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!