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Measurement of pH - the pH meter Bronsted-Lowry definition of acids and bases- an acid is a proton donor- a base is a proton acceptor- conjugate acid/conjugate base pairs- relationship of K a of a conjugate acid and K b of a conjugate base - PowerPoint PPT Presentation
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Acids and Bases - the Three Definitions
1. Measurement of pH - the pH meter
2. Bronsted-Lowry definition of acids and bases - an acid is a proton donor - a base is a proton acceptor - conjugate acid/conjugate base pairs
- relationship of Ka of a conjugate acid and Kb of a conjugate base
3. Lewis definition of acids and bases- a base is an
electron pair donor - an acid is an electron pair acceptor
- some examples of Lewis acids and Lewis bases
Ionization Constants (1)
Ionization Constants (2)
Exercise on Acid/Bases Strength
For each conjugate acid/base pair,
(1) Write the reactions defining Ka
and Kb.
(2)Find the values of pKa, pKb, and
Kb.
(3)Which species is the strongest conjugate acid, which is the strongest conjugate base?
nitrous acid: HNO2 / NO2
-
oxalic acid (2): HC2O4- /
C2O42-
arsenic acid (2): HAsO42- /
AsO43-
carbonic acid (1):H2CO3 / HCO3
-
Hydrolysis Reactions
Which salts undergo hydrolysis? Is the resulting solution acidic, basic, or neutral? Write the hydrolysis reaction (if any). Calculate the pH of a 0.10 M solution.
1. sodium acetate
(basic (pH=8.88), acetate (pKb=9.25)
hydrolyses to produce OH-)
2. ammonium chloride
(acidic (pH=5.12), ammonium (pKa=9.25)
hydrolyses to produce H3O+)
3. calcium chloride
(neutral, no hydrolysis)
4. sodium monohydrogen phosphate
(basic (pH=10.12), HPO42- (pKb2=6.79)
hydrolyses to produce OH-) (you need to consider two conjugate acid/base pairs..)
pH and % Dissociation of a Monoprotic Weak Acid
CH3COOH = CH3COO- + H+ Ka = 1.75 x 10-5
0.1 - x x x (We let x = [H+])
What is the pH of 0.10 M CH3COOH?
[CH3COO-] [H+]
[CH3COOH]
x2
0.1 - x
1.75 x 10-5
Ka = = =
Approximation Method: Since Ka <<1, assume x<<0.1x2 = 0.1 * 1.75 * 10-5 = 1.75 x 10-6
(x << 0.1)
and x = [H+] = 0.0013 M
Calculating % Dissociation and the pH
CH3COOH = CH3COO- + H+
% dissociation = 100 *
[CH3COO-]=
x
0.1
= 1.3%[CH3COOH]ini
t
[H+] = 1.3 x 10-3 M
pH = - log10[H+] = 2.88
Measurement of pH: the pH Meter
pH varies linearly with output voltage and can be
measured over the range pH 0 to
pH 14
Ka and Acid Strength
The stronger the acid, the larger the Ka and the smaller the pKa:
CH3COOH (aq) = CH3COO- (aq) + H+ (aq) Ka = 1.76 x 10-5
HCN (aq) = CN- (aq) + H+ (aq) Ka = 6.17 x 10-10
pKa = 4.75
HNO2 (aq) = NO2- (aq) + H+ (aq)
Ka = 4.6 x 10-4 pKa = 3.34
pKa = 9.21
stronger
weaker
Weak Acids
=
=
=
=
=
=
=
weaker
stronger
Kb and pKb
Arrhenius bases liberate OH- in solution.
Kb is the equilibrium constant for this reaction.NH4OH (aq) = NH4
+ (aq) + OH- (aq)
Kb =
[NH4+] [OH-]
[NH4OH]
= 1.76 x 10-5
pKb = - log10 Kb (definition)
pKb = - log10 (1.8 x 10-5) = 4.74
Kb and Base Strength
The stronger the base, the larger the Kb and the smaller the pKb:
NH4OH (aq) = NH4+ (aq) + OH- (aq)
Kb = 1.8 x 10-5 pKb = 4.74
stronger
weaker
PO43- (aq) + H2O (l) = HPO4
2- (aq) + OH- (aq) Kb = 4.5 x 10-2 pKb =
1.34
Conclusion: phosphate anion is a stronger base than NH4OH.
Kb's of weak bases Strength (Ranked)
2
1
6
7
3
5
4
Acids and Bases - Three Definitions
Arrhenius Definition:
Acids: increases [H+] in aqueous solution
Bases: increases [OH-] in aqueous solution
Bronsted-Lowry Definition: (based on proton transfer reactions)
Acids: proton (H+) donor
Bases: proton (H+) acceptor
Lewis Definition:
Acids: electron pair acceptor
Bases: electron pair donor
Some Lewis Acids and Bases
Lewis bases are characterized by having an available lone pair. Examples are:
O-H-..
.
.:N-HH
H: O-H
H:..I.
.
.
.: :
.
.
:S :
.
.2-
hydroxide iodide ammonia water sulfide
Lewis acids are electron deficient - i.e., electron pair acceptors
Examples are:
H+ Zn2+ Hg2+ Ag+ BF3metal cations electron
deficient compounds
Lewis Acids/Bases - the Most General Definition
Least general definition Most general definition
Arrhenius Bronsted-Lowry Lewis
H+(aq) + :OH-(aq) = H2O(l)
Electron pair acceptor electron pair donor
The Lewis definition generalizes the acid/base concept:
Every Arrhenius acid/base is also a Lewis acid/base.
Every Bronsted acid/base is also a Lewis acid/base.
Example: A strong acid reacts with a strong base:
The Lewis Acid-Base ReactionLewis definition:
Acid: electron pair acceptor Base: electron pair donor
HNO2 + ClO2
- = HClO2 + NO2-
Bronsted-Lowry: acid1 base2 acid2 base1
The Lewis acid is H+ (the electron deficient species)
There are 2 bases (NO2- and ClO2
-), which compete for the acid
The lone pairs donated by these bases are on oxygen atoms:
O=N- O....
.
.
.
. :.. O-N= O
.
..
...
.
.:..[ ]-O=Cl-
O
.
..
...
.
. :..
H..
O=N- O....
.
.
.
. :H.. O=Cl-
O
.
..
...
.
. : -O-Cl= O
.
...
:[ ]..
.
.
.
.
.
...
.
.+
+
Lewis Acids and Bases
The acid/base concept is further generalized by the Lewis acid/base definition. The driving force is the donation of an electron pair to electron-deficient atom.
Lewis acid - an electron pair acceptor
Lewis base - an electron pair donor
H+ + :O-H- = H-O
Lewis acid Lewis base
.
..
.
.
. :
H
F - B + N - H = F - B : N - H
F
F:
F
H
H
F H
H
Lewis acid Lewis base
Complex Ions in Solution
One example of Lewis acid-base neutralization involves the stepwise complexes formed between Hg(II) and I- ion.
There are four stepwise reactions, each of which is an acid/base neutralization by the Lewis definition:Hg2+(aq) + I-(aq) = HgI+
(aq)HgI+(aq) + I-(aq) = HgI2(s) (red-brown ppt)HgI2(s) + I-(aq) = HgI3
-
(aq)HgI3
-(aq) + I-(aq) = HgI4
2-(aq)Identify the Lewis acid and Lewis base in each reaction.
Complex Ions and Solubility
Stepwise Lewis acid/base complexes form between Al3+(aq) and OH- ion. All charged species are soluble in aqueous solution. Only the uncharged Al(OH)3(s) forms a white precipitate.
There are four stepwise reactions, each of which is an acid/base neutralization by the Lewis definition:
Identify the Lewis acid and Lewis base in each reaction.
Al3+ + :OH- = AlOH2+
AlOH2+ + :OH- = Al(OH)2
+
Al(OH)2+ + :OH- =
Al(OH)3 (s)Al(OH)3 + :OH- = Al(OH)4
-
white precipitate