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GenChem Ch 162013/03/03 TMHsiung 1/60
Chapter 16
Acids and Bases
GenChem Ch 162013/03/03 TMHsiung 2/60
Contents in Chapter 16
16-1 Arrhenius Theory: A Brief Review16-2 Brønsted–Lowry Theory of Acids and Bases16-3 Self-Ionization of Water and the pH Scale16-4 Strong Acids and Strong Bases16-5 Weak Acids and Weak Bases16-6 Polyprotic Acids16-7 Ions as Acids and Bases16-8 Molecular Structure and Acid–Base Behavior16-9 Lewis Acids and Bases
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HCl(g) → H+(aq) + Cl-(aq)
NaOH(s) → Na+(aq) + OH-(aq)H2O
H2O
Na+(aq) + OH-(aq) + H+(aq) + Cl-(aq) → H2O(l) + Na+(aq) + Cl-(aq)
H+(aq) + OH-(aq) → H2O(l) • Arrhenius theory did not handle non OH– bases such as
ammonia (NH3).• Neutralization reaction: Combination of hydrogen ions
(H+) and hydroxide ions (OH–) to form water.
16-1 Arrhenius Theory: A Brief Review Acid: A substance that provides H+ in aqueous solution,
e.g., HCl Base (alkalis): A substance that provides OH– in
aqueous solution, e.g., NaOH
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16-2 Brønsted–Lowry Theory of Acids and Bases*** Definition:
Acid: the substance act as H+ donor Base: the substance act as H+ acceptor
Conjugatebase of NH4
+
ConjugateAcid of NH3
Conjugatebase of H2O
Conjugateacid of OH−
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CH3COOH + H2O CH3COO– + H3O+
COOH][CH]COO][CHO[H
K3
33a
Conjugate acid
Conjugate acidConjugate base
Conjugate base
NH3 + H2O NH4+ + OH–
][NH
]][OH[NHK
3
4b
Conjugate base
Conjugate baseConjugate acid
Conjugate acid
Ionization constants*** Acid ionization constant
Base ionization constant
*H2O is an amphiprotic (amphoteric) substance,act as either an acid or a base.
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Strengths of conjugate acid-base pairs The stronger an acid, the weaker its conjugate
base. (The stronger a base, the weaker its conjugate acid.)
An acid-base reaction is favored in the direction from the stronger member to the weaker member of each conjugate acid-base pair.
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Brønsted–Lowry acid–base reaction: weak acid
• CH3COOH is only slightly ionized.• Reverse reaction proceeds to a greater extent than
does the forward reaction.• H3O+ is a stronger acid than CH3COOH and
CH3COO− is a stronger base than H2O.
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Brønsted–Lowry acid–base reaction: strong acid
• HCl is essentially completely ionized.• The forward reaction proceeds almost to completion.• H3O+ is a weaker acid than HCl and Cl− is a much
weaker base than H2O.
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More about strengths of acid and bases
Ka (Kb) values are used to compare the strengths of weak acids (bases).
Leveling (solvent) effect: The solvent's ability to level the effect of a strong acid (or strong base) dissolved in it. e.g., HI and HBr are leveled to the same acidic strength in H2O.
Differentiating (solvent) effect: The solvent's ability to differentiate the acidic (or basic) strength.Example:
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16-3 Ionization of Water and the pH Scale Self ionization of water
Ion-product of water, KW:KW = [H3O+][OH–]At 25oC, KW = 1x10–14
KW applies to all aqueous solutions—acids, bases, salts, and nonelectrolytes—not just to pure water.
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p function: –log pH = –log [H3O+] [H3O+] = 10–pH
pOH = –log [OH–] [OH–] = 10–pOH
Acidic solution: [H3O+] > [OH–] pH < pOHBasic solution: [H3O+] < [OH–] pH > pOH
Aqueous solution at 25oC:pKW = pH + pOH = 14.00
Aqueous solution at 25oC:acidic solution: [H3O+] > 1.0×10–7 pH < 7.00basic solution: [H3O+] < 1.0×10–7 pH > 7.00neutral solution: [H3O+] = 1.0×10–7 pH = 7.00
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The pH scale and pH values of some common materials
• pH = −1 ([H3O+] ≈ 10 M) and pH = 15 ([OH−] ≈ 10 M) are possible.
• The pH scale is useful only in the range 2 < pH < 12, because the molarities of H3O+ and OH− in concentrated acids and bases may differ from their true activities.
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16-4 Strong Acids and Strong Bases The contribution due to the self-ionization of water
can generally be ignored (unless the solution is extremely dilute), i.e., for strong acids and bases, dissociated completely. Therefore,[H3O+] CHCl CHCl: initial concentration of HCl[OH–] CNaOH CNaOH: initial concentration of NaOH.
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For extremely dilute solution of a strong acid and strong base, 1.0 x 10–8 M HCl for example:
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16-5 Weak Acids and Weak Bases Identifying Weak Acids and Bases
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Percent Ionization (A weak acid HA for example)
HA + H2O H3O+ + A-
Degree of ionization =[H3O+] from HA
[HA] originally
Percent ionization =[H3O+] from HA
[HA] originally 100%
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KH O A
HAa3
+
[ ][ ]
[ ]
]O[HC
]O[HK
+3HA
2+3
a
Equilibrium of monoprotic acid, HA
HA + H2O H3O+ + A–
CHA – x x x
Therefore
Assume CHA – x CHA
HA
HA a
x C100 5% about 100
C K
* CHA – x CHA, using:a
acbbx
242
5% rule:
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Equilibrium of monobasic base, B
Assume CB – x CB
B
b b
x C100 5% about 100
C K
* CB – x CB, using:a
acbbx
242
5% rule:][OHC
][OHK
B
2
b
B + H2O HB+ + OH–
CB – x x x
Therefore
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16-6 Polyprotic Acids
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Diprotic acid
A][H
]][HAO[HK
2
+3
a1
][HA
]][AO[HK
2+3
a2
][A
]][HA[OHK 2b1
][HA
A]][H[OHK 2
b2
H2A + H2O HA– + H3O+
HA– + H2O A2– + H3O+
For conjugate base:A2– + H2O HA– + OH–
HA– + H2O H2A + OH–
Ka1 x Kb2 = Kw Ka2 x Kb1 = Kw
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Triprotic acid
Ka1 x Kb3 = Kw Ka2 x Kb2 = Kw Ka3 x Kb1 = Kw
Phosphoric acid for example:
H3PO4 + H2O H2PO4– + H3O+
H2PO4– + H2O HPO4
2– + H3O+
HPO42– + H2O PO4
3– + H3O+
3
43
42+
3a1 101.7
]PO[H
]PO][HO[HK
8
42
24
+3
a2 103.6]PO[H
]][HPOO[HK
+ 3133 4
a3 24
[H O ][PO ]K 4.2 10
[HPO ]
Ionization constants for polyprotic acid progressively decrease:Ka1 > Ka2 > Ka3 > …..
Except in very dilute solutions, essentially all of the H3O+ ions come from the first ionization step alone.
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For a 3.0 M H3PO4, calculate (a) [H3O+]; (b) [H2PO4–];
(c) [HPO42–]; (d) [PO4
3–].
Solve
(a) assume that all H3O+ the forms in the first ionization step
H3PO4 + H2O H2PO4– + H3O+
Initial conc. 3.0 - - Change –x + x + x Equilibrium (3.0 – x) x x
x2 = 0.021 x = [H3O+] = 0.14 M Check: (x/CH3PO4
) x 100% = 4.7% < 5%, OK!!
EXAMPLE 16-9 Calculating Ion Concentrations in a Polyprotic Acid Solution
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(b) H3PO4 + H2O H2PO4– + H3O+
Equilibrium (3.0 – x) x x
[H2PO4–] [H3O+] = 0.14 M
+ 283 4
a22 4
[H O ][HPO ]K 6.3 10
[H PO ]
[HPO42–] = 6.3 x 10–8 M
(c) H2PO4– + H2O HPO4
2– + H3O+
+ 283 4
a22 4
[H O ][HPO ]K 6.3 10
[H PO ]
Since [H2PO4–] [H3O+], therefore
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Since [H3O+] = 0.14 M and [HPO42–] = 6.3 x 10–8 M
+ 3 3133 4 4
a3 2 84
[H O ][PO ] 0.14 [PO ]K 4.3 10
[HPO ] 6.3 10
(d) HPO42– + H2O PO4
3– + H3O+
[PO43–] = 1.9x 10–19 M
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2
4
24
+3
a2 101.1][HSO
]][SOO[HK
(1)H2SO4 + H2O HSO4– + H3O+ Ka1 103
CH2SO4 CH2SO4 CH2SO4
(2)HSO4– + H2O SO4
2– + H3O+
CH2SO4 - x x x
Concentrated solutions (>0.5 M H2SO4): H3O+ is predominated by first ionization step. e.g., 1.00 M H2SO4, [H3O+] 1.00 M.
Very dilute solutions (< 0.001 M H2SO4): both ionization steps are nearly completely dissociated, e.g., 0.001 M H2SO4, [H3O+] 0.002 M, [SO4
2–] 0.001 M. Intermediate concentrations (0.001 M< CH2SO4 <0.5 M), first
ionization step is completely dissociated, the second ionization step is partially dissociated.
A Somewhat Different Case: H2SO4
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Check: (x/CH2SO4) x 100% = 2.2% < 5%, OK!!
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16-7 Ions as Acids and Bases Hydrolysis: The reaction between an ion and water.
(1) NaCl(aq) Na+(aq) + Cl–
(aq) Neutral
Na+(aq) + H2O
Cl–(aq) + H2O
(2) NH4Cl(aq) NH4+
(aq) + Cl–(aq) Acidic
NH4+
(aq) + H2O NH3(aq) + H3O+(aq)
Cl–(aq) + H2O
(3) CH3COONa(aq) Na+(aq) + CH3COO–
(aq) Basic
Na+(aq) + H2O
CH3COO–(aq) + H2O CH3COOH(aq) + OH–
(aq)
(4) CH3COONH4(aq) NH4+
(aq) + CH3COO– (aq) ?????
NH4+
(aq) + H2O NH3(aq) + H3O+(aq)
CH3COO–(aq) + H2O CH3COOH(aq) + OH–
(aq)
XX
X
X
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Type of Salt Example
Ion That
Undergo
Hydrolysis
pH of
Solution
Cation from strong base/
Anion from strong acid
NaCl, KI, KNO3,
RbBr, BaCl2
None 7
Cation from strong base/
Anion from weak acid
CH3COONa,
KNO2, NaOCl
Anion >7
Cation from weak base/
Anion from strong acid
NH4Cl, NH4NO3 Cation <7
Cation from weak base/
Anion from weak acid
NH4NO2, NH4CN,
CH3COONH4
Anion and
Cation
<7 if Kb<Ka
7 if KbK
a
>7 if Kb>Ka
Cation is small, highly charged/
Anion from strong acid
AlCl3, Fe(NO3)3 Hydrated
Cation
<7
The pH of Salt Solutions
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(a) NaOCl(aq) Na+(aq) + OCl–
(aq) Basic
Na+(aq) + H2O
OCl–(aq) + H2O HOCl(aq) + OH–
(aq)
X
(b) KCl(aq) Na+(aq) + Cl–
(aq) Neutral
K+(aq) + H2O
Cl–(aq) + H2O
XX
(c) NH4NO3(aq) NH4+
(aq) + NO3–(aq) Acidic
NH4+
(aq) + H2O NH3(aq) + H3O+(aq)
NO3–(aq) + H2O X
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16-8 Molecular Structure and Acid–Base Behavior Strengths of Binary Acids Homolytic dissociation vs. Heterolytic dissociation• Homolytic dissociation
HX H + X D(H–X)• Heterolytic dissociation
HX H+ + X– D(H+X– ) Bond dissociation energy for the gas phase ionization
reaction
D(H+X– ) = D(H–X) + IE(H) + ΔHea
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• Bond dissociation energies (kJ/mol ) and Ka values for some binary acids
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Comparing binary acids of X in the same row*
Molecule CH4 NH3 H2O HFΔEN 0.4 0.9 1.4 1.9Acidity: CH4 < NH3 < H2O < HF
The higher polarity of the bond (the larger ΔEN (electronegativity difference)), the stronger acid.• Small ΔEN has more covalent character• Large ΔEN has more ionic character
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Comparing binary acids of X the same group
The larger bond length (the larger X radius, the weaker H—X bond) the stronger acid.
Molecule HF HCl HBr HIBE (kJ/mol) 565 431 364 297Anion radius (pm) 136 181 195 216Ka 6.6x10–4 ~106 ~108
~109
Acidity: HF < HCl < HBr < HI
Other example:Acidity: H2O < H2S < H2Se <
H2Te
*** HF(aq) is a weak acid
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Strengths of Oxoacids (H–O–EOn)
Molecule H–OI H–OBr H–OClEN 2.5 2.8 3.0Ka 2.3x10–11 2.5x10–9 2.9x10–8
Acid strength HOI < HOBr < HOCl
More examples:Acid strength: H2SeO3 < H2SO3
HBrO4 < HClO4
Comparing the EN of EThe larger EN (electronegativity) of E, the weaker H–O bond, the stronger acid.
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Comparing the n of On
The more number (n) of terminal O, the weaker protonated H–O bond, the stronger acid.
*O is the element has second higher electronegativity.
More examples:Ka1(H2SO3) < Ka1(H2SO4)
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Strengths of R–COOH vs. R–OH
• Ethoxide ion is a much stronger base than is acetate ion.• The stronger the conjugate base, the weaker the
corresponding acid.
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Strengths of carboxylic acids (R–COOH) Comparing electron-donating ability of R
• Electron-donating ability: –C2H5 > –CH3 > –H• The higher electron-donating ability, the weaker
acid strength.ExampleAcidity: CH3CH2–COOH < CH3–COOH < H–COOHKa: CH3CH2–COOH < CH3–COOH < H–COOHpKa: CH3CH2–COOH > CH3–COOH > H–COOH
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Comparing electron-withdrawing ability of R• Electron withdrawing ability: Cl > Br > I• The higher electron withdrawing ability, the
stronger acid strength Example
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Strengths of amines (–NH2) as bases Effect of electron-withdrawing (halogen, –OH, and –
NO2) group• Electronegative group withdraws electron
density from the N atom, i.e., the lone-pair electrons cannot bind a proton as strongly, and the base is weaker.
Example
Base strength: weaker
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Effect of electron-donating group• Electron-donating ability: –C2H5 > –CH3 > –H• The higher electron-donating, the stronger base
strength
Example
Base strength: Stronger
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Effect of aromatic group• π electrons in the benzene ring of an aromatic
molecule are delocalized and can involve the N’s lone-pair electrons in the resonance hybrid.
• Aromatic amines are much weaker bases than aliphatic amines.
Example 1
Example 2
The Weaker base
The Weaker base
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16-9 Lewis Acids and Bases Glossary and Definition
Lewis acid: The species act as electron-pair acceptor. Lewis base: The species act as electron-pair donor.
* In organic chemistry:Lewis acids called electrophiles (electron-loving)Lewis bases called nucleophiles (nucleus-loving)
Example:
Lewis base
Lewis acid
Complex(adduct)
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Example: CaO(s) reacts with SO2(g) to form CaSO3(s)
Ans: O2– act as Lewis base, SO2 act as Lewis acid
Ca2+ O2– + SO
O
Ca2+ O SO
O2–
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Summary of acid-base definitionAcid Base
Arrhenius H+ donor OH– donorBronsted-Lowry H+ donor H+ acceptorLewis e– acceptor e– donor
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Hydrated metal ion
Al3+ + 6H2O Al(H2O)63+
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Hydrolysis of hydrated metal ion
Lewis base
Lewis acid
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ionic chargeρ = charge density = ionic volume
The higher charge density, the higher acid strength.
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End of Chapter 16