Chapter 18
Equilibria Involving
Acids & bases
ARRHENIUS THEORY for ACIDS and BASES
 ACIDS: produce hydrogen ions (protons), H+, in
solution
 BASES: produce hydroxide ions, OH-,in solution,
 NEUTRALIZATION: H+ + OH-  H2O
 Problems with Arrhenius Theory
* H3O+: Hydronium ion rather than H+
* OH(H2O)3- present in solution, not OH* Other substances also have acidic or basic
properties
H+ surrounded by four H-bonded H2O molecules
H9O4+
OHsurrounded by
three H-bonded
H2O molecules
OH(H2O)3-
Bronsted-Lowery Theory of Acids and Bases
 Acid – any substance donating a proton, H+
 Base – any substance accepting a proton
 Conjugate Acid-Base Pairs:
e.g. HF + NH3  NH4+ + Facid 1
base 2
acid 2
base 1
 AMPHOTERIC substances have both acidic and
basic properties.
 Mono-, di-, tri-,………. to polyprotic acids.
 Acidic versus nonacidic H atoms in compounds.
For each of the following reactions, identify the acid, the
base, the conjugate base, and the conjugate acid
1.
2.
3.
4.
5.
6.
7.
8.
H2O + H2O  H3O+ + OHH2PO4- + H2PO4-  H3PO4 + HPO42H2SO4 + H2O  H3O+ + HSO4H2PO4- + H2O  H3PO4 + OHCO2 + 2H2O  HCO3- + H3O+
H2PO4- + H2O  HPO42- + H3O+
Fe(H2O)63+ + H2O  Fe(H2O)5(OH) 2+ + H3O+
HCN + CO32- CN- + HCO3-
Graphic
representations
of strong and
weak acid
equilibria
Strong Acid:
100% Dissociation
into ions
HA(aq) + H2O(l)  H3O+(aq) + A- (aq)
Equilibrium Favors undissociated acid

-
[H3O ] [A ]
Ka 
[HA]
Weak Acid:
Very little
Dissociation
into ions
Acid
strength
versus
conjugate
base
strength
Bronsted-Lowery Theory: Acid and Base Strengths
Proton transfers occur from a
• Strong acid to a strong base
e.g. HCl + NaOH  H2O + NaCl
• Weak acid to strong base
e.g. CH3COOH + NaOH  ??
•
Weak base to a stronger base
e.g. HSO41- + HSO31-  ???
• Will a Reaction occur between…..
a. HS1- and F 1- ??
b. HCl and ClO2 1- ??
c. HCl and ClO4 1- ??
d. HCl and HNO3??
Relative strengths of some Bronsted-Lowry acids and their conjugate bases
Strongest
Acids
Weakest
Acid
Acid
HClO4
H2SO4
HI
HBr
HCl
HNO3
H3 O +
HSO4H2SO3
H3PO4
HNO2
HF
CH3CO2H
H2CO3
H2 S
NH4+
HCN
HCO3HSH2 O
NH3
OH-
Base
ClO4HSO4IBrClNO3H2 O
SO42HSO3H2PO4NO2FCH3CO2HCO3HSNH3
CNCO32S2OHNH2O2-
Weakest
bases
Strongest
bases
Bronsted-Lowery Theory: Acid and Base Strengths
LEVELING EFFECT of SOLVENTS:
The strongest acid in a solvent is the conjugate acid of the
solvent. The strongest base is the conjugate base.
Acid
H3O+ in water
Base
OH- in water
Autoionization of Water and the pH Scale
+
H2O(l)
H2O(l)
+
H3O+(aq)
OH-(aq)
Autoionization of Water
H2O(l) + H2O(l)  H3O+(aq) + OH-(aq)
Kw = [H3O+] [OH-] = 1.0 x 10-14 at 25oC
At equilibrium
[H3O+] = [OH-] = 1.0 x 10-7
Kw changes with temperature
but [H3O+] = [OH-]
pH and pOH SCALES
pH = - log [H3O+]
pOH = - log
[OH ]
pH, pOH CALCULATIONS
pH = - log [H3O+]
pOH = - log [OH-]
Kw = [H3O+] [OH-] = 1.0 x 10-14
+
so p[H3O ] + p[OH ] = 14.00
or
pH = 14.00 – pOH
CALULATE SOME pH and pOH VALUES
Basic
Neutral
Acidic
[H+]
10-14
10-13
10-12
10-11
10-10
10-9
10-8
10-7
10-6
10-5
10-4
10-3
10-2
10-1
1
pH
14
13
12
11
10
9
8
7
6
5
4
3
2
1
0
1 M NaOH
Ammonia
(Household
Cleaner)
Blood
Pure Water
Milk
Vinegar
Lemon juice
Stomach acid
1 M HCl
Calculate the pH of each solution:
a.
[H+] = 1.4 x 10-3 M
e. [OH-] = 8 x 10-11 M
b.
[H+] = 2.5 x 10-10 M
f. [OH-] = 5.0 M
c,.
[H+] = 6.1 M
g. pOH = 10.5
d.
[OH-] = 3.5 x 10-2 M
h. pOH = 2.3
Calculate [H+] and [OH-] for each solution:
a.
pH = 7.41 (the normal pH of blood)
b.
pH = 15.3
c.
pH = -1.0
e. pOH = 5.0
d.
pH = 3.2
f. pOH = 9.6
 How many significant figures are there in the numbers: 10.78, 6.78,
0.78? If these were pH values, to how many significant figures can you
express the [H+]? Explain any discrepancies between your answers to
the two questions.
Values of Kw as a function of temperature are as follows:
Temp (oC)
Kw
0
1.14 x 10-15
25
1.00 x 10-14
35
2.09 X 10-14
40
2.92 x 10-14
50
5.47 x 10-14
a. Is the autoionization of water exothermic or
endothermic?
b. What is the pH of pure water at 50oC?
Values of Kw as a function of temperature are as follows:
Temp (oC)
Kw
0
1.14 x 10-15
25
1.00 x 10-14
35
2.09 X 10-14
40
2.92 x 10-14
50
5.47 x 10-14
a. Is the autoionization of water exothermic or
endothermic?
b. What is the pH of pure water at 50oC?
c. Restate your answers to water at 50oC. Which of the
three criteria for neutrality is most general?
d. From a plot of ln(Kw) versus 1/T (using the Kelvin
scale), estimate Kw at 37oC, normal physiological
temperature.
e. What is the pH of a neutral solution at 37oC?
-28
-29
Y = -9.2338 – 6870.6x R^2 = 0.999
-30
-31
-32
-33
-34
-35
0.0028
0.0030
0.0032
1/T
0.0034
0.0036
pH MEASUREMENT

Indicators: colored weak acids and bases
 pH Meters: Glass membrane with a voltage (potential)
difference across the glass.
pH and BODY CHEMISTRY
Normal pH 7.3 to 7.5
Acidosis pH < 7.3
Alkalosis pH > 7.45
Body Chemistry is “buffered” with bicarbonates (HCO3-)
dihydrogenphosphates (H2PO4-) and proteins which help
to maintain a constant pH
WEAK ACIDS
IONIZATION CONSTANTS
HA(aq) + H2O(l)  H3O+(aq) + A- (aq)

-
[H3O ] [A ]
Ka 
[HA]
Ka values at 25oC are known and tabulated for a
large number of weak acids.
Graphic
representations
of strong and
weak acid
equilibria
Strong Acid:
100% Dissociation
into ions
HA(aq) + H2O(l)  H3O+(aq) + A- (aq)
Equilibrium Favors undissociated acid

-
[H3O ] [A ]
Ka 
[HA]
Weak Acid:
Very little
Dissociation
into ions
Values of Ka for Some Common Monoprotic Acids
Name
Value of Ka*
HSO4HClO2
HC2H2ClO2
HF
HNO2
HC2H3O2
[Al(H2O)6]3+
HOCl
HCN
NH 4+
HOC6H5
Hydrogen sulfate ion
1.2 x 10-2
Chlorous acid
1.2 x 10-2
Monochloroacetic acid
1.35 x 10-3
Hydrofluoric acid
7.2 x 10-4
Nitrous acid
4.0 x 10-4
Acetic acid
1.8 x 10-5
Hydrated aluminum (III) ion 1.4 x10-5
Hypochlorous acid
3.5 x 10-8
Hydrocyanic acid
6.2 x 10-10
Ammonium ion
5.6 x 10-10
Phenol
1.6 x 10-10
*The units of Ka are mol/L, but are customarily omitted.
Increasing acid strength
Formula
Write the dissociation reaction and the
corresponding equilibrium expression for each of
the following acids in water.
a. H3PO4
b. H2PO41c. HCO31d. HCN
e. Glycine, H2NCH2COOH
f. Acetic acid, CH3COOH (HC2H3O2)
g. Phenol, C6H5OH
h. Benzoic acid, C6H5COOH
Write the reaction and the corresponding Kb
equilibrium expression for each of the following
substances acting as bases in water.
a. PO43g. Glycine, NH2CH2COOH
b. HPO42h. Ethylamine, CH3CH2NH2
c. H2PO4I. Aniline, C6H5NH2
d. NH3
j. Dimethylamine, (CH3)2NH
e. CNf. Pyridine, C5H5N
WEAK ACID CALCULATIONS
HA(aq) + H2O(l)  H3O+(aq) + A-(aq)
2H2O(l)  H3O+(aq) + OH-(aq)

[H3O ]
% Ionization 
x 100%
[HA]
To simplify calculations, if % ionization is < 5%,
then CHA  [HA] OR [HA] = CHA – [H3O+],
Set up pH equilibria calculations in tables as in
previous equilibria problems.
Solving Weak Acid Equilibrium Problems
1.
List the major species in the solution
2.
Choose the species that+ can produce H+, and write balanced equations for the
reactions producing H
3.
Using the values of the equilibrium constants for the reactions you
have
written, decide which equilibrium will dominate in producing H+
Write the equilibrium expression for the dominant equilibrium.
4.
5.
List the initial concentrations of the species participating in the dominant
equilibrium.
6.
Define the change needed to achieve equilibrium; that is, define x..
7.
Write the equilibrium concentrations in terms of x.
8.
Substitute the equilibrium concentrations into the equilibrium expression.
9.
Solve for x the “easy” way; that is , by assuming that [HA]0-x  [HA]0.
10. Use the 5% rule to verify whether the approximation is valid.
11. Calculate [H+] and pH.
1.
2.
3.
Ka Problems
For trichlorophenol (HC6H2Cl3O),
Ka = 1 x 10-6, Calculate the concentrations of
all species and the pH of a 0.05 M solution of
trichlorophenol in water.
A solution is prepared by dissolving 0.56 g
benzoic acid (C6H5CO2H), Ka = 6.4 x 10-5) in
enough water to make 1.0 L of solution.
Calculate [C6H5CO2H]. [C6H5CO2-], [H+], [OH], and the pH in this solution.
Calculate the pH of a solution containing a
mixture of 0.050 M HNO3 and 0.50M HC2H3O2.
WEAK BASES IONIZATION CONSTANTS
B(aq) + H2O(l)  BH+ (aq) + OH- (aq)

_
[BH ] [OH ]
Kb 
[B]
Kb values at 25oC are tabulated or may be
calculated from Kw and Ka
Kw = (Ka)(Kb) so Kb = Kw/Ka
Where Ka is the conjugate acid constant
1.
2.
Kb Problems
Thallium (Tl) hydroxide is a strong base used in
the synthesis of some organic compounds.
Calculate the pH of a solution containing 2.48 g
TlOH per liter.
For the reaction of hydrazine (N2H4) in water.
H2NNH2 + H2O  H2NNH3+ + OHKb is 3.0 x 10-6. Calculate the concentrations of
all species and the pH of a 2.0 M solution of
hydrazine in water.
Two Weak Acids in Solution
A solution of 0.100 M HClO, Ka =3.5 x 10-8
and 0.100 M Formic acid, HCO2H, Ka = 1.8 x
10-4 are mixed in equal proportions. What is
the resulting solution pH?
Polyprotic Acid Equilibria
What is the pH of a solution 0.100
M sulfurous acid, H2SO3,
Ka1 = 1.5 x 10-2 and Ka2 = 1.0 x 10-7?
ACID-BASE PROPERTIES OF THE OXIDES
(PART I)
ACID BASE PROPERTIES OF THE OXIDES
(PART II)
ACID-BASE PROPERTIES OF THE OXIDES
(PART III)
ACID BASE PROPERTIES OF THE OXIDES
(PART IV)
ACID-BASE PROPERTIES OF THE OXIDES
(PART V)
ACID BASE PROPERTIES OF THE OXIDES
PART (VI)
ACID BASE PROPERTIES OF THE OXIDES
PART (VII)
ACID-BASE PROPERTIES OF THE OXIDES
PART (VIII)`
BRONSTED-LOWRY
THEORY
OXIDES, HYDROXIDES, ANHYDRIDES
Acid, base reactions
Dehydrations (formation of anhydrides)
Hydration of oxides
ACID STRENGTHS
of WEAK ACIDS
Oxoacids:
Ka  up as the central atom
oxidation state  up.
Ka  up as central atom of same oxidation state
moves left to right in the periodic table.
Ka  up as the central atom of the same oxidation
state moves UP in the same Group or family.
Polyprotic acids:
Binary acids:
Ka decreases by approximately 105 for each
successive H+ ionized.
(only H and another element)
Within a period, Ka  up as electronegativity of
the other element.
Within a group, Ka  up going down the group
to higher mass and larger size.
SnO2 +
? H2O  ?
?HCrO4-  ?Cr2O72- + ?
?HMnO4-  ? Mn (VI) compound + ?
HYDROLOYSIS OF IONIC SALTS
The pH of each type salt in
solution depends on the Ka or Kb of
the hydrolyzing ion(s).
Salt Derived
From:
Ions Undergoing
Hydrolysis
Strong base
strong acid
pH
Examples
Neither
Neutral
pH = 7
NaCl, KNO3,
BaCl2, CaBr2
Strong base
weak acid
Anion
Basic
pH > 7
LiCN, KNO2, CaF2
NaCH3CO2
Weak base
strong acid
Cation
Acidic,
pH < 7
NH4Cl, Al(NO3)3,
(CH3)3NHBr
Weak base,
weak acid
Both
Acidic
if Kb< Ka;
Neutral
if Kb = Ka;
basic
if Kb > Ka
NH4NO2
NH4CH3CO2
NH4CN
Arrange the following 0.10 M solutions in order
from most acidic to most basic
KOH, KBr, KCN, NH4Br, NH4CN, HCN
Given that the Ka value for acetic acid is 1.8 x 10-5
and the Ka value for hypochlorous acid is 3 x 10-8,
which is the stronger base, OCl- or C2H3O2-
Acid-Base Equilibria
What is the pH of a solution of 0.150 M sodium
nitrite, NaNO2?
HNO2, Ka = 4.0 x 10-4.
What is the pH of a solution of 0.150 M
hydrazinnium chloride, H2NNH3+?
H2NNH3+, Kb = 3.0 x 10-6
Calculate the pH of each of the following solutions.
a. 0.10 M CH3NH3Cl
b. 0.050 M NaCN
c. 0.20 M Na2CO3 (consider only the reaction )
CO32- + H2O  HCO3- + OHSodium azid (NaN3) is sometimes added to water to
kill bacteria. Calculate the concentration of all
species in a 0.010 M solution of NaN3. The Ka
value for hydrazoic acid (HN3) is 1.9 x 10-5.
LEWIS THEORY
ACIDS and BASES
Lewis base - an electron pair donor
Lewis acid - an electron pair acceptor
Lewis acid + Lewis base  Adduct
(coordination compound)
e.g. Cu2+(aq) + 4CN-(aq)  Cu(CN)42-(aq)
Look at Lewis Dot structures for lone pairs.