17.1 – The Common-Ion Effect The Common Ion Effect = the suppression of the ionization of a weak acid or base by the presence of a common ion from a strong electrolyte. (because of LeChatelier’s Principle) *for math problems with this, you will start with a concentration of both the acid (or base), and the ion (see sample exercises 17.1 and 17.2)

Buffers Buffer Solutions– only change pH slightly when an acid or base is added They have to be either: a soln of a weak acid and its salt (conjugate base), or a soln of a weak base and its salt (conjugate acid) **the acid component neutralizes added OH -, and the basic component neutralizes added H 3 O +. Ex. For a buffer made up of a weak acid and its conjugate base: if OH - ion are added, OH - + HA → H 2 O + A - and if H 3 O + ions are added, H 3 O + + A - → HA + H 2 O

Usually buffers are prepared by mixing a weak acid or base with a salt of that acid or base To calculate the pH of a buffer solution, we can use the Henderson-Hasselbalch equation: Another way to calculate the pH is to treat it as a common ion problem See sample exercises 17.3 – 17.5

Buffer Capacity = the amount of acid or base the buffer can neutralize before the pH begins to change to an appreciable degree. It depends on the amount of acid and base used to prepare the buffer. pH range of a buffer = the pH range over which the buffer acts effectively. A buffer is most effective if the concentrations of the buffer acid and conjugate base are equal. Then, the pH = the pK a. This is the optimal pH of any buffer, so the acid chosen for a buffer should have a pK a close to the desired pH. Buffers usually have a pH range of ±1 pH unit.

Calculating how the pH of a buffer responds to the addition of a strong acid or base: 1. Do stoichiometry with the appropriate neutralization reaction (H 3 O + + A - → HA + H 2 O if an acid is added, or OH - + HA → H 2 O + A - if a base is added) to calculate the new values of [HA] and [A - ]. These neutralization reactions go to completion, so the strong acid or base is completely consumed in the reaction. 2. Use the Henderson-Hasselbalch equation with the new values of [HA] and [A - ] to calculate the pH See sample exercise 17.6

17.3 – Acid-Base Titrations Terms: Equivalence point = when the acid and base are in exact stoichiometric proportions (neither is in excess). An indicator can be used to detect this. End point = when the indicator changes color. This should be matched to the equivalence point. However, any indicator that changes color in the vertical part of the titration curve can be used. Titration curve = a graph of pH vs. volume of titrant added Titrant = the soln added from the buret

Titrating a strong acid with a strong base: The pH starts very low and ends very high At the equivalence point, the pH is 7.0

The pH can be calculated at various stages of the titration: 1)Initial pH – before any base is added, the pH is determined by the initial concentration of the strong acid 2)Between Initial pH and equivalence point – the pH is determined by the concentration of acid not yet neutralized (the OH - is the limiting reactant, so use stoichiometry to get the conc. of the acid) 3)Equivalence point – the pH is 7 4)After the equivalence point – the pH is determined by the concentration of excess NaOH in the solution. See sample exercise 17.7

Titration of a Strong Base with a Strong Acid: It looks like you “flipped over” the strong acid being titrated by a strong base. Start with a high pH (basic solution); the pH = 7 at the equivalence point; low pH to end.

Titration of a Weak Acid with a Strong Base: the initial pH is higher At the half neutralization point, pH = pK a (because the [HA] =[A - ] at this point, and using the H-H equat., the log[HA]/[A - ] = 0) At the equivalence point, the pH>7 (because the conjugate base of the weak acid hydrolyzes) The steep portion of the curve is shorter The same volume of base is required to reach the equivalence point, and the curve after the equivalence point is the same

To calculate the pH at different points: 1)Initial pH – use K a of the weak acid 2)Between initial pH and equivalence point – there is now a buffer, so use the procedure from section 17.2 (neutralization reaction to get new values of [HA] and [A - ], then H-H equation). See sample exercise )Equivalence point – only the conjugate base is present as the acid is completely neutralized by the strong base. To calculate the [A - ], use the following: initial moles HA = moles of A - at equivalence point, then moles of A - / vol. of solution at equivalence point = [A - ] at equivalence point. Then use the K b to get the pH. See sample exercise )After the equivalence point -- the pH is determined by the concentration of excess NaOH in the solution. (just like for a SA/SB titration)

Comparison of different acids:

Titration of a weak base with a strong acid: The graph starts lower than for a strong base titration The pH at the equivalence point <7

Titrations of Polyprotic Acids: When a polyprotic acid is titrated with a base, there is an equivalence point for each dissociation. Using the Henderson– Hasselbalch equation, we can see that half way to each equivalence point pH = pK a for that step.