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Slide 1
Chapter 14Chapter 14
•Aqueous Equilibria: Acids and Bases
•Aqueous Equilibria: Acids and Bases
Slide 2
Acid–Base Concepts 01Acid–Base Concepts 01
Arrhenius Acid: A substance which dissociates to form hydrogen ions (H+) in solution.
HA(aq) H+(aq) + A–(aq)
Arrhenius Base: A substance that dissociates in, or reacts with water to form hydroxide ions (OH–).
MOH(aq) M+(aq) + OH–(aq)
Slide 3
Acid–Base Concepts 02Acid–Base Concepts 02
• Brønsted–Lowry Acid: Substance that can donate H+
• Brønsted–Lowry Base: Substance that can accept H+
• Chemical species whose formulas differ only by one proton are said to be conjugate acid–base pairs.
Slide 4
Strong vs. Weak acids 03Strong vs. Weak acids 03
Slide 5
Hydrated Protons and Hydronium IonsHydrated Protons and Hydronium Ions
H1+(aq) + A1-(aq)HA(aq)
[H(H2O)n]1+
For our purposes, H1+ is equivalent to H3O1+.
n = 4 H9O41+
n = 1 H3O1+
n = 2 H5O21+
n = 3 H7O31+
Due to high reactivity of the hydrogen ion, it is actually hydrated by one or more water molecules.
Slide 6
Acid–Base Concepts Acid–Base Concepts
Slide 7
Lewis Acid–Base ConceptsLewis Acid–Base Concepts
Slide 8
• A Lewis Acid is an electron-pair acceptor. These are generally cations and neutral molecules with vacant valence orbitals, such as Al3+, Cu2+, H+, BF3.
• A Lewis Base is an electron-pair donor. These are generally anions and neutral molecules with available pairs of electrons, such as H2O, NH3, O2–.
• The bond formed is called a coordinate bond.
Acid–Base Concepts 05Acid–Base Concepts 05
Slide 9
Acid–Base Concepts 06Acid–Base Concepts 06
- +
Slide 10
Lewis Acids and BasesLewis Acids and Bases
Lewis Base: An electron-pair donor.
Lewis Acid: An electron-pair acceptor.
Slide 11
Lewis Acids and BasesLewis Acids and Bases
Lewis Base: An electron-pair donor.
Lewis Acid: An electron-pair acceptor.
Slide 12
Acid–Base Concepts 07Acid–Base Concepts 07
• Write balanced equations for the dissociation of each of the following Brønsted–Lowry acids.(a) H2SO4 (b) HSO4
– (c) H3O+
• Identify the Lewis acid and Lewis base in each of the following reactions:
(a) SnCl4(s) + 2 Cl–(aq) æ SnCl62–(aq)
(b) Hg2+(aq) + 4 CN–(aq) æ Hg(CN)42–(aq)
(c) Co3+(aq) + 6 NH3(aq) æ Co(NH3)63+(aq)
Slide 13
Dissociation of Water 01Dissociation of Water 01
• Water can act as an acid or as a base.
H2O(l) æ H+(aq) + OH–(aq)
• Kc = [H+][OH–]
• This is called the autoionization of water.
H2O(l) + H2O(l) æ H3O+(aq) + OH–(aq)
Slide 14
Dissociation of Water 02Dissociation of Water 02
• This equilibrium gives us the ion product constant for water.
Kw = Kc = [H+][OH–] = 1.0 x 10–14
• If we know either [H+] or [OH–] then we can
determine the other quantity.
Slide 15
Dissociation of Water 03Dissociation of Water 03
• The concentration of OH– ions in a certain household
ammonia cleaning solution is 0.0025 M. Calculate the
concentration of H+ ions.
• Calculate the concentration of OH– ions in a HCl
solution whose hydrogen ion concentration is 1.3 M.
Slide 16
pH – A Measure of Acidity 01pH – A Measure of Acidity 01
• The pH of a solution is the negative logarithm of the
hydrogen ion concentration (in mol/L).
pH = –log [H+], [H+] = 10-pH
pH + pOH = 14
Acidic solutions:[H+] > 1.0 x 10–7 M, pH < 7.00
Basic solutions: [H+] < 1.0 x 10–7 M, pH > 7.00
Neutral solutions: [H+] = 1.0 x 10–7 M, pH = 7.00
Slide 17
pH – A Measure of Acidity 02pH – A Measure of Acidity 02
• Nitric acid (HNO3) is used in the production of fertilizer, dyes, drugs, and explosives. Calculate the pH of a HNO3 solution having a hydrogen ion concentration of 0.76 M.
• The pH of a certain orange juice is 3.33. Calculate the H+ ion concentration.
• The OH– ion concentration of a blood sample is 2.5 x 10–
7 M. What is the pH of the blood?
Slide 18
pH – A Measure of Acidity 04pH – A Measure of Acidity 04
Color of Tea: Polyphenols, Thearubigins
Color of Red Cabbage: Anthocyanin
Slide 19
pH – A Measure of Acidity 04pH – A Measure of Acidity 04
Slide 20
Strength of Acids and Bases03Strength of Acids and Bases03
HClO4
HI
HBr
HCl
H2SO4
HNO3
H3O+
HSO4–
HSO4–
HF
HNO2
HCOOH
NH4+
HCN
H2O
NH3
ClO4–
I–
Br –
Cl –
HSO4 –
NO3 –
H2O
SO42–
SO42–
F –
NO2 –
HCOO –
NH3
CN –
OH –
NH2 –
ACID CONJ. BASE ACID CONJ. BASE
Incr
easi
ng A
cid
Str
engt
h
Incr
easi
ng A
cid
Str
engt
h
Slide 21
Strength of Acids and Bases04Strength of Acids and Bases04
• Stronger acid + stronger base
weaker acid + weaker base
• Predict the direction of the following:
HNO2(aq) + CN–(aq) æ HCN(aq) + NO2–(aq)
HF(aq) + NH3(aq) æ F–(aq) + NH4+(aq)
Slide 22
Acid Ionization Constants 01Acid Ionization Constants 01
• Acid Ionization Constant: the equilibrium constant for the ionization of an acid.
HA(aq) + H2O(l) æ H3O+(aq) + A–(aq)
• Or simply: HA(aq) æ H+(aq) + A–(aq)
[HA]]][A[H
aK
Slide 23
Conjugate Base Ionization ConstConjugate Base Ionization Const
[HA] [OH−][A-]
Kb =
A- + H2O(l) HA(aq) + OH−(aq)
Ka Kb = Kw
[HA] [OH−][A-]
Kb =Ka [HA]
]][A[H
= Kw
Slide 24
Acid Ionization Constants 02Acid Ionization Constants 02
7.1 x 10 –4
4.5 x 10 –4
3.0 x 10 –4
1.7 x 10 –4
8.0 x 10 –5
6.5 x 10 –5
1.8 x 10 –5
4.9 x 10 –10
1.3 x 10 –10
HF
HNO2
C9H8O4 (aspirin)
HCO2H (formic)
C6H8O6 (ascorbic)
C6H5CO2H (benzoic)
CH3CO2H (acetic)
HCN
C6H5OH (phenol)
F–
NO2 –
C9H7O4 –
HCO2 –
C6H7O6 –
C6H5CO2 –
CH3CO2 –
CN –
C6H5O –
ACID Ka CONJ. BASE Kb
1.4 x 10 –11
2.2 x 10 –11
3.3 x 10 –11
5.9 x 10 –11
1.3 x 10 –10
1.5 x 10 –10
5.6 x 10 –10
2.0 x 10 –5
7.7 x 10 –5
Slide 25
Strength of Acids and Bases03Strength of Acids and Bases03
(a) Arrange the three acids in order of increasing value of Ka.
(b) Which acid, if any, is a strong acid?(c) Which solution has the highest pH, and which has the
lowest?
(42/2) = 8 12/5= 0.2 Very LargeK =
Slide 26
HA æ H+ + A
(M): 0.50 0.00 0.00 (M): –x +x +x
Equilib (M): 0.50 –x x x
Acid Ionization Constants Determine the pH of 0.50 M HA solution at 25°C. Ka = 7.1 x 10–4 05
Acid Ionization Constants Determine the pH of 0.50 M HA solution at 25°C. Ka = 7.1 x 10–4 05
• Initial Change Equilibrium Table:.
InitialChange
(aq) (aq)-(aq)
Slide 27
What is the pH of a 0.50 M Citric acid solution (at 250C)?
HA (aq) H+ (aq) + A- (aq) Ka =[H+][A-][HA]
= 7.1 x 10-4
HA (aq) H+ (aq) + A- (aq)
Initial (M)
Change (M)
Equilibrium (M)
0.50 0.00
-x +x
0.50 - x
0.00
+x
x x
Ka =x2
0.50 - x= 7.1 x 10-4
Ka x2
0.50= 7.1 x 10-4
0.50 – x 0.50100•Ka < Co ?100 x 7.1 x 10-4
= 0.071 < 0.5x2 = 3.55 x 10-4 x = 0.019 M
[H+] = [A-] = 0.019 M pH = -log [H+] = 1.72
[HA] = 0.50 – x = 0.48 M
Slide 28
Acid Ionization Constants 06Acid Ionization Constants 06
• pH of a Weak Acid (Cont’d):
1. Substitute equilibrium concentrations into equilibrium
expression.
2. If 100•Ka < Co then (C0 – x) approximates to (C0).
3. The equation can now be solved for x and pH.
4. If 100•Ka is not significantly smaller than Co the
quadratic equation must be used to solve for x and pH.
Slide 29
Acid Ionization Constants 07Acid Ionization Constants 07
• The Quadratic Equation:
• The expression must first be rearranged to:
• The values are substituted into the quadratic and
solved for a positive solution to x and pH.
aacbb
x2
42
02 cbxax
Slide 30
Acid Ionization Constants 09Acid Ionization Constants 09
• Percent Dissociation: A measure of the strength of an acid.
• Stronger acids have higher percent dissociation.
• Percent dissociation of a weak acid decreases as
its concentration increases.
100[HA]
][HonDissociati %
H1+(aq) + A1-(aq)HA(aq)
Slide 31
Percent dissociation of a weak acid decreases as its concentration increases
Percent dissociation of a weak acid decreases as its concentration increases
• Concentration Dependence:
Slide 32
Weak Bases: Base Ionization Constants 01Weak Bases: Base Ionization Constants 01
• Base Ionization Constant: The equilibrium constant for the ionization of a base.
• The ionization of weak bases is treated in the same
way as the ionization of weak acids.
B(aq) + H2O(l) æ BH+(aq) + OH–(aq)
• Calculations follow the same procedure as used for
a weak acid but [OH–] is calculated, not [H+].
Slide 33
Base Ionization Constants 02Base Ionization Constants 02
5.6 x 10 –4
4.4 x 10 –4
4.1 x 10 –4
1.8 x 10 –5
1.7 x 10 –9
3.8 x 10 –10
1.5 x 10 –14
C2H5NH2 (ethylamine)
CH3NH2 (methylamine)
C8H10N4O2 (caffeine)
NH3 (ammonia)
C5H5N (pyridine)
C6H5NH2 (aniline)
NH2CONH2 (urea)
C2H5NH3+
CH3NH3+
C8H11N4O2+
NH4+
C5H6N+
C6H5NH3+
NH2CONH3+
BASE Kb CONJ. ACID Ka
1.8 x 10 –11
2.3 x 10 –11
2.4 x 10 –11
5.6 x 10 –10
5.9 x 10 –6
2.6 x 10 –5
0.67
Note that the positive charge sits on the nitrogen.(caffeine)
Slide 34
Base Ionization Constants 03Base Ionization Constants 03
• Product of Ka and Kb: multiplying out the
expressions for Ka and Kb equals Kw.
Ka Kb = Kw
Slide 35
pH of Basic SolutionspH of Basic Solutions
What is the pH of a 0.15 M solution of NH3?
[NH4+] [OH−]
[NH3]Kb = = 1.8 10−5
NH3(aq) + H2O(l) NH4+(aq) + OH−(aq)
[NH3], M [NH4+], M [OH−], M
Initially 0.15 0 0
At Equilibrium 0.15 - x x x
Slide 36
pH of Basic SolutionspH of Basic Solutions
(1.8 10−5) (0.15) = x2
2.7 10−6 = x2
1.6 10−3 = x2
(x)2
(0.15 - x )1.8 10−5 =
100 x Kb < C0 ?
1.8 10−3< 0.150.15 –x = 0.15
Slide 37
pH of Basic SolutionspH of Basic Solutions
Therefore,X = [OH−] = 1.6 10−3 MpOH = −log (1.6 10−3)pOH = 2.80pH = 14.00 − 2.80pH = 11.20
Slide 42
Diprotic & Polyprotic Acids 01Diprotic & Polyprotic Acids 01
• Diprotic and polyprotic acids yield more than one hydrogen
ion per molecule.
• One proton is lost at a time. Conjugate base of first step is
acid of second step.
• Ionization constants decrease as protons are removed.
H2SO4
H3PO4
Slide 43
Diprotic & Polyprotic Acids 02Diprotic & Polyprotic Acids 02
Very Large1.3 x 10 –2
6.5 x 10 –2
6.1 x 10 –5
1.3 x 10 –2
6.3 x 10 –8
4.2 x 10 –7
4.8 x 10 –11
9.5 x 10 –8
1 x 10 –19
7.5 x 10 –3
6.2 x 10 –8
4.8 x 10 –13
H2SO4
HSO4–
C2H2O4
C2HO4–
H2SO3
HSO3–
H2CO3
HCO3–
H2SHS–
H3PO4
H2PO4–
HPO42–
ACID Ka CONJ. BASE Kb
HSO4 –
SO4 2–
C2HO4–
C2O42–
HSO3 –
SO3 2–
HCO3–
CO3 2–
HS–
S 2–
H2PO4–
HPO42–
PO43–
Very Small7.7 x 10 –13
1.5 x 10 –13
1.6 x 10 –10
7.7 x 10 –13
1.6 x 10 –7
2.4 x 10 –8
2.1 x 10 –4
1.1 x 10 –7
1 x 10 –5
1.3 x 10 –12
1.6 x 10 –7
2.1 x 10 –2
Slide 44
Molecular Structure and Acid Strength 01Molecular Structure and Acid Strength 01
• The strength of an acid depends on its tendency to
ionize.
• For general acids of the type H–X:
1. The stronger the bond, the weaker the acid.
2. The more polar the bond, the stronger the acid.
• For the hydrohalic acids, bond strength plays the
key role giving: HF < HCl < HBr < HI299 kJ/mol for HI567 kJ/mol for HF
Slide 45
Molecular Structure and Acid Strength 02Molecular Structure and Acid Strength 02
• The electrostatic potential maps show all the hydrohalic
acids are polar. The variation in polarity is less
significant than the bond strength which decreases
from 567 kJ/mol for HF to 299 kJ/mol for HI.
Slide 46
(pm)
Slide 47
Molecular Structure and Acid Strength 03Molecular Structure and Acid Strength 03
• For binary acids in the same group, H–A bond strength decreases with increasing size of A, so acidity increases.
• For binary acids in the same row, H–A polarity increases with increasing electronegativity of A, so acidity increases.
Slide 48
Molecular Structure and Acid Strength 04Molecular Structure and Acid Strength 04
• For oxoacids bond polarity is more important. If we consider the main element (Y):
Y–O–H
• If Y is an electronegative element, the Y–O bond will
pull more electrons, the O–H bond will be more polar
and the acid will be stronger.
Slide 49
Molecular Structure and Acid Strength 05Molecular Structure and Acid Strength 05
• For oxoacids with different central atoms that are from the same group of the periodic table and that have the same oxidation number, acid strength increases with increasing electronegativity.
Slide 50
Polar Covalent Bonds 02Polar Covalent Bonds 02
Pauling ElectronegativitiesPauling Electronegativities
Detailed List of Electronegativity; http://environmentalchemistry.com/yogi/periodic/electronegativity.html
Slide 51
Molecular Structure and Acid Strength 07Molecular Structure and Acid Strength 07
• Oxoacids of Chlorine:
Slide 52
Molecular Structure and Acid Strength 08Molecular Structure and Acid Strength 08
• Predict the relative strengths of the following groups of oxoacids:
a) HClO, HBrO, and HIO.
b) HNO3 and HNO2.
c) H3PO3 and H3PO4.
Slide 53
Acid-Base Properties of SaltsAcid-Base Properties of Salts
Slide 54
Strong basesStrong bases
• Strong bases:• The following metals make strong hydroxy base
• Alkali metal cations of group 1A • Alkaline earth metal cations of group 2A
except for Be
Slide 55
Acid–Base Properties of Salts 01Acid–Base Properties of Salts 01
• Salts that produce neutral solutions are those
formed from strong acids and strong bases.
• Salts that produce basic solutions are those formed
from weak acids and strong bases.
• Salts that produce acidic solutions are those
formed from strong acids and weak bases.
Slide 56
The pH of an ammonium carbonate solution, (NH4)2CO3, depends on the relative acid strength of the cation and the relative base strength of the anion.
Is it acidic or basic?
Salts That Contain Cation from a Weak Base and anion from a Weak Base
Slide 57
Acid-Base Properties of SaltsAcid-Base Properties of Salts
Salts That Contain Acidic Cations and Basic Anions
HCO31-(aq) + OH1-(aq)CO3
2-(aq) + H2O(l) Kb
H3O1+(aq) + NH3(aq)NH41+(aq) + H2O(l) Ka
(NH4)2CO3:
Three possibilities:• Ka > Kb: The solution will contain an excess of
H3O1+ ions , Acidic solution, (pH < 7).• Ka < Kb: The solution will contain an excess of
OH1- ions, Basic solutions, (pH > 7).• Ka ≈ Kb: The solution will contain approximately
equal concentrations of H3O1+ and OH1- ions (pH ≈ 7).
Slide 58
Salts That Contain Cation from a Weak Acid and anion from a Weak Base
HCO31-(aq) + OH1-(aq)CO3
2-(aq) + H2O(l) Kb
H3O1+(aq) + NH3(aq)NH41+(aq) + H2O(l) Ka
(NH4)2CO3:
= 1.8 x 10-4
5.6 x 10-11
1.0 x 10-14
Kb for CO32- =
Ka for HCO31-
Kw
=
= 5.6 x 10-10
1.8 x 10-5
1.0 x 10-14
Ka for NH41+ =
Kb for NH3
Kw
=
Basic, Ka < Kb
Slide 59
Acid-Base Properties of SaltsAcid-Base Properties of Salts
Slide 60
Hydrated Cation of Al3+
Slide 61
Acid–Base Properties of Salts 03Acid–Base Properties of Salts 03
• Metal Ion Hydrolysis:
Slide 62
Acid–Base Properties of Salts 04Acid–Base Properties of Salts 04
• Calculate the pH of a 0.020 M Al(NO3)3 solution
Ka = 1.4 x 10-5.
• Predict whether the following solutions will be
acidic, basic, or nearly neutral:
(a) NH4I (b) CaCl2 (c) KCN (d) Fe(NO3)3