Chapter 16 Acid-Base Equilibria. Dissociation of water Autoionization or autoprotolysis Ion-product...
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Transcript of Chapter 16 Acid-Base Equilibria. Dissociation of water Autoionization or autoprotolysis Ion-product...
Chapter 16
Acid-Base Equilibria
Dissociation of water
Autoionization or autoprotolysis
Ion-product constantAutoprotolysis constant
-(aq)(aq)(l)2 OHHOH
]][OHH[K
O]H[
]][OH[HK
-w
2
-
constant
Kw = [H+][OH-] = 1.0x10-14
When [H+] = [OH-] neutral. Doesn’t usually happen.
As one increases, the other decreases; the product must equal 1.0x10-14.
When
[H+] > [OH-] acidic
[OH-] > [H+] basic
H+ is a proton with no electrons.In water:
H
HOH
Hydronium ion
Bronstead-Lowry Acid-Base
Acid - Can donate a proton
Base - Can accept a proton
*Doesn’t have to be in H2O. Can be in other solvents.
Conjugate Acid-Base Pairs
(aq)3-(aq)(l)2(aq) OHAOHHA
conj base
conj base
conj acid
conj acid
(aq)(aq)(l)23(aq) OHNH4OHNH
(aq)3-2(aq)(l)22(aq) OHNOOHHNO
The stronger an acid, the weaker its conjugate base.
The weaker an acid, the stronger its conjugate base.
pH scale
pH = -log [H+]
Remember
Kw = (1x10-7)(1x10-7) = 1.0x10-14
pH = -log [H+] = -log (1x10-7)
pH = 7 (neutral)
[H+] pH
acidic > 1.0x10-7 < 7.00
basic < 1.0x10-7 > 7.00
You can also speak in terms of [OH-]
pOH = -log [OH-]
= 14 - pH
Because
pH + pOH = -log Kw = 14
Measure pH by
pH meter
Acid-base indicators
Litmus
red = pH < 5
blue = pH > 8
Figure 16.7 shows several acid-base indicators and their ranges
Strong Acids and Bases
Strong electrolytes
Completely ionize
HA + H2O A- + H3O+
Bases form hydroxides in solvent
In H2O, Alkali metal hydroxides
Alkaline earth metal
Hydroxides (except Be)
Many are insoluble
Also, substances that will abstract a H+ from H2O.
O2- + H2O 2OH-
Na2O or CaO would do this. O2-, H-, N3- bases that would do this.
Weak acids
-(aq)(aq)(aq) AH HA
Only partially ionize
[HA]
]][A[HK
-
a
Acid dissociation constant
Larger Ka means stronger acid. ex.
N
O
C - O - H
O=
0.020M solutionpH = 3.26? Ka
pH = -log [H+] = 3.26[H+] = 5.50x10-4
N
O
C - OH
O=
N
O
C - O
O=
+ H+
HA A- H+
1:1
a
24-5-
4-
need stillthis
4-4-
a
-
a
K(0.0195)
)(5.5x101.55x10
[HA]0.0195MM)(5.5x10-M020.0
? [HA]
]][5.5x10[5.5x10K
[HA]
]][A[HK
Can calculate pH in same manner if you have Ka and concentration of solution.
Let’s use niacin again.
N
O
C - OH
O=
N
O
C - O
O=+ H+
HA A- H+
x)-(0.010
[HA]][H
(x)
][A
(x)
1.5x10
pH ?solution 0.010M1.5x10K
-5-
-5a
x)-(0.010
x1.5x10
25
** Simplifying Assumption **
x is very very small compared to 0.010M
sooooooooo,
ignore x in denominator
4-7-2
7-5-2
5-2
3.9x10x1.5x10x
1.5x10)(0.010)(1.5x10x
x105.10.010
x
pH = -log [H+]
x = [H+] = 3.9x10-4
pH = 3.41
What percent of niacin molecules ionized?
3.9%100x 0.010
x109.3 -4
Polyprotic Acids
ex. H2SO4 H3PO4 H2SeO4
H2SO4 H+ + HSO4-
Ka1 = 1.7x10-2
HSO4- H+ + SO4
2-
Ka2 = 6.4x10-8
Ka1 always larger than Ka2
If Ka1/ Ka2 103, can estimate pH by Ka1 only.
Weak Bases
ex. Amines
“an organic substituted ammonia”
ammoniaNH3
N
H
HH NH
CH3
H
methyl amine
N
H
CH3 + H2O H N
H
CH3 + OH-H
H
ClO- + H2O HClO + OH-
Kb = 3.3x10-7
][ClO
][HClO][OHx103.3
-
-7-
Can use this in the same manner in which you used Ka.
Anions of weak acids
Ka and Kb
How are they related?
][NH
]][OH[NHK
][NH
]][H[NHK
OHNHOHNH
HNHNH
3
-4
b4
3a
-(aq)4(aq)(l)23(aq)
(aq)3(aq)4(aq)
-(aq)(aq)(l)2
-(aq)4(aq)(l)23(aq)
(aq)3(aq)4(aq)
OHHOH
OHNHOHNH
HNHNH
1)
2)
3)
When two reactions are added together, the equilibrium constant for the third reaction is given by the product of equilibrium constants of equations 1 and 2.
K1 x K2 = K3
rxn 1 rxn 2 rxn 3
w-
3
-4
4
3ba
K]][OH[H
][NH
]][OH[NH
][NH
]][H[NHKK
Special Case
Ka x Kb = Kw
For conjugate acid-base pairs.
Bond polarity and Bond strength effect on Acid-base behavior: In binary acids
polarity(across a row) acidity
bond strength(in a group) acidity
stability of conj. base acidity
Metal hydrides are basic or show no acid/base properties in H2O.
Nonmetal hydrides are acidic or show no acid/base properties in H2O (except NH3)
Acidity increases moving down a group.
Oxyacids
H O S
O
O
O
HHave unprotonated and protonated oxygens.
Y O H H3PO4
• As electronegativity of Y increases, acidity increases.
• As number of unprotonated oxygens increases, acidity increases (effect of formal charge and oxidation number)
•Ex. HClO, HClO2, HClO3, HClO4
Carboxylic Acids
RC
OH
O COOH = Carboxyl group
R = H or an organic group.
The more electron withdrawing R is, the greater the acidity (this stabilizes anion and weakens O-H bond)
ex.
CH C
H
H O
OH
Acetic acidKa = 1.8x10-5
CF C
F
F O
OH
Trifluoroacetic acidKa = 5.0x10-1
Lewis Acids and Bases
This is a completely different definition for acid/base chemistry than what you have seen thus far!!!
Lewis acid = electron pair acceptor
Lewis base = electron pair ‘donor’
Not giving them away, just has them available to ‘share’.
H+ Bronstead-Lowry acid
also a Lewis acid
H+ electron pair acceptor
OH -
Electron pair donorLewis basealso Bronstead-Lowry base
B
H H
H
BH3 not a Bronstead-Lowry acid, but it’s a Lewis acid
Incomplete Octet
N
H
H
H
Lewis Basehas an electron pair available to attack an area that is e- deficient
Transition metal ions are often Lewis Acids. They have vacant d orbitals. (s and p also)
3
6
_3
NCFe
NC6Fe
H
HO
O = C = O Can be a Lewis Acid because e- density around the C is bound in just 2 directions.
H
HO
=
=
O
O
C
H
HO
=
=
O
O
C
H
H O
=
=
O
O
CCarbonic acid
Hydrolysis of metal ions
Metal ions have positive charge so they attract the lone e- pair on H2O molecules
6 ofthese
HO
H Fe3+H
OH
HO
H
HO
H
HO
H
H
O
H
H
O
H
Fe
3+
Because the metal is (+), e- density of H2O moves toward the metal. When this happens, there is less e- density in water’s O-H bonds, so H+ can come off easier… pH will drop.
The higher the charge density of the metal ion, the greater the acidity of its aqua complex.
OHin acidity stronger radius ionicsmaller
charge )(greater soooo
radius ionic
chargedensity charge
2