
However, because we have effectively removed some of the free Hydrogen from the
electrolyte (from H2SO4 to NaHSO4) and lowered its' total acidic concentration from
5M to our current max. solubility of Sodium Sulphate, yielding us a max
concentration of 1.37M, we will have reduced both the voltage potential across the
terminals as well as the total amp-hour capacity of the battery cell.
I go into great detail about this at the end of Part 2, under the chapters “Solubility
and Molarity” and “The Role That Hydrogen Plays”. I present a chart looking at solubility,
molar concentrations, and the different stripping power of sulphate compounds to
theorise what will work the best for us, since what we want is a combination of highmolar concentrations of H+ ions, but also an equilibrium of free metal ions in the
sulphate solution as well. This chart helps us to see that we can achieve higher
Hydrogen concentrations in the different Bisulphate compounds, but also balances them
with the other features of Pro's + Con's of some of the different various sulphate
compounds that we could use to replace the standard Sulphuric Acid Electrolyte.
As a final close to this Part 1, and to give some food for thought to anyone who just
read the previous paragraphs and thought “But hey, If I had up to 5x concentration of
Hydrogen with sulphuric acid, and now I only have 1.37x concentration of Hydrogen
with Sodium Bisulphate, is this even worth it since now I might only have 27% of the
original Amp-Hours of the battery. Would my 12.7v 100Ah batt, suddenly only become a
11.5v 27Ah batt?. We need to consider how much of the original 100Ah was available to
us under the Sulphuric Acid model. If we take the standard DOD of 30% for max. long
life cycles, then this effectively means that we can only draw 30% of that 100Ah
anyway, making our Sulphuric Acid 12v 100Ah battery, really, a 12v 30Ah battery for all
intents and purposes. Now when we look at them side by side it makes more sense.
(6x cell H2SO4) @ 12.7v 30Ah (381Watt-Hours) compared to (6x cell NaHSO4) @ 11.5v
27Ah (310.5Watt-Hours). And so for loosing 3Amp-hours in capacity, we have gained
the ability to discharge our batteries to near flat without causing any sulphation
problems, extended the life cycles of our batteries by up to 5x times (so while our
Sulphuric Acid deep cycle batt might have a life cycle count that runs up to an average
of 7 years, our Sodium Bisulphate deep cycle batt could live for up to 30+ years life
span). To compensate for the voltage drop, we also just add another cell in series, of the
same original Ah capacity (100Ah), to give us a (7x cell NaHSO4) @ 13.41v 27Ah
(362.07Watt-Hours). Now we are closer to the original Watt-Hour capability of the
Sulphuric Acid batt, but without the killing problems, and a huge long life span! Bonus.
Part 2----Alternative Liquid Electrolytes--------It suddenly dawned in me one day, in 2013, that the Redox reaction with the lead-acid
battery was using the H+ ion from the sulphuric acid (H2SO4) to reduce the PbO2
Lead Dioxide on the + positrode, and that was the source of
chemical-->electrical energy for the battery as the Lead on the positrode + was being
reduced from Pb4+ to Pb2+ (A change in oxidation state of the Lead), and requiring 2
electrons to do so. The 2 electrons come from the Negatrode -, where the solid Lead
Pb(s), is being oxidised (becoming Pb2+) and releasing 2 electrons, these electrons race
through our wire connecting the two poles, and we ultimately tap this energy off for our
needs during it's journey. The sulphate ion (SO42-) itself, and the ensuing PbSO4
Lead Sulphate was potentially a by-product of the reaction, and that the
production of Lead Sulphate on the plates was not necessary for the Battery to
work like it does.
This got me thinking about other possible electrolytes which only provided more
Hydrogen (H2) and Oxygen (O2) to the mix, such a Alkaline Hydroxides (Sodium
Hydroxide NaOH, Potassium Hydroxide KOH). However, initial testing of using
Hydroxides at high concentrations (5M), as a de-sulphating agent did not work, and
instead, reduced the PbO2, into a number of other undesirable oxidation states, such as
PbO, and decompositions into the potential Pb(OH)2 (which might not even exist)
Going back to the Activity Series of metals, one can discover a well tested and tried
chart of most metals and their reactivity compared to each other, listed in a table of
least reactive to most reactive.
The table gives us information to help estimate the outcomes of a reaction.
On the table, Pb (Lead) is quite low in reactivity, and above it, in order of lesser
reactive to more reactive, are some metals such as:
Tin (Sn), Nickel (Ni), Iron (Fe), Zinc (Zn), Manganese (Mn), Titanium (Ti),
Aluminium (Al), Magnesium (Mg), Calcium (Ca), Lithium (Li), Sodium (Na) and
Potassium (K)
This tells us that under normal conditions, these metals will not plate onto Lead, and are
more likely to lose electrons easily to form ions (floating metal potentials ready to bond
with other non-metal elements). They are stronger reducing agents (that is, they donate
their electrons more easily) and they oxidise (or corrode) more easily.
Looking at this table, it became obvious that I could experiment with any one
of these metals in a sulphate solution as the electrolyte, and I was fairly certain
that the metal ion in the sulphate compound would not come out of the liquid, become
solid metal, and cover the lead plates (thus ruining my little plan), unlike a metal that is
less reactive than lead, such as Copper, which will easily plate onto the lead, and not do
the job that I have for it later on during the discharge process. :)
Some of the metals though, simply create sulphate solutions that will not work, for
example: Calcium sulphate is even more insoluble in Lead sulphate, and so will not form
a liquid electrolyte for us to use, and Manganese sulphate readily reacts with PbO2 to
form compounds such as permanganates (MnO4), and thus not help us by staying in the
sulphate form during a full charged state.
Later in the PDF, at the end of Part 2, is a chart where I explore the outcomes of many
of these sulphate solutions to find ones that will work.
It also just so happens that metal sulphates are really, really common
chemicals.
Some are sold at garden centres for fertilising the ground and re-mineralising soils.
Some can be found at pool shops, some are bath salts, Calcium sulphate, for example,
is just plaster-of-paris (though we cant use Calcium Sulphate for our alternative
electrolytes due to its' total lack of solubility), but the majority of these sulphates are
low in toxicity or non-toxic to humans and low in environmental impact (some actually
helping the soil as a fertiliser, such as Potassium Sulphate)
------------------MAGNESIUM SULPHATE / BISULPHATE THEORY-------------------The first one to try was going to be the easiest one I could get my hands on locally.
So I drove to the supermarket and bought myself 1 Kg of the classically common and
available compound called Magnesium Sulphate, or “Epsom Salts”.
It's powdered chemical formula is MgSO4.7H2O, and when dissolved in water it forms
ions of Mg2+ and So42-.
I began at once and made up a 1M (unit of concentration) solution of MgSO4.7H2O (This
is called Magnesium Sulphate Pentahydrate, which is the common Epsom Salt form, and
thus has a higher molecular weight than the theoretical MgSO4 by itself)
(246.5g / Ltr H2O), and made 4 litres of the stuff. Emptied out the sulphuric acid inside
one of the junk 100amp 12v deep-cycle batteries that I had lying around, filled it
instead with the MgSO4 solution, and popped on to the charge from the solar panels.
This particular battery came flying back to life, with some altered characteristics, but a
lower total voltage.
Some exploring of the battery led me believe that one of the cells had discharged way
back when, and then had it's polarity reversed, and was pulling down the total voltage,
so even though the battery had not returned to 12v, things were looking promising.
-------------------------------------I replicated the experiment again on a different 12v 100a battery, which I knew was in
better condition, and the battery worked better than it had for years.
It produced approx. 11.5v total resting voltage when fully charged, and each cell
showed 1.91v full, instead of the usual Lead-Acid 2.11v (hence 6x cells in series give an
approx. 12.7v total charge)
It packed a punch with discharge amperage. It provided a much smaller spark (now
theorised to be due to the reduction of Hydrogen concentration and overall acidity), but
would climb to discharges of 30amp very easily.
I knew that, in the lead-acid chemistry, a fully charged battery has all the sulphate
(SO42-) in the liquid solution, and not on the plates, so I figured that this battery would,
theoretically, have this net charged equation state at full.
Pb(s) + PbO2 (s) + Mg2+ + 2SO42- + 3H+ + OHThen I thought that potentially a few things might occur during discharge, and I was
hoping that my plan would work to put the Magnesium ions to work on the sulphate
ions.
Remember, during discharge of a standard Lead-Acid battery PbSO 4 (lead
sulphate) is formed on both the – and + electrodes, and it stays there usually
until the battery is recharged.
My idea was to use the free metal ions (in this case, Mg2+) which are more reactive than
Pb Lead, and hope that the magnesium ions would bond with the Sulphate ions (SO42-)
to form a compound with the Mg2+ again (which might be highly soluble and would
dissolve back into the liquid), rather than the problematic PbSO4.
Here is how I thought that might occur, starting from the “first” CHARGE reaction
when using Magnesium Sulphate (Epsom Salts) in a battery that already has some
sulphation. These following equations are called half-equations, and are to look at what
might happen at both electrodes of the battery:
FIRST CHARGE OF SULPHATED BATTERY USING MAGNESIUM SULPHATE
@ - Negatrode:
(Discharged State)
(Charged State)
PbSO4(s) + MgSO4 + 2H+ + 2e- ==> Pb(s) + Mg(HSO4)2
@ + Positrode:
(Discharged State)
(Charged State)
PbSO4(s) + 2H2O ==> PbO2 + 3H+ + HSO4- + 2eTHEN – Possibly spontaneously react with other Magnesium Sulphate in solution:
3H+ + HSO4- + MgSO4 ==> Mg(HSO4)2 + 2H+
Note: I am assuming the H+ presence at the Negatrode during the CHARGE cycle, since
we are inputting electrical current into the battery at a voltage of 2.3v (or at least,
greater than 1.5v, which is the lowest voltage required to begin water electrolysis),
providing us with free Hydrogen from the water electrolysis at the Negatrode for as long
as there is a >1.5v charge across the terminals.
At first I thought that maybe the MgSO4 would simply be replaced intact during either
charge or discharge, and so I expected that a fully charged cell, or a fully discharged cell
would have a pH value of 6 or 7 (the pH of MgSO4), but when I tested the liquid when
fully charged it showed a pH of approx. 1 - 2, and when I tested the liquid fully
discharged (0.1v) it showed a pH of approx. 1 - 2 also. (Using pH indicator “Wide Range
Indicator”, colour was vibrant red/pink)
This perplexed me until I came across the chemistry idea of the “Bisulphate” compound.
This is where the metal ion and the sulphate ion bond together along with a Hydrogen
ion, and my equations now subsequently use the Bisulphate compound theory.
While I have been able to find good data on the specs of Sodium Bisulphate and
Potassium Bisulphate, I have found it difficult to find good data (mainly solubility) on the
other Bisulphate compounds, namely, Magnesium Bisulphate, Aluminium Bisulphate,
Zinc Bisulphate and Potassium Aluminium Sulphate (Cooking Alum) (My online Wikipedia
research has shown that the word “Alum” just means a double-metal sulphate salt, with
the general formula of AM(SO4)2.12H2O, where (A) is a monovalent cation (1+ metal
such as Sodium or Potassium) and (M) is a trivalent cation (3+ metal such as
Alumunium).
For metal ions that have a valency (charge) of 1+, such as Sodium (Na+) and Potassium
(K+), then the bisulphate compound would simply be NaHSO4, or KHSO4 respectively.
These compounds exhibit a pH of approx. 1 - 2, so this seemed to give an answer to my
ponderings. They can be prepared outside of the battery by mixing the metal sulphate
compound with Sulphuric Acid, to form the acid-salt (Bisulphate). Some examples, with
Sodium Sulphate and Potassium Sulphate, are:
Na2SO4 + H2SO4 ==> 2 NaHSO4 (Sodium Bisulphate)
K2SO4 + H2SO4 ==> 2 KHSO4 (Potassium Bisulphate)
These equations have a mix ratio of simply 1:1, since 50% is the metal
sulphate and 50% is the Sulphuric Acid, i.e., this happens if you mix 1 Litre of
1M Sodium Sulphate with 1 Litre of 1M Sulphuric Acid.
-------------------------------------For Magnesium however, the compound would be slightly more complex because of
Magnesium's 2+ charge. I theorised the compounds formula to be:
Mg2+ + 2H+ + 2SO42- ==> Mg2+(HSO4-)2
And that it could be produced by mixing a 1:1 ratio of Magnesium Sulphate with
Sulphuric Acid, of the same molarity, e.g.
MgSO4 + H2SO4 ==> Mg(HSO4)2 (Magnesium Bisulphate)
These equations above seemed to explain what I was observing given the pH readings.
They explain how the initial magnesium sulphate gets turned into magnesium bisulphate
after the first charge / discharge cycle, and from then on, perhaps it operates as a
bisulphate on-goingly and re-establishes itself during both the charge + discharge
cycles, for example:
After the battery plates have been desulphated, if we assume that we have plates of
only Lead (Pb) (Negatrode) and Lead Dioxide (PbO2) (Positrode), then the half equations
for our lead-acid battery using our Magnesium Bisulphate compound might work like
this:
SUBSEQUENT DISCHARGE OF MAGNESIUM BISULPHATE ELECTROLYTE IN
LEAD-ACID BATTERY WITH PLATES WITH NO LEAD SULPHATE ON THEM
@ - Negatrode:
(Charged State)
(Discharging State)
Pb(s) + Mg(HSO4)2 ==> PbSO4(s) + MgSO4 + 2H+ + 2eTHE IONIC EQUATION BREAKDOWN LOOKS LIKE:
(Pb(s) - 2e-) + (Mg2+ + 2HSO4-)
Pb2+ + Mg2+ + (2H+ + 2SO42-) ==> (PbSO4) + (MgSO4) + 2H+ + 2eTHEN, AFTERWARDS, SPONTANEOUSLY (AND SECONDARILY) ?:
PbSO4(s) + MgSO4(aq) + 2H+ ==> Pb(s) + Mg(HSO4)2(aq)
(Charged State)
PbO2(s)+ 3H+ + HSO4- + 2e-
@ + Positrode:
(Discharging State)
==> PbSO4 + 2H2O
THE IONIC EQUATION BREAKDOWN LOOKS LIKE:
PbO2(s) + 3H+ + Mg(HSO4)2 + 2ePbO2(s)+ 3H+ + HSO4- + Mg2+ + 2ePbO2(s)+ 3H+ + HSO4- + 2e((Pb4+ + 2e- )+ 2O2-) + (4H+ + SO42-)
(Pb2+ + 2O2-) + (4H+ + SO42-)
2+
2+
(Pb + SO4 ) + (4H + 2O2-) ==> (PbSO4) + (2H2O)
THEN, Possibly spontaneously react with Magnesium Sulphate in solution, in
the presence of the + charged Lead Electrode:
PbSO4 + 2H2O + MgSO4 ==> Mg(HSO4)2(aq) + PbO(s) + H2O
@ Negatrode, Pb is being Oxidised during Discharge (giving up electrons) turning into
Pb2+ + 2e-.
@ Negatrode, the Magnesium bisulphate is already dissociated being in an aqueous
solutions, effectively having the Bisulphate anion floating in a loose orbit to the
Magnesium, so we can assume the Mg is already in an ionised state (Mg2+) and the
bisulphate anion is loosely bonded in its' own ionised state 2(HSO4-)
@ Negatrode, the bisulphate anion (HSO4-) is being oxidised (loosing hydrogen) turning
into H+ + SO42-.
During Discharge, the Negatrode is undergoing OXIDATION
@ Positrode, PbO2 (Lead Dioxide, also called Lead (IV) oxide) is being Reduced during
discharge (gaining electrons) turning from Tetravalent Lead (Pb4+) to Divalent Lead
(Pb2+), by receiving electrons.
@ Positrode the Pb is loosing the Oxygen to the Hydrogen
@ Positrode, I am assuming that the Mg2+ itself does not take part in the reaction given
the + charge on the electrode, and the + charge of the Magnesium ion, so I am using
the standard Lead-Acid Positrode reduction equation of assuming the presence of the
HSO4- anion.
@ Positrode I theorise that the Lead Sulphate formed during reduction is then further
reacted with the Magnesium Sulphate in solution (from the initial dissociation of Mg +
HSO4 at the positrode) to consume 1 water molecule and reduce the Lead to Lead Oxide
(PbO) (also called Lead (II) Oxide), which is an orange/yellow/pastel coloured oxide
powder. I have observed this colour forming on the Positrodes in my test batteries.
During Discharge, the Positrode is undergoing REDUCTION
--------------THOUGHTS ON OTHER SULPHATE COMPOUND OPTIONS---------------The example above is looking at the chemistry of using Magnesium Sulphate to
desulphate sulphated lead acid batteries, and using Magnesium Bisulphate to continue
on the normal reaction after desulphation, or when using in new/clean plate lead-acid
batteries.
But what about the other sulphate compounds?, We have many different options in the
sulphate compounds that we could choose;
Some are highly toxic, some completely benign, some environmentally friendly, some
environmentally disastrous. Some insoluble, some highly soluble, some based on Metal
ions, some based on Non-metal ions, some that will work, some that will not.
To get an understanding of how I calculated what Sulphates would work, and then
narrowed down the list to the few that I use today, let's take a look at the basic features
and characteristics of some of the different common sulphates:
Sulphate Name
Quick Comments
Could we use it?
Sodium Sulphate
Cheap, non-toxic, soluble
YES
Potassium Sulphate
Cheap, non-toxic, soluble
YES
Calcium Sulphate
Not soluble
NO
Copper (II) Sulphate
Less reactive than Lead, will plate instead
NO
Copper (I) Sulphate
Less reactive than Lead, will plate instead
NO
Aluminium Sulphate
Easy to make, dangerous to environ.
YES
Magnesium Sulphate
Cheap, non-toxic, soluble
YES
Lithium Sulphate
Difficult to obtain, but has high solubility in
cold water
YES
Iron (II) +(III) Sulphate
Cheap, available, lowly-toxic, soluble, but
may precipitate easily
YES
Nickel (III) +(II)Sulphate Toxic, Dangerous for environment
NO
Zinc Sulphate
Cheap, non-toxic, highly soluble
YES
Tin (IV) Sulphate
Easy to make, non-toxic
YES
Ammonium Sulphate
Might dissociate into gases?, highly soluble
YES
Cobalt (III) Sulphate
Toxic and dangerous to environment
NO
Chromium (III) Sulphate
Almost non-soluble, dangerous to environ.
NO
Manganese (III) Sulphate Forms permanganate with PbO2
NO
Manganese (II) Sulphate
NO
Forms permanganate with PbO2
Iron Ammonium Sulphate Soluble, Irritant
YES
Titanium (IV) Sulphate
Difficult to obtain, soluble
YES
Barium Sulphate
Not Soluble, radioactive,
NO
Potassium Aluminium
Sulphate
Used for food production, soluble
YES
Cadmium Sulphate
Highly toxic, Dangerous to environ.
NO
Silver Sulphate
Barely Soluble, dangerous to environ, toxic
NO
------------------A DEEPER LOOK AT SOLUBILITY AND MOLARITY-------------------After experimenting with some different sulphate compounds, it became apparent to me
that some of them were working much better than others. Magnesium sulphate, for
example, was shown to strip lead-sulphate far more effectively than Sodium, and that
Potassium sulphate barely worked at all. At first I thought this might just be differences
in the old cells I was using, non-uniform charge rates, temperature, etc.. but then I
decided to really understand it properly.
Knowing that what I was wanting to occur, was to have the Sulphate compound convert
into its' Bisulphate form, as this was my theoretical explanation for how the sulphates
were de-sulphating the old batteries. Looking at a few of the sulphates, I knew I could
calculate the total stripping power by seeing how many SO42- anions the Bisulphate
would “consume” so to speak, per mol of concentration:
Sodium Sulphate – Sodium Bisulphate Conversion
Na2SO4 + 2H+ + PbSO4 ==> 2NaHSO4 + Pb(s)
(can strip 1mol PbSO4)
Potassium Sulphate – Potassium Bisulphate Conversion
K2SO4 + 2H+ + PbSO4 ==> 2KHSO4 + Pb(s)
(can strip 1mol PbSO4)
Magnesium Sulphate – Magnesium Bisulphate Conversion
MgSO4 + 2H+ + PbSO4 ==> Mg(HSO4)2 + Pb(s)
(can strip 1mol PbSO4)
Aluminium Sulphate – Aluminium Bisulphate Conversion
Al2(SO4)3 + 6H+ + 3PbSO4 ==> 2Al(HSO4)3 + 3Pb(s)
(can strip 3mol PbSO4)
Potassium Aluminium Sulphate (Alum) – Potassium Aluminium Bisulphate Conversion
KAl(SO4)2 + 4H+ + 2PbSO4 ==> 2KAl(HSO4)4 + 2Pb(s)
(can strip 2mol PbSO4)
Looking at this short list, we can see that 1mol Sodium, Potassium and Magnesium all
have the ability to strip only 1mol of SO42- from the Lead (Pb) during the conversion
from Sulphate compound to Bisulphate compound. They also consume Hydrogen during
the process (this is important, as it suggests that we need to have free hydrogen in the
solution, or a negative – charge on the electrode to produce H+ ions to enable to process
to happen).
We can also see that for each 1mol Aluminium Sulphate to Bisulphate conversion it has
the ability to strip 3mol of SO42- from the Lead (Pb). Alum can strip 2mol, etc. (One day I
will complete a chart of all known sulphate compounds and calculate their features for
this use).
At first glance then, it looks like using Aluminium would be the clear winner, but
remember, this is looking at the ability of the chemistry to strip SO42-, and is entirely
based on an assumption that we have equal molar concentrations of the metal sulphate
compound, and the Lead Sulphate. So the next thing to do is to calculate the total
solubility of the metal sulphate compounds to see how many mol of each we can
actually dissolve into 1-litre of water and make available:
This next chart is looking at each of our example sulphate compounds, and lists their
max. solubility in water @ 20*c (since we are using our batteries in every day
environments, lets use a standard average temp in Australia to calculate this data). It
also lists their molar weight (the total amount of the compound we would need to get a
concentration of 1M in 1-litre of water), then lists the compounds “SO42- stripping power
as a molar value. (This is based on using the molar weight of SO4 which is calculated to
be 96.06g, therefore a concentration of 1M of SO4 will be 96.06g/L H2O.), then finally
multiplies the stripping potential of SO42- expressed as p/M value by the total mol of
metal sulphate that we can obtain in 1-L H2O @ 20*c.
This total info gives us a clear value that we can use, which tells us the exact theoretical
amount of SO42- that a maximum concentration p/L H2O of that metal sulphate
compound has the ability to strip.
-----------SO4 STRIPPING POTENTIAL OF ALT. SULPHATE COMPOUNDS-----------For comparison sake, I have included data for both @ 20*c and @ 40*c, though I am
basing this PDF any my main calculations on the use at 20*c, since this could represent
an average day temp here in Australia. Even though we might get higher solubility at
higher temps, once the temp cools down, the excess sulphate compounds would then
precipitate out of the solution back into powder, and this could build up on the bottom of
the cell and short out the plates, so let's avoid that happening! ^_^
Na2SO4 (Sodium Sulphate)
@ 20*c
max. solubility p/L H2O
@ 40*c
195g/L
488g/L
molar weight (to make 1M)
142.04g/L
142.04g/L
Max. concentration p/L H2O
1.37mol
3.43mol
96.06g/L (1M)
96.06g/L (1M)
131.87g/L
329.48g/L
24
Stripping Potential of SO
Total SO42- that can be stripped p/L
K2SO4 (Potassium Sulphate)
max. solubility p/L H2O
@ 20*c
@ 40*c
111g/L
148g/L
molar weight (to make 1M)
174.26g/L
174.26g/L
Max. concentration p/L H2O
0.63mol
0.85mol
96.06g/L (1M)
96.06g/L (1M)
61.18g/L
81.65g/L
Stripping Potential of SO42Total SO42- that can be stripped p/L
MgSO4 (Magnesium Sulphate)
max. solubility p/L H2O
@ 20*c
@ 40*c
351g/L
447g/L
molar weight (to make 1M)
120.36g/L
120.36g/L
Max. concentration p/L H2O
2.91mol
3.71mol
96.06g/L (1M)
96.06g/L (1M)
280.1g/L
356.38g/L
Stripping Potential of SO42Total SO42- that can be stripped p/L
Al2(SO4)3 (Aluminium Sulphate)
@ 20*c
@ 40*c
max. solubility p/L H2O
364g/L
458g/L
molar weight (to make 1M)
342g/L
342g/L
Max. concentration p/L H2O
1.06mol
1.34mol
288.18g/L (3M)
288.18g/L (3M)
306.5g/L
386.16g/L
Stripping Potential of SO42Total SO42- that can be stripped p/L
KAl(SO4)2 (Alum)
@ 20*c
max. solubility p/L H2O
@ 50*c
140g/L
368g/L
molar weight (to make 1M)
258.19g/L
258.19g/L
Max. concentration p/L H2O
0.54mol
1.42mol
192.12g/L (2M)
192.12g/L (2M)
104.17g/L
273.80g/L
Stripping Potential of SO42Total SO42- that can be stripped p/L
ZnSO4 (Zinc Sulphate)
max. solubility p/L H2O
@ 20*c
@ 40*c
538g/L
705g/L
molar weight (to make 1M)
161.45g/L
161.45g/L
Max. concentration p/L H2O
3.33mol
4.36mol
96.06g/L (1M)
96.06g/L (1M)
320.10g/L
419.46g/L
Stripping Potential of SO42Total SO42- that can be stripped p/L
For comparison sake, let's also take a look at Sulphuric Acid, H2SO4.
From my research, the majority of modern flooded Lead Acid Batteries for standard
applications, use a concentration of Sulphuric Acid at between 4-5M. So I will over
estimate and use the idea of a 5M H2SO4 concentration.
The molar weight of SO4 is 96.06g, this means that we will expect to have 96.06g of
SO4 per L of H2O for a 1M concentration of SO4.
If we use our 5M battery acid statistic, then we can assume that we will have:
5M x 96.06g = 480.30g/L of SO4 dissolved into the water.
Given that, when I have worked on old batteries, even heavily sulphated ones, the liquid
that comes out, or if dried out, the liquid that comes out once you fill the battery with
water and slosh it around, is still highly acidic, indicating to me, that even in heavily
sulphated Lead-Acid batteries, not all of the SO4 will come out of the solution and onto
the Lead Plate as PbSO4.
Lets theorise then, even a conservative idea of what if ½ of the available SO4 were to
come out of the solution and form PbSO4 on the Lead plates? So;
2.5M x 96.06g = 240.15g/L of SO4 dissolved into the water.
Now, with this data, looking back at our table that shows us the total potential stripping
ability, in g/L (or M), we can deduce that a single treatment of Sodium Sulphate or Alum
(Potassium Aluminium Sulphate) in a heavily sulphated battery cannot fully desulphate
the cell (therefore it would require a few treatments to do so).
Potassium Sulphate has an even lower stripping potential, and theoretically Magnesium
and Aluminium sulphates have the potential to strip all of the SO4 away from
the Lead in a single treatment.
I currently have various tests setup to confirm this. Magnesium sulphate definitely is
effective at stripping sulphate from the lead plates in sulphated batteries.
My experimenting with Potassium Sulphate showed me that it barely worked at all, I did
not notice much reduction in the total Lead Sulphate with a single treatment of it.
My experimenting with Sodium shows that it works much better than Potassium, but
less totally than Magnesium (for just one treatment), however, it seems to desulphate at
a quicker rate?,
Lastly, my experiments with Aluminium sulphate definitely showed me that it
provides a larger spark than the others upon short-circuit with high amp
cables, and that it produced higher discharge amp-ratings than the others.
Batteries that have less SO4 on the Lead plates might be able to be recovered by using
Sodium Sulphate or Alum, but it is hard to tell what amount in grams (or M) of SO4 is on
the lead plates just by looking at them, not to mention hard to look into the whole cell
of closed plastic batteries.
In the future I would like to calculate a way to estimate potential SO4 on the Lead plates
by looking at the g/cm2 or g/cm3. This would also help to estimate total charge potential
(in Amp-Hours) using a formula that tells us total Ah/cm2.
There is a lot more research to be done. Glaringly obvious things that need more
understanding such as could the Magnesium Sulphate/Bisulphate cause any solid
precipitate under certain conditions? Such as overcharging, etc..?
How does the total solubility of the Bisulphate compounds differ to their sulphate
compounds, and how can I use this to get the maximum amount of concentration into
solution.
Does the presence of a free metal ion in the sulphate solutions (such as Alum
“Potassium Aluminium sulphate) allow the SO4 stripping process to occur even without a
charge applied to the cell? (So that it would not be stripping SO4 from the Pb by forming
a Bisulphate compound, but rather, simply forms a separate sulphate compound, i.e.
2KAL(SO4)2 + PbSO4 + H+==> Pb(s) + K2SO4 + Al2(SO4)3 + HSO4- ?
Any many many more...
---------BUT WHAT ABOUT ONCE THE BATTERY HAS BEEN DESULPHATED--------------------------------THE ROLE THAT HYDROGEN IS PLAYING------------------------Now that we have an idea of which sulphate compound will be the best contender for
helping us to desulphate our old lead-acid battery based on the depth of its' previous
sulphation, and the max. solubility and molarity of our chosen sulphate compound, we
will be able to choose an alternative sulphate compound for use as our de-sulphating
electrolyte.
But what about its' properties as a normal everyday electrolyte once the desulphating
process has finished?
How does it compare to Sulphuric Acid in terms of amp-hour production and voltage?
Gathering that the process of discharging goes as follows (using the Magnesium
Bisulphate example), we can see that we require H+ ions in order to facilitate the
discharge process, by reducing the Oxygen on the Positrode +
@ - Negatrode:
(Charged State)
(Discharging State)
Pb(s) + Mg(HSO4)2 ==> PbSO4(s) + MgSO4 + 2H+ + 2eTHEN, AFTERWARDS, SPONTANEOUSLY (AND SECONDARILY) ?:
PbSO4(s) + MgSO4(aq) + 2H+ ==> Pb(s) + Mg(HSO4)2(aq)
(Charged State)
PbO2(s)+ 3H+ + HSO4- + 2e-
@ + Positrode:
(Discharging State)
==> PbSO4 + 2H2O
THEN, Possibly spontaneously react with Magnesium Sulphate in solution, in
the presence of the + charged Lead Electrode ?:
PbSO4 + 2H2O + MgSO4 ==> Mg(HSO4)2(aq) + PbO(s) + H2O
So what we are mainly wanting to look at here is the available Hydrogen as H+ ions,
free to reduce the PbO2 (Lead Dioxide) and bond with the Oxygen O2- and form a
molecule of water H2O.
Let's compare the total max. available Hydrogen in the Bisulphate solutions per Litre
based on molar concentration and solubility
Again, we have to go back to our molarity of the compounds to get an understanding
and come up with yet another chart. :)
Bisulphate Name
Bisulphate
Per/
Compound molecule
Formula
H+ count
Max. Molar
Available H+
concentration p/L
per
based on solubility L/electrolyte
@ 20*c
Sulphuric Acid
H2SO4 (for
reference)
2
5M (reference of
most used conc.)
10g/L
Aluminium Bisulphate
Al2(HSO4)3
3
1.45M
* 4.37g/L
Sodium Bisulphate
NaHSO4
1
2.06M
2.06g/L
Potassium Bisulphate
KHSO4
1
3.56M
3.56g/L
Magnesium
Bisulphate
Mg(HSO4)2
2
2.04M
Potassium Aluminium
Bisulphate
KAl(HSO4)4
4
0.52M
* 4.09g/L
* 2.11g/L
Note: Hydrogen has a molecular weight of 1 gram (rounded down from 1.008g)
For this chart, I also needed to look into the solubility for the new Bisulphate
compounds, as they are different from the Sulphate compounds. The molarity also
changes. Here are the compounds above, listed with Solubility and Molar weight data for
both Sulphate and then the Bisulphate.
* Some Bisulphate compounds, such as Aluminium Bisulphate, Magnesium
Bisulphate and Potassium Aluminium Sulphate I have not yet found solubility
data online for them despite my research, so I am considering making them and
precipitating them by drying and measuring the solubility data for myself, and I will then
add the data to this PDF in future versions. What I have done for the purpose of this
chart, is assume a greater solubility than their Sulphate partners, but kept it to a
conservative increase in solubility by 100g/L, I have no idea if this is accurate
or not.
Compound Name
Formula
Molar weight
Max. Solubility @
20*c p/L H2O
Aluminium Sulphate
Al2(SO4)3
342.15g/mol
364g/L
Aluminium Bisulphate
Al(HSO4)3
318.19g/mol
? est. (464g/L)
Sodium Sulphate
Na2SO4
142.04g/mol
195g/L
Sodium Bisulphate
NaHSO4
138.07g/mol
285g/L
Potassium Sulphate
K2SO4
174.25g/mol
111g/L
Potassium Bisulphate
KHSO4
136.17g/mol
486g/L
Magnesium Sulphate
MgSO4
120.36g/mol
351g/L
Magnesium Bisulphate
Mg(HSO4)2
220.42g/mol
? est. (450g/L)
Potassium Aluminium
Sulphate
KAl(SO4)2
258.19g/mol
140g/L
Potassium Aluminium
Bisulphate
KAl(HSO4)4
454.31g/mol
? est. (240g/L)
Fascinating Notes: Notice how some of the Bisulphate versions of these solutions have
remarkably higher solubility rates?, Potassium shows itself to not be very good for
desulphating batteries (since it lacks the stripping power to work on the SO4, however,
for use as mixing up into a pre-made Bisulphate, it has a high solubility and this will
mean that we can get a much higher molar concentration of Potassium Bisulphate than
we can for Potassium Sulphate.
The two charts above help us to calculate which of the sulphate solutions I have so far
experimented with will be the best contenders for making Bisulphate compounds that
could compare with the use of high molar concentration Sulphuric Acid as an alternative
liquid electrolyte for Flooded Lead Acid batteries.
We can clearly see the following as the stand out Bisulphate compounds in terms of H+
availability:
Magnesium Bisulphate could provide 4.09g/L of available H+
Potassium Bisulphate could provide 3.56g/L of available H+
Aluminium Bisulphate could provide 4.37g/L of available H+
Sodium Bisulphate could provide 2.06g/L of available H+
------------WRAP UP SUMMARY OF THE PDF AND RECOMMENDATIONS------------Lastly on this topic, I am making specific assumptions that, for the majority of the Amp
capacity of the battery, it is the availability of the dissolved Hydrogen in the
solution that is the limit for the amp-hour capacity of the cell, not the Oxygen
attached to the Positrode, nor specifically just the plate surface area. (though
the plate surface area will determine the amount of Oxygen available at the Positrode,
and the amount of H+ generated under charging at the Negatrode)
I need to confirm this with more testing, but I am fast coming to this conclusion.
So, to wrap this PDF all together, if we take the 5M sulphuric acid concentration as an
example, and we have calculated that we could have up to 10g/L electrolyte of available
Hydrogen, then we could theoretically reduce up to 5g/L of oxygen from the Positrode
(given that the water molecule is 2xH+ and 1xO2- = “H2O”).
Given my idea (it may be wrong, but I am interested in the H+/amp-hour
relationship and am open to being more enlightened on the subject), let's say
that in our 12v 100Ah battery, that 100Ah is a representation of the 10g/L H+ from the
5M sulphuric acid, so theoretically, every 1g H+ could be used to represent every 10Ah
of capacity.
Given that, in most Sulphuric Acid based Flooded Lead Acid Batteries, we want to only
consume a total of 30% to 50% maximum capacity, before recharging, to keep to our
max. DOD for long(ish) life cycles. Then we can consider this instead to be a total
amount of “useable” H+ of 3g/L to 5g/L. Effectively making the total usable AmpHour capacity of our Lead-Acid battery to be 12v 30Ah or 12v 50Ah respectively.
If we now consider our alternative sulphate compounds, we get that:
Magnesium Bisulphate = 4.09g/L of available H+(12v @ 40.9Ah) = 100% DOD
Potassium Bisulphate = 3.56g/L of available H+ (12v @ 35.6Ah) = 100% DOD
Aluminium Bisulphate = 4.37g/L of available H+ (12v @ 43.7Ah) = 100% DOD
Sodium Bisulphate = 2.06g/L of available H+
(12v @ 20.6Ah) = 100% DOD
And for Comparison:
Sulphuric Acid = 3g/L of “useable” H+
Sulphuric Acid = 5g/L of “useable” H+
(12v @ 30.0Ah) = 30% DOD
(12v @ 50.0Ah) = 50% DOD
The battery would arrive at this 0% SOC at a lower per/cell voltage than only using 30%
of the sulphuric acid based battery, but, we might not do any damage to the plates
through sulphation by doing so, and we can routinely discharge to flat, as well as
extending the life-cycle of the battery possibly up to 5x times.
Through my research, I feel that it becomes very evident that there are numerous
options available to use to use as alternative sulphate compounds to run inside our
Flooded Lead Acid Batteries.
Some have features that are more desirable than others, and this research has only just
scratched the surface. New alt. Sulphate electrolytes that account for below freezing
temperature needs to be done, for our friends in the snowy countries, and a look at the
numerous remaining sulphate compounds which are as yet un-researched, needs to be
done.
It's also possible that Aluminium bisulphate, with its' 3x reduction potential of Lead
Sulphate, and higher availability of Hydrogen, could make the best CCA (cold cranking
amps) starter battery?
Personally, I will never be running my lead-acid batteries on sulphuric acid
ever again. If I need more Ah capacity, I simply get more old batteries and add
them in parallel to my system. In the case of car-batteries, I simply find
discarded car batteries which are larger than the size needed for my car, and
that will at least come up to 12v on a quick charge for an hour, and then work
on desulphating them.
This ends the main body of the PDF, and what follows on from this is a collection of
my notes and various test ramblings. This PDF is by no far exhaustive, and so I will be
continuing to test various compounds and run workshops and collecting data until,
ideally, I can report on my research of all known available non-toxic sulphate
compounds.
I think that we can alter the current Lead-Acid Battery technology to help us head
towards a cleaner and more healthy future, while still using what is massively available
to us on a world scale.
As toxic as lead is, Lead Acid Batteries are not going to vanish overnight, and while
recycling rates might be high for countries like Australia, I have heard many stories of
children in India pulling old discarded lead-batteries apart to get the lead out by burning
and melting. Groundwater pollution from the lead scum that accumulates at the bottom
of the case and flows out of the battery if spilled or tipped upside down, is a problem,
not to mention general soil contamination of lead and heavy metals.
While this PDF and my research doesn't solve the problem, or provide a brand new nontoxic battery, I do feel that it helps us head in a better direction by helping us to extend
the life span of Lead Acid Batteries massively (reducing the number in landfill or
dumped), make them generally safer to handle, reduce explosion potential with reduced
Hydrogen gassing, reduce acidity of the electrolyte in case of spills and splashes, but
still maintain usable Amp-capacity, and empower the people of Earth to take ownership
of a possible highly available energy storage medium.
How large an off-grid energy system could you create if you can reuse old discarded
batteries that you can obtain for free?. Think about that. In Love ^_^
For the assistance of all the people of Planet Earth
- Sylph Dominic Hawkins
-----------------MY EXPERIMENT NOTES AND FURTHER COMMENTS---------------This section will be just a random summary of my notes (taken from my
workbooks and summarised here for you). It will not be intended to be
presented in an explained way, so this section is mainly as a dump-place for
my personal thoughts and experiments, in a hope that it may shed more light
on the process for others, and further add to the anecdotal evidence of this
alternative electrolyte.
1st battery conversion was with Magnesium Sulphate (Epsom salts)
This is technically Magnesium Sulphate Pentahydrate (MgSO4.7H2O)
Cell voltage = 1.9v,
Short-circuit spark is much much less than normal (using H2SO4), is this possibly
because there is less H2 available to reduce the O2 from the Positrode?, effectively
reducing the total amount of redox reaction that can occur in each moment with the
given surface area?
The lowest voltage I have drawn this MgSO4 12v @ 100aH battery to is 7.6v under load
over 48hrs, and then recharged happily to settle again at 11.7/11.6v.
Battery continues to maintain strong voltage under load even after many 70/80% DOD
discharges and beyond.
6 cells in series add to give me a theoretical 11.4v battery (it seems to be happy
returning close to this after charge has stopped (night time) (~11.7v).
Cells seem to be happy charging at ~2.2v (6 cells therefore charges at 13.2v)
I have found that above this, there is a substantial spontaneous increase in the amount
of liquid produced. Perhaps car-alternator charge regulators / solar regulators / etc., will
need to be adjusted to lower the cut-off voltage from ~14.4v, to ~13.2/13.5v
My solar charge controller was set to cut out at 14.4v (standard sulphuric acid lead
acid), and so was potentially overcharging the batt with new electrolyte. Batt settled
back to 12v from overcharging.
Was liquid created as a product from the chemistry due to the excess electrons on the
positrode after the surface area of the positrode had been fully turned to PbO2
(charged).
When tested, the MgSO4.7H2O has a pH of 7,
when put into the battery cells, fully charged and tested, the solution has a pH of ~1,
and then when drained and tested, the solution also has a pH of ~1.
Bisulphate compound is assumed.
I have not yet investigated into taking specific gravity measurements of the liquids at
both states in order to get an understanding of what compounds are forming, and to
confirm whether or not the Bisulphate is indeed being produced during both the Charge
and Discharge cycle (this theory is very important, since it provides an explanation for
how the sulphate compound can strip and “desulphate” batteries, as well as provide a
charge and Hydrogen ions for the discharge process)
Charging produces very little Hydrogen gas compared to sulphuric acid electrolyte
(visibly) (is this because it is being bonded with the sulphate compound to produce
more Bisulphate, and is not becoming Hydrogen gas?)
Bisulphate is assumed to be the vehicle that provides Hydrogen ions for the PbO2 to be
reduced with. Otherwise it would be needing to split the H2O, and while I am aware that
water theoretically exists as H+ + OH-, this seems less likely than for a H+ ion to be
available more readily in the Bisulphate compound, and would also require a greater
water loss (I imagine) than what I have observed.
This would also mean that testing the fully charged electrolyte would still show a pH of
~7 if it was returning to simply MgSO4.
The bisulphate compound of Potassium Bisulphate is described as exhibiting features as
if K2SO4 + H2SO4 were living side by side in solution, both performing their individual
features, so perhaps this is also true of Magnesium, Sodium, Aluminium, and other
sulphate compounds?
Great discharge of the battery (up to 80% DOD and more) has shown that not only does
the battery happily bounce back up to it's settled post-charge no-load voltage of ~11.7v
(1.95v per cell) (drops to 11.4v under-load (1.9v per cell), but it actually performs
better the next time! Voltage stays more stable under greater load, and this trend has
continued for 12months. Each cycle use is causing the battery to be able to deliver more
Ah.
Most tests have been performed at average rates of discharge and charge such as 0.1C
(10% of the total Ah capacity of the cell, so in our 12v 100Ah battery example, this
would equate to a discharge and charge rate of 10amps).
I have done higher amp-draw tests on both my Magnesium and Sodium sulphate battery
cells, discharging of upwards of 40amps from 100Ah cells, and charging at 30amps for
100Ah cells.
Under high amp discharge, the voltage drop is more stable than traditional lead-acid
using sulphuric acid, with a 1.9v Sodium Sulphate cell, it could discharge 40amps over 5
mins, and it's voltage drop was to ~1.85v.
(This cell was desulphated and then used with the same electrolyte still in the battery,
so therefore I could not calculate the exact molarity of the assumed Sodium Bisulphate,
as I have not found a way to calculate the grams of Sulphate attached to the Lead in
lead Sulphate of sulphated batteries, so I always overestimate to theorising a possible
5M Sulphuric acid concentration, and then work with that for understanding the amount
of SO42- that I may have to strip from the plates, when mixing my sulphate solutions.
This process is cumbersome, and realistically, once the cells are completely cleaned of
all lead-sulphate crystals, I should dump out the electrolyte, and then refill the cell with
a specific molarity of pre-made Bisulphate (by pre mixing equal molar concentrations of
Sodium Sulphate and Sulphuric Acid). Then I can easily calculate the total available
hydrogen to take part in the reaction, and will probably even see much better Amp-Hour
ratings and capacity in my converted batteries.
I would like to see a 5 year test study done with 80% DOD cycles, to see more longterm effects this has on the battery. How will other factors come into play, such as
shedding of positrode material?
----------------Aluminium Sulphate is very interesting. Theoretically its chemistry shows the potential
of being able to reduce 3x amounts of PbO2 in the same given reaction.
Would this potentially result in a larger faster discharge rate?, i.e., would it spark larger,
more like a conventional lead-acid?, would it reduce the positrode faster?, would a
combination of Al2(SO4)3 and other sulphate compounds produce a middle ground of
larger instantaneous discharge (kind of like a CCA rating), with the higher reactivity and
efficiency of sulphate-stripping of something like Sodium or Potassium?.
Aluminium does react with water, but much more slowly than Sodium or Potassium.
Aluminium sulphate can however, be easily created in a DIY setup by electrolysing
aluminium saucepans + scrap in buckets of sulphuric acid from waste batteries.
----------------Na and K should always spontaneously form their associated Hydroxide if, in some
circumstance, they were to become solid metals from their ions (possibly in an
overcharge state with the battery?). My theory is that they would automatically react
with the water to form NaOH, or KOH, and then continue doing their job, and be
dissolved in the solution.
The battery electrolyte has to be a compound that will remain in liquid form over a
range of temperatures (note to self, test these chemistries and solubility at freezing
temps -0*c and 40*c+ temps). It needs to be a compound that will not precipitate out
of solution, and will not form any particles that could short out between plates.
It needs to be fully reversible, and the re-charge process should return our active
compound automatically.
Another experiment will be with Potassium Sodium Sulphate (if it exists?)
K3Na(SO4)2
and Potassium Magnesium Sulphate (if it exists?)
K2Mg2(SO4)3
Would these co-crystalline sulphate compounds provide a free metal ion (K+) into the
solution, and therefore be able to act on already heavily-sulphate lead plates in second
hand batteries? Without needing any charge input to split the sulphate and form
bisulphate?
If so, then perhaps this accounts for the success of using Potassium Aluminium
Sulphate, and it could possibly strip sulphate ions away from the Lead even without
being put through charge cycles, just through being there and left for a time?
It could be a great way to get Potassium ions into the solution, while getting around
Potassium Sulphate's low solubility in water problem.
(@ 120g/L H2O solubility at 20*c , it is difficult to prepare a concentration of Potassium
Sulphate high enough for our tests (1M concentration), since 1-mol K2SO4 requires
174g / L H2O)
Sodium and Potassium
The metal ions will never exist as solids in water, and will always form hydroxides, so
there is no chance that they will precipitate out of solution.
They are massively reactive, far far greater than Lead Pb, and I imagine they would be
more ionically available to strip the SO4- due to the fact that they more readily dissolve
to form ions than many metals.
Their sulphate compounds are non-toxic to humans, and easily obtainable, cheap and
non-environmentally toxic.
They can be dissolved into distilled water without producing heat or vapours, and
therefore be mixed in plastic containers and without special tools.
Unfortunately, Potassium simply lacks the stripping power of the SO42-, due to it's low
solubility and difficulty to obtain high molar concentrations. However, its' bisulphate
solubility is much higher, making it a possible contender for use as a pre-made
bisulphate compound, but not for the desulphating process.
------------------Comments on Lead as well as its possible Oxide compounds
that we are dealing with
Lead could potentially form 3 different forms of oxides, all highly toxic, in lead-acid
battery chemistry:
- PbO (lead(II)oxide) – a red / yellow / pastel pink - rusty coloured powder
- Pb3O4 (lead(II/IV)oxide, also called “Red Lead”) - Bright Red/Orange powder
(formula is: Pb3O4, but actually is [2PbO].[PbO2] )
- PbO2 (lead(IV)oxide, also called Lead Dioxide) - dark brown powder, insoluble in H2O
Lead Dioxide dissolves in “Strong” alkali (potentially >40% concentration ?) to form the
hydroxyplumbate ion, Pb(OH)62−
PbO2 + 2NaOH + 2H2O → Na2[Pb(OH)6]
Lead Hydroxide (Pb(OH)2) is noted as being a particularly strangely unstable compound ,
and that Lead (II) Oxide (PbO) or Basic Lead Carbonate (PbCO3·2Pb(OH)2) is usually
found in place where Lead Hydroxide is expected.
(This note is directly from: https://en.wikipedia.org/wiki/Lead(II)_hydroxide)
Maybe this accounts for the automatic and relatively spontaneous reaction of Pb(OH)2 in
the discharge reaction of the lead-acid battery chemistry theory?
Chart tables of Voltage characteristics of some different sulphate compounds
relative to their SOC % -------------------------------------------------------------------
SOC %
P/Cell Sulphuric P/cell Al2(SO4)3 P/cell NaSO4
Acid H2SO4
(1.85v
(1.91v
(Baseline)
nominal)
nominal)
P/cell MgSO4
(1.91v
nominal)
Ideal charging
voltage (0.3v above
nominal)
2.41v
2.15v
2.21v
2.21v
Float charge voltage 2.21v
(0.1v above
nominal)
1.95v
2.01v
2.01v
100
2.12v
1.85v
1.91v
1.91v
90
2.10v
80
2.08v
70
2.06v
60
2.04v
50% SOC (0.1v
below nominal)
2.01v
1.75v
1.81v
1.81v
40
1.99v
30
1.96v
20
1.94v
10% SOC (0.2v
below nominal)
1.91v
1.65v
1.71v
1.71v
Interesting to note that regardless of the different sulphate compounds used as
electrolytes, on average, the potential difference in voltage of the cells from a 100%
charged state to an “empty” or 10% charged state, remains around the 0.2v (200mV)
range, suggesting that the lead-acid chemistry, regardless of the lower voltage potential
due to the presence of the metal ions in the sulphate solutions, operates around a 0.2v
range from full to empty.
They also exhibit similar charge voltage ranges, of 0.3v above resting “nominal” charge.
I have observed an average range of 0.5v between ideal charging voltage and “empty”
voltage.
------------Recommended 7x cell series arrangement for alt. sulphates-----------This chart below is the same chart setup, but using values of a normal 6x cell
arrangement (such as in a “standard” 12v car battery), and also an alternative 7x
cell arrangement. Notice the average range of 2.9v (that is 0.49v per cell) between
ideal charging voltage and 10% remaining SOC (“empty”).
SOC %
6x cells
Sulphuric
Acid H2SO4
(Baseline)
6x cells
Al2(SO4)3
(1.85v
nominal)
7x cells
Al2(SO4)3
(1.85v
nominal)
6x cells
NaSO4
(1.91v
nominal)
6x cells
MgSO4
(1.91v
nominal)
7x cells
MgSO4
(1.91v
nominal)
Ideal charging
14.4v
voltage (1.7v above
nominal)
12.90v
15.05v
13.26v
13.26v
15.47v
Float charge voltage 13.3v
(0.6v above
nominal)
11.70v
13.65v
12.06v
12.06v
14.07v
100
12.72v
11.10v
12.95v
11.46v
11.46v
13.37v
90
12.62v
80
12.50v
70
12.37v
60
12.24v
50% SOC
(0.6v below
nominal)
12.10v
10.50v
12.25v
10.86v
10.86v
12.67v
40
11.96v
30
11.81v
20
11.66v
10% SOC (1.2v
below nominal)
11.51v
9.90v
11.55v
10.26v
10.26v
11.97v
This next chart gives us an overview of the different standard off-grid lead-acid battery
voltage setups, 12v, 24v, 36v, 48v, and shows us the new number of cells needed with
the alt. chemistries, as well as the relatively expected charge/float/SOC values.
SOC %
12v Sulphuric
Acid H2SO4
(Comparison)
6x cells
12v Mg(HSO4)2
/ NaHSO4 (1.91v
nominal)
7x cells
(1 extra cell)
24v Sulphuric
Acid H2SO4
(Comparison)
12x cells
24v Mg(HSO4)2 /
NaHSO4 (1.91v
nominal)
13x cells
(1 extra cell)
Ideal charging volt 14.40v
15.47v
28.80v
28.73v
Float charge volt
13.30v
14.07v
26.60v
26.13v
100% SOC
12.72v
13.37v
25.40v
24.83v
90
12.62v
25.24v
80
12.50v
25.00v
70
12.37v
24.74v
60
12.24v
24.48v
50% SOC
12.10v
40
11.96v
23.92v
30
11.81v
23.62v
20
11.66v
23.32v
10% SOC
11.51v
11.97v
23.02v
22.23v
SOC %
36v Sulphuric
Acid H2SO4
(Comparison)
18x cells
36v Mg(HSO4)2
/ NaHSO4 (1.91v
nominal)
20x cells
(2 extra cells)
48v Sulphuric
Acid H2SO4
(Comparison)
24x cells
48v Mg(HSO4)2 /
NaHSO4 (1.91v
nominal)
26x cells
(2 extra cells)
Ideal charging
voltage
43.20v
44.20v
57.76v
57.46v
Float charge
voltage
39.90v
40.20v
53.20v
52.26v
100% SOC
38.10v
38.20v
50.88v
49.66v
90
37.86v
50.48v
80
37.50v
50.00v
70
37.11v
49.48v
60
36.72v
48.96v
50% SOC
36.30v
40
35.88v
47.84v
30
35.43v
47.24v
20
34.98v
46.64v
10% SOC
34.53v
12.67v
36.20v
34.20v
24.20v
48.40v
46.04v
23.53v
47.06v
44.46v
Please take nothing that I discover as fact, experiment with it for yourself.
I have provided the basis to doing all of your own experimenting of this process, and no
doubt others will be able to test other compounds that I have not even thought of yet.
Let's share this information together, rather than arguing about expected outcomes
based on information that we receive from battery companies.
The best thing to do is to experiment. If you can get your hands on any second hand
lead-acid wet cell batteries that have probably only worn out due to sulphation, not
mechanical stress, and ideally not shorted out cells due to dendrils and shedding, then
this could be a brilliant way to rejuvenate deep-cycle batteries, making them work
better than designed in the first place, and power your solar / wind / off-grid system for
almost no cost.
Alternatively, if you were to purchase a brand new lead-acid wet-cell battery, you could
immediately dump out the sulphuric acid into a tub and neutralise it using Sodium
Bicarb or Sodium Hydroxide (so that it can then be safely poured down the
sewer/drain), and then fill your new battery with one of these pre-made bisulphate
solutions to protect against sulphation occurring right from the start.
I'd hedge my bets that it would extend the life span of an average car starter battery
from 3 years to over 7 years, and for deep-cycle spiral wound batteries, maybe it would
take their life span up to over 20-30years, up from the nominal 10-14yrs if only 30%
DOD is drawn.
I drive around quite happily with my 2.5L diesel van, and both starter batteries have
been emptied and filled with my Sodium Sulphate solution (to desulphate, then I simply
left the liquid in there, and have not yet dumped it out and refilled with a specific
Bisulphate concentration, but it works. :), and the engine starts happily every time.
Voltage drop is much less than it was with sulphuric acid under the starter motor pull.
The batteries themselves probably have less Amp-hours available now, given that
Sodium Bisulphate generates less Hydrogen per mol and I haven't calculated the specific
concentration of the bisulphate that might be in there, or dumped it out and refilled),
but still, for testing, when I replace the fuel filter on the van, and put in a new one, I
need to crank the engine for up to 40seconds to pump the diesel through the filter and
get the engine going again, and routinely, the starter batteries drop from about 11.7v to
around 9v and stay there the whole time, not dropping further under the load like
sulphuric acid batts do.
The van also has a 2nd battery 12v @ 100Ah charged from roof-top solar, and it is filled
with my Magnesium Sulphate solution (again, I have not dumped out the liquid and
refilled with an exact concentration of pre-made Magnesium Bisulphate, but it still works
well), and it continues to run all the lights/fridge/phone chargers/ipad/electric
blanket/inverter.
Sometimes being discharged down to 8v, and then charged back up again to ~11.7v
(Keep in mind that my intention is to see if I can destroy the battery, and test the
durability under the new electrolyte conditions, I don't normally recommend drawing
your batteries down to 8v, even with a different electrolyte, mainly to reduce shedding
of the active material (lead) from the positrode, but you know, just saying. ^_^)
Best of luck!, thanks for reading pages and pages of chemistry and my rambling
thoughts. ^_^
This information and research is for the benefit of
all free and sovereign Humans
Lots of Love,
© Sylph Dominic Hawkins
March 2016
Australia
*1st Version was originally published November 2015*
To get in contact:
email:
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