Chapter 6 Lecture 1 Acid-Base Concepts

I. Unifying ConceptsA. The Acid-Base Concept

1) There are many acid-base definitions, each at times useful

2) Acid-Base concepts are not facts or even theories, but are useful generalizations for classification, and organization

3) Acid-Base concepts are powerful ways to explain data and predict trends

B. Arrhenius Concept

1) An acid forms H+ in water; a base forms OH- in water

2) Applicable to aqueous solutions only

3) HCl + NaOH H+ + OH- + Na+ + Cl-

C. Bronsted-Lowery Concept

1) Acid is a proton donor; Base is a proton acceptor

2) Conjugate acid/base pairs differ only by a proton

3) Reactions proceed to produce the weakest acid and base

4) H3O+ + NO2- H2O + HNO2

5) Includes non-aqueous systems

NH4+ + NH2

- 2 NH3

D. Solvent System Concept

1) Useful for aprotic, non-aqueous systems

2) Applies to any solvent that can dissociate to cation (acid) and anion (base)

3) For water: 2 H2O H3O+ + OH-

a) Any solute increasing [H3O+] is an acid

HCl + H2O H3O+ + Cl-

b) Any solute increasing [OH-] is a base

NH3 + H2O NH4+ + OH-

4) Aprotic, non-aqueous example: 2 BrF3 BrF2+ + BrF4

-

a) Acid: SbF5 + BrF3 BrF2+ + SbF6

-

b) Base: F- + BrF3 BrF4-

5) The acid-base reaction: acid + base = solvent (reverse of the ionization)

a) Arrhenius: acid + base = salt + water

b) Bronsted: acid1 + base2 = base1 + acid2

6) pKion = -log[acid][base]

a) pKW = -log[H3O+][OH-] = -log[10-7][10-7] = 14

b) pKH2SO4 = -log[H3SO4+][HSO4

-] = 3.4

c) The smaller the number, the more dissociation has occurred

E. Lewis Concept

1) Acid = e- pair acceptor; Base = e- pair donor

2) Includes metal ions and non-aqueous systems; encompasses other concepts

3) We will use this concept throughout the rest of the chapter and course

4) Metal ion Example: Ag+ + 2 NH3 [Ag(NH3)2]+

a) Acid-Base product is called an adduct

b) If the acid is a metal ion, it is also called a coordination compound or a coordination complex or a complex ion

5) A non-metal example: BF3 + NH3 H3N:BF3 (or BF3 • NH3)

II. Acid-Base StrengthA. Thermodynamic Measurement

1) We can easily measure pH, but that doesn’t really tell us about acid strength

2) HA + H2O H3O+ + A-

3) Go = -RTlnKa = H - TS

a) G = free energy

b) H = enthalpy

c) S = entropy

4) Solving for Ka:

Ka

[HA]

]][AO[HK 3

a

R

ΔS

RT

ΔHlnKa

B. Binary Hydrogen Compounds

1) Acidity increases down a column of the periodic table

a) H2Se > H2S > H2O

b) HI > HBr > HCl > HF

c) Conjugate bases of larger ions have lower charge density, thus a smaller attraction for H+

2) Acidity increases from left to right of the periodic table

a) NH3 < H2O < HF

b) The more electronegative the conjugate base is, the easier is it for H+ to dissociate

C. Inductive Effects (electron pulling/pushing through sigma bonds)

1) Electronegative substituents increase acidity and decrease basicity

Basicity: :PF3 < :PH3

2) Electron Donating substituents decrease acidity and increase basicity

Basicity: NMe3 > NHMe2 > NH2Me > NH3

3) Oxyacids: the more unprotonated Oxygens, the stronger the acid

a) Acidity: HOClO3 > HOClO2 > HOClO > HOCl

b) The electronegative O’s pull e- away from the H—O bond

c) The electronegative O’s stabilize the conjugate base

D. Cations in Aqueous Solution

1) Cationic metal ions are generally Lewis acids in water solutions

2) Example: [Fe(H2O)6]3+ + H2O [Fe(H2O)5(OH)]2+ + H3O+

3) Large charge and small radii increase acidity

a) Alkali metals are not acidic (Na+); Alkaline Earths are weakly acidic (Ca2+)

b) 2+ Transition Metals are weak acids; 3+ Transition Metals are strong acids

c) All 4+ or higher metals are very strong acids MxOy

4) The stronger acid the cation is, the less soluble the hydroxide complex is. OH- can’t dissociate to dissolve because of strong charge attraction.

We can use this property to estimate the acid strength of the cation

E. Steric Effects

1) Steric bulk can repel an acid-base partner, modifying the acid-base strength

a) F = front strain = direct steric interference at the site of interaction

b) B = back strain = bulky groups interfere opposite the interaction site upon binding as the molecule adjusts its VSEPR geometry

2) The order of basicity can scramble depending on bulk of the acid

N

CH3

CH3

N

4

3

2

1

4

2

1

3

4

2

1

3

F. Solvation

1) Solvation is interaction with solvent molecules

2) Basicity in water: NHMe2 > NH2Me > NMe3 > NH3

a) By induction, the more substituted amine should be the most basic

b) This amine has less H’s to interact with water

G. Non-aqueous Solvents

1) The Leveling Effect: the strongest acid possible in a solvent is the solvent cation; the strongest base possible in a solvent is the solvent anion.

a) H2SO4 + H2O H3O+ + HSO4- (100% dissociation)

b) Na2O + H2O 2 Na+ + 2 OH- (100% dissociation)

c) H2SO4 + HOAc H2OAc+ + HSO4- (< 100%)

d) NH3 + HOAc NH4+ + OAc- (100%)

2) HNO3, H2SO4, HClO4, HCl are all equally acidic in water (H3O+)

3) HClO4 > HCl > H2SO4 > HNO3 in HOAc

4) Hydrocarbon Solvents don’t level acids or bases

H. Superacids = acids stronger than H2SO4

1) Hammet Acidity Function = Ho

B = nitroaniline indicator used as the base

2) Lewis Superacids are often made by protonating an already strong acid

a) This is often done using HF as the acid to be protonated

b) It requires a very stable anion to make the reaction proceed

c) 2 HF + 2 SbF5 H2F+ + Sb2F11- (Fluoroantimonic acid)

[B]

][BHlogpKH BHo