Chapter 15: Applications of Aqueous Equilibria

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# Chapter 15: Applications of Aqueous Equilibria - PowerPoint PPT Presentation

Chapter 15: Applications of Aqueous Equilibria. Buffers Common Ion Effect Henderson-Hasselbalch Equation Buffer Capacity Acid-Base Titrations &amp; Titration Curves Strong Acid-Strong Base Titrations Weak Acid-Strong Base Titrations Strong Acid-Weak Base Titrations Acid-Base Indicators

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Chapter 15: Applications of Aqueous Equilibria
• Buffers
• Common Ion Effect
• Henderson-Hasselbalch Equation
• Buffer Capacity
• Acid-Base Titrations & Titration Curves
• Strong Acid-Strong Base Titrations
• Weak Acid-Strong Base Titrations
• Strong Acid-Weak Base Titrations
• Acid-Base Indicators
• Solubility Equilibria
• Calculating Solubility
• Common Ion Effect
• Selective Precipitation
• Qualitative Analysis
• Complex Ion Equilibria
Example 1

Find the pH of a solution that is 0.20M

KCNO and 0.10 M HCNO.

Ka for HCNO is 3.5 x 10-4.

Example 2

The ratio of [HCO3–] to [H2CO3] in

human blood is 20:1. Find the pH of an

aqueous solution with this composition.

Example 3

Find the pH of a buffer that is made of

0.20M KCNO and 0.10 M HCNO. (Note:

This is the same as Example 1, but here

we will double check our answer using

the Henderson-Hasselbalch equation.)

Ka for HCNO is 3.5 x 10-4.

Example 4

What pH change would result from the

addition of 5.0 mL of 0.10M HCl to 50.0

mL of a buffer containing 0.10 M NH3

and 0.10 M NH4Cl?

How much would the pH of 50.0 mL of

pure water change if the same amount

of acid was added to it?

Example 5

How many grams of Na2CO3 should be

added to 1.5 L of 0.20 M NaHCO3 to

make a buffer of pH = 10.00?

Ka2 of H2CO3 = 5.6 x 10-11

Example 6

What is the pH of a buffer made by

adding 5.0 mL of 0.20 M NaOH to 25.0

mL of 0.10 M HC2H3O2?

Example 7

Calculate the pH change that occurs if

1.0 mL of 0.10 M HCl is added to 50.0

mL of a buffer containing:

• 0.30 M pyruvic acid (HC3H3O3) and 0.30 M potassium pyruvate? Ka for pyruvic acid is 1.4 x 10-4.
• 0.0030 M pyruvic acid and 0.0030 M potassium pyruvate?

Example 8

A 75.0 mL sample of 0.200 M HBr is

titrated with 0.100 M KOH to a

phenolphthalein endpoint. How much

KOH solution is needed to reach the

equivalence point? What is the pH of

the solution at the equivalence point?

Example 9

When a 50.0 mL sample of 0.250 M

nitrous acid is titrated with 0.100 M

NaOH, what volume of NaOH is needed

to reach the equivalence point? What is

the pH at the equivalence point? What

is the pH at the halfway point?

Ka for nitrous acid is 4.0 x 10-4.

Example 10

Find the volume of 0.100 M HCl needed

to reach the equivalence point in the

titration of 25.0 mL of 0.100 M NH3.

Also find the pH of the solution in the

• Initially
• After 10.0 mL of HCl have been added
• At the halfway point
• At the equivalence point
• After 35.0 mL of HCl have been added

Kb for NH3 is 1.8 x 10-5

Example 11

A 100.0 mL sample of a weak,

monoprotic acid with a concentration

of 0.200 M is titrated with 0.100 M

NaOH. After 10.0 mL of NaOH have

been added, the pH is 5.79. What is

Ka for this acid?

Example 12

Silver bromide has a solubility of 0.133

mg per 1.00 L of water. Find Ksp for

silver bromide.

Example 13

Mercury (I) chloride has a Ksp of

1.3 x 10-18. Find its solubility in units of

mole/L and g/L.

Example 14

Which of the following ionic compounds

is more soluble in water? (i.e. Which

will dissolve more moles per liter?)

• CaSO4 or CaCO3
• CaSO4 or Ca(OH)2

Ksp for CaSO4 = 6.1 x 10-5

Ksp for CaCO3 = 8.7 x 10-9

Ksp for Ca(OH)2 = 1.3 x 10-6

Example 15

Calcium oxalate has a solubility of

6.1x10-3 g/L in water. Find its solubility

in 0.20 M CaCl2.

Ksp for CaC2O4 = 2.3x10-9

Example 16

One type of kidney stones is made of

calcium phospate. If [Ca2+] in urine is

0.080 g/L, what is the minimum

molarity of phosphate that will cause

kidney stones to form?

Ksp for calcium phosphate = 1.3x10-32

Example 17

A 65.0 mL sample of 0.010 M Pb(NO3)2

was added to a beaker containing 40.0

mL of 0.035 M KCl. Will a precipitate

form?

Example 18

What percentage of Ca2+ ions remain in

solution after CaCO3 precipitates when

25.0 mL of 0.10 M CaCl2 is added to

25.0 mL of 0.10 M Na2CO3?

Ksp for CaCO3 is 8.7 x 10-9.

Example 19

When 1.0 M AgNO3 is slowly added to a

solution containing 0.015M Cl- and

0.015M CrO42-, what percent of Cl-

remains in solution when the Ag2CrO4

begins to precipitate? (i.e. What is the

maximum separation of Cl- from CrO42-

that can be achieved?)

Example 20

A solution contains 0.10 M Cd2+ and 0.10 M

Ni2+. What concentration of S2- will precipitate

a maximum amount of one cation without

precipitating the other?

Ksp NiS = 3.0 x 10-21

Ksp CdS = 1.0 x 10-28

Example 21

How much Zn2+ ion remains in solution

in a mixture that is 0.010 M Zn(NO3)2

and 0.10 M NH3?

Kf Zn(NH3)42+ = 2.9 x 109

Example 22

Calculate the solubility of AgI in:

• 0.10 M KCN [Kf Ag(CN)2- = 5.6 x 1018]
• water

Ksp for AgI = 1.5 x 10-16

Example 23

Will nickel (II) hydroxide precipitate from

in a solution that is 0.0020 M NiSO4,

0.010 M NaOH, and 0.10 M NH3?

Kf for Ni(NH3)62+ = 5.6 x 108

Ksp for Ni(OH)2 = 2.0 x 10-15