# Why doesn't a weak acid/base act as a buffer on its own?

1. May 12, 2013

### sungholee

Hello. I am a high school level student and I had a question about weak acids/bases and why they wouldn't work as a buffer on its own.

If the dissociation of a weak acid is $HA \rightleftharpoons H^{+} + A^{-}$, then when we add $OH^{-}$, it will react with the $H^{+}$ to form water. Because the concentration of the products had been decreased, more $HA$ will dissociate. It's very simple for me upto this point.

But the problem is when we add $H^{+}$, it will react with the $A^{-}$ to form $HA$. Then, according to Le Chatelier's principle, wouldn't the $HA$ dissociate, because its concentration is greater than the products? What I have been told is that now no more $H^{+}$ can be neutralized because all of the $A^{-}$ have been used up, and that's why we need a salt of the acid, to provide that $A^{-}$. However, I fail to see why the $A^{-}$ would work differently to $H^{+}$.

Of course, I can pretend I understand it and simply memorise it, but I'll be studying chem in uni so I want to have a strong understanding of the basics.

Thanks.

Last edited: May 12, 2013
2. May 12, 2013

### mishima

Hi, are you asking for example why a weak acid alone would not act as a buffer? ...As opposed to the typical buffer being a weak acid and conjugate base mixture?

3. May 13, 2013

### sungholee

In essence, yes I am.

4. May 13, 2013

### Staff: Mentor

It is a matter of how effective a solution is at maintaining the constant pH. If it resist pH changes we call it a buffer, when it doesn't - we don't. Resistance to pH changes is called a buffer capacity. And it is highest for pH=pKa, goes down the further from the pKa we are.

See derivation and plot here: buffer capacity.

5. May 13, 2013

### mishima

Well, you can't really have a weak acid by itself, as far as I know. It will always partially dissociate into its conjugate base, reaching equilibrium. This is in contrast with a strong acid, which dissociates essentially completely.

I think maybe more related to what you are asking though is something called percent ionization. This is a simple ratio of the concentration of H+ over the concentration of HA (expressed as a percent). Weak acids have small percent ionizations, for example 0.2 M HNO3 would only have a percent ionization of about 4.8% (so like 5 molecules out of 100 dissociate).

Now maybe sort of an unintuitive thing happens, when you increase the concentration of a weak acid, the percent ionization actually decreases. In other words a smaller fraction of the weak acid molecules dissociate. So like a 2.5 M HNO3 solution has a percent ionization of 1.4% (now only like 1 out of 100 dissociate). (I am just getting these vaules from a text) Its still the case that the concentration of H+ is greater for the higher molarity, but the (percent) amount of dissociation that occurs is less.

So in other words the dissociation of HA isn't a linear relationship with concentration, as (if I am reading you right) you might be thinking.