chemistry 232
DESCRIPTION
Chemistry 232. Reaction Equilibria in Nonideal Systems. The Equilibrium Constant. For a nonideal system, the nonstandard Gibbs energy of reaction is written. The Equilibrium Condition. If we apply the equilibrium conditions to the above equation. The Autoionization of Water. - PowerPoint PPT PresentationTRANSCRIPT
The Equilibrium Constant
For a nonideal system, the nonstandard Gibbs energy of reaction is written
J
eqJrrJaRTGG ,
The Equilibrium Condition
If we apply the equilibrium conditions to the above equation
J
eqJrJaRTG0 ,
J
oJJ
JeqJr
JaRTG ,
J
eqJo JaK
,
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 amphoteric.
The Autoionization Equilibrium
From the equilibrium chapter
22
3
2
OHaOHa OHa
or
OHaOHa( Ha
K =
)()()(
)())(
But we know a(H2O) is 1.00!
The Defination of Kw
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 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 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
Equilibria in Aqueous Solutions of 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?
Equilibria of Weak Acids in Water: The Ka 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 Water
For the dissolution of HF(aq) in water.
HF (aq) H+ (aq) + F- (aq)
4a 10x17
HFaFa Ha
K
.
The Nonelectrolyte Activity
HF (aq) ⇌ H+ (aq) + F- (aq) The undissociated HF is a nonelectrolyte
a(HF) = (HF) m[HF] m[HF]
(HF) 1
4a 10x17
HFmFa Ha
K
.
Equilibria of Weak Bases in Water
Calculate the percentage dissociation of a weak base in water (and the pH of the solutions)
CH3NH2 (aq) + H2O ⇌ CH3NH3+(aq) + OH-
(aq)
The Kb 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
BmOHaBa
Kb
Calculating the pH of Solutions of Strong Acids
For the dissolution of HCl, HI, or any of the other seven strong acids in water
HCl (aq) H+ (aq) + Cl- (aq) 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+])
Calculating the pH of Solution of Strong Bases
For the dissolution of NaOH, Ba(OH)2, or any of the other strong bases in water
NaOH (aq) Na+ (aq) + OH- (aq)
pOH = -log ( (base) m[OH-])
Calculating the pH of a Weak Acid 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.
The Definition of a Buffer
Buffer a reasonably concentrated solution of a weak acid and its conjugate base
Buffers resist pH changes when an additional amount of strong acid or strong base is added to the solutions.
The Generalized Buffer Equation
The pH of the solution determined by the ratio of the weak acid to the conjugate base.
Henderson-Hasselbalch equation often used for buffer calculations!
acid weakm
base conjapKpH a
).(log
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!!
][log
AmBm
pKpH a
Solubility Equilibria
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
The 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 an 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 Effect
AgCl (s) ⇌ Ag+ (aq) + Cl- (aq). Designate the solubility of the salt in
the absence of the inert electrolyte as so = m(Ag+) = m(Cl-) at equilibrium.
2o2
102
sp
s
10x81ClmAgm
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.
o2sp
sssK
Reaction Equilibria in Nonideal Gaseous Systems
For a nonideal system gaseous, the nonstandard Gibbs energy of reaction is written
J
eqJrrJfRTGG ,
JoeqjJ
JoeqJo
JJ
P
Px
P
fK
,,
The Equilibrium Condition
Calculate the equilibrium composition from the fugacity coefficients from compression factor data
Jo
eqJ
JJ
o J
J P
PxK
,