Chapter 18: Aqueous Ionic Equilibrium (Exam 3)

Question 1

Buffer

Answer 1

• Resists pH change by neutralizing added acid or base
  1) Significant amounts of both weak acid and conjugate base
  2) Significant amounts of both weak base and conjugate acid

Question 2

Common Ion Effect

Calculate pH?

Answer 2

• When solution contains two substances that share common ions, solution ionizes less than it normally would
• Write balanced equation, subtract x from reactants and add x to products; substitute into equilibrium concentrations, use x is small approximatio, and determine H₃O⁺ from calculated value of x
• Substitue into pH equation pH = log[H₃O⁺]

Question 3

Henderson-Hasselbach Equation

Answer 3

• pH = pK_a + log [base]/[acid]
• Remember, pK_a = -logK_a
• Use when x is small

Question 4

Generic Approach to Calculating pH in Buffer Solution

Answer 4

1) Stoichiometry calculation in which we calculate the addition changes relative amounts of acid and conjugate base.
2) Equilibrium calculation in which we calculate pH based on new amounts of acid and conjugate base.

–> Stoichiometry: As added acid is neutralized, it converts the stoichiometric amount of base into conjugate acid through neutralization reaction

Question 5

1) Effect of adding small amount of strong acid to buffer solution?
2) Effect of adding small amount of strong base to buffer?

Answer 5

1) Converts stoichiometric amount of base to conjugate acid and LOWERS pH of the buffer.
2) Converts stoichiometric amount of acid to conjugate base and RAISES pH of the buffer.

Question 6

What does an effective buffer accomplish?

Answer 6

• Neutralizes small to moderate amounts of added acid or base

Question 7

What is a buffer’s effectiveness influenced by?

Answer 7

• Relative amounts of acid and conjugate base
• Absolute concentrations of acid and conjugate base
• Most effective when concentration of acid/base are equal and high; range must be 0.10-10

Question 8

How to prepare a buffer?

Answer 8

• Calculate ratio of conjugate base to acid required to attain desired pH
• Plug in pH, pK_a, and use [base]/[acid] = 10ⁿ
• where “n” is equal to pH - pK_a
• (Use H-H equation)

Question 9

Buffer Capacity

Answer 9

• Amount of acid or base that can be added to a buffer without changin the pH significantly

Question 10

Acid-Base Titration

Answer 10

• Basic (or acidic) solution of unknown concentration reacts with acidic (or basic) solution of known concentration

Question 11

Indicator

Answer 11

• Substance whose color depends on pH; monitors titration

Question 12

Equivalence Point

Answer 12

• Point in titration in which number of moles of base is stoichiometrically equal to number of moles of acid

Question 13

Titration: Strong Acid-Strong Base

Answer 13

• Before equivalence point, H₃O⁺ is in excess
• Calculate [H₃O⁺] by substracting number of moles added of OH- from initial moles of H₃O⁺ and dividing by total volume
• Equivalence point is pH = 7
• Beyond equivalence point, OH- is in excess

Question 14

Titration: Weak Acid-Strong Base

• Calculate initial
• Calculate half equivalence point
• Calculate beyond equivalence point
• pH at equivalence point

Answer 14

• Initial pH is that of weak acid solution to be titrated
• Between iitial pH and equivalence point, solution becomes a buffer (calculate with reaction stoichiometry to calculate amounts of each buffer component)
• Halfway to equivalence point: Buffer components are exactly equal (pH = pK_a) –> (calculate: equilibrium problem for ionization of water by ion acting as weak base; # moles ion / volume at equivalence point)
• Beyond equivalence point, OH- is in excess (calculate: OH- by subracting initial moles weak acid from number moles added OH- and divide by total volume)
• Equivalence point: pH = 8

Question 15

Titration: Weak Base-Strong Acid

• Calculate pH?
• Equivalence point?

Answer 15

• Use Henderson-Hasselbach equation
• pK_a = conjugate acid being titrated
• Equivalence point: pH = 5.26

Question 16

Titration: Polyprotic Acid

Answer 16

• pH curve has two equivalence points
• Volume required to reach first equivalence point is identical to volume required to reach second

Question 17

Endpoint

Answer 17

• Occurs at equivalence point; point in which indicator changes color
• Indicators are typically weak organic acids

Question 18

Solubility-Product Constant (K_sp)

Answer 18

• Equilibrium constant for a chemical equation representing the dissolution of an ionic compound
• Meausre of solubility of a compound
• Written as concentration of products [C]^c[D]^d

A + B –> cC + dD

Question 19

Molar Solubility

Answer 19

• Solubility in units of moles/liter (calculated from K_sp)
• Calculate with ICE box

Question 20

Effect of pH on solubility

Answer 20

• Solubility of an ionic compound with strongly basic or weakly acidic anion increases with increasing acidity

Question 21

Q < K_sp

Answer 21

• Solution is unsaturated; more solid ionic compound can dissolve in solution

Question 22

Q = K_sp

Answer 22

• Solution is saturated; holds equilibrium amount of dissolved ions
• Added solid does not dissolve ins olution

Question 23

Q > K_sp

Answer 23

• Solution is supersatruated
• Under most circumstances, excess solid precipitates out of supersaturated solution

Question 24

Predicting Precipitation Reactions

Answer 24

• Use reaction quotient to compare to K_sp

Question 25

Selective Precipitation

Answer 25

• Process involving the addition of a reagent that forms a precipitate with one of the dissolved cations (but not the others)

Question 26

Complex Ion

Answer 26

• Contains central metal ion bound to one or more ligands

Question 27

Ligands

Answer 27

• Neutral molecule or ion that acts as a Lewis base with central metal ion

Question 28

Formation Constant

Answer 28

• K_f
• Equilibrium constant associated with reaction for the formation of a complex ion
• Written the same as a reaction quotient

Question 29

Effect of Complex Ion Equilibrium on Solubility

Answer 29

• Solubility of an ionic compound containing a metal ion increases in the presence of Lewis bases that form complex ions with the cation

Question 30

Effect of solution pH on solubility of amphoteric metal hydroxides?

Answer 30

• Metal hydroxides become more soluble in acidic solutions; can acts as base and react with H₃O⁺
• Some metals can also act as an acid (amphoteric)