Chapter 16

Question 1

Buffer

Answer 1

• Resists pH change by neutralizing added acid or added base.
• Contains significant amounts of both a weak acid and it’s conjugate base (or weak base and conjugate acid)

Question 2

Buffer Characteristics

Answer 2

• Buffers resist pH change.
• A buffer contains significant amounts of both a weak acid and its conjugate base.
• The weak acid neutralizes added base.
• The conjugate base neutralizes added acid.

Question 3

Common Ion Effect

Answer 3

The tendency for a common ion to decrease the solubility of an ionic compound or to decrease the ionization of a weak acid or weak base.

Question 4

Henderson-Hasselbalch Equation

Answer 4

An equation used to calculate the pH of a buffer solution from the initial concentrations of the
buffer components.

pH = pk_a + log( [base]/[acid] )

Question 5

Calculating pH changes in buffer solution after adding small amounts of a strong acid or base

Answer 5

1. Stoichiometry—use the stoichiometry of the neutralization equation to calculate the changes in the amounts (in moles) of the buffer components upon addition of the acid or base.
2. Equilibrium—use the new amounts of buffer components to work an equilibrium problem to find pH. (For most buffers, this can also be done with the Henderson–Hasselbalch equation.)

Question 6

Buffer Effectiveness

Answer 6

Buffer is most effective when:

• Concentrations of acid and conjugate base are equal and high.

Question 7

Buffer Range

Answer 7

• Lowest pH for effective buffer occurs when the base is one-tenth as concentrated as the acid.
• Highest pH for effective buffer occurs when the base is 10 times as concentrated as the acid.

Question 8

Buffer Capacity

Answer 8

The amount of acid or base that you can add to a buffer without causing a large change in pH.

Question 9

Acid-Base Titration

Answer 9

A basic (or acidic) solution of unknown concentration is reacted with an acidic (or basic) solution of known concentration, in order to determine the concentration of the unknown.

Question 10

Equivalence Point

Answer 10

The point in a titration at which the added solute completely reacts with the solute present in the solution; for
acid–base titrations, the point at which the amount of acid is stoichiometrically equal to the amount of base in solution.

Question 11

Titration of a Strong Acid with a Strong Base

Answer 11

• The initial pH is simply the pH of the strong acid solution to be titrated.
• Before the equivalence point, H₃O⁺is in excess. Calculate the [H₃O⁺] by subtracting the number of moles of added OH^-from the initial number of moles of H₃O⁺ and dividing by the total volume.
• At the equivalence point, neither reactant is in excess and the pH=7.00.
• Beyond the equivalence point, OH^- is in excess. Calculate the [OH^-] by subtracting the initial number of moles of H₃O⁺ from the number of moles of added OH^- and dividing by the total volume.

Question 12

Titration of a Weak Acid with a Strong Base

Answer 12

• The initial pH is that of the weak acid solution to be titrated. Calculate the pH by working an equilibrium problem using the concentration of the weak acid as the initial concentration.
• Between the initial pH and the equivalence point, the solution becomes a buffer. Use the reaction stoichiometry to calculate the amounts of each buffer component and then use the Henderson–Hasselbalch equation to calculate the pH.
• Halfway to the equivalence point, the buffer components are exactly equal and pH = pK_a.
• At the equivalence point, the acid has all been converted into its conjugate base. Calculate the pH by working an equilibrium problem for the ionization of water by the ion acting as a weak base . (Calculate the concentration
  of the ion acting as a weak base by dividing the number of moles of the ion by the total volume at the equivalence point.)
• Beyond the equivalence point, OH^- is in excess. Ignore the weak base and calculate the [OH-] by subtracting the initial number of moles of H₃O⁺ from the number of moles of added OH^- and dividing by the total volume.

Question 13

Solubility Product Constant (K_sp)

Answer 13

• Represents the dissolution of a slightly to moderately soluble ionic compound.
• Example:

K_sp = [Ca²⁺][F^-]²

Question 14

Molar Solubility

Answer 14

The solubility of a compound in units of mol/L

Question 15

Solubility of an Ionic Compound

Answer 15

• Lower in a solution containing a common ion than in pure water.
• With a strongly basic or weakly basic anion the solubility increases with increasing acidity (decreasing pH).

Question 16

Q < K_sp

Answer 16

• The solution is unsaturated.
• More of the solid ionic compound can dissolve in the solution.

Question 17

Q = K_sp

Answer 17

• Solution is saturated.
• Additional solid will not dissolve in the solution.

Question 18

Q > K_sp

Answer 18

• Solution is supersaturated.
• Excess solid will precipitate out of a supersaturated solution.

Question 19

Selective Precipitation

Answer 19

A process involving the addition of a reagant to a solution that forms a precipitate with one of the dissolved ions but not the others.

Question 20

Qualitative Analysis

Answer 20

• Way to determine the ions present in an unknown solution

Question 21

Quantitative Analysis

Answer 21

• Concerned with quantity.
• Way to determine the amounts of substances in a solution or mixture.

Question 22

Complex Ion

Answer 22

Contains a central metal ion bound to one or more ligands.

Question 23

Ligand

Answer 23

A neutral molecule or ion that acts as a Lewis base with the central metal ion.

Question 24

Formation Constant, K_f

Answer 24

Kf = [products] / [reactants]

• Equilibrium constant for the formation of a complex ion.

Question 25

Molar Solubility

Answer 25

The number of moles of salt dissolved in one liter of its saturated solution.