Ch 17 Aqueous Ionic Equilibrium

Question 1

What is a buffer? Give an example.

Answer 1

Buffer - A solution that resists changes in pH by neutralizing added acid or added base.
A buffer contains either a significant amount of a weak acid and its conjugate base or a significant amount of a weak base and its conjugate acid.
For example - Blood is a buffer that contains a mixture of H2CO3 and HCO3 and maintains a pH between 7.35 and 7.45.

Question 2

What happens if a strong base or strong acid is added to an acidic buffer solution?

Answer 2

1) If a strong base is added, it is neutralized by the weak acid in the buffer.
2) If a strong acid is added, it is neutralized by the conjugate base in the buffer.

Question 3

Explain the Common Ion Effect.

Answer 3

In a weak acid solution, adding a salt containing the anion NaA, the conjugate base of the acid (known as as the common ion), shifts the position of equilibrium to the left.
This lowers the H3O+ ion and increase pH.

Question 4

What’s the significance of the Henderson-Hasselbalch Equation?what is the equation?

Answer 4

It’s an equation derived from the Ka expression that allows us to calculate the pH of a buffer solution.
The equation calculates the pH of a buffer from the pKa and initial concentrations of the weak acid and salt of the conjugate base, as long as the “x is small” approximation is valid.
pH = pKa + log ([base]/[acid])

Question 5

The Common Ion is also known as what?

Answer 5

The Conjugate Base

Question 6

Define Buffer Capacity and Buffer Range

Answer 6

1) Buffer Capacity - The amount of acid or base that can be added to a buffer without causing a large change in pH value and destroying its effectiveness.
2) Buffer Range - The pH range in which the buffer can be effective.

Question 7

The effectiveness of a buffer depends on what factors?

Answer 7

1) The absolute concentrations of buffer acid and base.

2) The relative amounts of buffer and base.

Question 8

When is a Buffer most effective/resistant to pH changes?

Answer 8

1) When the [acid] and the [base] are large, since they can neutralize more added acid or base.
2) When [acid] = [conjugate base]

Question 9

How do you calculate the maximum and minimum pH at which the buffer will be creative?

Answer 9

1) Lowest pH: pH = pKa + log(0.10) aka pKa - 1
2) Highest pH: pH = pKa + log(10) aka pKa + 1
Therefore, the effective pH range of a buffer is pKa +/- 1.

Question 10

How should you choose an acid to make a buffer?

Answer 10

Choose one whose pKa is closest to the pH of the buffer.

Question 11

Define Acid-Base Titration.

Answer 11

When a solution of known concentration (titrant) is slowly added to a solution of unknown concentration in order to determine the concentration of the unknown.

Question 12

Define Equivalence Point

Answer 12

The point in an acid-base titration in which the amount of acid is stoichiometrically equal to the amount of base in a solution. Aka, the End point.

Question 13

Define an Indicator

Answer 13

A chemical that changes color when the pH changes, and is used in titration to indicate the equivalence point (end point).

Question 14

What are the characteristics of a Titration Curve?

Answer 14

1) It is a plot of pH versus the amount of added titrant
2) The inflection Point of the curve is the equivalence point of the titration.
3) Before the equivalence point, the original solution in the flask is in excess, so the the pH is closest to its pH.
4) The pH of the equivalence point depends on the pH of the salt solutions.
5) Beyond the equivalence point, the unknown solution in the burette is in excess, so the pH approaches its pH.

Question 15

What are the two Indicators used for acid-base titrations, that were mentioned in class?

Answer 15

1) Methyl Red - Pink when acidic and yellow when basic.

2) Phenopthalein - Most common indicator used; Clear when acidic and pink when basic.

Question 16

Define Solubility Product.

Answer 16

The equilibrium constant for the dissociation of a solid into its aqueous ions. aka ksp

Question 17

Define Solubility and Molar Solubility

Answer 17

1) Solubility - The amount of solute that will dissolve in a given amount of solution at a particular temperature.
2) Molar Solubility - The number of moles of solute that will dissolve in a liter of solution. (The molarity of the dissolved solute in a saturated solution).
To compare Ksp values, the compounds must have the same dissociation stoichiometry.

Question 18

What is the effect of Common Ion on Solubility?

Answer 18

Addition of a a soluble salt that contains one of the ions of the “insoluble” salt decreases the solubility of the “insoluble” salt.
For example, addition of NaF to the solubility equilibrium of CaF2 decreases the solubility of CaF2.
CaF2 Ca2 + 2F-
Addition of F- shifts the equilibrium to the left, and decrease the solubility

Question 19

What is the effect of pH on Solubility?

Answer 19

1) For insoluble ionic hydroxides, the higher the pH, the lower the solubility of the ionic hydroxide.
2) The lower the pH, the higher the solubility

Question 20

How can we determine if Precipitation will occur?

Answer 20

By comparing the reaction quotient, Q, for the current solution concentrations to the value of Ksp.

1) Q = Ksp, the solution is saturated (no precipitation)
2) Q < Ksp, the solution is unsaturated (no precipitation)
3) Q > Ksp, the solution is above saturation (the salt will precipitate)

Question 21

Define precipitate

Answer 21

A solid, insoluble ionic compound that forms in, and separates from, a solution.

Question 22

Define Selective Precipitation and Qualitative Analysis.

Answer 22

1) Selective Precipitation - A process involving the addition of a reagent to a solution that forms a precipitate with one of the dissolved ions but not the others.
2) Qualitative Analysis - An analytical scheme that utilizes selective precipitation to determine the amounts of substances in a solution.

Question 23

What are the characteristics of Complex Ion Formation?

Answer 23

1) Complex Ions are ions formed by combining a cation with several anions or neutral molecules.
2) Transition metals tend to be good Lewis acids.
3) They often bond to one or more H2O molecules to form a hydrated ion.
4) The ions or molecules that act as Lewis bases are called ligands. i.e. H2O and NH3.

Question 24

Define Ligand

Answer 24

A neutral molecule or an ion that acts as a Lewis Base with the central metal ion in a complex ion.

Question 25

What happens when a ligand is added to solution that forms a stronger bond than the current ligand?

Answer 25

It will replace the current ligand. For example:
Ag (H2O)2 + 2NH3(aq) Ag(NH3)2(aq) + 2H2O(l)
Generally, H2O is not included because its complex ion is always present in aqueous solution.
AG+(aq) +2NH3(aq) Ag(NH3)2+(aq)
The equilibrium constant for the formation reaction is called the formation constant (Kf).
Kf = [Ag(NH3)2]/[Ag]*[NH3]^2

Question 26

What is a Complex Ion?

Answer 26

An Ion that contains a central metal cation bonded to one or more molecules or ions.

Question 27

What is the effect of Complex Ion Formation on Solubility?

Answer 27

The Solubility of an ionic compound that contains a metal cation that forms a complex ion increases in the presence of aqueous ligands.

Question 28

What are the characteristics of the Solubility of Amphoteric Metal Hydroxides?

Answer 28

1) Some metal hydroxides also act as acids, becoming more soluble in basic solution.
2) Some cations that form amphoteric hydroxides includes: AL3+, CR3+, ZN2+, Pb2+, and Sn2+

Question 29

How do you calculate pKa, given pKb?

Answer 29

pKa * pKb = 14

Question 30

How do you calculate pKb or pKa, given Kb or Ka?

Answer 30

pKb = - log (Kb)
pKa = -log (Ka)

Question 32

What happens at low pH, neutral pH, and high pH for aluminum Hydroxide?

Answer 32

1) Low pH - Al(H2O)63 dissolves
2) Neutral pH - Al(H2O)3 precipitates
3) High pH - Al(H2O)2OH4 dissolves

Question 33

What is the product for the following reactions:

1) NH3 + H2O —>
2) NH3 + HNO3 —>

Answer 33

1) NH3 + H2O —> NH4+ + OH-

2) NH3 + HNO3 —> NH4+ + NO3-

Question 34

How to solve buffer questions?

Answer 34

Try to calculate the concentrations with the total volume. If wrong, try without the total volumes.

Question 35

How to solve titration problems.

Answer 35

Try to get the difference in concentrations when either acid or base is used up. Then divide that concentration by total volume and that gives you [H]. You can take the -log of [H] to find pH.

Question 36

What’s the relationship between Q and also for the following solutions:
1) unsaturated
2 saturated
3) precipitate

Answer 36

1) unsaturated = Q < Ksp
2 saturated = Q = Ksp
3) precipitate = Q > Ksp