Topic 19: Ionic Equilibria in Aqueous Systems Flashcards
Acid-Base Buffer
Solution that lessens the impact on the pH of the addition of acid/base
Components of conjugate acid-base buffer
Conjugate acid-base pair
a) WA + CB
b) WB + CA
How does an acid-base buffer neutralize acid/base?
a) A large amount of the acidic component of the buffer consumes small amounts of added OH-
b) A large amount of the basic component A- consumes small amounts of added H3O+
Definition of buffer capacity
Measure of the ability to maintain the pH following the addition of strong acid or base
Factors affecting buffer capacity
a) Absolute concentration
b) Relative concentration
How does absolute concentration affect buffer capacity?
a) The more concentrated the buffer components, the greater the capacity
b) Amounts of the components must be greater than the amount of H3O+ or OH- added
How does relative concentration affect buffer capacity?
a) The closer the component concentrations are to each other, the greater the capacity
b) A buffer whose pH is equal to or near the pKa of its acid has the highest capacity for a given concentration
Definition of buffer range
pH range over which the buffer is effective
What is the usable range of a buffer?
Within ±1 pH unit of the pKa of the acid component
How is a buffer prepared?
a) Choose conjugate acid-base pair
b) Calculate the ratio of buffer component concentrations
c) Determine the buffer concentration
d) Mix the solution and correct the pH / Partial neutralization
Conditions for common-ion effect
a) An ion is added to an equilibrium mixture that already contains that ion
b) The position of equilibrium shifts away from forming it
Consequence of the common ion effect
Common ion A- suppresses the dissociation of HA
Henderson-Hasselback Equation
pH=pKa+log([base]/[acid] )
Strong Acid + Strong Base Titration Curve (Major species in each step)
a) Initial pH = pH of strong acid
(Cl- / H3O+)
b) pH changes only gradually until equivalence (Cl- / H3O+ / Na+)
c) Very sharp jump in pH at equivalence (6 – 8 units) (Na+ / Cl-)
d) After equivalence the curve flattens out at the pH of the strong base (Na+ / Cl- / OH-)
Weak Acid + Strong Base Titration Curve (Major species in each step)
a) Initial pH = pH of weak acid (HPr)
b) pH stays relatively constant until equivalence (Pr- / HPr / Na+)
c) Jump in pH at equivalence (Pr- / Na+)
d) After equivalence, the curve flattens out at the pH of the strong base (Pr- / Na+ / OH-)
What is the half-equivalence point?
Position where half of an acid has been neutralized by a base and converted into a salt, while the other half in the flask remains unreacted
Strong Acid and Weak Base (Major species in each step)
a) Initial pH = pH of weak base (NH3)
b) pH stays relatively constant through the buffer region to equivalence (NH3 / NH4+ Cl-)
c) Jump in pH at equivalence (Cl- / NH4+)
d) After equivalence the curve flattens out at a pH of weak base (Cl- / NH4+ / H3O+)
Strong base and Polyprotic acids (Major species in each step)
a) The same amount of base is required per mole of H+
b) There are two equivalence points and two buffer regions. The pH at the midpoint of each buffer region is equal to the pKa of the acidic species present then
c) The pH of the first equivalence point is below 7.00 due to weak acid HSO3-
d) The pH of the second equivalence point is above 7.00 due to weak base SO3-
Definition of an indicator
Weak acid or a weak base where the components of the conjugate acid-base pair have different colors
Formula that explain the function of an indicator
Equilibrium response to a change in the pH of the medium
HIn(aq) ⇌ H+(aq) +In-(aq)
When does an indicator change color (end point)?
When the pH is equal to their pKa
When is an indicator effective at signaling the equivalence point?
When its end-point (pH at which it changes color) coincides with the pH at the equivalence poin
What is the pH range of an indicator?
pH values between which the indicator has intermediate colors
a) +- 1 pH on either side of the pKa
Solubility-product constant (Ksp)
Equilibrium value between solid solute and aqueous ions at small amounts
a) How far the dissolution occurs?
b) Ksp=[Pb2+][F-]^2
Effect of a common ion on solubility
Adding a common ion decreases the solubility of a slightly soluble ionic compound
Assumption of Ksp
Small dissolved amount of compound dissociates completely into separate ions
Effect of acidity on solubility
If a slightly soluble ionic compounds contains the anion of a weak/strong acid, addition of H3O+ increases/does not change the compound’s solubility
What would happen if… ?
a) Qsp = Ksp
b) Qsp < Ksp
c) Qsp > Ksp
a) No change
b) No precipitate will form
c) Precipitate will form
Explanation of selective precipitation
A solution of one precipitating ion is added to a solution of two ionic compounds until the Qsp of the more soluble compound is almost equal to its K_sp
Definition of complex ions
Adduct of a central metal ion covalently bonded to two or more ligands (anions/molecules)
Formation of a complex ions
When a salt dissolves in water, a complex ion forms with water as ligands around the metal ion
To what extent complex ions form from its hydrated ion?
Complex ions form readily from the hydrated ion
Relationship between complex ion and solubility of precipitates
A ligand increases the solubility of a slightly soluble ionic compound if it forms a complex ion with the compound’s cation.
Dissolution of complex ions of amphoteric hydroxides
Dissolve very little in water but to a much greater extent in both acidic and basic solution
a) Reaction of H3O+ with OH- anion
b) Formation of the complex ion with 4OH-. Insoluble hydroxide might form, which is further dissolved by addition of OH-
