### Introduction to buffers

This is the definition of a buffer:

a solution that contains a weak acid (or a base) and the salt of that weak acid (or that base)
Here are the two classic examples:
1) a solution of acetic acid (CH3COOH) and sodium acetate (CH3COONa)

2) a solution of ammonia (NH3) and ammonium chloride (NH4Cl)

The first example is a weak acid and its salt. The second is a weak base and its salt.

Buffers have a very important chemical property. They can resist changes in pH better than solutions which are not buffers.

If you add some NaOH to pure water (pH = 7), the pH changes by several pH. However, when you add the same amount of NaOH to a buffer with the same pH of 7, the change in pH will be much, much less.

For example, let us assume we have some pure water at pH = 7. Then we add 0.1 mole of NaOH. Since the [OH¯] now equals 0.1 M, the pOH = 1 and the pH = 13. When the pH = 13, this means that the [H+] = 1.0 x 10¯13 M. This means the [H+] decreased by a factor of one million with the addition of 0.1 mole of NaOH, since the water at pH = 7 had [H+] = 1.0 x 10¯7 M.

I'm going to reserve a discussion of the change in pH of a buffer until discussing the actual calculations you will be called upon to perform. However, let me assure you, the pH change in a buffer (when some acid or base is added) will be on the order of tenths of a pH unit rather than 6 pH units, as in the pure water example just above.

Just to be emphatic: buffers will change their pH as an acid or a base is added to the buffer. It is just that a buffer changes pH much slower than an unbuffered solution. Why am I saying this? Because people being introduced to buffers sometimes get the misconception that buffers never change their pH. This is not correct.

More Buffer Examples

The various phosphates form several different buffer combinations; phosphate buffers are popular as test questions.

Here are the 4 possible phosphate compounds:

H3PO4 - phosphoric acid
NaH2PO4 - sodium dihydrogen phosphate
Na2HPO4 - sodium hydrogen phosphate
Na3PO4 - sodium phosphate

Here are the possible phosphate buffer combinations:

H3PO4 / H2PO4¯
H2PO4¯ / HPO42¯
HPO42¯ / PO43¯

There is one carbonate buffer that you could see used:

HCO3¯ / CO32¯

This carbonate buffer, H2CO3 / HCO3¯, is seen from time to time. It is seen much less than the one just above.

In all buffer examples, the two substances differ by only a proton:

two examples: H3PO4 / H2PO4¯ or NH3 / NH4+

Let's say that again:

In a buffer solution, the two substances differ by only one proton.

By the way, you might see that these two phosphate combinations can exist:

H3PO4 / HPO42¯

and

H2PO4¯ / PO43¯

They do not differ by one proton, they differ by two. These types of solutions are not normally discussed in an introductory chemistry class. This is because the chemistry would be too complex. As soon as some base or acid were to be added, a third substance would be introduced (H2PO4¯ for the first example and HPO42¯ for the second).

There may exist somewhere some instructions on how to deal with solutions of this sort. You will not see them on the ChemTeam.

Note that, in the carbonate and phosphate buffers, I have eliminated the cation. This is because the cation has no effect on the buffer action and so it is ignored. It simply does not matter if the salt is the sodium salt (as above) or the potassium salt (KHCO3, for example).

The same is also true of the ammonium used at the top of this tutorial. It does not matter if it is ammonium chloride or ammonium nitrate.

This ignoring of certain cations and anions also extends to the strong acids and bases you will see: sodium hydroxide and potassim hydroxide will be mentioned, but only the hydroxide concentration ([OH¯]) will commonly be discussed. With strong acids, you will see hydrochloric acid or nitric acid, but only the hydrogen ion ([H+]) concentration will be important.

The key equation to use in buffer calculations is the Henderson-Hasselbalch Equation.