Tech Talk: What is this Alkalinity? Part I

In Pumproom Press 21 we looked at the different meaning of the words alkali, alkaline and alkalinity. But we really still don't know is what "alkalinity" is, how it is defined, and why it is so important for pools.

Before the concept of alkalinity can be explained, we need to take another look at acids. Remember, an acid was defined by the amount of H+ ions in water (PrP 20). The more H+ ions the solution contains (measured in moles per liter), the lower the pH and the stronger the acid.

We will now define something called "weak acids." Strong acids are completely split – the chemists call it dissociated – into the H+ ions, and the leftover part, called the anions. Some examples for strong acids are:

Depending on the actual concentration, these acids are usually completely dissociated at a pH of less than 1. And now we need another definition: Acids, where more than one hydrogen atom can split off, are called multi-protic acids. Sulfuric acid is di-protic (2 H+ released), phosphoric acid (H3PO4), which can release three H+ ions, is tri-protic. What is important is how easily those H+ ions separate themselves from the anion?

All that leads us to the definition of strong and weak acids. Strong acids dissociate easily and completely. Weak acids are happy to hang onto their hydrogen atoms, and give them up (dissociate) only reluctantly. The only way to force them let go is to increase the pH. You are all familiar with the dissociation curve of one weak acid: hypochlorous acid, the regular pool chlorine. Only when the pH is raised above 5 does it start to split up:

At pH 6 the dissociation is about 5%, at pH 7 it's a little less than 30%, the mid-point is at pH 7.52; at pH 8 it reaches 80%, and finally at pH 9.5 the dissociation is complete. This dissociation is important, because only the undissiated acid is a strong sanitizer. HOCl works more than 100 times better than OCl- in killing germs – the good vs. the bad chlorine in you AFO textbook.

Some other weak acids are: acetic acid (CH3·COOH, vinegar); oxalic acid (HOOC·COOH, found naturally in rhubarb, its salts being the major constituents of kidney and bladder stones); phosphoric acid (H3PO4, the salt of which is trisodium phosphate (TSP), a common cleaner used around pools. Also phosphates are excellent fertilizers, and promote algae growth). Finally there is cyanuric acid (also a tri-protic acid), used often as a so-called stabilizer.

We now come to the crux of the matter: why do we want a weak acid in pool water? If you have been with me so far (not always easy?) you have learned that weak acids fall apart – dissociate – with increasing pH. How do we increase the pH? We add something that swallows up the excess hydrogen ions (H+). The easiest way to do that is by adding a base, with lots of OH- ions. Remember that water itself can split up:

By adding a lot of OH- ions, we can tie up some or most of the H+ ions; consequently the pH rises. It's really all a question of how much of each of the individual constituents is in the water.

What we will do now is look at the amount of weak acids in our pool water, and call all of them together "alkalinity". If the water has low alkalinity, it has only a small amount of these weak acids in solution. What happens with low alkalinity water, when you add a strong acid? Well, the poor few anions floating around are quickly gobbled up by the excess H+ ions from the strong acid, and the pH plummets. If, however, you have lots and lots of these weak acid anions in the water, they'll bind the H+ ions from the strong acid, and consequently the pH will drop only a little bit.

If you hate math, just skip the calculations below and trust me! If you have fond memories of powers and logarithms, stay with it. You must remember from PrP 20 that you get the concentration of H+ ions by changing the sign in front of the pH, and then use it as a power of ten. If the pH is 7.5, then the concentration of H+ ions is 10-7.5 (mol/L).

Let's look at an example (it'll be in metric units, because it's easier, but doing the calculations in gallons would also work). You have a 100,000 L (about 25,500 gallon) pool, at a pH of 7.5. That means you have:

This means you have approximately 0.003 moles of H+ in your whole pool!

Now you add 10 liters of acid with a pH of zero (0) to the water. With 100 = 1, we get: 10 L x 100 mol/L (from the pH) = 10 x 1 = 10 mol.

So your total amount of H+ ions in the pool water is now 10.003 moles; 0.003 mol from the original water, and 10 mol from the added acid. Will you believe it when we say the 0.003 may be safely ignored, and we can round the number to 10?

Let's find out what the resulting pH is (which means we must go back to moles/liter): We have 10 moles in 100,000 liters.

So the pH that you get when you add that amount of acid to a pool with no alkalinity drops from 7.5 to 4! That is much lower than you like. Left there for long and you can kiss your plaster, your heater piping and copper heat exchangers good-bye.

The same calculation for a pool with an alkalinity of 100 mg/L (or ppm) is a bit more tricky. If you are interested, get a college chemistry textbook and look up the Henderson-Hasselbach equation. With it you can calculate the reaction outcome. Or you can simply say: "I believe" and trust the result. The outcome is astounding. If you add the very same 10 liters of pH-zero acid to the same 100,000 liter pool, starting at pH = 7.5 and with an alkalinity of 100 ppm, the pH will drop only to 7.1.

So that's the reason we want alkalinity in the pool water! Chemists call it a buffer, for the effect it has. It helps avoiding huge fluctuations in pH – the famous pH-bounce. With low alkalinity the pH will shoot up when you add a base, and drop dangerously low when you add a strong acid. The alkalinity smoothes out or "buffers" those "yo-yo-like" fluctuations, and keeps things in check.

We'd like to continue this alkalinity-concept discussion ad infinitum – and that's about how long it would take to cover it at the pool-college level – but we need to break here and pick up in the next issue. We will continue it on the website too, so if you're really into alkalinity, www.ppoa.org will be the place to get it sooner. Mostly we'll be wrapping up by going into CO2 in some depth. See you then – meanwhile, measure the pH in a can of soda and see a weak acid in action.

~wh