I’ve always thought alkalinity was the opposite of acidity, and that an alkaloid was simply a chemical with a pH above 7.0. However, a recent swimming pool thread here suggests otherwise, with no corrections in the thread. The Wikipedia article on it seems to agree with me, but isn’t perfectly clear. I’ve been using the term to refer to chemicals with a pH above 7, since base and basic can have several meanings even when referring to chemicals. I’d hate to think I’ve been wrong about something. What’s the truth about this?

I was confused as well when my parents got a hot-tub.

From the web:

  1. The quantitative capacity of water to neutralize an acid; that is, the measure of how much acid can be added to a liquid without causing a significant change in pH.

So based on this definition, a highly alkaline solution could either have a high pH or be a highly buffered solution of pH 7.

Now I’m curious as to what indicators they are using.

you are right. Alkaline ususally means basic. i.e. pH greater than 7. “Alkaloids” are often basic, but can just refer to a large class nitrogenous natural chemicals that may or may not be basic.

Yes, high alkalinity means a tendency or ability to resist acidity…

So, in pool terms, think like this and it will seem clear:

You have clear water and the pH is where it should be. It rains. The pH is now WAY off. Your pools alkalinity sucks. The water cannot resist a swing in pH.

Adjust the alkalinity (add baking soda) and the pool chemicals and rain will not be able to swing the pH very easily.

This is still confusing me. I had a chemist tell me that pH could be described as bacicity. How is that different from alkalinity? It seems clear they are different properties but what makes them different?

Water is H[sub]2[/sub]0. In pure water, a very small proportion is seperated into H[sup]+[/sup] and OH[sup]-[/sup].

pH is a measure of how much H[sup]+[/sup] you have in your water. pH 7 means that you have a ten-millionth (10[sup]-7[/sup]) of a mole of H[sup]+[/sup] per litre of water. pH 6 means that you have a millionth (10[sup]-6[/sup]) of a mole per litre - ten times as much. Each step down the pH ladder means you have multiplied the amount of H[sup]+[/sup] by 10.

Going the other way has the same effect - pH 8 means ten times less H[sup]+[/sup] than pH 7 and so on.

Now for the complicated bit. In pure water, the amount of H[sup]+[/sup] and OH[sup]-[/sup] will be the same. But if you artificially add H[sup]+[/sup], some of that H[sup]+[/sup] reacts with the OH[sup]-[/sup] already present to form water, so the amount of OH[sup]-[/sup] decreases. Adding H[sup]+[/sup] reduces OH[sup]-[/sup] and vice-versa. Push one up and the other goes down. Acidity and alkalinity are therefore tied together, so pH is a measure of both or either.

Strong acids (low pH) therefore have high concentrations of H[sup]+[/sup], and strong alkalis (high pH) necessarily have high concentrations of OH[sup]-[/sup]. Both these species are quite reactive, and are therefore best prevented from getting in your eyes, on your skin or on your favourite metal possession, especially if it’s aluminium, which is attacked by both.

People tend to think of acids as destructive, and this is generally true. Weak acids are more corrosive to metals than plain water, although they tend to be harmless to people (HF being the notable exception, but that is due to the fluoride ion rather than the H[sup]+[/sup].) Strong acids can have various nasty effects - sulphuric has a tremendous affinity to water and will consume it when it finds it with the production of a lot of heat. Very concentrated nitric acid is sufficiently oxidising to ignite paper and cotton when it encounters it. Note however that both these effects are due in part to the sulphate and nitrate ions in these acids, and are not universal to all acids.

Strong alkalis are caustic, and are great for breaking up organic molecules. This is why they are used to get the baked-on crud off your oven (oven cleaner is basically strong hydroxide). This is also why they are nasty if they sit on your skin for any length of time. Paradoxically, strong alkali doesn’t attack steel at all.

So how can the ability of water to resist acidity change without changing the pH?


Alkalinity and Buffering Capacity

This is a textbook definition of a buffer, not a base; Your web site is mistaken.

Alkinity is almost like the opposite of acidity. One definition of basicity is the concentration of hydroxyl ions (OH[sup]-[/sup] while aciditiy is the concentration of Hydronium ions (H[sup+[/sup]).

The way they’re related is simple: pH + pOH=14.

If you have a pH of 3, then the pOH is 11 (the smaller the number, the more concentrated)

The swimming pool guys, and the marine scientists (see my second link) seem to be using the term alkalinity in a nonstandard (as far as I’m concerned) way, which includes buffering capacity in the definition. So in that context, alkalinity is not identical to basicity.

So I understand the concept of alkalinity now. What about the term alkaloid though? Does it refer to basic chemicals, buffers, or both?

alkaloids are organic compounds with an organic nitrogen (Nitrates don’t count) in them. As nitrogen has an extra pair of elecrons it can either donate them (lewis base) or accept a proton (Bronstead base). They are usually mildly basic compounds so they make good buffers.

Nitpick: Isolated H+ ions (protons) are far too reactive to exist freely in aqueous solution. What actually happens is that the free H+ will group with a water molecule to form a H3O+ (hydronium) ion. The equilibrium is between H3O(+) + OH(-) <—> 2H2O.

pH is defined as the negative base-10 logarithm of the [H3O+] concentration.

Your first statement is more or less correct, (though the concentration of free hydrogen ions is very low, it is never zero). Your second statement is not entirely correct, however, IIRC. Recent research suggests that a free proton groups with one or more water molecules, resulting in a variety of species besides hydronium. This being the case, it is no more “correct” to prefer the use of hydronium ion to hydrogen ion. Use of the latter has the advantage of simplicity, though, when working with equations.