Buffers1.

Titrations and pH curves2.

Solubility equilibria3.

Topics in this chapter:

Buffers resist changes to the pH of a solution.

Consider a normal weak acid equilibrium:

HA + H2O ⇌ H3O+ + A−

Which direction will the reaction shift if more A− is

added?

What happens to the acidity?

What happens to the % ionization of HA?

The equilibrium concentrations of HA and A− will be very

close to their initial concentrations:

Ka = [H3O+] =

Chapter 16: Aqueous ionic equilibrium

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pH = pKa + log [base]

[acid]

This only works for a buffer solution

(high concentrations of both parts of a

conjugate acid-base pair)

�

Don't flip the fraction! The conj. base is on top!�

Use the pKa of the conjugate acid.�

But be careful…

What is the pH of a solution of 0.50 M acetic acid and

0.50 M sodium acetate?

Henderson-Hasselbalch Equation

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What is the pH of a solution of

0.10 M ammonium chloride and 0.15 M ammonia?

Buffer pH calculation

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Buffers resist changes in pH because they contain both

an acid and a base.

What happens when you add H3O+ to the buffer?

In a buffer of HA and A−:

What happens when you add OH− to the buffer?

Adding a strong acid or base to a buffer

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Calculate the pH when 0.0050 mol strong acid is added to

a buffer which contains 0.010 mol NH4+ and 0.015 mol

NH3?

Stoichiometry calculation1.

Equilibrium calculation (H-H equation)2.

Adding a strong acid or base to a buffer

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A buffer contains 0.47 M HNO2 and 0.61 M NaNO2. What

is the pH of this buffer?

What is the pH when 8.7 mL of 0.128 M NaOH is added

to 243 mL of the buffer solution?

Adding a strong acid or base to a buffer

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Relative amounts of acid and base:

A buffer is most effective when the ratio of A− to HA is

between 0.1 and 10. What pH range does this

correspond to?

So, when creating a buffer, make sure the pKa of the

acid is as close to the desired pH as possible.

What is the pH range of a buffer created by formic

acid (pKa = 3.74) and its conjugate?

What ratio of the base to acid would be needed to

make a pH 4.00 buffer?

Absolute concentrations of acid and base:

A buffer is effective when the concentrations of A−

and HA are at least 10x greater than the

concentration of an added strong acid or base.

Buffer capacity and buffer range

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Known concentration

Unknown concentration

Titration curve:

graph of pH vs volume

of solution added

Strong acid titrated with strong base

pH

Vol NaOH added (mL)

Titrations and titration curves

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50.0 mL HCl of unknown concentration is titrated with

0.150 M NaOH. The titration curve's only inflection point

occurs at 40.0 mL of added base. What was the initial HCl

concentration?

What was the initial pH of the HCl solution?

What is the pH after 60.0 mL NaOH is added?

Strong acid titrated with strong base

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pH

Vol NaOH added (mL)

HA + OH− →

Buffer region:

At the midpoint (half-equivalence pt):

Weak acid titrated with strong base

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40.0 mL of 0.0788 M HA are titrated with 0.100 M NaOH.

What is the volume of NaOH added at the equivalence

point?

If the pH = 4.88 at the midpoint, what is Ka and what is

the pH at the endpoint?

Weak acid titrated with strong base

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What is the pH after 10.0 mL 0.100 M NaOH is added to

40.0 mL of 0.0788 M HA (the titration on the last page)?

Weak acid titrated with strong base

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What is the initial acid concentration?�

What is the Ka?�

What is the pH initially and at the equivalence

point?

�

75.0 mL of a weak acid are titrated by 0.125 M NaOH.

The endpoint occurs at 62.1 mL of added base, and the

pH at the midpoint is 3.12.

Weak acid titrated with strong base

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Weak base titr by strong acid

pH

Vol HCl added (mL)

Polyprotic acid titr by strong base

pH

Vol NaOH added (mL)

Weak base and polyprotic acid titrations

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Make sure the color range of your indicator is within the

most vertical part of the titration curve!

pH Indicators

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Compounds we called "insoluble" before, are actually

very slightly soluble, governed by an equilibrium:

CaF2(s) ⇌ Ca2+(aq) + 2 F−(aq)

Solubility product constant, Ksp for CaF2 =

Molar solubility, S = mol solute (when saturated)

L solution

If the molar solubility, S, of PbF2 is 2.61 x 10−3 M, what is

the Ksp of PbF2?

PbF2(s) ⇌

What are the ion concentrations?

Solubility equilibria

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If Ksp for Mg(OH)2 is 2.06 x 10−13, calculate the molar

solubility, S.

Calculating molar solubility

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Direct comparison of Ksp values only works between

compounds that dissociate into the same number of

ions.

Which is the most soluble?

BaF2 Ksp = 2.45 x 10−5

CaF2 Ksp = 1.46 x 10−10

Fe(OH)2 Ksp = 4.87 x 10−17

PbCl2 Ksp = 1.17 x 10−15

PbBr2 Ksp = 4.67 x 10−6

Ksp and relative solubility

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How is solubility affected when more ions are added to

the solution?

Ca(OH)2(s) ⇌ Ca2+(aq) + 2 OH−(aq); Ksp = 4.68 x 10−6

To this solution, we add Ca(NO3)2

The reaction shifts:

A common ion in solution will _________ a compound's

solubility.

Na+, K+, NH4+,

NO3−, C2H3O2

−

make compounds

totally soluble

Common Ion Effect

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What is S for Ca(OH)2? Ksp = 4.68 x 10−6

What is S for Ca(OH)2 in 0.10 M Ca(NO3)2 solution?

Common ion effect calculation

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If a compound's dissociation produces a basic anion

(usually OH−, CO32−, S2−), solubility is dependent on pH.

CaCO3(s) ⇌ Ca2+(aq) + CO32−(aq)

Addition of strong acid to this solution…

A similar effect occurs when base is added to a solution

which contains an acidic cation.

Effect of pH on solubility

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To calculate whether or not a precipitate will form, take

the given ion concentrations and calculate Q (the

reaction quotient).

If Q = Ksp:

If Q > Ksp:

If Q < Ksp:

A newly mixed solution has these ion concentrations:

[Pb2+] = 0.25 M, [Cl−] = 0.017 M. Will a precipitate form?

Ksp (PbCl2) = 1.17 x 10−5.

Precipitation

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