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Why and how does lime work? A chemical discussion of lime and pH in soil.
Posted 12 January 2011 - 05:22 PM
As a chemist, I was very curious as to why lime (dolomite lime) works in soils. At first glance, it didn't seem like it should, as I will explain below. Note that this explanation is necessarily chemical in nature, but I will try to explain things as I go.
Before going further, I want to briefly explain chemical equilibria. If you're fluent in this area, you can skip this paragraph, otherwise, read on. Most chemical reactions are reversible, which means that they can go either way (forward or backward). This reversibility is repeatable, meaning a given chemical reaction will always perform the same way. Chemical reactions are described by a chemical equation, and the equilibrium constant tells the chemist which "side" of the reaction is more favored. The equilibrium constant is defined for a given reaction by the ratio of the concentration (or activity, or partial pressure, depending on your measurement system) of products of a reaction, multiplied by each other and taken to the power of their coefficient in the reaction, divided by the concentration (activity, partial pressure, etc.) of the reactants. An equilibrium constant is just that - a constant, and is unchanged at a given set of conditions. This means that a chemical reaction will do whatever is necessary to achieve the equilibrium defined by its equilibrium constant. This is important later on . What the equilibrium constant allows one to do is figure out what will happen when a given set of chemicals are put together. The equilibrium constant is usually given the symbol Keq ("eq" is subscript); for acids and bases, the equilibrium constant is known as Ka (the "a" should be subscript), which describes the "acidity" or "basicity" of the given chemical. The larger the Ka, the more acidic. Back to the main discussion...
The idea behind lime is that it takes acidity in the soil and neutralizes it, and the result is that the pH becomes neutral, or about 7. Some people think lime "buffers" the soil. This false belief will be explained below.
Limestone is composed primarily of magnesium and calcium carbonate, along with some other minerals. These have the formulas: Mg(CO3) and Ca(CO3). Note that when dissolved, these will create carbonate (CO3) (2-) and the respective 2+ metal ions.
However, when you look up the Ka2 for carbonic acid, the Henderson Hasselbalch equation tells us that the pH should be in the 10.25 range (Ka2 is 5.6E-11; (-)log of this value is 10.2518) (note, this value is a bit different in a 50/50 mix of carbonate and bicarbonate in this publication: http://www.biochemj....245/0390245.pdf, however the pH of the initial water will affect this value). If lime is used in soil, how could it be that the pH of the soil would end up as "neutral?" Reaction of acid with the carbonate would yield hydrogen carbonate (aka bicarbonate, or the anion of baking soda) so that the buffer system would be around 10.25 assuming equal parts bicarbonate and carbonate.
CO3 (2-) + H+ <--> HCO3 - Keq = 1/Ka2 (carbonic acid) = 1/5.6E-11
[An interesting sidenote: the buffer solution I bought with my pH meter uses this particular system as the pH = 10.0 buffer! This served as some of the motivation for this post.] If the system is composed of carbonate and bicarbonate, we will have a buffer, and at way too high of a pH!
But wait...the bicarbonate is still basic, i.e. it can accept a proton (Arrhenius base; in the same way that baking soda is a base!). However, because of the large separation of Ka values, this will happen only AFTER carbonate has been made into bicarbonate, i.e. no acid will react with bicarbonate when there is carbonate still around.
HCO3 - + H+ <---> H2CO3 Keq = 1/Ka1 (carbonic acid) = 1/4.3E-7
Thus, acid will convert carbonate to bicarbonate, then bicarbonate to carbonic acid.
The key to this system is that carbonic acid (our final product here) spontaneously decomposes into carbon dioxide and water.
H2CO3 --> CO2 + H2O
So as the carbonate is converted to bicarbonate and the bicarbonate is converted to carbonic acid, it is effectively "evaporated" from the soil!
So why not just add bicarbonate (baking soda) to the solution or to the soil right off the bat? If bicarbonate is the species that is created by lime reacting with natural soil acidity, why bother with lime and just instead use good ol' baking soda?
The answer is: you could, but there are two problems. The first is that you'll be introducing a lot of sodium into your system (soil). That's not good. The second is that, in the absence of acid, you would get a pH of soil around 8.2, which is far too high. Like I mentioned above, bicarbonate (we call sodium bicarbonate "baking soda") is a base, and will raise your pH.
But wait...if this is true, wouldn't carbonate be even more basic? Shouldn't it be worse? So...why does lime work?
The issue here is the solubility of Calcium/Magnesium carbonate. They have respective solubilities of 4.5E-9 (Ca) and 3.5E-8 (Mg); this is very low. In other words, there is not "free" carbonate in solution, the magnesium/calcium carbonate don't dissolve very well and as a result, CO3 (2-) is only released when it encounters acid. [Calcium carbonate is also "chalk," have you noticed that chalk doesn't dissolve very well in water? But the acid rain we now have around the world does eventually remove the sidewalk art that children draw, or outlines of dead bodies if you want to be dark...]
When you use calcium carbonate, and react with acid, you get a complex solubility problem; i.e. two equilibria are taking place simultaneously...dissolution and acid base reaction:
CaCO3 <--> Ca2+ and CO3 (2-)
CO3 (2-) + H+ <---> HCO3-
Keq (total) = Ksp(=4.5E-9)* 1/Ka2 =~ 80.
This means this reaction will go (though 80 is not that huge of a Keq); this is because the calcium carbonate doesn't WANT to dissolve, but it is immediately "eaten up" when it does dissolve that, by Le Chateliers principle, by "removing" carbonate in the first equation, the reaction will continue to shift to the right to ensure chemical equilibrium, so the calcium carbonate will continue to dissolve. Thus, carbonate is liberated and turned to bicarbonate, but note that the acidification of bicarbonate to make carbonic acid is much more likely to happen next (because of the small Keq here and the much larger Keq for the acidification of bicarbonate). So what will happen is that as you dissolve carbonate, you create bicarbonate, this in turn becomes carbonic acid, and then the carbonic acid will dissipate naturally.
Thus, a quick fix for a very low pH soil seems like it would be to add a baking soda solution, since this will raise the PH. However, you can easily "overshoot" and end up with a high pH, and you are unnecessarily adding sodium to your soil!
On the other hand, unless there is acidity, the lime won't do anything. People who say they have a high pH value because of lime added are blaming the wrong suspect...the lime is insoluble in basic solutions, so it will not raise the pH. These are first cousins of people who say that lime "buffers" the soil, which it does not do (from a strictly chemistry point of view). A buffer is created when there is a sufficient amount of a base and conjugate acid, or acid and conjugate base. In this case, we would be talking about a roughly equal proportion of CO3/HCO3 (which made my buffer solution, above!) or HCO3/H2CO3. This will not happen in a lime solution, no matter how much acid you have, because the acid will react with the carbonate to make bicarbonate, and when there is any bicarbonate around, the acid would much rather react with it until it is all converted to carbonic acid...which, as we know, will naturally degrade. So you cannot have both carbonate and bicarbonate in any appreciable amounts together, nor can you have bicarbonate and carbonic acid in appreciable amounts. Thus, lime does not "buffer" the pH, but it prevents it from getting too low. A buffer would protect against both positive and negative changes in pH, and lime will not prevent pH from raising from some other factor (e.g. you're stoned and put "quicklime" or, essentially, hydroxide [a strong base] into the soil!).
I hope this discussion explains the use of lime and helps guide people with their soil mixes
Some other questions: can you use too much lime?
Definitely. Not so much from an acid/base point of view, but we can't forget about our "spectator" ions...Mg and Ca. As the soil continues to be buffered, carbonate used up, etc., Magnesium and Calcium ions will be left in the soil. The result is a salt buildup of these ions, which can mess up the soil! Some Mg and Ca is good for the plant, but too much can be a problem.
How can I get rid of excess salt?
This question is tricky. One way is to add a complexing/chelating agent, such as EDTA, which will bind up the ions. This is hard for a normal person to come by, but it would work, and prevent "salt buildup" in a classic sense (the ions are no longer free to be absorbed by the plant, etc.). If present in the ionic form, you can wash them out of the soil, but of course in the process you will wash out nutrients, etc. The main point is that you want to use the minimum amount of lime you can get away with to prevent acidification of the soil without adding so much as to overload the system with ions. This is why watering with a neutral pH water is important... you don't want to liberate that lime you added, freeing too much Mg and Ca!
I welcome your comments, corrections, and discussion!
Posted 12 January 2011 - 06:40 PM
Could you explain hydrated lime a bit? I'm at a loss with the "why" when I tell people (especially inexperienced growers), to avoid it, other than sad experience. LOL
Posted 13 January 2011 - 12:23 AM
+rep for you! Great info. This would make a great sticky.
Could you explain hydrated lime a bit? I'm at a loss with the "why" when I tell people (especially inexperienced growers), to avoid it, other than sad experience. LOL
Thanks very much.
Hydrated lime is calcium hydroxide. Limestone (from which dolomite lime comes from) is heated and converts to calcium oxide, CaO. If you wanted to, you could take your bag of dolomite lime and heat it to make calcium oxide. Adding this to water would yield calcium hydroxide, Ca(OH)2, so people who have dolomite lime could make some hydrated lime if they were so inclined. I would NOT recommend doing this because of dangerous side reactions that might occur because dolomite lime is not pure 100% calcium carbonate. Do not try this at home!
Hydroxide is a strong base. Calcium hydroxide is much more basic because instead of dissolving in solution to give the calcium cation (the metal ion) and carbonate, it dissolves to give the calcium cation and the hydroxide anion.
Hydroxide is the strongest base that can exist in water. Anything that is "basic" is so because it can abstract (remove) a proton from a water molecule, leaving hydroxide:
H20 + A- <---> HA + OH-
Generally, the A- is a weak base, so the equilibrium here lies to the left side, so not much hydroxide is made. However, it doesn't take a lot of hydroxide to dramatically affect the pH. I'll take a step back to explain this...if you are up to snuff with your pH chemistry, you can skip the next 2 paragraphs.
pH is a measure of the concentration of H+ in water (n.b. there is much discussion and disagreement in the chemical literature as to the nature of protons dissolved in water...generally, the blanket term "hydronium ion" is used and represented as H3O+, which would be a water molecule with an extra proton attached. The theory is that free protons don't exist in water, and that they are tied to a molecule, and there is some thought that these water molecules are actually part of a larger moiety, held together by noncovalent bonds involving multiple water molecules [up to 9? not sure here ]. Just think of it as protons, it makes life easier). The pH scale is defined as the negative logarithm of the concentration of H+ in solution. So if you have 1E-04 moles of H+ in a 1 liter volume of water (0.0001 moles/L, 0.0001M), you have (-)log(1E-04) = pH = 4. It is clear that it doesn't take a large concentration of H+ before you are getting very acidic. What's wild, and a lot of fun to think about, is that the pH scale is a measure of the hydroxide concentration as well. This is because water has an autoionization constant:
H2O <---> H+ + OH-
Which is another chemical equilibrium! This is Kw (w is subscript), which is 1E-14. Remember earlier I said that chemical reactions always are trying to reach equilbria. So when you have a given concentration of H+ in solution, this means that there has to be a corresponding amount of OH- in solution. To solve, we set Kw = [H+][OH-] (brackets indicate "concentration of") and we find that the concentration of OH- is 1E-10. This solution therefore has an analogous pOH, which here would be 10. pH = 14-pOH. Note that this reaction lies far to the left side, so most free hydroxide in solution will instantly react with protons and make water. This is why hydroxide is such a strong base (really, the only base in water), and why it only takes a small bump in the concentration of hydroxide to have a large affect on pH!
So calcium hydroxide will release OH- into solution, dramatically raising the pH. However, like it's cousin calcium carbonate, quicklime (or slaked lime, aka calcium hydroxide) it is poorly soluble (though more soluble than calcium carbonate = dolomite lime) in water. What does dissolve will dramatically raise the pH, because free hydroxide will be present in solution (rather than the carbonate reacting with water to liberate small amounts of hydroxide).
The kicker is that the rest of the quicklime will not dissolve right away, and will, like lime, slowly leach out over time. This is why just rinsing the soil with quicklime in it will not do anything (assuming you're using regular pH neutral water). Remember that we are dealing in logarithm units...so let's illustrate.
Suppose you have a 1 liter container (let's make the numbers easy) of soil. You put some quicklime in there, who knows how much You get a pH reading that is 9. Whoa! This means that the pOH = 5, which means that the concentration of hydroxide in the container is 1E-05. OK, you say, let's just get rid of this quicklime by washing through. You wash through with 2 liters of water (which would, if this was soil, remove a lot of good things with it!). Let's assume that you were smart and saw that the solublility of quicklime in water is 0.173g/L. Which means you added 0.173g in the soil (a very small amount...probably about the size of 1/2 a dime as a powder). Therefore, it's all dissolved. However, when you wash through with 2L of water, you're not perfectly washing through, there is mixing that occurs, so really you're doing some combination of washing through (i.e. completely displacing the water present) with mixing and diluting the water present. Again, let's be optimistic and let's say you managed to decease the concentration 10 fold with your 2L wash. Now you concentration is 1E-05 / 10 = 1E-06 --> pH = 8. This is still no good.
In real life, you would use MUCH more than 0.173g in a liter because it is difficult to measure such a small amount (and consequently, you can see that if you need to use quicklime to make drastic pH changes, using the smallest amount you can is best!). Furthermoe, you wouldn't have 1L of water when there is soil in the container, so not all of the 0.173g would dissolve right away anyways! If more than can dissolve is present, it means that as you wash your quicklime from the soil, more would dissolve to take its place, and you have to keep diluting and diluting. Secondly, you probably wouldn't regularly wash through with water that is the volume of your container (times 2!). So in this case, dilution is not the solution to pollution! The solution in a dire strait such as this would be to use some acid to neutralize the quicklime until the pH came to a proper range, then go straight to the store to buy the right lime!
Hope this provides some further insights. Again, questions and comments encouraged!
Posted 21 January 2011 - 08:25 PM
I'm planning to mix 5 gallons of soil (for a single plant grow) as follows:
12 qts foxfarms ocean forest, which is "buffered" to pH 6.5 - 6.8 (I understand it's not really buffered, but adjusted),
6 qts worm castings
2 qts perlite
The FFOF has some dolomitic lime already in it, but I don't know how much. What would be an appropriate amount of dol. lime to add to this mix?
Another question -- I usually add a couple tablespoons of bone meal for the phosphorous and calcium. Is this inadvisable in conjunction with the lime? I'm not sure how available the calcium in bone meal is, or how long it takes to break down. I feed with organic nutes that have P, so I'm thinking I should skip the bone meal if I add dolomitic lime to the mix.
Posted 21 January 2011 - 09:44 PM
Since FFOF has 'some' dolomite in there, perhaps 1/2 this amount? FFOF does seem to run out of lime in flowering.
Posted 21 January 2011 - 10:07 PM
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