inorganic

Acid-Base Concepts I.

Unifying Concepts A.

B.

C.

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

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-

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

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4)

H3O+ + NO2-

5)

Includes non-aqueous systems NH4+ + NH2-

D.

H2O + HNO2 2 NH3

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 a)

Any solute increasing [H 3O+] is an acid HCl + H2O

b)

H3O+ + OHH3O+ + Cl-

Any solute increasing [OH -] is a base NH3 + H2O

NH4+ + OH-







6) pK ion = -log[acid][base]

E.

a)

pK W = -log[H3O+][OH-] = -log[10 -7][10-7] = 14

b)

pK H2SO4 = -log[H3SO4+][HSO4-] = 3.4

c)

The smaller the number, the more dissociation has occurred

Lewis Concept 1)

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

2)

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







5)

A nonn-me meta tall ex exam amp ple le:: BF3 + NH3

H3 N:BF3 (or BF3 • NH3)







II.

Acid-Base Strength A.

Thermodynamic Measurement 1) 2) 3)

4)

We can easily measure pH, but that doesn t really tell us about acid strength 

HA + H2O !Go

+

K a

= -RTlnK a =

a)

G = free energy

b)

H = enthalpy

c)

S = entropy

Solving for K a:

K a

H3O+ + A!H

lnK a

- T!S

=

! !H RT

+

!S

R

=

[H3O ][A! ] [HA]







B.

Binary Hydrogen Compounds 1)

2)

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+

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 i t 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, Oxygens, the stronger the acid a)

Acidity: HOClO3 > HOClO2 > HOClO > HOCl







D.

Cations in Aqueous Solution 1)

Cationic metal ions are generally Lewis acids in i n water solutions

2)

Example: [Fe(H2O)6]3+ + H2O

3)

Large charge and small radii increase acidity

[Fe(H2O)5(OH)]2+ + H3O+

a)

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

b)

2+ Transition Metals are weak acids; 3+ Transition Metals are







4)

The stron

acid the cation cation is, the the less less soluble soluble the hydro hydroxide xide

lex i







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 CH3 N

N







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 Solvents











H.

Superacids = acids stronger than H 2SO4 1)

Hammet Acidity Function = H

+

H

pK

log

[BH ]

inorganic

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