Chemical Equilibrium is a special property of the reactions and solutions of many substances. Whenever you see a chemical equation with arrows pointing both forward and backward, it is an equilibrium reaction. In an ordinary reaction, the starting materials simply react until one of them is all gone, and then it stops dead. But an equilibrium reaction does not automatically go to completion. Instead, it grinds to a stall (achieves equilibrium: the reverse reaction equals the forward reaction) at some characteristic point, even though there may still be plenty of starting materials left. Every reaction has its own regular balance-point: one reaction may stall after only 0.0000001% of the starting materials have been used up; another may go to the point of being 99.999999% complete. The endpoint characteristic to each equilibrium reaction is defined by a number called the Equilibrium Constant, KEq, which never changes. Whenever the concentration of any the reactants or products is changed (by reacting or adding or removing something), all of the other substances in the equation change in compensation, and they change in such a way as to preserve the value of the KEq. For example, consider the most important equilibrium of all, which is also the simplest:

The Water Equilibrium describes the dissociation of water molecules into their component parts, hydrogen ions and hydroxide ions (H+ and OH):

H2O <=> H+  +  OH KEq  =  10-14

10-14 is a very small number, which means the reaction doesn’t “go very far,” and only about one molecule in ten million is actually dissociated at any instant. In pure water with nothing added, H+ and OH- will be identical in concentration, equal in this simple case to the square root of the KEq, or 10-7. This reaction and relationship are so important that a special shorthand notation has been developed, called the pH system: use the exponent only, drop the sign and use “p” as a symbol, and you get pH = 7.0 for pure water. (The pOH = 7.0, also.) All this is important, because it’s water we’re talking about, but also because H+ is very important to water chemistry. The hydrogen ion is the smallest and most chemically active piece of ordinary matter known. Hydrogen is the smallest of the elements – a single proton with a single electron whizzing around it. Remove the election, and you have a single, lone, naked proton in solution that is highly reactive, invasive, and corrosive. (It’s not really naked and alone; sometimes it is shown attached to a water molecule as H3O+.) H+ defines acidity; H+ is acid. So, to a water equilibrium defines acidity: pH less than 7 is acidic; pH greater than 7 is alkaline or “basic.” If either H+ or OH is increased artificially, by adding or removing acid or alkali, the other changes in the same proportion tain the value of the KEq, which is chiseled in stone.” Thus, if acid is added to pure water to make the pH = 3, the equilibrium shifts so that the pOH = 11 and the KEq is preserved as 10-14.

The Carbonic Acid-Calcium Carbonate Equilibrium is the other “most important” equilibrium in water chemistry. When water falls as rain it absorbs carbon dioxide gas from the air, forming carbonic acid, which instantly dissociates into bicarbonate and carbonate ions:




















When this slightly acidic water soaks into the ground and contacts limestone (calcium carbonate) or dolomite (mixed calcium and magnesium carbonate), it dissolves some of the rock, making “hardness” and alkalinity”:




Ca+2 + Mg+2


(which dissociates As shown above)





“Hardness” is the sum of all ions that react with soap to inhibit lathering and precipitate soap scum or “bathtub ring”. It happens that they are all metals with more than one “+” charge-mostly Ca+2 and Mg+2 in most water supplies, but also including Fe+2, Cu+2, Zn+2, Mn+2, etc., if present.

“Alkalinity” is a confusing term because it came into use before the chemistry was understood. As a word, alkalinity just means the opposite of acidity, or something that consumes or neutralizes acidity. When it was realized that both of the carbonic acids and their ions in the two equations just above are the same thing, “alkalinity” became the term in water chemistry to represent the sum of CO2 + H2CO3 + HCO3 + C03-2, and thus the total buffering power of the water to resist changes in acidity. All the double arrows show that the reactions go both forward and backward, which means that limestone dissolved in one place may deposit as lime scale in another place later, if the equilibrium shifts. The equilibrium can be shifted by adding or consuming any of the constituents involved. Remember that the relationship between reactants and products – the KEq – is chiseled in stone, and any change from the outside will be compensated for, instantly, in a way that preserves the value of the KEq.

The values of the three KEq for the three carbonic acid equilibrium (one for each arrow) are intentionally absent from this discussion, to avoid getting you bogged down in numbers. What matters is that you get a “feel” for the concept of equilibrium. It makes sense that adding more reactants (on the left side of an equation) would make a reaction “go” further (to the right). But with equilibrium, you can also make a reaction go to the right by removing some of the final products. It’s as if the removal creates sort of a chemical vacuum which demands to be filled. The reverse is also true: to make a reaction reverse (go to the left), you can either load up the right side (add more final product) or lighten the left side (remove reactants). It is as though the reactants and products are physically connected by the arrow, and you can push or pull on the reaction from either end. And in chemical systems with several reactions connected in series, like the first one, reprinted here, the “connection” goes through all of them. Note that one of the final products of the third reaction (far right) is H+, or acid:

CO2 + H2O <=> H2CO3 <=> H+ + HCO3 <=> 2H + CO3-2

If you add more acid, the entire three-reaction continuum will reverse (go to the left), causing some CO2 to bubble away and be lost forever, thus reducing the total alkalinity of the system. This is called dealkalization.

About Buffering: Just above, it was said that the presence of alkalinity “buffers” water against changes in pH. Buffers are chemical stabilizers that use equilibrium to establish and maintain a chemical balance. The common “mineral” acids such as hydrochloric, sulfuric, and nitric acid do not work as buffers because their dissociation to produce H+ ions is not governed by equilibrium: they all become virtually 100% dissociated the instant they are dissolved. Because of this, they are called “strong” acids. Only “weak” acids and their salts like carbonic acid and sodium carbonate can be buffers. The KEq of carbonic acid is about 10-7, which means that normally, only one out of millions of carbonic acid molecules is split into H+ and HCO3 ions at any instant. Note that the value of the KEq is close to the value of the pH (and also the pOH) when acidity is “neutral:” pH = 7 = pOH. That means carbonic acid is good for buffering water near pH 7. However, note that the bicarbonate ion, HCO3, still has one hydrogen to lose, and bicarbonate ion therefore has its own KEq, which is about 10-11. That means sodium bicarbonate would be a good buffer for water at pH values near pH 11. For comparison, the KEq of acetic acid (vinegar) is about 10-5, which means that acetic acid would be good for buffering water in the vicinity of pH 5.

So, how do buffers work, exactly? They work by having half of their capability inactive, in reserves, so to speak, from both the acidic and alkaline point of view: When the pH is equal to the exponent on the KEq, 50% of the weak acid molecules in solution (carbonic acid in this case) are whole and 50% are dissociated into H+ and HCO3. If extra H+ is artificially added to the solution, it will instantly be incorporated into the carbonic acid-bicarbonate-carbonate “alkalinity continuum,” with very little effect on the pH. Likewise, if some OH ion is artificially added, it will instantly react with and be neutralized by a H+ ion, but that H+ will instantly be replenished from the alkalinity, again with very little impact on the pH. As long as any total alkalinity remains in the water, it will be able to “absorb” either H+ or OH in the vicinity of pH 7 without changing the pH very much. Incidentally, it is possible to buffer a solution against changes in ions other than H+ – CI, for example.