Prentice Hall ©2004 Chapter 14 Aqueous Equilibria: Acids and Bases.
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Transcript of Prentice Hall ©2004 Chapter 14 Aqueous Equilibria: Acids and Bases.
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Chapter 14Chapter 14
Aqueous Equilibria: Acids and Bases
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Acid–Base Concepts 01Acid–Base Concepts 01
Arrhenius Acid: A substance which dissociates in water to form hydrogen ions (H+) in solution.
HA(aq) + H2O(l) H3O+(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)
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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.
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Acid–Base Concepts 03Acid–Base Concepts 03
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Acid–Base Concepts 04Acid–Base Concepts 04
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• 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
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Acid–Base Concepts 06Acid–Base Concepts 06
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Acid–Base Concepts 07Acid–Base Concepts 07
• Problems 14.1,14.2• Write balanced equations for the dissociation of each of the following Brønsted–Lowry acids.(a) H2SO4 (b) HSO4
– (c) H3O+
• Problems 14.27• Identify the Lewis acid and Lewis base in each of the following reactions:
(a) AlCl3(s) + Cl–(aq) æ AlCl4–(aq)
(b) SO2 (aq) + OH–(aq) æ HSO3–(aq)
(c) Ag+(aq) + 2 NH3(aq) æ Ag(NH3)2+(aq)
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Dissociation of Water 01Dissociation of Water 01
• Water can act as an acid or as a base.
H2O(l) æ H+(aq) + OH–(aq)
• This is called the autoionization of water.
H2O(l) + H2O(l) æ H3O+(aq) + OH–(aq)
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Dissociation of Water 02Dissociation of Water 02
• This equilibrium gives us the ion product constant for water.
Kw = Kc = [H3O+][OH–] = 1.0 x 10–14
If we know either [H3O+] or [OH–] then we can
determine the other quantity.
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Dissociation of Water 03Dissociation of Water 03
• Problem 14.6 The concentration of OH– in a sample of seawater is 5.0 × 10-6 M. Calculate the concentration of H3O+ ions, and classify the solution as acidic, neutral, or basic.
• Problem 14.7 At 50°C the value of Kw is 5.5 × 10-14.
What are the concentrations of H3O+ and OH– in a
neutral solution at 50°C?
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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 [H3O+]
pOH = –log [OH– ]
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
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pH – A Measure of Acidity 02pH – A Measure of Acidity 02
• Problem 14.8 Calculate the pH of each of the following
Problem 14.9 Calculate the concentrations of H3O+ and OH– in each of the following solutions:
(a) Human blood (pH 7.40)(b) A cola beverage (pH 2.8)
Problem 14.10 Calculate the pH of(a)0.050 M HClO4 (b) 6.0 M HCl (c) 4.0 M KOH (d) 0.010 M Ba(OH)2
Problem 14.11 Calculate the pH of a solution prepared by dissolving 0.25 g of BaO in enough water to make 0.500 L of
solution
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pH – A Measure of Acidity 03pH – A Measure of Acidity 03
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Strength of Acids and Bases01Strength of Acids and Bases01
• Strong acids and bases: are strong electrolytes
that are assumed to ionize completely in water.
• Weak acids and bases: are weak electrolytes that
ionize only to a limited extent in water.
• Solutions of weak acids and bases contain ionized
and non-ionized species.
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Strength of Acids and Bases02Strength of Acids and Bases02
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
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Strength of Acids and Bases03Strength of Acids and Bases03
• Stronger acid + stronger base
weaker acid + weaker base
• Problems 14.4 Predict the direction of the following:
HF(aq) + NO3– (aq) æ F–(aq) + HNO3(aq)
NH4+(aq) + CO3
–2 (aq) æ HCO3– (aq) + NH3(aq)
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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)
[HA]
]][A[H3O
aK
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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
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CalculatingEquilibriumConcentrationinSolutions ofWeak Acids
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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 04Acid Ionization Constants 04
• Initial Change Equilibrium Table: Determine the pH
of 0.50 M HA solution at 25°C. Ka = 7.1 x 10–4.
InitialChange
(aq) (aq)-(aq)
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Acid Ionization Constants 05Acid Ionization Constants 05
• pH of a Weak Acid (Cont’d):
1. Substitute new values into equilibrium expression.
2. If Ka is significantly (>1000 x) smaller than [HA] the expression
(0.50 – x) approximates to (0.50).
3. The equation can now be solved for x and pH.
4. If Ka is not significantly smaller than [HA] the quadratic
equation must be used to solve for x and pH.
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Acid Ionization Constants 06Acid Ionization Constants 06
• 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
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Acid Ionization Constants 08Acid Ionization Constants 08
• 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]
][H3OonDissociati %
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Base Ionization Constants 01Base 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+].
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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.
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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.
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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
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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 < HI
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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.
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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.
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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, or in a high
oxidation state, the Y–O bond will be more covalent
and the O–H bond more polar and the acid stronger.
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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.
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Molecular Structure and Acid Strength 06Molecular Structure and Acid Strength 06
• For oxoacids having the same central atom but different numbers of attached groups, acid strength increases with increasing central atom oxidation number.
• As shown on the next slide, the number of oxygen atoms increases the positive charge on the chlorine which weakens the O–H bond and increases its polarity.
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Molecular Structure and Acid Strength 07Molecular Structure and Acid Strength 07
• Oxoacids of Chlorine:
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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.
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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.