Oxoacids

Simple oxoacids: the simplest oxoacids are the mononuclear acids that contain one atom of the parent element Ee.g. H2CO3, HNO3, H3PO4 and H2SO4

E(O)p(OH)q

p = 0 H-O-Cl

p = 1

O

PHO OH

OH

OH

SiHO OH

OH

OH

TeHO OH

OH

OH

HO HO BOHOH

O ClOH

OH SeO OH

OH

p = 2 O NOH

OS

O OHOH

O

p = 3 ClO O

OH

O

Cl+O O-

OH

O

Cl+

O O-OH Cl

O OOH

O-

I+HO OH

OH

OH

HOO

IHO OH

OH

OH

HO

107.2 7.8, 11.2 9.1

2.1, 7.4, 12.7 3.6 1.6, 7.0 2.6, 8.0

-1.4-2.0, 1.9 -1.0

-10

pKa

The first dissociation constant of oxoacids pKa(1)pKa(1) decreases as the number of “hydrogen free” oxygen atoms (p) increases

-1.02Cl(O)2(OH)HClO3

2.01Cl(O)(OH)HClO2

≈-83Cl(O)3(OH)HClO4

7.20Cl(OH)HOCl

≈-22S(O)2(OH)2H2SO4

1.821S(O)(OH)2H2SO3

-1.402N(O)2(OH)HNO3

3.371N(O)(OH)HNO2

pKa(1)nE(O)p(OH)qFormula of acid

•After the first dissociation, the number of “hydrogen free” O atoms increases to p+1

•The greater the number of the oxygen atom in the conjugated base[E(O)p+1(OH)q-1]-, the more the negative charge of the anion can be spread out.

e.g. the negative charge of nitrate ion is delocalized uniform over via the π-bonding system

O

O ON

-1/3

-1/3-1/3

Empirical rules devised by Linus Pauling

•For the oxoacids EOp(OH)q. pKa ≈ 8 – 5p

•The successive pKa values of polyproticacids (those with q > 1) increases by 5 units for each successive deprotonation

Aquated Cations as Brønsted Acids – Aqua Acids

•When a metal salt dissolves in H2O, the cation and anion are hydrated

e.g. Aqueous solution of NaCl

Na+

O

O

O

O

O

O

HHH

H

H

HH H

H

H

H

H

δ-

δ-

δ-

δ-

δ-

δ-

Ion-dipoleinteractions

First hydrationshell

For many transition metal salts, the cations Mn+ form covalent bonds with H2O

Mn+

O

O

O

O

O

O

HHH

H

H

HH H

H

H

H

H

Covalent bond

Coordination/dative/metal-ligand bond

Cu2+

O

O

O

O

O

O

HHH

H

H

HH H

H

H

H

H

Hexaaquacopper(II) ion

Distorted octahedralComplex/adduct

The metal-H2O complexes can act as BrønstedAcids by loss of H+ from a coordinated H2O molecule:

[M(H2O)6]n+(aq) + H2O (l)

[M(H2O)5(OH)](n-1)+(aq) + H3O+

(aq)

The position of the equilibrium depends on the degree of the O-H bonds are polarized

Hydrolysis/deprotonation

M

OH H

When H2O coordinates to M with n+ positivecharge, charge is withdrawn towards the metal center, leaving the H more δ+ than in bulk water

δ+ δ+

e- density

OH

H

Mn+

OH

OH

HH

+

n+(n-1)+

Hydroxide ion

13.30.030132Sr2+

11.40.04786Mg2+

12.80.035114Ca2+

13.50.027149Ba2+

13.60.01190Li+14.20.009116Na+

14.50.007152K+

pKa(a)RadiusIon

pKa Values of Metal Cations

(a) = Z2/r

a pKa

Harnessing the Acidity of Metal Cations in Biology

e.g. Carbonic anhydraseAn enzyme located in RBC that catalyzes conversion of CO2to bicarbonate ion

CO2 + H2O HCO3- + H+

Carbonic anhydrase

•In the absence of enzyme, this reaction is very slow: ~ 10-2 s-1

in the presence of enzyme: 106 s-1

•The active site of this enzyme is composed of a Zn2+ ion bonded to the backbone of the enzyme

bicarbonate

(aq) (aq) (aq)(l)

• One of the reasons for the slow rate constant is the poor nucleophlicity of H2O. Hydroxide is a strong nucleophile but its concentration at pH = 7 is only 10-7 M

HO

H

O

C

O

δ+

δ-

δ--H+ O

C

O-HO

Metal-OH has stronger nucleophilicity

Coordination of a water molecule to Zn2+ in the enzyme increases its acidity to 107 times!

Zn2+

O

HH

δ+δ+Coordinated water

pKa = 7H

ydrophobic region

Hydrophobic region – binding of CO2

Active site

Zn

O

HH

O

C

O Zn

OH

O

C

O

δ+δ+

O

HH

Zn

OH

O

C

O

HCO3-

H2O

H3O+δ+

δ-

δ-δ+

δ-

δ-

(b) Lewis Concept

•Lewis Acid – electron pair acceptor

•Lewis Base – electron pair donor

Lewis Acid + Lewis Base Acid – base adduct

N H Cl+ N H

+

Cl-+

AdductLewis base Adductacid base

Lewis base

H+ + OH- H2Oe--acceptor e--donor Adduct

N +4 Ni2+

N

Ni2+

N

NN

2+

Lewis base Lewis acid Adduct/complex

Metal ion

Examples of Lewis Acids and Bases

•Anions i.e. OH-, O2-

Halides, S2-, Se2-, NH2-,

S-R-, O-R-, alkyl & aryl ions

•pseduohalides,•Neutral molecules i.e.

H2O, amines, ammonia, phosphines, arsines,

ethers, thioethers, heterocyclics, (many of them contain group 15, 16 or 17 elements as

their donor atoms)

•Alkali metal ions•Alkaline earth metal ions

•Transition metal ions(e.g. Fe3+, Cu2+, Zn2+, Ni2+, Au+1,

Ti4+, Co3+)•Main group metal ions (e.g. Tl+,

Pb2+, Al3+) •Neutral molecules(e.g. BF3, BeCl2)

BaseAcid

Examples of Lewis Acids: Transition Metals

Different Oxidation States

e.g. Fe0, Fe+1, Fe+2, Fe+3, Fe+4

Different ionic radii

e.g. Fe0 > Fe+1> Fe+2> Fe+3> Fe+4

Ru+2 > Fe+2; Au+1> Au+3

Cu2+

O

O

O

O

O

O

HHH

H

H

HH H

H

H

H

HCu2+ + 6 H2O

Ag+ + 2 NH3 H3N Ag NH3

2+

+

Cu2+ + 4 Cl- CuCl42-

Lewis acid Lewis base Adduct/complex

Acid-Base Reactions Involving Metal Ions

H3N

Cl Cl

NH3

Pt S Au ClH3C

H3C

WCl

Cl Cl

Cl

Cl

Cl

WH3C

H3C CH3

CH3

CH3

CH3

Metal-carbon bond

HexamethyltungstenAn organometallic complex

Identify the acids and bases

Examples of Lewis Acids: Group 2

Cl Be Cl 4-e-

BeCl2 +

Beryllium dichloride

2 Cl-

CH2CH3

O

CH2CH32

Cl

BeCl Cl

Cl

2-

O

BeO Cl

ClEt

Et

Et Et

Examples of Lewis Acids: Group 13

X

X XB

6-e-X = halides, alkyl

X

X XAl

X

X XGa

X

X XIn

X

X XTl

F

F FB

CH3

CH3

CH3N

+

N

BF F

F

CH3

CH3H3C

+ CH2CH3

OCH2CH3

trimethylamine

diethyl ether

O

BF F

F

CH2CH3H3CH2C

Acid-Base Reactions Involving Group 13 Acids

Al

Cl

Al

ClCl

Cl

Cl

Cl+ 2 CN CH3 2

AlCl

Cl

Cl

NC

CH3

H

H HGa

+ NMe3

GaH

H

H

Me3N

+ 2 NMe3 H GaH

HNMe2

NMe3

Acid-Base Reactions Involving Group 13 Acids