Chapter 17

Page 1

17.1 The Common-Ion

Effect

We have a weak acid,

acetic acid (CH3COOH) and soluble salt of that acid, sodium acetate (CH3COONa).

Both of which share the

common ion CH3COO-

Sodium acetate is a soluble

ionic compound therefore a strong electrolyte.

It dissociates completely in

aqueous solution to form Na+ and Ch3COO- ions

CH3COONa is a weak

electrolyte that ionizes only partially

Equation

17.1

CH3COOH(aq) <----> H+ (aq) + CH3COO- (aq)

Ka=

1.8 *10 -5

The presence of the added

acetate ion cause the acetic acid to ionize less that it normally would

The

equilibrium constant itself does not change; it is the relative

concentrations of products and reactants in the equilibrium expression

that change.

Whenever a

weak electrolyte and a strong el4ectrolte containing a common ion are

together in solution, the weak electrolyte ionizes less than it world if

it were alone in the solution----common ion effect.

17.2 Buffers

Buffers: contain high concentrations of weak

conjugate acid-base pair and that resist drastic changes in pH when small

amount of strong acid or strong base are added to the solution.

Composition and

Action of Buffers

A buffer resists

changes in pH because it contains both an acid to neutralize added OH- ions and

a base to neutralize added H+ ions.

They however cannot consume

each other through a neutralization reaction

You must have a weak

acid-base conjugate pair.

The key is to have roughly

equal concentrations of both the weak acid and its conjugate base

2 WAYS TO MAKE A BUFFER

Mix weak acid or

weak base with a salt of that acid or base.

Make a conjugate acid or

base from a solution of weak base or acid by addition of strong acid or

base.

Neutralization reaction have

very large equilibrium constants so the amount of acetate formed in CH3COOH + OH- ----> CH3COO- + H2O , will be limited by the relative

amounts of the acid and strong base that are mixed.

By choosing

appropriate components and adjusting their relative concentrations, we can

buffer a solution at virtually any pH.

HA (aq) <---->

H+ (aq) + A- (aq)

Corresponding

acid-dissociation-constant expression is:

Calculating the pH of a Buffer

Henderson-Hasselbalch equation: pH= pKa + log [base]/[acid]

* you have to assume to use the initial concentrationfggvvgfgs of the acid and base components in the buffer are equal to the equilibrium concentrations

Buffer Capacity and pH Range

*Buffer capacity: is the amount of acid or base the buffer can neutralize before the pH begins to change to an appreciable degree

the buffer capacity depends on the amount of acid and base used to prepare the buffer

*pH range: of any buffer is the pH range over which the buffer acts effectively

buffers most effectively resist a change in pH in either direction when the concentrations of weak acid and conjugate base are about the same

*when the concentrations of weak acid and conjugate base are equal, pH=pKa.

because log of 10 = 1, buffers usually have a usable range within + or - pH unit of pKa (range of pH= pKa +-1)

Addition of Strong Acids or Bases to Buffers

* neutralization reactions between strong acids and weak bases proceed essential to completion, as do those between strong bases and weak acids

*this is because water is a product of the reaction, and when you have an equilibrium constant of 1/Kw=10(power of 14)

as long as you do not exceed the buffering capacity of the buffer, you can assume that the strong acid or strong base is completely consumed by reaction with the buffer

Ex: a buffer that contains a weak acid HA and its conjugate base A-. when a strong acid is added to this buffer, the added H+ is consumed by A- to produce HA; thus [HA] increase and [A-] decreases.

*Ex: upon addition of a strong base, the added OH- is consumed by HA to produce A-; in this case [HA] decreases and [A-] increases.

To calculate pH of the buffer responds to addition of a strong acid or a strong base

1. the acid-base neutralization reaction and determines its effect on [HA] and [A-]. this is a limiting reactant stoichiometry calculation

2. dUse the calculated values of [HA] and [A-] along with Ka to calculate [H+]. this step is an equilibrium calculation and is most easily done when using the Henderson Hasselbalch equation (if the concentrations of the weak acid-base pair are very large compared to Ka for the acid).

Examples on page 736

17.3 Acid-Base Titrations

titrations are procedures in which one reactant is slowly added into a solution of another reactant, while equilibrium concentrations along the way are monitored.

pH titration curve: a graph of the pH as a function of the volume of titrant added. the shape of the curve makes it possible to determine the equivalent point. also a suitable indicator to determine the Ka of the weak acid or the Kb of the weak base being titrated

Strong Acid-Strong Base Titrations

1. initial pH: the pH of the solution before the addition of any base is determined hy the initial concentration of the strong acid

.1M HCL, [H+] =.1 M and pH=-log(.1) =1.0 thus the initial pH is low

2. Between initial pH and equivalence point: as NaOH is added, the pH increases slowly at first and then rapidly in the vicinity of the equivalency point. the pH before the equivalency point is determined by the concentration of acid not yet neutralized. (ICE table)

3. Equivalence point: at the point an equal number of moles of NaOH and HCl have reacted, leaving only a solution of their salt, NaCl. the Ph of the solution is 7 because the cation of a strong base and the anion of the strong acid are neigher acids nor bases and, have no appreciable effect on pH.

4. After equivalence point: the pH of the solution after the equivalence point is determined by the concentration of excess NaOH in the solution

Weak Acid-strong base titrations

1. Initial pH: we use Ka to calculate this pH.

2. Between initial pH and equivalence point: prior to reaching the equivalence poit, the acid is being neutralized, and its conjugate base is being formed

*2 steps for calulating the pH in this regions

1. consider the neutraliztion reaction between the acid and base to determine the concentrations of both

2. next we calculate the pH of this buffer pair

3. Equivalence point: pH at the equivalence point is always above 7 in a weak acid-strong base titration because the anion of the salt formed is a weak base.

4. after equivalence point (excess base): in this region, [OH-] from te reaction of CH3COO- with water is negligible relative to [OH-] from the excess NaOH. thus, the pH is determined by the concentration of OH- from the ecess NaOH

Titrations of polyprotic acids

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17.4 Solubility Equilibria