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Chemistry Chemistry 232232
Applications of Aqueous Equilbria
The BrThe Brøønsted Definitionsnsted Definitions
Brønsted Acid proton donor
Brønsted Base proton acceptor
Conjugate acid - base pair an acid and its conjugate base or a base and its conjugate acid
Example Acid-Base ReactionsExample Acid-Base Reactions
Look at acetic acid dissociatingCH3COOH(aq) CH3COO-(aq) + H+(aq)
Brønsted acid Conjugate base
Look at NH3(aq) in waterNH3(aq) + H2O(l) NH4
+(aq) + OH-(aq) Brønsted base conjugate acid
Representing Protons in Aqueous Representing Protons in Aqueous SolutionSolution
CH3COOH(aq) CH3COO-(aq) + H+(aq)
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)
HCl (aq) Cl-(aq) + H+(aq)
HCl(aq) + H2O(l) Cl-(aq) + H3O+(aq)
What is HWhat is H++ (aq)? (aq)?
+H O
H
HH3O+
HO
H
H +OH
H
H5O2+
+H
HO
OH2H2O
H2O
+
H
H
H+
H9O4+
Representing ProtonsRepresenting Protons
Both representations of the proton are equivalent.
H5O2+ (aq), H7O3
+ (aq), H9O4+ (aq) have
been observed.
We will use H+(aq)!
The Autoionization of Water The Autoionization of Water
Water autoionizes (self-dissociates) to a small extent
2H2O(l) H3O+(aq) + OH-(aq)
H2O(l) H+(aq) + OH-(aq) These are both equivalent definitions of the
autoionization reaction. Water is acting as a base and an acid in the above reaction water is amphoteric.
The Autoionization The Autoionization EquilibriumEquilibrium
From the equilibrium chapter
)OH(a
)OH(a )OH(a or
)OH(a
)OHa( )H(a = K
2
3
2c
But we know a(H2O) is 1.00!
The Defination of KThe Defination of Kww
Kw = a(H+) a(OH-)
Ion product constant for water, Kw, is the product of the activities of the H+ and OH-
ions in pure water at a temperature of 298.15 K
Kw = a(H+) a(OH-) = 1.0x10-14 at 298.2 K
The pH scaleThe pH scale Attributed to Sørenson in 1909
We should define the pH of the solution in terms of the hydrogen ion activity in solution
pH -log a(H+)
Single ion activities and activity coefficients can’t be measured
Determination of pHDetermination of pH
What are we really measuring when we measure the pH?
pH -log a(H+)
a (H+) is the best approximation to the hydrogen ion activity in solution.
How do we measure a(H+)?
For the dissociation of HCl in water
HCl (aq) Cl-(aq) + H+(aq) We measure the mean activity of the acid
a(HCl) = a(H+) a(Cl-)
a(H+) a(Cl-) = (a(HCl))2
Under the assumption
a(H+) = a(Cl-) We obtain
a´(H+) = (a(HCl))1/2 = a(HCl)
Equilibria in Aqueous Solutions Equilibria in Aqueous Solutions of Weak Acids/ Weak Basesof Weak Acids/ Weak Bases
By definition, a weak acid or a weak base does not ionize completely in water ( <<100%). How would we calculate the pH of a solution of a weak acid or a weak base in water?
To obtain the pH of a weak acid solution, we must apply the principles of chemical equilibrium.
Equilibria of Weak Acids in Equilibria of Weak Acids in WaterWater: The K: The Kaa Value Value
Define the acid dissociation constant Ka
For a general weak acid reaction
HA (aq) H+ (aq) + A- (aq)
HAa
Aa HaKa
Equilibria of Weak Acids in Equilibria of Weak Acids in WaterWater
For the dissolution of HF(aq) in water.
HF (aq) H+ (aq) + F- (aq)
HFa
Fa HaKa
The small value of Ka indicates that this acid is only ionized to a small extent at equilibrium.
The Nonelectrolyte ActivityThe Nonelectrolyte Activity
HF (aq) H+ (aq) + F- (aq) The undissociated HF is a nonelectrolyte
a(HF) = (HF) m[HF] m[HF]
(HF) 1
4a 10x1.7
HFm
Fa HaK
Equilibria of Weak Bases in Equilibria of Weak Bases in WaterWater
To calculate the percentage dissociation of a weak base in water (and the pH of the solutions)
CH3NH2 (aq) + H2O CH3NH3+(aq) + OH- (aq)
We approach the problem as in the case of the weak acid above, i.e., from the chemical equilibrium viewpoint.
The KThe Kbb Value Value
Define the base dissociation constant Kb For a general weak base reaction with water
B (aq) + H2O (aq) B+ (aq) + OH- (aq)
Ba
OHa BaKb
For the above system
Bm
OHaBaKb
Examples of Acid-Base Examples of Acid-Base CalculationsCalculations
Determining the pH of a strong acid (or base solution).
Determining the pH of a weak acid (or base solution).
Calculating the pH of Solutions Calculating the pH of Solutions of Strong Acidsof Strong Acids
For the dissolution of HCl, HI, or any of the other seven strong acids in water
HCl (aq) H+ (aq) + Cl- (aq) HI (aq) H+ (aq) + I- (aq) % eq = 100%
The pH of these solutions can be estimated from the molality and the mean activity coefficient of the dissolved acid
pH -log ( (acid) m[H+])
For the dissolution of NaOH, Ba(OH)2, or any of the other strong bases in water
NaOH (aq) Na+ (aq) + OH- (aq) Ba(OH)2 (aq) Ba2+ (aq) + 2OH- (aq)
% eq = 100%
Calculating the pH of Solution of Calculating the pH of Solution of Strong BasesStrong Bases
The pH of these solutions is obtained by first estimating the pOH from the molality and mean activity coefficient of the dissolved base
pOH -log ( (NaOH) m[OH-])
pOH -log{ (Ba(OH)2) 2 m[Ba(OH)2]}
pH = 14.00 - pOH
Calculating the pH of a Weak Calculating the pH of a Weak Acid SolutionAcid Solution
The pH of a weak acid solution is obtained via an iterative procedure.
We begin by making the assumption that the mean activity coefficient of the dissociated acid is 1.00.
We ‘correct’ the value of (H+) by calculating the mean activity coefficient of the dissociated acid.
Repeat the procedure until (H+) converges.
Measuring the pH of SolutionsMeasuring the pH of Solutions
Because the activity of a single ion cannot be measured, we can only measure our ‘best approximation’ to the hydrogen ion activity.
Let’s assume that we are going to couple a hydrogen electrode with another reference electrode, e.g., a calomel reference electrode.
A Cell for ‘Measuring’ the pHA Cell for ‘Measuring’ the pH
Half-cell reactions.HgCl2 (s) + 2e- Hg (l) + 2 Cl- (aq)
E(SCE) = 0.2415 V
2 H+ (aq) + 2e- H2 (g)
E (H+/H2) = 0.000 V Cell Reaction
HgCl2 (s) + H2 (g) Hg (l) + 2 H+ (aq) + 2Cl- (aq)
Ecell = (0.2415 - 0.0000 V) = 0.2415 V
Pt H2 (g), f=1 H+ (aq) HgCl2 Hg Cl- (aq), 3.5 M Pt
The Nernst EquationThe Nernst Equation
For the above cell
2
cell Hf
Cla Haln
F 2
RTV2415.0E
Note since the concentration of the KCl on one side of the liquid junction is so large, the magnitude of the junction potential should be small!
The Practical Problem The Practical Problem
The activity of the Cl- ion in the cell is not accurately known.
We try to place the cell in a reference solution with an accurately known pH (solution I).
Next place the solution whose pH we are attempting to measure into the cell (solution II).
Assuming that the ELJ and the a(Cl-) are the same in both cases,
I,HalnF
RTII,Haln
F
RTEE III
cellcell
Substituting the definition of the pH into the above expression,
Icell
IIcell
III EETR2.303
FpHpH
Standard SolutionsStandard Solutions
Generally, two solutions are used as references. Saturated aqueous solution of sodium hydrogen
tartarate, pH = 3.557 at 25C. 0.0100 mol/kg disodium tetraborate, pH =
9.180 at 25C.
The Glass ElectrodeThe Glass Electrode
The glass electrode has replaced the hydrogen electrode in the operational definition of the pH.
Glass ElectrodesGlass Electrodes
Measuring the pH – the glass electrode is immersed in the solution of interest.
Inner solution – solution is generally a phosphate buffer with a sufficient quantity of Cl- (aq).
Silver-silver chloride electrode is sealed within the cell and a calomel electrode is used as the reference electrode.
Glass Electrodes and pHGlass Electrodes and pH
The potential difference across the special glass membrane arises to equilibrate the hydronium ions inside the membrane with those outside the membrane.
The Definition of a Buffer The Definition of a Buffer
Buffer a reasonably concentrated solution of a weak acid and its conjugate base that resists changes in the pH when an additional amount of strong acid or strong base is added to the solutions.
How would we calculate the pH of a buffer solution?
HCOOHm
HCOOa HaKa
HCOOHm
HCOOa HalogKlogpK aa
HCOOHmlogHCOOalogHalogpKa
note pH = -log a(H+)
Define pKa = -log (Ka )
The Buffer EquationThe Buffer Equation
HCOOHm
)HCOO(alogpHpKa
Substituting and rearranging
HCOOHm
)HCOO(alogpKpH a
The Generalized Buffer The Generalized Buffer EquationEquation
The pH of the solution determined by the ratio of the weak acid to the conjugate base. This equation (the Henderson-Hasselbalch equation) is often used by chemists, biochemists, and biologists for calculating the pH of a solution of a weak acid and its conjugate base!
acid weakm
)base .conj(alogpKpH a
Note: The Henderson-Hasselbalch equation is really only valid for pH ranges near the pKa of the weak acid!
Buffer CH3COONa (aq) and CH3COOH (aq))
CH3COOH (aq) ⇄ CH3COO- (aq) + H+ (aq)
The Equilibrium Data Table
n(CH3COOH) n(H+) n(CH3COO-)
Start A 0 B
Change -eq + eq +eq
m (A-eq) (eq) (B+ eq)
The pH of the solution will be almost entirely due to the original molalities of acid and base!!
]A[m
BmlogpKpH a
This ratio will be practically unchanged in the presence of a small amount of added strong acid or base
The pH of the solution changes very little after adding strong acid or base (i.e., it is buffered)
How does the pH change after the addition of strong acid or base?
Example of Buffer Example of Buffer CalculationsCalculations
How do we calculate the pH of a buffer solution?
The pH of a Buffer SolutionThe pH of a Buffer Solution
The major task in almost all buffer calculations is to obtain the ratio of the concentrations of conjugate base to weak acid!
Using the Ka of the appropriate acid, the pH of the solution is obtained from the Henderson-Hasselbalch equation.
Adding Strong Acid or Base to Adding Strong Acid or Base to Buffer Solutions Buffer Solutions
To obtain the pH after the addition of a strong acid or base, we must calculate the new amount of weak acid and conjugate base from the reaction of the strong acid (or base) to the buffer system.
The pH of the solution may again be calculated with the Henderson-Hasselbalch equation.
Solubility EquiSolubility Equilibria libria
Examine the following systems
AgCl (s) ⇌ Ag+ (aq) + Cl- (aq)
BaF2 (s) ⇌ Ba2+ (aq) + 2 F- (aq)
Using the principles of chemical equilibrium, we write the equilibrium constant expressions as follows
10
sp 10x8.1Cla AgaK
1AgCla note
AgCla
Cla AgaK
622sp 10x0.1Fa BaaK
Calculate the solubility of a solid in the presence of a common ion.
Examples of KExamples of Kspsp Calculations Calculations
Calculate the solubility of a sparingly soluble solid in water.
Calculate the solubility of a solid in the presence of an inert electrolyte.
Solubility of Sparingly Soluble Solubility of Sparingly Soluble Solids in WaterSolids in Water
AgCl (s) ⇌ Ag+ (aq) + Cl- (aq) We approach this using the principles of
chemical equilibrium. We set up the equilibrium data table, and calculate the numerical value of the activity of the dissolved ions in solution.
The Common Ion EffectThe Common Ion Effect
What about the solubility of AgCl in solution containing NaCl (aq)?
AgCl (s) ⇌ Ag+ (aq) + Cl- (aq)
NaCl (aq) Na+ (aq) + Cl- (aq)
AgCl (s) ⇌ Ag+ (aq) + Cl- (aq)
Equilibrium is displaced to the left by LeChatelier’s principle (an example of the common ion effect).
Solubility in the Presence of Solubility in the Presence of an Inert Electrolytean Inert Electrolyte
What happens when we try to dissolve a solid like AgCl in solutions of an inert electrolyte (e.g., KNO3 (aq))?
We must now take into account of the effect of the ionic strength on the mean activity coefficient!
The Salting-In EffectThe Salting-In Effect
AgCl (s) ⇌ Ag+ (aq) + Cl- (aq). We designate the solubility of the salt in the
absence of the inert electrolyte as so = m(Ag+) = m(Cl-) at equilibrium.
2o
2
102
sp
s
10x8.1ClmAgm
Cla AgaK
For a dilute solution
2osp sK1
Designate s as the solubility of the salt in the presence of varying concentrations of inert electrolyte.
o2
sp
sssK