Organic 3 Unit 2
Organic acids
The slides are made from the info obtain in:
• McMurry Organic chemistry Edition 7 and 8
• Study guide
• Chapter 5 from R.F. Daley and S. J. Daley
www.ochem4free.com
Please consult the text book for more complete info and
explanations. You will understand much better if you read the text
book since the slides are just summaries and does not contain all
information
ACID-BASE: three theories
• Arrhenius theory: about H+ and OH– Acid: the hydrogen ion, proton H+ (H3O+) producer in water
– Base: OH- producer in water
• Bronsted-Lowry theory: about protons, H+
– Acid: the proton, H+ donor
– Base: the proton, H+ acceptor
• Lewis Theory: about electrons, e– Acid :electron pair acceptor
– Base :electron pair donor
Inorganic 2 revision
Very important
ACID-BASE: three theories
Arrhenius
Bronsted-Lowry
Lewis
All Bronsted –Lowry acids can also be defined as Lewis acids
Not all Lewis acids can be defined as Bronsted-Lowry acids
Bronsted-Lowry theory
• Bronsted-Lowry theory: about protons, H+
– Acid: the proton, H+ donor
– Base: the proton, H+ acceptor
McMurry ed 7 p49
Bronsted-Lowry theory
Bronsted-Lowry theory: about protons, H+
Acid: the proton, H+ donor Base: the proton, H+ acceptor
McMurry ed 7 p50
Questions
• Problem 2:11 in McMurry
Organic 2 revision
•
Nitric acid (HNO3) reacts
with
ammonia (NH3)
to yield
ammonium
nitrate.
Write
the
reaction, and identify the acid, the
base, the conjugate acid product, and
the conjugate base product
Acid base strength
Acid strength: magnitude of the equilibrium constant,
Keq or acid constant Ka
HA + H2O ⇌ H3O+ + A-

H O A 
Keq 



H O A 
Keq H O  

3
2
3
HAH 2O

H O A 
Ka 

3
HA


HA
pKa = - log Ka
A stronger acid (larger Ka) has a smaller pKa,
and a weaker acid (smaller Ka) has
a larger pKa.
Stronger acids have the equilibria towards the right
and therefore have larger Ka (Acid Constants)
Ka values
The common inorganic acids such as H2SO4, HNO3, and HCl
have Ka’s in the range of 102 to 109, while organic acids
generally have Ka’s in the range of 10-5 to 10-15.
Ka of 102 pKa -2
Ka of 109 pKa -9
Strong acid: small pKa
Ka of 10-5 pKa 5
Ka of 10-15 pKa 15 Weak acid: large pKa
pKa values
•
Chapter 5 from R.F. Daley and S. J. Daley
www.ochem4free.com
McMurry ed 7 p51
pKa values
A strong acid has a weak conjugate base, and a weak acid has a strong conjugate base
Questions
• Problem 2:12 and 2:13 in McMurry
• Which one is the weaker acid and why?
– Phenylalanine , pKa = 1,83
– Tryptophan, pKa = 2,83
Organic Acids
But what is an organic acid?
Organic acids are characterised by the presence of
a positively polarised hydrogen atom
Two main kinds:
1. Hydrogen atom bonded to an electronegative
oxygen atom (O-H).
2. Hydrogen atom bonded to a carbon atom next
to a C=O bond (O=C-C-H).
Organic acids: polar covalent bonds
Organic Acids
• Moreover, the conjugate base of acetic acid is
stabilised by resonance
Organic Acids
Organic Bases
Characterised by the presence of an atom
with a lone pair of electrons that can bond to
H+
 Nitrogen containing compounds are the most
common, while oxygen containing compounds
may also act as bases when reacting wit
sufficiently strong acids.
Oxygen-containing compounds may also act as
both acids and bases, depending on circumstances
Organic Bases
Lewis acids and bases
• Lewis Theory: about electrons, e– Acid :electron pair acceptor
– Base :electron pair donor
McMurry ed 7 57
Lewis acids
All Bronsted-Lowry acids are also Lewis Acids, but
the reverse is not necessarily true
McMurry ed 7 p58
Lewis bases
Others can act either as an acid or base (Alcohols and carboxylic acids)
act as acids when they donate an H+ but as
bases when their oxygen atom accepts an H+.
McMurry ed 7 59
Lewis acids and basis
electrophile
McMurry ed 7 58
nucleophile
Lewis acids and basis
Acid base strength
Acid strength: magnitude of the equilibrium constant,
Keq or acid constant Ka
HA + H2O ⇌ H3O+ + A-

H O A 
Keq H O  

H O A 
Keq 



3
3
HAH 2O
2

H O A 
Ka 

3
HA
McMurry ed 7 p51
HA

pKa = - log Ka

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 Bronsted-Lowry acid H-A is dissolved in
water, an acid-base reaction occurs, and an
equilibrium constant can be written for the rxn.
Factors that determine acid strength
• Stability of conjugated base A influences the
strength acidity
• Higher stability makes acid more acidic
Factors influencing acidity:
– Resonance effects
– Induction effects
– Hybridisation effects
– Charge effects
– Elemental effects
Method to compare two acids
• Draw the conjugated base
• Determine the relative stabilities of
conjugated bases
• Higher stability of conjugated base makes acid
more acidic
Resonance effects
• Resonance effects makes conjugated base more
stable, therefore acid has higher strength
no resonance structures
resonance structures thus stable
conjugated base
Qs: why is acetic acid more acidic than methanol and ethanol?
Qs: which of these two structures are more acidic?
Anion,
the conjugated
base has
resonance
structures thus
more stable
Anion
without
resonance
structures
Resonance structures
http://chemwiki.ucdavis.edu/Core/Organic_Chemistry/UMM_chemwiki_project/Acid%2F%2FBase_Reactions/Resonance_effects_o
n_acidity
Resonance effects: Alcohols
• Ethanol (CH3CH2OH) pKa= 16.00
– no resonance structures
• Phenol pKa= 9.89
– many resonance structures thus stable conjugated base
Substituted phenols acidity
http://chemistry.tutorvista.com/organic-chemistry/phenol.html
http://www.chem.ucalgary.ca/courses/351/Carey5th/Ch24/ch24-1.html
http://chemistry.stackexchange.com/questions/19205/acidity-of-substituted-phenols
Inductive
effects
 Atom’s ability to polarize a bond
 Shifting of electrons in a sigma bond in response to the electronegativity of nearby
atoms.
 Metals, such as lithium and magnesium, inductively donate electrons, whereas
reactive non-metals, such as oxygen and nitrogen, inductively withdraw electrons.
 Stabilises the conjugated base (anion) leading to a stronger
acid
http://chemwiki.ucdavis.edu/Core/Organic_Chemistry/UMM_chemwiki_project/Acid%2F%2FBase_Reactions/I
nductive_effects_on_acidity
Inductive effects: Alcohols
• Phenol pKa= 9.89
• Electron donating groups like amino (-NH2), alkyl (-R) decreases the
acidity of phenol leading to a less stable conjugated base (anion)
– p-Methylphenol pKa= 10.17
• Electron withdrawing groups like Cl and NO2 increases the acidity
of phenol, leading to a more stable conjugated base (anion)
– p-Chlorophenol pKa= 9.38
– p-Nitrophenol pKa= 7.15
Hybridisation effects
• The higher the percent of s-character of the
hybridized orbital, the closer the lone pair is held to
the nucleus, and the more stable the conjugate base.
Hybridisation effects
http://www.masterorganicchemistry.com/2012/04/25/walkthrough-of-acid-base-reactions-3-acidity-trends/
Elemental effects
The more electronegative an atom, the better it is able to bear a negative charge.
Less stable
Weaker acid
More stable
Stronger acid
Thus, the ethoxide
anion is the most stable
(lowest energy, least
basic) of the three
conjugate bases, and
the ethyl anion is the
least stable (highest
energy, most basic).
http://chemwiki.ucdavis.edu/Core/Organic_Chemistry/UMM_chemwiki_project/Acid%2F%2FBase_Reactions/I
nductive_effects_on_acidity
Elemental effects
Stronger
base
Weaker
acid
More
stable
Stronger
acid
Less
stable
Weaker
base
Acidity of alpha () hydrogen
• Why are the protons adjacent to the carbonyl
group acidic?
McMurry ed 8 718
Question
• Will a ketone (like acetone) or an alcohol (like
propanol) be more acidic?
• The conjugated base of the ketone (anion) has resonance structures
• Thus more stable
• Thus ketone more acidic
Charge effects
• The more negative charge: the more basic
• The more positive charge: the more acidic
– H3O+ : most acidic
– H3O+  H+ + H2O
– H2O neutral and stable
– H2O : middle
– H2O  H+ + OH– OH- less stable than water thus H2O less acidic than H3O+
– OH- :least acid (most basic)
– OH-  H+ + O2– O2- unstable, OH- more stable
Acidity of alpha () hydrogen
Qs: Which molecule is more acidic? Aldehyde, ketone or ester. Explain why.
Qs: Explain why is an aldehyde is more acidic than a ketone?
Qs: Explain why a ketone is more acidic than an ester?
Stability
http://www.chem.ucalgary.ca/courses/350/Carey5th/Ch21/ch21-2.html
Acidity of alpha () hydrogen
Competing resonance
leading to destabilisation
The difference is in the nature of the group attached to the carbonyl group.
Aldehyde: H
Ketone: alkyl (e.g. CH3)
Ester: alkoxy group (e.g. OCH3)
Acidity of alpha () hydrogen
•
The difference between the 3 systems is in the nature of the group attached to the
common carbonyl. The aldehyde has a hydrogen, the ketone an alkyl- group and
the ester an alkoxy- group.
•
H atoms are regarded as having no electronic effect : they don't withdraw or
donate electrons.
•
Alkyl groups are weakly electron donating, they tend to destabilise anions. This is
because they will be "pushing" electrons towards a negative system which
is unfavourable electrostatically. Hence, the anion of a ketone, where there are
extra alkyl groups is less stable than that of an aldehyde, and so, a ketone
is less acidic.
•
In the ester, there is also a resonance donation from the alkoxy group towards the
carbonyl that competes with the stabilisation of the enolate charge. This makes
the ester enolate less stable than those of aldehydes and ketones so esters are
even less acidic.
http://www.chem.ucalgary.ca/courses/350/Carey5th/Ch21/ch21-2.html
Basicity of amines
Amines are much stronger bases than alcohols
and ethers, their oxygen containing
Analogs……..Why ?
Lone pair of electrons on the nitrogen making it basic and nucleophilic
McMurry ed 8 p949 ed 7 p921
Effects that decreases the electron
density on Nitrogen
Lone pair of nitrogen will be less available for
protonation and the amine less basic, if:
The lone pair is involved in maintaining the
aromaticity of the molecule
Nitrogen attached to an electron-withdrawing
group
The lone pair is conjugated with an electronwithdrawing group
The lone pair is in an sp or sp2 hydridized orbital
McMurry ed 8 950
Basicity of amines
McMurry ed 8 950
Basicity of amines
Aromatic—A cyclic, planar, completely conjugated compound with 4n + 2  electrons.
Look at aromaticity electron count to know if the lone pairs are part of the aromatic
system or not
Basicity of amines ?
• NH3< primary amine ~ tertiary amine <
secondary amine
• Electron donation effects
– reason why 2° amine more basic than 1° amine
• Steric hindrance
– alkyl group would hinder the attack of a hydrogen
atom
– reason why 3° amine less basic than 2° amine
• Solvation
– Less hydrogen bonding for 3° amines
– reason why 3° amine less basic than 2° amine
http://www.tutorvista.com/chemistry/basicity-of-amines
Basicity of amines
Electron withdrawing group from N – Less basic (smaller PKa)
Electron donating group on the N –Increases the basicity (Larger pKa)
Arylamines are less basic than alkyl amines
Arylamines
• Arylamines are less basic than alkyl amines
• The lone pair of electrons on the nitrogen of
aniline are conjugated to the -electrons of
the aromatic ring and therefore less available
for acid-base chemistry
• Protonation disrupts the conjugation
Electron donating
groups, stabilise
conjugated acid
Electron withdrawing
groups, destabilise
conjugated acid
Pyridine vs pyrrole
Protonation disrupts the
conjugation
Lone pair part of aromatic sextet: less available thus less basic
No longer aromatic
http://chemwiki.ucdavis.edu/Textbook_Maps/Organic_Chemistry_Textbook_Maps/Map%3A_Organic_Chemistry_(McMurray)/Unit_24_Amines_and_Heterocycles/
24.03_Basicity_of_Amines
Lone pair NOT part of aromatic sextet: more available thus more basic
Aromatic—A cyclic, planar, completely conjugated compound with 4n + 2  electrons.
Look at aromaticity electron count to know if the lone pairs are part of the aromatic
system or not