acids and bases - manu's...

Acids and BasesDr. Sapna Gupta

Arrows in Organic Chemistry

synthesis (yield)

equilibrium

resonance

retrosynthesis (backward)

transfer of two electrons

transfer of one electron

Dr. Sapna Gupta/Acids and Bases 2

Acids and Bases

• A Brønsted acid donates a hydrogen cation (H+)

• A Brønsted base accepts the H+

• “proton” is a synonym for H+ - loss of an electron from H leaving the bare nucleus—a proton

• Full headed arrows indicate transfer of electrons.

Example


Conjugate Acid Base Pairs

• Conjugate base: formed from an acid when it donates a proton to a base. A strong acid gives a weak conjugate base and vice versa.

• Conjugate acid: formed from a base when it accepts a proton from an acid. A strong base gives a weak conjugate acid and vice versa.

HCl( aq) H2 O( l) Cl-( aq) H3 O+ ( aq)+ +

WaterHydrogen

chloride

Hydronium

ion

Chloride

ion

(base)(acid) (conjugate

acid of H 2O)

(conjugate

base of HCl)

conjugate acid-base pair


NH4+CH3 COOH CH3 COO-

NH3+ +

Acetic acid Ammonia

(acid)


Acetate

ion

Ammonium

ion

(base) (conjugate base

acetic acid)

(conjugate acid

of ammonia)



Examples of Acids and Bases

• There are inorganic (mineral) and organic acids and bases.


Solved Problems1) What is conjugate acid of NH3?

a) NH2

b) NH2+

c) NH2-

d) NH4

e) NH4+

2) What are the conjugate bases in the reaction below?

CO32 + HSO4

HCO3 + SO4

2

a) HCO3 and HSO4

b) HSO4 and CO3

2

c) CO32 and OH

d) SO42 and HSO4

e) CO32and SO4

2

3) For the reaction below which two substances which are both acids

CH3NH3+ + H2O CH3NH2 + H3O+

a) H2O and H3O+

b) CH3NH3+ and H2O

c) CH3NH3+ and CH3NH2

d) CH3NH3+ and H3O+

e) CH3NH2 and H2O

4) A strong acid leads to a

a) weak conjugate acidb) strong conjugate basec) weak conjugate based) strong basee) pure water


Acid Strength and pKa

• Acid strength is the tendency of an acid to donate a proton.

• The more readily a compound donates a proton, the stronger an acid it is.

• Acidity is measured by an equilibrium constant.

• When a Brønsted-Lowry acid H—A is dissolved in water, an acid-base reaction occurs, and an equilibrium constant can be written for the reaction.

• It is more convenient to use “pKa” values than Ka values.


pKa Table


Factors Determining Acid Strength

1. Electronegativity: Across the row, e.g. H2O vs CH4 (the more electronegative the element the more acidic it is)

2. Size of the anion: down the group, e.g. HF, HCl, HBr, HI (the larger the anion the more stable it is)

3. Number of oxygen, e.g. HNO2 vs HNO3( more oxygen cause more electronegativity hence easier for H to leave)

4. Inductive effects, e.g. CH3COOH vs CH2ClCOOH

5. Resonance stabilization of conjugate base, e.g. CH3COOH vsCH3OH

6. Acidity of hydrocarbons due to delocalization of e- in conjugate base, e.g.CH=CH, CH2=CH2, CH3CH3

7. Solvent effect: more polar solvents will support more bronstead-lowry type acids because of ions formed during reactions.


1) Electronegativity in the Row

Within a row, the greater the electronegativity of the atom bearing the negative charge stabilizes the anion and the acid is stronger.


2) Size of Anion

• Within a column of the Periodic Table, acidity is related to the size of the the atom bearing the negative charge.

• Atomic size increases from top to bottom of a column.

• The larger the atom bearing the charge, the greater its stability.

S HCH3 CH3 O – CH3 S

–O HCH3

Methanethiol

pKa 7.0

(stronger acid)

Methoxide

ion

(stronger base)

Methanethiolate

ion

(weaker base)

Methanol

pKa 16

(weaker acid)

+ +


3) Number of Oxygen Atoms and Electronegativity

• In all Bronsted acids the proton that dissociates is bonded to oxygen.

• Compare the Lewis structures of phosphorous acid (H3PO3) and phosphoric acid (H3PO4). Phosphoric acid is triprotic and all three protons are bonded to oxygen whereas phosphorous is diprotic as it has only 2 protons bonded to oxygen, the third one, bonded to P, does not ionize. The bond between oxygen and hydrogen is more polar than between hydrogen and phosphorous due to electronegativity difference.

• For mineral acids more oxygen atoms means they are the more acidic e.g. HNO3 > HNO2 and H2SO4 > H2SO3

• For organic acids CH3COOH > CH3OH


4) Inductive Effect

• Electronic effects that are transmitted through space and through the bonds of a molecule due to the electronegativity of an adjacent covalent bond.

• Stabilization by the inductive effect falls off rapidly with increasing distance of the electronegative atom from the site of negative charge in the conjugate base.

• We also see the operation of the inductive effect in the acidity of alchoholsand acids and also the halogen substituted alcohols and carboxylic acids.

C-CH2O-H

H

H

H

C-CH2O-H

F

F

F

Ethanol

pKa 15.9

2,2,2-Trifluoroethanol

pKa 12.4

Butanoic acid

pKa 4.82

4-Chlorobutanoicacid

pKa 4.52


pKa 3.98


pKa 2.83

OH

O

OH

O

OH

O

OH

O

Cl

Cl

Cl


5) ResonanceResonance delocalization of charge in A-. The more stable the anion, the more the position of equilibrium is shifted to the right.

• Compare the acidity alcohols and carboxylic acids.

• Ionization of the O-H bond of an alcohol gives an anion for which there is no resonance stabilization.

• Ionization of a carboxylic acid gives a resonance-stabilized anion.

• The pKa of acetic acid is 4.76

• Carboxylic acids are stronger acids than alcohols as a result of the resonance stabilization of the carboxylate anion.

CH3CH2O-H H2 O CH3CH2O -

H3 O+

+

An alcohol An alkoxide ion

+ pKa = 15.9

equivalent contributing structures;

the carboxylate anion is stabilized by

delocalization of the negative charge.

+CH3 C

O

O

C

O

O

CH3

O H

C

O

CH3 H3 O+

+ H2 O


6) Hybridization

• For anions differing only in the hybridization of the charged atom, the greater the % s character to the hybrid orbital of the charged atom, the more stable the anion, therefore more acidic.

• Consider the acidity of alkanes, alkenes, and alkynes (given for comparison are the acidities of water and ammonia).

CH3 CH2 -H CH3 CH2–

CH2 = CH- H CH2 = CH–

H2 N-H H2 N–

HC C H HC C–

HO- H HO–

Weak

Acid

Alkyne

Alkene

Alkane

Water

25

44

51

15.7

Conjugate

Base pKa

Incr

easi

ng

aci

dit

y

Ammonia 38

More acidic

More basic


7) Effect of Solvent

• Acidity values in gas phase are generally very low

• It is difficult to separate the product ions without solvent molecules to stabilize them

• Acetic acid has pKa of 130 in the gas phase

• A protic solvent is one in which hydrogen is attached to a highly electronegative atom such as oxygen or nitrogen e.g. water

• Solvation of both acetic acid and acetate ion occurs in water although the acetate is more stabilized by this solvation


Strength of Bases

• Bases are opposite of acids in strength; a strong acid will give a weak conjugate base and a weak acid gives a strong conjugate base.

• E.g. HCl is a strong acid so Cl- is a weak conjugate base; methanol (CH3OH) is a weak acid so methoxide (CH3O-) is a strong conjugate base.

• It is best to remember acidity rules to remember bases also.


Lewis Acids and Bases

• Lewis acid: A molecule/ion that can accept a pair of electrons.

• Lewis base: A molecule/ion that can donate a pair of electrons.

• There is no pKa scale for these acids and bases. Their effectiveness is determined by how well they donate or accept electrons.

• Examples of Lewis Acids • Group 3A elements, such as BF3 and AlCl3, are Lewis acids because they

have unfilled valence orbitals and can accept electron pairs from Lewis bases

• Transition-metal compounds, such as TiCl4, FeCl3, ZnCl2, and SnCl4, are Lewis acids

• In case of organic compounds, any carbocation (carbon with a positive charge) would be a Lewis acid.


Organic Compounds as Lewis Bases

• Any organic compound containing an atom with a lone pair (O,N) can act as a base


Solved Problems

Which of the following is a Lewis base?a) BCl3

b) Cu2+

c) SH-

d) Mn2+

e) NH4+

Iodine trichloride, ICl3, will react with a chloride ion to form ICl4

-. Which species, is the Lewis base this reaction?a) ICl4

-

b) ICl3

c) Cl-

d) the solvent

Which one of the following is a Lewis acid but not a Brønsted-Lowry acid?a) Fe3+

b) H3O+

c) HSO4-

d) NH3

Which of these species will act as a Lewis acid?a) NH3

b) NH4+

c) H2O d) BF3

e) F


Lewis Acid Base Reactions

Examples

• For writing Lewis acid base reactions look for electron poor and electron rich sites.

• Electron poor sites are cations and incomplete valence shell.

• Electron rich sites are anions, lone pair of electrons, double and triple bonds.

• NOTE: all the red arrows indicate an electron pair is being transferred.

Lewis base

Lewis acid

+ BA new covalent bondformed in this Lewisacid-base reaction

:+-

A B

:

: :

Br -CH3 -C C-CH3

H H

H

CH3 -C C-CH3

H H

H Br

Bromide

ion

sec- Butyl cation

(a carbocation)

2-Bromobutane

++

: :

::

+ -

Diethyl ether(a Lewis base)

Boron trifluoride (a Lewis acid)

O

CH3 CH2

CH3 CH2

F

F

O

CH3 CH2

CH3 CH2

B-F

F

F

+ B

A BF3-ether complex

F:: :

CH3 -CH= CH-CH3 + H- Br CH3 -C- C-CH3

H

H

+

2-Butene sec- Butyl cation

(a 2° carbocation)

H

Br

(a Lewis base)Dr. Sapna Gupta/Acids and Bases 21

Examples of Some Acid-Base Reactions

• Any base stronger than hydroxide will be converted to hydroxide in water

• Sodium amide can be used as a strong base in solvents such as liquid NH3

• Alkyl lithium reagents in hexane are very strong bases


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