Lab Cat

9 Mar 2007

Buffers

Filed under: Basic Concepts, Chemistry — Cat @ 12:07 pm

In my last basic concepts post, I discussed strong and weak acids. This is also true of bases (alkalines). When dissolved in water, weak bases and acids do not fully dissociate, causing an equilibrium to exist between the acid (HA) and acid ion (A) (also known as the conjugate base):

HA + H2O ⇌ H3O+ + A

ref (1)

Equilibrium reactions work to keep the concentration of reactants and products in balance. So if more of the compounds on the left are added, the reaction is said to move to the right. So the addition of more acid causes the formation of more H3O+ and A. Also if compounds from the right are removed from the solution, somehow, the more of them will form. This is a key point.

Before we discuss that key point, here is a recap. In my post on acids and bases, I stated the fact that H3O+ is the equivalent of H+ and in my pH post, I explained that pH = -log[H+]. So where are we?

H3O+ ~= H+

Thus if:

pH = -log[H+]

pH ~= -log[H3O+]

Thus, the higher the concentration of H3O+ the lower the pH since a lower pH implies a more acidic solution. If more acid is added, more H3O+ is formed and the pH is lowered.

So back to weak acid dissociation.

If H3O+ is removed, more will form by more acid dissociating. Thus, as long as the acid is present in ample supply, the pH will remain constant even as H3O+ is removed as more H3O+ is formed to keep the reaction in balance.

How can H3O+ be removed?

The typical way would be for it to react with base (HOB) to form water and conjugate acid:

H3O+ + HOB —> B+ + 2H2O

In fact, the conjugate acid (B+) would most likely react with the conjugate base (A) to form a conjugate salt, AB.

The removal of both H3O+ and A from the system force the reaction to move to the right, causing more acid to dissociate and more H3O+ and A to form. So as long as there is sufficient acid and not too much base, the pH of the weak acid solution resists change and maintains the pH. This is known as buffering.

Buffers work best with a solution within plus & minus one pH unit of the pKa of the weak acid. Where the pKa is:

pKa = −log10 Ka

and Ka is the acid dissociation constant (ref 2):

Ka = [H3O+][A]/[HA]

The Ka, pH and knowledge of the equilibrium reaction behavior, were used to develop the Henderson-Hasselbalch equation. This allows chemists to calculate the pH of a solution if the concentration of the weak acid and conjugate base are known. It also allows us to calculate how much acid or base are needed in a system to get a certain pH.

Next up: Why are buffers important?


References

(1) I have been teaching this for so long that I forget my original sources. It certainly wasn’t Wikipedia. Probably an A-level chemistry textbook.

(2) When the solution pH is equal to the pKa, there is an equal concentration of acid (AH) and of acid ion (A)

(3) Another useful reference on Henderson-Hasselbalch is at ChemBuddy

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2 Comments »

  1. Maybe start off with a definition of Bronsted acids (which is what you’re talking about above…)

    Plus the idea of basicity, Kb and pKb?

    Then move alittle onto Lewis acids / bases… ?

    Otherwise – it’s a nice little section – I imagine very useful for food chemists!

    Comment by Mark C R UK — 9 Mar 2007 @ 5:19 pm

  2. Hi Mark

    This is part of a series I have written on basic concepts. Check the link to acids and bases. That discusses the different definitions of acids and bases

    I admit to ignoring basicity as it is less important for food science than acidity.

    Thanks for your comments.

    Comment by Cat — 10 Mar 2007 @ 2:07 pm


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