Buffers, Titrations. & pH curves

Common Ion Effect

shift in equilibrium position that occurs because of the addition of an ion already involved in the equilibrium reaction; shifts away from the ion added to equilibrium; an application of Le Chatelier’s Principle

Key Points about Buffered Solutions

Resist changes in pH; they’re weak acids or bases containing a common ion in solution; after addition of strong acid or base, deal with stoichiometry first, then the equilibrium (double RICE charts)

2 Ways to Make a Buffer

1. Direct mixing: mixing a weak acid with its conjugate base salt (e.g. acetic acid and sodium acetate)

2. Partial Neutralization: reacting a weak acid with a strong base (e.g. acetic acid + NaOH) or a weak base with a strong acid (e.g. ammonia + HCl) to create the conjugate pair in solution

Henderson-Hasselbalch Equation

pH = pKa + log [A-]/[HA] where…[A-] = conjugate base, [HA] = acid

For particular buffering systems (conjugate acid-base pair), all solutions that have the same ratio [A-]/[HA] will have the same pH; equation is only used for buffered solutions

If more base than acid is present in a buffered solution, then…

the buffer is better against acids

If more acid than base is present in a buffered solution, then…

the buffer is better against bases

What concentration of bases and acids create the best buffer?

Equal concentrations of base and acid

Buffered Solution Characteristics

Buffers contain relatively large amounts of weak acid and corresponding conjugate bases; added H+ reacts to completion with the conjugate base; added OH- reacts to completion with the weak acid; the pH in the buffered solution is determined by the ratio of the concentrations of the weak acid and weak base; as long as this ratio remainds virtually constant the pH will remain virtually constant

Buffer Capacity

The amount of protons or hydroxide ions the buffer can absorb without a significant change in pH; determined by the magnitudes (size of numerical value) of [HA] and [A-]; a buffer with large capacity contains large concentrations of the buffering components;

What is the optimal buffering situation?

When [HA] = [A-]; it is for this condition that the ratio [A-]/ [HA] is most resistant to change when H+ or OH- is added to the buffered solution

Choosing a Buffer

pKa of the weak acid to be used in the buffer should be as close as possible to the desired pH

Titration Curve

Plotting the pH of the solution being analyzed as a function of the amount of titrant added; possible to switch the substance being added (acid or base) —> flipped “s” curve

Equivalence Point

point in the titration when enough titrant has been added to react exactly with the substance in solution being titrated; mol of acid = mol of base

Weak Acid - Strong Base Titration Steps

1. A stoichiometry problem (reaction is assumed to run to completion) then determine remaining species

2. An equilibrium problem (determine position of weak acid equilibrium and calculate pH) —> watch for additive volumes (find new volume)

Acid-Base Indicators

Marks the end point of a titration by changing color; the equivalence point is NOT the same as the end point; the equivalence point is quantitative (mol of acid = mol of base) and the end point is qualitative (color change)

Indicator Selection

Ideally, the end point will be as close as possible to the actual equivalence point

Solubility Equilibria

Solubility product (Ksp) - equilibrium constant; has only one value for a given solid at a given temp; Solubility - an equilibrium position

Does a higher Ksp value always mean a higher solutbility?

No—> certain salts can’t be compared because they produce different amounts of ions; Ksp values can be used to compare solubility only if they have the same number of ions

How do you compare the solubilities of a salt in different solutions?

If there is no common ion effect, a possible acid-base reaction that could happen, or a temp. change, then their solubilities will be the same