I'm no chemist, but what helped me get over the hump was:

There are a number of chemical reactions that may occur depending on what bumps into what.

Read up on the Arrhenius theory (Edit: try this), that may help.

ReEdit: Sorry about the bad link. Instead, go to Wikipedia and search:

Acid–base reaction

It's Wiki, but it's a very useful explanation. Especially on the the Arrhenius theory.

Daniel, here is the equation I am looking at (from Carrow's book Turfgrass Soil Fertility and Chemical Problems):

CaCO3 + H2O leads to Ca+2 + HCO3- + OH-

(don't seem to find an arrow on the keyboard so "leads to" substitutes for an arrow)

Then he says Ca+2 can replace H+ or Al+3 while OH- reacts with H+ to form water or with Al+3 to form Al(OH).

You have a different explanation: H+ being soaked up into HCO3-. Do you have an equation?

Good point on what happens to H+ when magnesium is added depends on the salt the Mg is in. So magnesium sulfate does not raise pH because the sulfate SO4 and H+ won't combine. Is it the same with potassium sulfate?

Thanks for trying to explain this stuff to me. It has been a long time since high school chemistry and I'm sure it's changed a lot since then.

You have a different explanation: H+ being soaked up into HCO3-. Do you have an equation?

Daniel,

Would explain why the hydrogen atom freed up from the reaction between CaCO3 and H2O combines with the CO3 to create HCO3 and doesn't just recombine with the OH to reform into water making Ca, CO3 and H2O. Why does the OH "wait" for the Ca to free up an H atom to combine with and form water? As long as I'm asking, why does the H combine with OH and not the HCO# to form H2CO3- carbonic acid?

Becky and yard: Carrow's equation might be technically correct, yet at the same time irrelevant. First of all, calcium carbonate is essentially insoluble in pure water, so his reaction isn't going to happen in any quantity worth mentioning. Second, it would have to be done in a lab with a nitrogen atmosphere above the solution because CO2 will interfere (more below).

The bigger picture is that carbonate chemistry is unique in nature from the standpoint that a similar molecule, carbon dioxide, is in the atmosphere everywhere and is very soluble in water. This means that you have to consider all the reactions that happen at the same time. The ones that "win out" are those that have the greatest driving force. In other words, some reactions "want" to happen, while others don't. This is a measurable quantity and is recorded as a solubility product constant, or an acid or base constant. (been awhile, had to google some of this!). So if you look at the solubility constants, you find that calcium carbonate really doesn't want to dissolve in water, but at the same time calcium bicarbonate very much wants to form. In fact, I was reading that calcium bicarb is so soluble that it doesn't even exist as a solid. So what I'm saying is that calcium carbonate would love to exist in solution in the bicarbonate (HCO3) form, but it can't get there because it can't dissolve in water. However, it can make an end-around play to get there. CO2 will dissolve in water and form carbonic acid. Now we have the supply of H+ from the H2CO3 that came from the CO2 and H2O. This H+ is able to break down the calcium carbonate and get it into solution as calcium bicarbonate. So what we end up with, like I said above, is lots of species all in equilibrium, from CO2 to carbonic acid, to bicarb and carbonate, but I think more HCO3 bicarbonate ion in solution than anything else.

So if we look at the H+ attached to the clay particle we know that is a static attraction, and not a strong bond (compared to an actual "covalent" molecular bond). I'm going to assume therefore that the H attached to the clay's cation site acts like a weak acid. So now in addition to all the CO2 chemistry above, we have calcium carbonate reacting with another acid, and not just with carbonic acid from the atmosphere. Call it "clay acid" I guess. So since limestone likes to react with acids to form the bicarb ion, the reaction might look like this:

Clay-H+ + CaCO3 ---> Clay-Ca+ + HCO3-

I realize the equation is not electrically balanced but I just woke up. I'm not sure if addressed the question adequately but I'll have to read this again later today. The wife is kicking me off the pc, lol.

Also, Becky, yes all else being equal it could be Mg or K sulfate. It is the sulfate that matters when looking at how it combines (or doesn't) with H+.

Here is another explanation of how lime works. What do you think of this one? He also answers the question about over liming: saying that adding lime to alkaline soil doesn't raise the pH but could mess up the soil with too much calcium or magnesium. Is that correct, that lime won't have the carbonate effect in alkaline soil? Is there anything that raises pH in an alkaline soil? Daniel, your explanation about calcium carbonate dissolving in carbonic acid (from CO2 and H2O) would seem to work whether the soil is acidic or alkaline.

Whatcha growing there, Becky? ;-)

Yes, I know. And my daughter and family are soon moving to the Humbolt area of California . I'm just trying to understand the chemistry of lime and pH changes and that came up in a search of how does lime work. More questions: what happens to the sulfate part of magnesium sulfate or potassium sulfate? If H+ doesn't combine with the sulfate, I guess the H+ off the exchange site will swim around in the soil solution as H3O until a proton is needed somewhere (or does it?), but what happens to the SO4? What happens when potassium chloride is used? Does the chloride combine with H+? Is there a pH change?

Becky,

You're already moving on to sulfates and chlorides? I'm still trying to digest daniel's discussion of carbonic acid and bicarbonates. NM, I was lost long ago.

Good luck.

daniel,

Thanks for taking the time to write up that explanation. I was lost a long time ago when "sometimes it acts as a particle, sometimes it acts like a wave, it can depend on whether it's being observed or not." :/

I don't understand everything either but I get the gist of what he's saying. it seems what is missing from Carrow's formula is the role of carbon dioxide. Nothing is going to happen to the lime until you get water and carbon dioxide there to react with it. I'm trying to digest the idea of hydrogen so easily migrating between the carbonate forms. And it's so little, just one measley proton. Yet in the lyotropic series, it's between Al and Ca in adsorption strength (Al+3 > H+ > Ca+2 > Mg+2 > K+ = NH4+ > Na+), at least in some forms of the series. I've also read that it's about the same adsorption strength as K+ on clay surfaces and more tightly held than Al+3 to oxygen below pK, whatever pK is. Quite a mysterious character.

Yet in the lyotropic series, it's between Al and Ca in adsorption
strength (Al+3 > H+ > Ca+2 > Mg+2 > K+ = NH4+ > Na+), at
least in some forms of the series.

Yes H+ seems to move around on different lyotropic series. Hydrated radius? One article explained that if there is enough H in solution they can gang up and displace any cation, including Al. And then there is the secondary cation layer...My search goes on for a layman's explanation, what an amendment will do, formulas for calculating the amount and what amendments in what quantities in conjunction with another may have adverse effects. I'd rather avoid the details at the nuclear level, although I know it would be advantageous to understand it.

I'll continue to follow this discussion and pick-up what I can.

Carrow's formula may be correct, but it is kind of like saying that a first date leads to a baby. You might eventually get there, but a lot of other stuff happens in between.

I can try again later to explain it in other terms. The link Becky provided was very good and explained it better than I did, since he knows the material better, for one, and is a chemist by trade, for another! I'm working off of memory and the internet.

I was lost a long time ago when "sometimes it acts as a particle, sometimes it acts like a wave, it can depend on whether it's being observed or not."

In order to measure or observe, you have to interact with the system one way or another - say by bouncing a photon off - and it is this interaction which "collapses" the wavefunction. This collapse will happen whether a human is "observing" or not. It will also happen if you use a measuring device that does not display any information, taking the "human consciousness" factor out of the equation.

You were lost and now you're found :-)

This is a bit theoretical, but explains why the ph can become greater when you add a basic material.

First pH is there is an equilibrium of the H and OH ions. At a pH of 1, the H concentration is .1 (10 to the -1 power) molar. At a pH of 14 the H ion concentration is (10 to the -14 power) molar The H and OH ions are in equilibrium when the pH is 7. As you add H ions the concentration of the H ions goes up. As you add OH ions the H ions combine with the OH to form H2O. and the H ion concentration becomes more dilute.

It has been 50 years since I studied Chemistry but as I remember the pH scale is based on what is called a equilibrium constant (10 to the 14th power), or the propensity for the molecule to dissociate. In the case H and OH. At a pH of 7 there are an equal number of H and OH ions.

I am sure there are those out there that have had or worked in chemistry since I have so maybe some one can give a better explanation.

This is the definition from a search (Their K is what I called a dissociation constant)

"Like any other equilibrium constant, the value of Kw varies with temperature. Its value is usually taken to be 1.00 x 10-14 mol2 dm-6 at room temperature. In fact, this is its value at a bit less than 25°C."

I became interested in this idea of calcium carbonate not being soluble in water, yet at the same time it is said to have a pH of 9.5 in aqueous solution. The two seemed contradictory. After about a 3 hour search I found a very interesting article on the extent to which calcium carbonate hydrolyzes, or picks up an H+ and leaves behind an OH- as in the Carrow equation Becky provided. I read a good deal of this paper, which contradicted some of what I have read, and also illustrated the complexity of something as simple as dissolving limestone in soil. Of course after all that reading I realized this paper was written in 1936! Somehow I think it still stands.

For anybody so inclined, click on the link, and then on the image of the paper at the right part of the screen to download.

http://arizona.openrepository.com/arizona/handle/10150/190519

Thanks for the paper. I read a good part of it, skimming where I would not understand it anyway. He did not discuss carbon dioxide as part of the cause of the reaction. Do you think that's because it was not part of the lime formula (which is the same as Carrow's)? There was not a lot of discussion about the changes in carbonate forms, but that was not the focus of the study and may be already well known. Anyway, I have answers to a couple of my questions. Calcium carbonate added to alkaline soil does not raise the pH. What happens to the HCO3 is that it transforms to water (gets another H+ from somewhere) and carbon dioxide. The carbon dioxide would go into the atmosphere and the water would stick around in the solution, leach away, be taken up by the plant, or evaporate.

By the way, I heard back from Pennington yesterday about my question about just how much of their lime could be used at a time. I gave them my pitiful numbers (pH 5.3, TEC 10.1). The answer was 18 lb/K annually, which could be done in either one application or two applications of 9lb/K. Still waiting to hear from Encap and SoluCal. I sent the questions to them again. Jonathan Greene's answer about MagiCal was 9 lb/K.

I wonder how much H+ the plant puts into the soil. Never seen any data on that. What might be other sources? As far as high PH soils, they all are sodic or saline soils. I'm sure that's a whole new can of worms.

That's a bit of heavy reading, daniel. I'll read that one when I have more time to dedicate.

The roots exude H+ in their normal operation. From what I read when I was trying to get my head around cation exchange a month ago, the H+ from roots blends into water to make H3O (another hydrogen migration!) and then the H+ can jump onto a cation site from the H3O, turning it back H2O. It seemed to me that H+ was sort of a placeholder on a cation site, essentially nothing there, until a more nutritious substance came along. What does it take to get H+ off a site? With lime, do you think maybe the OH- created in the process helps to attract H+ off a site, letting something else on? One interesting thing, among many, from the Carrow book is that increasing the pH in an acidic soil creates more cation sites. There are pH independent sites (what the soil has anyway) and pH dependent sites which exist depending on pH. In particular increasing pH of soils with Fe or Al oxides, allophane, kaolinite, or humus can raise CEC.

Not all pH soils are sodic or saline. Sodic and saline soils seem to be rare. Lots of people have high pH soil. Some is calcareous, some is not.

Not all pH soils are sodic or saline. Sodic and saline soils seem to be
rare. Lots of people have high pH soil. Some is calcareous, some is not.

My bad. I should have qualified. I was refering to >8-8.5.

As far as increase of CEC with increased PH, the only thing I've seen stated that a proton is removed. No explanation. That makes no sense to me.

Any chance what you read about roots and H+ mention anything about how much H+ is generated by roots?

What does it take to get H+ off a site? With lime, do you think maybe
the OH- created in the process helps to attract H+ off a site, letting
something else on?

My understanding is that that is the basis of the lyotropic series. Dependent in large part on the hydrated radius of each of the elements (and the strength of the charge of the cation--maybe due to the shell that the electron would occupy?) and its relation to the others. However, from what I've read, they can gang up on a an occupied site and displace (squeeze it out by the sum of their total force of attraction) whatever is already occupying the site.

becky said: Thanks for the paper. I read a good part of it, skimming where I would
not understand it anyway. He did not discuss carbon dioxide as part of
the cause of the reaction. Do you think that's because it was not part
of the lime formula (which is the same as Carrow's)?

Well, in plain terms, the real purpose of the study was to see how much calcium carbonate would dissolve in pure (ie, distilled) water with no CO2 present in the atmosphere above the liquid. They intentionally excluded CO2 and the whole carbonate balance, other than what came specifically from solid lime dissolving in the pure water. I was interested in this because I was curious how much increase in pH could be attributed to reactions like you mentioned, CaCO3 + H2O -->Ca + HCO3 + OH. This is the reaction they were studying, called hydrolysis (adding an H to the CO3). The other reason pH increases in acidic soil when you add lime is the reaction of the H+ already in the soil with the solid carbonate. This reaction off-gasses CO2 and turns the H+ into water, causing a rise in pH.

If you read the conclusions on pg 37 they explain how the theoretical, lab measured pH of 9.5 from just hydrolysis of calcium carbonate is about impossible to achieve without perfect conditions. They also found that naturally occurring carbonates were much less likely to hydrolyze than lab precipitated ones. (I've seen this happen first hand. When you make a solid in the lab, it can be gelatinous with very small colloidal particles. These solids dissolve much more readily than an equal solid take from nature.) As if that weren't enough, they also show how the lime gets coated with gunk in the soil and has an even harder time hydrolyzing (although this gunk didn't hinder acids from dissolving lime so much). So the bottom line is that even though lime in the lab will get to a 9.5 pH solely through hydrolysis, there is no chance of that happening in a real soil. So that takes us back to the idea that unless lime is allowed to react with an acid, it won't do much of anything to modify pH because it is otherwise too inert.

yard said: I wonder how much H+ the plant puts into the soil. Never seen any data on that. What might be other sources?

I'd bet morpheus has some information on that, if he's so inclined to chime in. Of course you also get H+ from atmospheric CO2 dissolved in the soil solution. I'm sure lots of CO2 is also generated by microb activity, which will result in more H+ as well.

becky said: From what I read when I was trying to get my head around cation
exchange a month ago, the H+ from roots blends into water to make H3O
(another hydrogen migration!) and then the H+ can jump onto a cation
site from the H3O, turning it back H2O. It seemed to me that H+ was sort
of a placeholder on a cation site, essentially nothing there, until a
more nutritious substance came along.

I'd be surprised if the bold part were happening. From what I've been reading on how cation exchange sites actually hold on to cations (weakly), It doesn't appear that the adsorption of H3O+ onto a site would involve the stripping of the H2O. I think the whole H3O+ molecule, typically abbreviated as H+, is adhered to the site.

I would think of it like this: atoms bond in different ways (covalent, ionic bonds) with each other to form strong bonds and become molecules. In aqueous solution all these dissolved molecules (and also individual atoms as cations) are going to be hydrated in one way or another. Water is polar and is going to stick to anything else that has any charge (or even a neutral molecule that has polarity). It is like the Mars rover landing where the rover is surrounded by big balloons that protect it while bounding along the surface until coming to rest.

Of course you also get H+ from atmospheric CO2 dissolved in the soil solution.

Remind me, How does CO2 produce H+? Forgot that roots also release CO2.

Last post for awhile, promise...

becky said:What does it take to get H+ off a site? With lime, do you think maybe
the OH- created in the process helps to attract H+ off a site, letting
something else on?

In an acidic environment, I think I've learned more precisely what is happening. If you have H+ on an exchange site, the soil is probably also acidic and you have free H+ in solution. When you drop a particle of CaCO3 in the solution what happens? Two of the H+ ions in solution pry apart the Ca and CO3 solid and surround, or hydrate them with water. In other words, acid dissolves lime into Ca2+ and CO3(2-). However, in this environment, the anion is immediately converted to H2CO3, but because this carbonic acid would much, much prefer to be CO2 gas and water instead, the acid off-gasses the CO2 and liberates a water molecule. So in essence adding lime had the effect of turning some of the H+ acidity into water, thereby increasing the pH. Meanwhile, the H+ on the cation exchange site is in equilibrium with the H+ in the soil solution... less H+ in solution means less H+ at the exchange site, so the Ca2+ liberated from the lime now has a place to go.

I think this is also the basis for your statement about higher pH's increasing the CEC. I'm not sure, but I don't think more sites are actually being created at higher pH's. I think you just have fewer H's occupying those sites, which are then freed up for mineral cations.

last, last post for awhile: :o)

yard said: Remind me, How does CO2 produce H+?

CO2 from the atmosphere dissolves in water. A small amount of that CO2 is converted to H2CO3 by reacting with H2O. The H2CO3 is a weak acid settling in around pH 5.6 in pure water. That provides the H+.

Here is Carrow's explanation of pH independent and pH dependent exchange sites:

"Negative charges arise during the formation of layer silicate clays when a cation is replaced by another cation of similar size but lower charge. This is termed isomorphic substitution. Typical substitutions are AL+3 for Si+4 and Mg+2 or Fe+2 for AL+3. This negative charge does not change when soil pH is changed; it is called pH independent.

"Some of the CEC sites (i.e., negative charges) do change as soil pH is altered. These are variable or pH-dependent charges and are associated with the “broken edges” of silicate clay layers. On these edges >Al-OH and >Si-OH groups are exposed. Under high acidity (i.e. low pH), the -OH group remains intact, but as pH increases, hydrogen dissociates from the -OH and a negative charge is formed.

>Si-OH becomes >SiO- +H+

"A second source of pH dependent charge as pH increases is by removal of positively charged complex aluminium hydroxy ions. Blockage of negative charges by AL(OH)2+ ions occurs at low pH; thereby, they are unavailable for cation exchange. As pH is increased, AL(OH)2+ ions react with OH- ions in the soil solution and form insoluble AL(OH)3, which free the negatively charged sites. Only about 5 to 10% of the CEC sites are pH-dependent in 2:1 clays. But 1:1 clays (kaolinite) may have 50 to 90% of the CEC sites as pH-dependent (see Table 4.4).”

Go to that chart and here are percentages of pH dependent sites (there is also total cmol kg-1 for these but I’m not typing that), expressed as % of total CEC
Vermiculite 5-10%
Smectites (montmorilionite) 5-10%
Fine-grained micas (illite) 5-10%
Chlorite 15-25%
Kaolinite 50-90%
Goethite (Fe-oxide) 100%
Gibbsite (Al-oxide) 100%
Allophane 100%
Humus 90-100%
Thatch 90-100%

Skipping ahead to another chapter there is this info:
“In summary, many soils have a CEC component that does not vary as pH changes (pH independent CEC) and a component that is altered as pH changes (pH dependent CEC). Soils with the most pH dependent CEC include Fe or Al hydrous oxides, allophane, kaolinite, or humus. Increasing pH of these soils from pH 4.0 to 5.0 up to pH 6.5 to 7.0 results in significant increase in CEC. Conversely, acidification of soils containing pH dependent CEC results in a significant reduction in CEC. Thus, excessively acid soils containing these colloids cannot retain as many nutrients. However, excessively acid soils with 2:1 clay types do not exhibit loss of CEC since their CEC is pH-dependent."
virginiagal Posts: 583Joined: January 18th, 2014, 4:26 pmLocation:Richmond VAGrass Type: tall fescue

Yardtractor, this webpage discusses roots giving off hydrogen as part of the nutrient exchange at the root.

Daniel, so H+ is actually just the extra H in H3O. That makes it easier to visualize--easier to visualize a water molecule clinging to a soil particle than a teeny tiny proton of a gas taking up space there. That's why I was thinking of it as a placeholder on a cation site, essentially nothing there. An acidic soil would be one where the soil solution has a lot of H3O. An alkaline soil would be one where there is more OH-. I can't visualize OH-. Is it wet like water?

Thanks for the links Becky and daniel. It appears that I'm going to be doing a lot of reading. Need to set aside some time.

becky: while H+ is considered to exist as H3O+, I don't believe there is anything like that for OH-. It exists in solution as is. However, the exact chemistry is complicated and difficult to figure out. They're always studying it to learn more. Nothing obvious, though. As a side note, H2O is shaped like a boomerang, with positive charges on the ends from the H+ and negative in the middle, at the "outside" part of the boomerang where the oxygen atom is. This is why water is polar. Water will surround whatever is in solution based on the charges. That 's why they say "like dissolves like." Polar solvents dissolve polar solutes and non polar in to non polar.

You can also look up Van der Waals forces if interested, but not really necessary.

I'm going to read more about your Carrow CEC discussion. I'm curious exactly how the cations are bound to the exchange site. ie, what kind of bond is it?

There's only so much a non chemist can absorb. What I think I understand is that OH- is often part of something and it can dissociate to be OH- by itself, until it meets up with H+ (H3O) to form plain water. Is that right? In a solution, there could be H20, H3O, OH-, in various combinations. If H3O and OH- are there in equal amounts, they will merge to plain H2O. Is that how it works?

The Carrow book can be found in Google books with pages deleted. Still a lot there. The pH dependent CEC pages were some of the deleted ones. Google pH dependent CEC and you should be able to find more about it. I don't think the kind of bond is different, just that more sites are available to receive cations in those kinds of soils when excess acidity is neutralized..

Yardtractor, there is a table in Carrow's book on acidity and basicity of fertilizers (p. 329, one of the deleted pages from Google books' copy). It's derived from a study by W. H. Pierre in 1933. I found this chapter in another book (from Google books) which talks about it. Chart on page 80. Oh, I just found Pierre's study. It seems that neither potassium chloride, potassium sulfate, nor potassium magnesium sulfate affect pH. Another of my questions answered.

Plants do produce a lot of H+ under many circumstances, but it does depend. They're also going to pump out negatively charged ions (which could be an OH-) under other circumstances.

Consider. A plant is busily going about its day (this is not going to be a deeply scientific post as I don't have time right now) and nibbling on passing sunlight. It feels the need for calcium, so absorbs a calcium ion, Ca++ (no, it's not that simple, either).

It would have just picked up two positive charges, which would make absorbing any other positive ions more difficult (and eventually result in a static discharge from the plant to the soil). Not good. So it has to jettison two positive charges to get back to neutral. Not to mention that this is how the cellular ion pumps work anyway.

Highly available and essentially valueless (a lot gets used, but there's boat tons around) hydrogen (H+) fits the bill nicely. So throwing away two hydrogen ions (two protons) will get you square with the charge situation.

But now consider that you're a plant that's absorbing nitrogen. Some plants can absorb urea (no charge) directly, also directly resulting in no net charge trading (it does happen later but that's another story).

Or you could absorb an ammonium ion (NH2+) and dump out a hydrogen ion to compensate.

Or a nitrate ion (NO3-) and...crud, that's negative. You have to dump another negative ion to get rid of that...so an OH- would fit the bill, but so would any number of other negative ions floating around.

That last case results in what could be an effective pH rise in the rhizosphere.

TL;DR...you can always look up the lime requirement of anything you're using in the garden, although those are more available for nitrogen fertilizers as most other things either do not cause a pH shift, or don't do so via interaction with the plant.

becky said: What I think I understand is that OH- is often part of something and it
can dissociate to be OH- by itself, until it meets up with H+ (H3O) to
form plain water. Is that right? In a solution, there could be H20, H3O,
OH-, in various combinations. If H3O and OH- are there in equal
amounts, they will merge to plain H2O. Is that how it works?

I can go over some basics. I'm not sure what you know and what you don't, so don't shoot me if I go too simple in my explanation. So in no particular order, addressing your comment above:

OH- is called the hydroxyl group. It exists as part of many compounds, or molecules. Most any substance that ends in an "-ol" has a hydroxyl group in its structure. In fact alcohols are called alcohols because of this OH group (I'm going to leave off the negative sign unless needed for simplicity). Methanol, ethanol, propanol, butanol (butyl alcohol) are some examples. Classes of compounds like phenols and glycols also feature the hydroxyl group. In some cases the hydroxyl group in these molecules acts similarly to what we have been discussing, ie, the hydroxyl dissociates and can be replaced by another anion like chloride. (There are cases where the OH is present in a molecule, but actually sheds the H and acts like an acid, such as in acetic acid, or any acid containing a carboxyl group, but don't let this confuse you. There are always exceptions to the rules).

In our discussion about soil chemistry we've been discussing inorganic chemistry and acid/base chemistry for the most part, and much of that kind of chemistry occurs in water, or usually termed "aqueous solution." The hydroxyl group in our situation is present in water and is also introduced to the soil from sources like NaOH, so we are talking about OH floating around free in solution. You say that if H and OH are present in equal amounts, that they will form water. Not necessarily, as follows:

To make my point I have to briefly discuss the dissociation of water and pH. (As a side note, you should look up the term, "stoichiometric ratio." When you say you have equal amounts of something, or if you say "2 H's will combine with 1 O to make water," you are describing a stoichiometric ratio. It is the exact amount needed of each reactant). OK, so if I have a glass of pure water it WILL have equal amounts of H and OH floating around, even though you think they should combine to make H2O. It is an extremely small amount of each, just 1x10^-7 mol/liter. The water ionization constant Kw is derived from the concentrations of H and OH by multiplying the two concentrations together, 10^-7 x 10^-7 = 1x10^-14. This is where the pH scale comes from. The 14 is the max on the pH scale. So the important thing to know is that when you multiply the concentrations of H and OH in water together, they must equal 1x10^-14. Depending on what else is in the water, sometimes you will have excess H and vice versa. If you have 10^-4 mol/l of H ions, then you know through algebra that you must have 10^-10 OH ions. So to answer your question, H and OH will not always react to form water when they see each other. There will always be some of each in aqueous solution, and this is the basis for the pH scale. Even though pH below 7 is dominated by H, there is still a corresponding amount of OH floating around (remember that 1x10^14 ionization constant). But you could say, "Yeah but 10^-7 is so small it doesn't even matter." Well it is a small amount, but the whole pH scale from 1 to 14 still deals with very dilute solutions, so it does matter.

I thought of another way to put it. Let's make it "Fisher Price" easy by thinking in terms of individual atoms. Lets say in a given hypothetical beaker I have 100 H ions and 100 OH ions floating around in water. Let's also say the ionization constant for water requires these 100 + 100 = 200 ions to coexist without combining to make water. What would happen if we then drop 20 molecules of HCl into the beaker? Well, you could say that the 20 H ions will try to react with the free OH ions to make 20 waters. OK, then we would have 100 original H and now only 80 OH, plus 20 Cl in solution. (Note: The solution always has to be electrically neutral and the 80 + 20 Cl anions balances the 100 H cations, like when morpheus said if grass soaks up a Ca2+ ion it has to kick out 2 H+ ions to remain neutral.) But now we only have 180 total free H and OH, and our Kw said we need 200 ions. So more water will ionize into H and OH until you have 200 of them again. That means 10 water molecules dissociate to make 10 H and 10 OH. That gives us a total of 110 H and 90 OH = 200. OK, so now our ionization constant is happy because we have 200 ions coexising again. The difference is that we have 20 more H than OH, causing the solution to become acidic. Hopefully this give some idea of how and why H and OH coexist in solution even though you think they should be reacting.

Let me know if that makes sense, or confuses you more!

Wonderful explanation! Thanks!

Thanks! I edited to say there are 20 more H than OH (not and) in the second to last sentence, just to be clear.

What I'm learning from this discussion is that if pH needs adjusting, it's the H+ and OH- and what happens to them after applications that matters, not the cations themselves. We could add lots of calcium via gypsum to an acidic soil and it won't help the pH rise. We need the lime for the carbonate action to make pH rise. So though it's nice to have cations balanced, it's not the balancing of cations that corrects pH, it's the attached substances that do (or don't do) the job. Adding potassium sulfate or potassium chloride or magnesium potassium sulfate will not affect pH one way or the other. Most nitrogen fertilizers will have an acidifying effect (ammonium sulfate the most acidifying) but some have an alkaline effect (some of the nitrate forms). Phosphorus fertilizer doesn't affect pH unless it has a nitrogen form mixed in. The reason to avoid lime in an alkaline soil is not because it will raise pH even higher but because you could end up with too much calcium and that's hard to get rid of. If these are not the correct conclusions, please set me straight.

We could add lots of calcium via gypsum to an acidic soil and it won't
help the pH rise. We need the lime for the carbonate actio to make pH
rise.

There's a typo there, but what I think you are saying is that lime will increase pH in an acidic soil because of the carbonate anion. (the positive cation always associates with the negative anion. Carbonates, nitrates, chlorides, sulfides, etc. all have negative charges and are all anions.) The acid in the soil reacts at the surface of the solid lime to liberate CO2 and H2O (after becoming H2CO3 briefly, first). This serves to remove H+ from soil. For extra credit, can you explain why CaSO4 will not impact pH?

The reason to avoid lime in an alkaline soil is not because it will
raise pH even higher but because you could end up with too much calcium
and that's hard to get rid of.

Yes. In addition, all that carbonate will buffer any attempt to lower the pH and make it that much harder to do.

You are correct in that all that matters in determining pH is the amount of H+ (and corresponding OH-) is in the soil solution. After reading up on soil chemistry, I fully understand why you are confused on this point. Like I was discussing with azdoctor, I think, soil chemists have their own, slightly incorrect, language. Read anything on the internet concerning "base cations" and they have you believing that calcium raises pH! I think what happens is that in regards to soil, calcium is pretty much always in carbonate or sulfate form, as is magnesium, sodium and potassium, so the short hand is just to say that these elements will raise pH. I don't know how many people understand that it is the anion, not the cation, that is responsible for interacting with H+ or OH-.

Regarding nitrogen fertilizers being acidifying or "alkalyzing":

In the case of ammonium salts like ammonium sulfate, it is the 4th H+ on the ammonium cation that is important. It acts just like the H+ in HCl in acidifying a solution. So when the nitrogen is part of the cation in a fertilizer it will form the ammonium cation. (I'm not a fertilizer expert but I imagine this is mostly correct). I'm not sure which ferts produce a basic solution due to a nitrate. The nitrate anion is like the sulfate anion, and so won't affect pH. (That's a hint for my quiz question above!)

I fixed my typo: action instead of actio.

I'm cheating by copying an answer to your test question (maybe I get credit for searching skills?): "Gypsum is the neutral salt of a strong acid and strong base and does not increase or decrease acidity. Dissolving gypsum in water or soil results in the following reaction: CaSO4·2H2O = Ca2+ + SO42- + 2H2O. It adds calcium ions (Ca2+) and sulfate ions (SO42-), but does not add or take away hydrogen ions (H+). Therefore, it does not act as a liming or acidifying material. The Ca2+ ions simply interact with exchange sites in soil and sulfate remains dissolved in soil water." Source is here.

The fertilizers listed as having a small basic effect are calcium nitrate, sodium nitrate, and potassium nitrate.

I've seen a number of bare statements that Potassium chloride raises pH.

Scroll down to potassium nitrate. Unfortunately, No explanation.

here

I also found this:

Nutrients affect PH in root zone

Edit: Many of the Reference links are active.

OK, you get partial extra credit! Actually it isn't totally fair when I know the answer I'm looking for but you don't. I bring this question up because I think it is important and interesting to understand.

In my question I asked why CaSO4 won't affect pH, while CaCO3, in an acidic environment, will. First off, it isn't a good question because I have found that CaSO4 won't dissolve in water, or I think even in acidic pH's. So let's go with NaSO4 which is soluble in water. I was going for the SO4 anyway, so the cation doesn't really matter.

The answer has to do with whether the anion (carbonate or sulfate in this case) came from a weak or strong acid. Not that the anion literally came from the acid, but is the anion associated with a weak or strong acid. For instance, if you had NaCl you could say that Cl in acid form would be HCl. Or, if you had CaOH you could say that OH in base form could be NaOH. Sulfate, SO4, comes from sulfuric acid, H2SO4, which is a strong acid. This means that if you have a salt of sulfuric acid, like NaSO4 (look up "conjugate base") then we can predict something about SO4- when it dissolves in water. Since sulfuric acid is completely dissociated in water, that tells us that SO4- has very little ability to hold on to an H+. That means it won't try to get H+ by extracting it from the free H+ floating around in water, and it WON'T raise the pH. On the other hand, we know that carbonic acid, H2CO3, is a weak acid. So if we have a salt like NaCO3 in solution, the CO3 WILL be strong enough to suck H+ out of water and form HCO3 and H2CO3. As mentioned earlier in this thread, removal of the H+ from water will allow free OH- to increase, which ups the pH. So the secret is understanding that salts from strong acids will not be able to capture a H, while salts from weak acids will be able to. (The reason they are weak acids is that they don't like to let go of their H.) If you strip the H by making a salt, and then dissolve it in water, it will go looking for the H in the pure water, raising the OH concentration.

Likewise for nitrates, I'm not sure why fertilizers raise the pH with them. Nitric acid is also a strong acid, so I'm not sure if they do raise pH. I did read one thing like what morpheus said: It said that if a root sucks up a nitrate, it will spit out an OH-, thereby increasing the local pH at the root. But that is more of an application beyond the basic chemistry and illustrates how interaction in the soil with plants and microbes can change what might be expected from theory.

Said more than enough, again. :) I guess it is fun relearning some of my long ago chemistry.

Yard, we were cross posting. It appears from your link that my final comments (from morpheus in part) were correct. It isn't that the fert in itself is acidic or alkaline, but its about what happens when the plant absorbs the nutrient. In the second link you provide, look at the paragraph heading: Nutrient Sources Affect Soil pH in Root Zones.

It explains it quite well. Did you mean to say KCl raises or lowers pH? According to the link, it lowers pH. The root takes up K+ but doesn't need very much Cl-. So it has to kick out an H to keep the electrical charge balanced. Likewise with potassium nitrate, the final pH near the root depends on whether the plant needs more K or NO3. If it takes up more positive charged K then it will become more acidic, and vice versa more basic if it takes up more NO3-.

Very timely contribution from my perspective. Thanks.

Did you mean to say KCl raises or lowers pH?

I'm saying that I've read a number of articles that state that KCl raises pH, but they don't give any explanation. It's often given as the rational for recommending potassium sulfate, but this article contradicts that. Have you or Becky comes across anything that explains why >8.5 soils are sodic? Everything I've seen implies that for a soil to have a ph greater than 8.5 they need to be sodic.

In the second link you provide, look at the paragraph heading: Nutrient Sources Affect Soil pH in Root Zones.

You found the exact section of the article that I found interesting.

Very timely contribution from my perspective. Thanks.

I owed you and Becky a link or two. :)

Edit: Considering the mass of turf roots in the top zone 0-4" of the soil, it would appear that the turf (plant) has a significant influence on pH and that the nutrient ratios would influence that--BPN.