GenChem Ch 162013/03/03TMHsiung 1/60 Chapter 16 Acids and Bases.

GenChem Ch 162013/03/03 TMHsiung 1/60

Chapter 16

Acids and Bases


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


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


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−


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.


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.


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.


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.


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:


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.


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


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.


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.


For extremely dilute solution of a strong acid and strong base, 1.0 x 10–8 M HCl for example:


16-5 Weak Acids and Weak Bases Identifying Weak Acids and Bases


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%


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:


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


16-6 Polyprotic Acids


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


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.


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


(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


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


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


Check: (x/CH2SO4) x 100% = 2.2% < 5%, OK!!


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+


CH3COO–(aq) + H2O CH3COOH(aq) + OH–

(aq)

XX

X

X


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/


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/


AlCl3, Fe(NO3)3 Hydrated

Cation

<7

The pH of Salt Solutions


(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+


NO3–(aq) + H2O X


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


• Bond dissociation energies (kJ/mol ) and Ka values for some binary acids


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


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


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.


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)


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.


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


Comparing electron-withdrawing ability of R• Electron withdrawing ability: Cl > Br > I• The higher electron withdrawing ability, the

stronger acid strength Example


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


Effect of electron-donating group• Electron-donating ability: –C2H5 > –CH3 > –H• The higher electron-donating, the stronger base

strength

Example

Base strength: Stronger


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


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)


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–


Summary of acid-base definitionAcid Base

Arrhenius H+ donor OH– donorBronsted-Lowry H+ donor H+ acceptorLewis e– acceptor e– donor


Hydrated metal ion

Al3+ + 6H2O Al(H2O)63+


Hydrolysis of hydrated metal ion

Lewis base

Lewis acid


ionic chargeρ = charge density = ionic volume

The higher charge density, the higher acid strength.


End of Chapter 16

GenChem Ch 162013/03/03TMHsiung 1/60 Chapter 16 Acids and Bases.

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Transcript of GenChem Ch 162013/03/03TMHsiung 1/60 Chapter 16 Acids and Bases.