From: Oz <Oz@upthorpe.demon.co.uk>
Newsgroups: sci.agriculture,rec.gardens,rec.gardens.edible,uk.rec.gardens
Subject: Re: Was: an experiment w/ pH, Now: an experiment w/ elemental sulphur
Date: Sat, 3 May 1997 06:59:16 +0100

In article <050297084803Rnf0.78@zephyr.manawatu.planet.co.nz>, Stuart
Brown <stuart@zephyr.manawatu.planet.co.nz> writes

>I recently asked a fertiliser salesperson about the risk of sulphate leaching
>from applying elemental sulphur just proir to the winter (a normally wet
>time of the year here). He told me that as soil temperature falls below
>about 8 deg Celicus (about 46 F) elemental S stops being oxidised
>(converted to soluble S). He said this was because bacteria are involved
>in the conversion process and their activity is reduced with low
>temperatures. He said you can chuck a handful in a bucket of cold water
>and it won't dissolve. Any comments ?

This is an interesting area. Quite a lot of work has been done on it but
I have never been able to locate any papers on the Net. If anyone has
any real data on the oxidisation rates of elemental sulphur wrt particle
size, temperature and soil water content I would be very grateful indeed
for the location and/or a copy of any relevent papers.

As I understand it there has been significant work done in New Zealand
(Aukland U.?) and at Newcastle U. in the UK. Doubtless work has also
been done elsewhere. Some at least is confidential, having been paid for
by sulphur suppliers.

I have been able to glean the following:

1) Elemental sulphur is indeed very insoluble in water.

2) Particle size is critical. Small particles break down very much
faster than large ones. For realistic rates of oxidation we are talking
micron sized particles. Large lumps are essentially inert.

3) I believe that it's breakdown is essentially a 'slow burn' and is not
mediated by bacterial activity.

4) Rates of oxidation at low temperatures and in waterlogged (ie low
oxygen levels) soils are very slow, even for small particles. This is
unsurprising as one would expect the rate of oxidation to be dependant
on temperature and oxygen concentration.

5) Plants can, and do, absorb very significant amounts of sulphur as
sulphur dioxide directly from the air, through their stomata, in the
same way they absorb CO2. Some experiments have indicated 80% of
requirement can be taken up in this way.

Elemental sulpur is probably the cheapest source of sulphur for plants
and it should be possible to select particle sizes so that an
application could be made in the autumn that would supply sulphur
throughout the following growing year. The problem is in spreading the
very fine dust, and some sort of pellet that will break down liberating
the powder is required. Alternatively an aqueous slurry could be sprayed
on.

Sulphur is positively hazardous if it were to be stored close to
ammonium nitrate and there is a significant risk of fire and/or
explosion.

Sprayed as a foliar spray in water, finely divided sulphur is a useful
(if not very effective) fungicide as well as providing sulphur for
plants. Presumably the low level of SO2 produced is the active
ingredient.

The reduction in sulphur emissions from industrial and domestic sources
is such that in the UK, and probably elsewhere, plants are becoming
deficient in sulphur and there is now a requirement to apply it as a
fertiliser for optimum plant growth. The quantities that will be
required worldwide are truly huge.

--
'Oz     "Is it better to seem ignorant and learn,
         - or seem wise and stay ignorant?"

From: Oz <Oz@upthorpe.demon.co.uk>
Newsgroups: sci.agriculture,rec.gardens,rec.gardens.edible,uk.rec.gardens
Subject: Re: Was: an experiment w/ pH, Now: an experiment w/ elemental sulphur
Date: Sun, 4 May 1997 08:03:44 +0100

In article <050497083254Rnf0.78@zephyr.manawatu.planet.co.nz>, Stuart
Brown <stuart@zephyr.manawatu.planet.co.nz> writes

>Reading from a NZ fertiliser newsletter (Quinphos) :

Forgive me for being a bit cautious on material produced by fertiliser
manufacturers. Unless one of the major names, and even then they can be
selective about which trials they use, I would want independent work to
back them up.

>To be effective,
>elemental sulphur must be less than 1mm in size.

Do you have lifetimes in soil for sulphur at the 1mm size? I would be a
bit suspicious of sizes this large. They may well be effective, but for
decades. Nothing wrong with this as such, so long as you know.

>It must also contain good
>proportions  of very fine sulphur ie below 0.15mm (150 microns). The
>proportion in this fine size range depends to which it's being put.

Could you illustrate some uses and proportions?

>They recommend for genral maintenence in the North Is of NZ 20-30% of the
>sulphur should be in this fine (<.15mm) fraction.

OK. How much is typically applied as a maintenance dressing, how much
for the initial dose, and on which crops. I presume mostly clover-rich
grassland.

>Ground elemental sulphur sells in NZ for around $NZ375/ton ($US260), this
>is 100% S.

Similar to the UK. Other forms of sulphur are significantly more
expensive (in general), but very much easier to buy and spread.

>Another fertiliser manufacturer in NZ (Ravensdown) produces a product
>they call Maxi Sulphur Super. This is made by spraying molten sulphur
>onto super phosphate granules. This achicves a fine particle size (indeed
>the opposition claim it's too fine).

Why *too* fine?

>>As I understand it there has been significant work done in New Zealand
>>(Aukland U.?) and at Newcastle U. in the UK. Doubtless work has also
>
>I couldn't see any references. I should think a trip to an agricultural
>university library should yield some info.

Not so easy for a rural peasant.

>>3) I believe that it's breakdown is essentially a 'slow burn' and is not
>>mediated by bacterial activity.
>
>Both Ravensdown and Quinphos claim it is significantly affected by
>bacterial activity.

I find it difficult to see a mechanism due to the aqueous insolubility
of elemental sulphur. On top of that it is known that S (presumably due
to the SO2 produced) is rather toxic to micro-organisms. On reflection I
suppose that S can be dissolved in various organic solvents and
microbial activity in the soil will produce some of these, albeit at low
level. OK, it's possible.

I presume that elemental sulphur is used in NZ mainly because it's high
concentration makes spreading bt plane that bit cheaper.

--
'Oz     "Is it better to seem ignorant and learn,
         - or seem wise and stay ignorant?"

From: Oz <Oz@upthorpe.demon.co.uk>
Newsgroups: sci.agriculture,rec.gardens,rec.gardens.edible,uk.rec.gardens
Subject: Re: Was: an experiment w/ pH, Now: an experiment w/ elemental sulphur
Date: Sun, 4 May 1997 12:46:42 +0100

In article <336C4908.1366@icx.net>, "David B. Hedrick"
<davidbhedrick@icx.net> writes
>
>Oz:
>
>       I've got a bit of info to add to your growing pile.

Great, keep it coming.

>> 2) Particle size is critical. Small particles break down very much
>> faster than large ones. For realistic rates of oxidation we are talking
>> micron sized particles. Large lumps are essentially inert.
>
>Since S has such a low solubility, the oxidation can only happen at the
>surface.  Smaller particles, more surface area.

Agreed. So if, say, it takes ten years for a 1mm sized piece of sulphur
to oxidise at room temperature then it will only take one year for a
0.1mm sized particle and 1/10 year for a 10 micron particle. (Number of
particles up as the cube, area per particle decreases as the square).

Hang on, we ought to be able to get an estimate from the micronised
sulphur used as a foliar spray. I'll look it up. Ahh, it has a particle
size of 6-10 microns. Probable lifetime at 15C, a few months? This would
suggest an oxidation rate of about 1.5 micron per month (from both
sides). At that rate a 100micron (0.1mm) would take about three years, a
1mm particle about thirty years and a 1cm cube about 300 years

>> 3) I believe that it's breakdown is essentially a 'slow burn' and is not
>> mediated by bacterial activity.
>
>S does not appreciably oxidize in air.

See above. It depends what you mean by 'appreciably'.

>It is bacterially mediated, and
>that is probably the source of the temperature dependence.

It may be bacterially mediated. However temperature dependence is not
solely a bacterial phenomena. Certainly oxidation rates are temperature
dependent also, if you don't believe me, light some sulphur! Typically
reactions change by a factor of 2 or 3 for a 10C temp change round room
temperature. As a result I would expect temperature dependence even if
the reaction was entirely one of inorganic chemistry. Of course in a wet
cold soil, with oxygen supply to the sulphur particles slowed by a water
layer, the reaction rates could be very significantly lower.

--
'Oz     "Is it better to seem ignorant and learn,
         - or seem wise and stay ignorant?"

From: Oz <Oz@upthorpe.demon.co.uk>
Newsgroups: sci.agriculture,rec.gardens,rec.gardens.edible,uk.rec.gardens
Subject: Re: Was: an experiment w/ pH, Now: an experiment w/ elemental sulphur
Date: Sun, 4 May 1997 22:11:31 +0100

In article <336da98f.1676166@nntp.ftc-i.net>, Wayne Richardson
<wrichardson@ftc-i.net> writes
>
>On Sun, 4 May 1997 12:46:42 +0100, in sci.agriculture you wrote:
>
>>In article <336C4908.1366@icx.net>, "David B. Hedrick"
>><davidbhedrick@icx.net> writes
>
>>
>>>> 2) Particle size is critical. Small particles break down very much
>>>> faster than large ones. For realistic rates of oxidation we are talking
>>>> micron sized particles. Large lumps are essentially inert.
>>>
>>>Since S has such a low solubility, the oxidation can only happen at the
>>>surface.  Smaller particles, more surface area.
>>
>>Agreed. So if, say, it takes ten years for a 1mm sized piece of sulphur
>>to oxidise at room temperature then it will only take one year for a
>>0.1mm sized particle and 1/10 year for a 10 micron particle. (Number of
>>particles up as the cube, area per particle decreases as the square).
>
>If you are assuming dissolution rate is rigorously proportional to
>surface area - 10 years dissolution time for a 1mm diameter ((0.524
>mm3 per particle) is equivalent to only 0.01 year for the 0.1 mm
>diameter particle (0.000524 mm3). - The surface area is a cubic
>function of the particle size.

Oh? Take a cube 10mm on a side, each side has an area of 10x10=100mm^2,
to give a total surface area is 800mm^2. Cut it into cubes 1mm on a
side. You have 1000 of them. Each cubelet has a surface area of 8mm^2 so
a total surface of 8000mm^2. You have increased the area by a factor of
ten, not 1000. The surface area is proportional to the particle size for
any given quantity. Sorry about this as it validates my original figures
and invalidates yours.

>More realistically, the rates of reaction are probably first order,
>dependent on the amount of surface remaining at any given time.

True. However we are actually more interested in how long it takes for
all the sulphur to be oxidised. This is likely to be most easily
expressed as a surface reaction rate expressed as mm per year (or
microns or whatever) which should be approximately constant. The amount
liberated as oxidised sulphur will not be linear of course, I haven't
bothered to work out the shape of the curve but for any given quantity
it will start high (lots of material, big surface area) and fall as the
particles become smaller and the surface area reduces. The rate of
change of volume will start small and grow bigger of course.

>I believe that - on this scale - other factors may become quite
>significant. See, e.g. Stumm and Morgan.  Solubility increases as a
>function of SA but so does sublimation in the case of sulfur.

I would expect sublimated sulphur to have a rather short life in an
oxygen rich environment, wouldn't you? Soon converted to SO2? This
explains it's fungicidal activity of course since SO2 is known to have
fungicidal (and bactericidal) effects. In high doses it is not very good
for plants either.

--
'Oz     "Is it better to seem ignorant and learn,
         - or seem wise and stay ignorant?"

From: Oz <Oz@upthorpe.demon.co.uk>
Newsgroups: sci.agriculture,rec.gardens,rec.gardens.edible,uk.rec.gardens
Subject: Re: Was: an experiment w/ pH, Now: an experiment w/ elemental sulphur
Date: Sun, 4 May 1997 22:19:15 +0100

In article <allyn-ya023080000405970202130001@news.u.washington.edu>,
Allyn Weaks <allyn@u.washington.edu> writes

>> 3) I believe that it's breakdown is essentially a 'slow burn' and is not
>> mediated by bacterial activity.
>
>Quoting from Killham: "Elemental sulfur, sulphides, and several other
>inorganic S-compounds can be oxidized in soil by purely chemical processes,

I would not disagree about S-**COMPOUNDS*, particularly where they are
soluble, or at least partly so. I probably wouldn't disagree if we were
talking about those strange species that live in sulphurous pools,
although most of these do seem to use H2S as a food source rather than
elemental S.

>but these are generally less important than microbial S-oxidation."

OK. I suppose I have to bow to the evidence. I would be more convinced
if I saw some experimental evidence of this, though.

>*  Thiobacilli are bacteria that get their energy by oxidizing sulfur
>(thio) compounds.  They tend to be the most acid tolerant species, which is
>good, since they tend to produce sulfuric acid as a waste product :-)

I wonder how well they would survive in a soil with pH 8? Typically acid
tolerant species are highly INtolerant of alkaline conditions. Put
another way, they die. Despite this elemental sulphur sprays carry no
warning about their use on alkaline soils.

>'Substrate' above is the sulfur compound eaten by the bacteria.  There are
>several possible pathways to get from S to SO4++; from Killham's figures,
>it looks like it takes a minimum of 2 steps, i.e. 2 species.

Whilst I would not disagree that many other soil microflora would get
into the act once the sulphur has been transferred to anything other
than elemental sulphur, I would not agree that each bacteria can only do
one step. Typically bacteria can do several steps though some might be
better than others at any particular step. Also bear in mind that
bacteria excrete the active chemicals since their rigid cell wall does
not allow them to ingest particles. They thus exist in a soup of enzymes
and other products produced by themselves and any other species that may
exist nearby.

>As for the rate at which S becomes
>available to the plants:  "With very few exceptions, the rate of supply of
>plant-available sulphate in soils is not limited by the rate of
>S-oxidation, but by the rate at which organic sulfur is mineralized into
>the inorganic S-pool."

Ah. So now we are getting into soil anaysis. This IS a valid point since
work done in the UK has had serious problems associating sulphur
deficiency in soils with sulphur deficiency in crops. As far as I can
see this could be due to at least three things.

1) As is documented, sulphur dioxide can be absorbed directly into
plants from the air which is not of course affected by soil anaysis.

2) Availability from the organic sulphur compounds in the soil may vary
widely from season to season, and indeed during a season.

3) Deep sulphur from older atmospheric deposition mat be accesses by
deeper rooting plants. People tend to forget that wheat roots get 2M
deep by harvest.

>In practice, I don't think particle size matters all that much (aside from
>gigantic lumps, perhaps).  Certainly, gardeners have been using straight
>flowers of sulfur (admittedly fairly finely divided, but I'm pretty sure
>it's a good bit bigger than micron sized) successfully for years to lower
>pH and add sulfur.

Hah! Unless you are going to apply tens (at least) of tonnes a year on
my soil you aren't going to get anywhere as far a pH is concerned.
Sulphates are capable of blocking the absorbtion of some essential trace
elements so I would suggest caution. I was always under the impression
that flowers of sulphur was mainly used in gardens as a fungicide, for
various rose diseases for example.

>But you need to be careful about when you apply it to foliage.  It's
>phytotoxic above about 80F, isn't it?

Odd statement to be made for an element that is 'almost completely
insoluble' and 'does not react directly with oxygen at normal
temperatures'? On the other hand I would expect sulphur dioxide to be
distinctly phytotoxic at high levels.  :-)

>From Killham again, in addition to the reduction of SO2 in air, the
>reduction of available sulfur in some soils is due to modern fertilizers
>not containing significant sulfur content,

Correct. Single superphosphate contained significant sulphur (12%, I
forget). However the phosphate content was very much lower and the price
was the same. Given the costs of haulage, I doubt this will change much
which makes it an expensive fertiliser. The same with sulphate of potash
and ammonia.

>He says that the
>most common agricultural addition of sulfur now is from pesticides!

I think rather tongue in cheek. Until very recently (in the UK at least)
NO sulphur was applied, except elemental sulphur as a fungicide.

>I
>doubt that sulfur deficiency is really a crisis though.  It's quite a
>common element, and obviously easy enough to add to soils in a way that
>keeps the soil ecology healthy,

Better to have it for free.   :-)
I suppose sooner or later the power station sulphur cleaning waste will
be processed (lots of heavy metals in them I understand) and sold back
to farmers to put on the land. Not remotely as efficient as the old way
that gave agricultural sulphur for free, and no sulphur processing
charge for your electricity. OK, somewhat in jest I admit but I doubt
anyone included the cost of having to use sulphur as a fertiliser when
they did the cost-benefit analysis.

>unlike pesticides and the fortunately
>diminishing air pollution.

The soil, in my part of the world, was not affected by acid rain. Quite
the contrary. It's pretty much neat chalk chippings.

--
'Oz     "Is it better to seem ignorant and learn,
         - or seem wise and stay ignorant?"

From: Oz <Oz@upthorpe.demon.co.uk>
Newsgroups: sci.agriculture
Subject: Re: Was: an experiment w/ pH, Now: an experiment w/ elemental sulphur
Date: Mon, 5 May 1997 10:30:11 +0100

In article <336e4c51.7468687@nntp.ftc-i.net>, Wayne Richardson
<wrichardson@ftc-i.net> writes
>On Sun, 4 May 1997 22:11:31 +0100, Oz <Oz@upthorpe.demon.co.uk> wrote:
>>Wayne Richardson <wrichardson@ftc-i.net> writes

>>Sorry about this as it validates my original figures
>>and invalidates yours.

>True, true - beat me if you must - I have it coming.

Having been in this situation too, I sympathise. The sudden realisation,
the cold sweat, the frisson of fear that runs up the spine. However take
heart from my .sig, it refers to *me* looking ignorant and I do it all
the time.

>>I would expect sublimated sulphur to have a rather short life in an
>>oxygen rich environment, wouldn't you? Soon converted to SO2?
>I have no idea of the activation energy of the oxidation process - do
>you?  Under ambient conditions, it may be possible that moisture is
>needed to promote (effect a pathway) for the oxidation???

Possibly. Ozone would be very active in the gas phase of course as would
any surface the isolated sulphur molecule got absorbed onto. I suspect
it's mean free time in soil would be rather short and little would
escape to the atmosphere unoxidised. On a leaf one might expect quite a
bit of it to bind locally, at least initially. Hmmm. Leaf cuticle is
typically waxy, it's not impossible that micron sized particles would at
least partly dissolve in the wax layer, particularly if a little wetter
in the formulation disrupted the wax somewhat.

Reprise: I have some figures for the vapour pressure of S down to 400C.
However extrapolating to room temp gives absurdly low vapour pressures
and another mode of action must be involved in sublimation, for which I
have no results. I suppose one could ask a vacuum specialist if they had
more appropriate figures for these temperatures.

>>This
>>explains it's fungicidal activity of course since SO2 is known to have
>>fungicidal (and bactericidal) effects. In high doses it is not very good
>>for plants either.

>If moisture is important  for the reaction --- You may be able to
>enhance the affinity of the sulfur for the water by addition of a
>small amount of a cationic surfactant. Elemental sulfur is
>electrokinetically negative - a cationic surfactant should adsorb onto
>the surface of the sulfur and may lower any barrier to dissolution.
>Once the sulfur is dissolved and reacted the cationic is free to bind
>with another one...there is at least one industrial process using
>elemental sulfur that makes good use of this method.

Interesting. I don't know what surfactants are used for micronised
sulphur, these tend to be trade secrets (OK all the competition know,
it's just the users that don't). However keeping 6u particles separate
and suspended for periods of over a year is probably the over-riding
factor.

I would, however, expect to find all sorts of natural surfactants in
soil. I also note that organic peroxides are a common weapon used by
many micro-organisms and it's not beyond the bounds of possibility that
these could be used to accelerate oxidation. The energy balance might be
somewhat suspect, however.


--
'Oz     "Is it better to seem ignorant and learn,
         - or seem wise and stay ignorant?"

From: Oz <Oz@upthorpe.demon.co.uk>
Newsgroups: sci.agriculture
Subject: Re: an experiment w/ pH
Date: Sun, 11 May 1997 22:48:26 +0100

In article <863364320.9078@dejanews.com>, yojimbo@ix.netcom.com writes
>In article <DzfpZbAxEEZzEw$m@upthorpe.demon.co.uk>,
>  Oz@upthorpe.demon.co.uk wrote:

>> Overliming can indeed cause problems, particularly when the soil is
>> naturally acid. Quite a lot of trace elements vary wildely in their
>> availability depending on pH and plants can easily become deficient in
>> an overlimed soil simply because they cannot absorb them from an
>> alkaline soil. Manganese is the best known example.
>>
>> Fortunately grass is rather tolerant of pH changes and rarely shows much
>> in the way of deficiency, so you needn't bother too much.
>
>Perhaps you are confusing dolomite (calcium/magnesium CARBONATE) with
>slaked lime (calcium hydroxide).

No, I am NOT confusing dolomite with slaked lime. No farmer would dream
of using slaked lime to increase his soil pH. Horribly expensive. My
chalky soil manages to have a pH between 7.8 and 8.2 just from the
natural level of chalk (chalk: calcium carbonate) and dolomite in
solution should achieve similar results.


>While slaked lime is VERY alkaline,
>dolomite is not. Dolomite mediates soil pH by acting as a buffer.

And buffers to what pH?? Eh?
Certainly NOT pH 6.5. At pH 6.5 you get these little bubbles of CO2
given off ......

>The dolomite in question is sold as a quick-acting dolomite, the particle
>size is stated to be guaranteed at aome very small figure which I can't
>recall, but is was about one-tenth the screen size of the "standard"
>crushed rock dolomite that I compared it to.

Since I have no need to lime my soil I am not (from memory) up to date
with the approved particle size for the ground chalk that is used for
this purpose. My memory tells me that 100 microns is about what is used
in agriculture, but it might be 80. Using anything larger is very
inefficient and requires a lot to alter the pH by much.

NB Agricultural history:
Slaked lime was used years ago because it was easier to make limestone
(a very hard rock often used in concrete mixes as the aggregate) into a
fine powder by converting it to lime and then slaking it (with water)
which converted it to a fine powder suitable for changing soil pH when
spread. This only required burning a few trees each time and in those
days trees were two a penny. Once steam engines appeared that could
power a crushing mill it fell out of favour and ground chalk (a soft
rock) was used instead. This is the derivation of the word 'liming'.
This is not, of course, to be confused with the derogatory word for
Englishmen (limey) that derived from the citrus lime fruit, used to fend
off scurvey in British ships about the time of the American Revolt.

--
'Oz     "Is it better to seem ignorant and learn,
         - or seem wise and stay ignorant?"