Can't wrap my head around buffers

[seratia]

Let's keep things simple. I have pure water. pH = 7.
I want to add a buffer. I add the weak acid. It dissociates little. So the pH is still close to neutral. So far so good.

The problem is, we are also adding conjugate base (the conjugate base is strong). Why doesn't the added base consume the H+ in the water, thus increasing the pH?

In other words, I am confused, because in the process of making a buffer, we are screwing with the original pH. Am I missing something?

 

[Borek]

You choose amounts of the acid and its conjugate base so that the final pH is what you need.

 

[Merlin3189]

... . I have pure water. pH = 7.
I want to add a buffer. I add the weak acid (such as H Actetate.) It dissociates little. So the pH is still close to neutral (or close to the pH of a weak ACID.) So far so good.

The problem is, we are also adding conjugate base (acetate ion) (the conjugate base is strong). Why doesn't the added base consume the H+ in the water, thus increasing the pH? ...

You only have Ac^- ions when HAc molecules have dissociated into H⁺ and Ac^- ions. So for every Ac^- ion in solution, you have already produced one extra H⁺ ion. If you say the Ac^- ions are mopping up the H⁺ ions, to form HAc molecules, then you are saying the HAc does not dissociate at all and leaves the water neutral.

When you add a weak acid to water, you don't get a nearly neutral solution, you get a weakly acid solution. The clue is in the name.
If, when you added HAc to water, it did (by some utterly improbable chance) completely dissociate, then it would be a strong acid and pH would zip down to 1. But Ac^- ions would, as you say, grab protons and hang onto them, bringing the equilibrium back to the point where pH is that of a weak acid.

When you have your weak acid equilibrium (HAc ⇔ H⁺ + Ac^- ), if you were to add a small amount of stronger acid, you add more H⁺ ions upsetting the equilibrium The Ac^- ions do grab some of the excess H⁺, so the pH doesn't get so acid as it would if the stronger acid had been added to pure water.
Equally, if you add some OH^- ions from some alkali, then they remove H⁺ ions again upsetting the acetate equilibrium, so more HAc dissociates giving H⁺ ions and stopping the pH from going as alkaline as it would have.
That's the role of the buffer, to put the brakes on and oppose changes in pH a little bit.

 

[James Pelezo]

When you have your weak acid equilibrium (HAc ⇔ H+ + Ac- ), if you were to add a small amount of stronger acid, you add more H+ ions upsetting the equilibrium The Ac- ions do grab some of the excess H+, so the pH doesn't get so acid as it would if the stronger acid had been added to pure water.
Equally, if you add some OH- ions from some alkali, then they remove H+ ions again upsetting the acetate equilibrium, so more HAc dissociates giving H+ ions and stopping the pH from going as alkaline as it would have.

With respect, you seem to be missing the conjugate base in your comment. Just depending on the shift of the weak acid (or, weak base) ionization on additions of H⁺ or OH^- will not provide the buffer effect. A common ion (or common ion system) must be in the weak electrolyte solution also so as to inhibit shifting of pH values due to weak electrolyte ionization limitations. By definition, a buffer is a weak electrolyte (wk acid or wk base) plus the salt of the weak electrolyte (providing the common ion).

Example: compare the change in pH of HOAc as a 0.10MHOAc(aq) + 0.01MHCl(aq) solution; Ka(HOAc) = 1.85 x 10^-5 @25^oC, to change in pH on addition of 0.01MHCl(aq) to a 0.10MHOAc(aq)/0.10MNaOAc(aq) solution under same conditions.

pH(0.10M HOAc@25^oC) = 4.73 ... If adding 0.01MHCl(aq) to 0.10MHOAc@25^oC => pH = 3.73 => ΔpH = 3.73 - 4.73 = -1.00

pH(0.10MHOAc/0.10MNaOAc)@25^oC = 4.73 ... If adding 0.01MHCl(aq) to this soln @25^oC => pH = 4.64 (limited ioniz'n) => ΔpH = 4.64 - 4.73 = -0.09

A drop of -1.00 without OAc^- ion vs -0.09 with OAc^- ion in HOAc solution. The limited change is due to the common ion effect, in this example, the acetate ion (or, conj. base). Without the common ion, pH changes are dependent on amount of H⁺ added. A shift in the ionization equilibrium toward HOAc does occur due to excess H⁺, but is NOT a new equilibrium for HOAC <=> H⁺ + OAc^- due to the chemistry of the common ion effect required in the buffering process.