Transcript Chapter 1--Title - Imperial Valley College

Chapter 3
An Introduction to Organic Reactions:
Acids and Bases
Chapter 3
An Introduction to Organic Reactions:
Acids and Bases
Organic Reactions and their Mechanisms
A reaction mechanism is a detailed description of the bonding changes as
a reaction proceeds. The reaction mechanism also includes the many
important principles of organic chemistry. A plausible reaction
mechanism must be consistent with the principles of organic chemistry.
Four General Categories of Organic Reactions
Organic reactions tend to fall into four categories:
substitutions, additions, eliminations, rearrangements.
Substitutions
In a substitution reaction, one atom or group replaces another in a
structure. This type of reaction is commonly observed in saturated
hydrocarbons and aromatics.
+ Na+ -OH
R-X
H2O
R-OH + Na+ X-
an alkyl halide
Ar-H
an aromatic
+ Br2
Fe
Ar-Br + HBr
The mechanisms of the above two substitution reactions are
completely different.
Additions
Addition reactions are found in organic compounds with multiple bonds:
alkenes, alkynes, carbonyl-containing compounds. In this reaction, thecomponent of the multiple bond is lost as new bonds are formed to the
carbon (or other atomic) centers.
H
H
C
H
C
+ Br2
H
H
H
H
C
C
Br Br
Bromine adds to the alkene (ethene).
H
Eliminations
These reactions are the reverse of addition reactions. In an elimination
reaction, a molecule loses atoms or groups from adjoining carbon (or
other atomic) centers, forming a multiple bond.
H
H
H
C
C
H
Br
H
H
C
H + KOH
H
+ K+ Br- + H2O
C
H
The above reaction is a dehydrohalogenation, loss of HBr, of an alkyl
halide to form an alkene.
Rearrangements
In rearrangement reactions, there is a reorganization of
the atoms or groups in a structure.
H+
In the presence of acid, the alkene on the left rearranges to the alkene
on the right.
Reaction Mechanisms and Chemical Intermediates
Reaction mechanisms are detailed descriptions of changes at the
molecular level as reactants become products. Often the reactions
involve a sequence of steps with one or more chemical species
called intermediates that are formed and consumed.
Chemical intermediates typically are not stable structures that can be
put in a bottle. Many exist for very short times (10-6 - 10-9 seconds).
We will explore how structural and electronic influences affect the
stability of chemical intermediates and, thereby, control the path that
reactions follow.
Bond Making and Bond Breaking Processes:
Heterolysis and Homolysis
A covalent bond may break by either of two different processes:
heterolysis or homolysis.
Heterolysis (Gr: hetero- "different" + lysis-"cleavage")
A:B
+
A + B
Double-barbed arrow
is used to show movement
of an electron pair.
ions
Homolysis (Gr: homo-"the same" + lysis)
A:B
A. + B.
radicals
Single-barbed arrow
is used to show movement
of a single electron.
Cleavage of Covalent Bonds
• Homolysis
• Heterolysis
9
• Heterolytic reactions almost always occur at
polar bonds
• The reaction is often assisted by formation of a new
bond to another molecule
10
Acid-Base Reactions
The Brønsted-Lowry Definitions
In 1923 the Danish chemist Johannes Brønsted (1879-1947) and the
English chemist Thomas Lowry (1874-1936) independently proposed
that acids and bases be defined in terms of their ability to give up or
accept a proton.
base
acid
(accepts H+)
(gives up H+)
+
:Cl :
: :
H-Cl :
:
+
: :
:
H-O:
H
+
H-O-H
H
conjugate acid conjugate base
of HCl
of H2O
• Example
• Aqueous hydrogen chloride and aqueous sodium
hydroxide are mixed
• The actual reaction is between hydronium and
hydroxide ions
12
Conjugate Pairs
The species formed when an acid loses a proton is the conjugate base.
The species formed when a base accepts a proton is the conjugate acid.
Conjugate pairs are
pairs of chemical
species that only differ
in H+ such as
H2O/H3O+
HCl/Cl-
Diprotic Acids
Diprotic acids have two acidic protons that may be released. An
example is sulfuric acid where the two H+ are sequentially released.
(1) H2SO4
+
H2O
H3O+ +
HSO4-
(2) HSO4-
+
H2O
H3O+ +
SO42-
Step 1 occurs completely while step 2 proceeds to only about 10% in water.
The Leveling Effect of Water
The "leveling effect of water" refers to the limitations imposed on acid
and base strengths in water because of the acid-base reactions of water.
The strongest acids and bases that can exist in water are H3O+ and HO-.
When stronger acids or bases are added to water, they immediately
react with water to produce these species.
H-A
+
H2O
fast
H3O+ + A-
fast
HO-
very strong acid
B:-
+ H2O
very strong base
+ H-B
chemical species in water
A practical consequence of the leveling effect of water is that most
organic reactions that involve very strong bases or acids are carried out
in nonaqueous solvents such as ethers or hydrocarbons.
Acids and Bases
The Lewis Definitions of Acids and Bases
In 1923, the same year the Brønsted-Lowry definitions were introduced,
G.N. Lewis (1875-1946) broadened the definitions of acids and bases:
An acid is a chemical species that can accept an electron pair.
A base is a chemical species that can donate an electron pair.
Lewis acid + Lewis base reactions
H+
+
:F:
+
:Br-Br:
NH4+
: :
: :
F
F-B-F
F
Br
- +
Br-Fe:Br-Br:
Br
: :
: :
:NH3
: :
F
F-B
F
Br
Br-Fe
Br
+
Note: Any electron
deficient atom can act as a
Lewis acid by accepting an
electron pair. This idea is
important in organic
reaction mechanisms.
Quiz Chapter 3 Section 2
In the reaction below, identify the acid, base, conjugate acid and
conjugate base.
CH3 C C H + NaNH2
acid
base
CH3 C CNa + NH3
conjugat e base conjugate acid
Solut ion
Use t he Brønst ed-Low ry definit ions. Look for t he species t hat
giv e up and accept t he proton..
• Lewis Definition of Acids and Bases
• Lewis Acid: electron pair acceptor
• Lewis Base: electron pair donor
• Curved arrows show movement of electrons to
form and break bonds
17
Opposite Charges Attract and React
• BF3 and NH3 react based on their relative
electron densities
• BF3 has substantial positive charge on the boron
• NH3 has substantial negative charge localized at the
lone pair
18
Carbon Bond Heterolysis Processes:
Carbocations and Carbanions
When a bond to carbon is broken heterolytically, the carbon may
carry either a positive (carbocation) or negative (carbanion) charge..
Modes of Heterolytic Bond Cleavage
Heterolysis of Bonds to Carbons: Carbanions and Carbocations
• Reaction can occur to give a carbocation or carbanion
depending on the nature of Z
• Carbocations have only 6 valence electrons and a positive
charge
20
Formal Charges
Carbocations have only six electrons in the valence or bonding level and
are electron-deficient. The charge on carbon can be determined by a
simple calculation.
Step One: Determine how many of the valence electrons
"belong" to carbon.
For each pair of bonding electrons, one "belongs" to carbon.
All nonbonding electrons in the valence level "belong" to carbon.
C+
a carbocation
In a carbocation, only 3 of the 6 bonding
electrons "belong" to the carbon. There
are no nonbonding electrons.
Step Two: Compare the number of "owned" electrons in the valence
level of the bonded state with the number in the atomic state.
Atomic carbon has 4 electrons in the valence level. Since in the
carbocation, only 3 electrons belong to carbon, there is a deficiency of
one electron (a formal charge imbalance at carbon). Therefore, there is a
formal charge of +1 on carbon.
Carbanions
Carbanions have 6 bonding and two nonbonding electrons in the valence
level, and have a formal charge of -1.
C:
a carbanion
Three of the bonding electrons
and the two nonbonding electrons
belong to carbon.
Calculation of Formal Charge
Since atomic carbon has 4 electrons in the valence level while 5 of the
8 valence electrons in the carbanion belong to carbon, there is a
surplus of one electron in the bonded state. The formal charge is -1.
• Carbanions have 8 valence electrons and a negative
charge
• Organic chemistry terms for Lewis acids and bases
• Electrophiles (“electron-loving” reagents ): seek
electrons to obtain a stable valence shell of electrons
– Are electron-deficient themselves e.g. carbocations
• Nucleophiles (“nucleus-loving” reagents): seek a proton
or some other positively charged center
– Are electron-rich themselves e.g. carbanions
23
The Reactivity of Chemical Intermediates
Carbocations as Electrophiles
Since carbocations are electron-deficient in the valence level, they are
strong Lewis acids. Carbocations react rapidly with Lewis bases,
species that are capable of donating electrons. Carbocations are called
electrophiles.
C+
+ :B
-
CB
C B
a carbocation
(an electrophile)
An Example
C+
+
O H
H
Lewis Base
H
-H+
H C O H
H H
H
H C O H
H
an alcohol
Carbanions as Nucleophiles
Carbanions are Lewis bases. They donate an electron pair to Lewis
acids such as H+ and other electropositive atoms and groups.
Carbanions are called nucleophiles.
-
C
+
carbanion
(a nucleophile)
C
-
carbanion
(a nucleophile)
 
H-A
C H +A
Lewis acid
+
 
C L
Lewis acid
C C
+ L-
The Curved Arrow Formalism
The above bonding changes are illustrated with a formalism called
curved arrows where the arrow shows the direction of electron flow
from nucleophile to electrophile.
The curved arrow begins at an electron pair (bonding or
nonbonding) and moves towards an electron-deficient atom or
group (Lewis acid).
direction of electron flow
nucleophile
electrophile
Examples
C C
nucleophile
+
 
H Cl
-
electrophile
nucleophile
: :
:O H
H
O
CH3C O : +
=
:
=
O  
CH3C O H +
electrophile
C C H + Cl
H
+
:O H
H
Quiz Chapter 3 Section 4
:
: :
In the reaction below, use the curved arrow formalism to show
movement of an electron pair.
+
+
C6H5CH2-OCH3
C6H5CH2 + CH3OH
H
electrophile nucleophile
Lewis acid
Lewis base
Identify the electrophile and nucleophile in the reaction.
Identify the Lewis acid and Lewis base in the reaction.
The Use of Curved Arrows in Illustrating Reactions
• Curved arrows show the flow of electrons in a reaction
• An arrow starts at a site of higher electron density (a
covalent bond or unshared electron pair) and points to
a site of electron deficiency
• Example: Mechanism of reaction of HCl and water
28
Acid and Base Strengths: Ka and pKa
Organic acids (carboxylic acids) typically are weaker acids than the
mineral acids (HCl, H2SO4). While the latter dissociate completely in
water, carboxylic acids, such as acetic acid, dissociate only to a small
degree.
+ H2O
O
CH3CO-
=
=
O
CH3CO-H
+ H3O+
In a 0.1 M solution of acetic acid in water at 25oC, only about 10% of
the acetic acid molecules are dissociated to the acetate and hydronium
ions.
Acidity Constant, Ka
An equilibrium is established for the dissociation of acetic acid in water
with an equilibrium constant, Keq, that is expressed as
Keq =
[H3O+] [CH3CO2-]
[CH3COOH] [H2O]
However, for dilute solutions, the concentration of water (~55.5 M)
does not change significantly during the reaction (the activity of water
a = 1) , and a new equilibrium constant is defined: the acidity
constant, Ka.
Ka = Keq[H2O] =
[H3O+] [CH3CO2-]
[CH3COOH]
Values of Ka are tabulated. For acetic acid, at 25oC, Ka = 1.76 x 10-5.
The units of Ka are mol/L, but they are usually omitted.
Magnitude of Ka and Acid Strength
For the general case,
HA + H2O
Ka =
[H3O+]
[A- ]
[HA]
H3O+ + AThe larger the magnitude of Ka, the
more the equilibrium is shifted to the
products side, and the greater the acid
strength of HA.
Strengths of Acids and Bases
• Ka and pKa
• Acetic acid is a relatively weak acid and a 0.1M
solution is only able to protonate water to the
extent of about 1%
• The equilibrium equation for this reaction is:
32
• Dilute acids have a constant concentration of water
(about 55.5 M) and so the concentration of water can
be factored out to obtain the acidity constant (Ka)
– Ka for acetic acid is 1.76 X 10-5
• Any weak acid (HA) dissolved in water fits the general
Ka expression
– The stronger the acid, the larger the Ka
33
• Acidity is usually expressed in terms of pKa
– pKa is the negative log of Ka
– The pKa for acetic acid is 4.75
• The larger the pKa, the weaker the acid
34
Acidity and pKa
Because Ka values range over many powers of 10, a logarithmic scale
is used where
pKa = -Log Ka
in analogy with
pH = -Log [H3O+]
For acetic acid at 25oC,
pKa = -Log Ka = -Log (1.76 x 10-5) = -(-4.75) = 4.75
Note the inverse relationship between the magnitude of Ka and the
magnitude of pKa because of the negative sign in the definition of
pKa. The smaller the magnitude of Ka, the larger the magnitude
of pKa.
A Quantitative Measure of Acid Strength:
Examples of Ka and pKa Values
acid
Ka
pKa
H-F
6.8 x 10--4
3.17
strongest acid
C6H5COOH
H-CN
6.5 x 10-5
4.9 x 10-10
4.19
9.3
weakest acid
"Acidity" in Organic Chemistry
In organic chemistry, the term "acidity" is broadly used and is a
fundamental property of all organic materials. Many reactions
depend on the relative acidities of organic compounds which direct
proton transfers, and the course of reactions. The range of acidities of
organic compounds is enormous, with the Ka values covering many
powers of 10.
Conjugate Pair Relationships
The acidity of a compound and the basicity of its conjugate base are
related by
Ka x Kb = Kw = 1.0 x 10-14
or
pKa + pKb = 14
for
conjugate pairs
This relationship means the stronger the acid strength of HA, the
weaker the base strength of A-. Also, if the value of Ka for HA is
known, the value of Kb for A- may be calculated.
Some Selected Acids and their Conjugate Bases
The defining equilibrium for acid strength (Ka or pKa) is
HA + H2O
H3O+ + A-
either from direct measurement or from indirect methods.
Clearly, the magnitude of Ka for the very weak "acids" in the table
on the next slide cannot be directly measured in water.
acid
H2SO4
HCl
approximate pKa
strongest
acid
C6H5SO3H
-10
-7
conjugate base
weakest
base
C6H5SO3-
-6.5
H3O+
HNO3
HSO4Cl-
-1.74
H2O
-1.4
NO3-
CF3COOH
0.18
CF3CO2-
CH3COOH
4.75
CH3CO2-
NH4+
H2O
CH3CH2OH
HC
CH
9.2
NH3
15.7
HO-
16
CH3CH2O-
25
HC
C:-
H2
NH3
35
38
HNH2-
CH2=CH2
44
H2C=CH-
CH3CH3
weakest
acid
>50
strongest
base
CH3CH2-
40
Self-Ionization of Water
Even in pure water, there are still finite concentrations of H3O+
and HO- because of the self-ionization of water:
H-O: + :O-H
H
H
acid
: :
:
:
:
+
H O-H + H-O:
H
base
In pure water, at 25oC, [H3O+] = [HO-] = 1.0 x 10-7 M. Since the
concentration of water in pure water is 55.5 M,,
Ka =
[H3O+] [HO-]
[H2O]
(10-7) (10-7)
=
(55.5)
and the pKa of water is 15.7.
= 1.8 x 10-16
The Acidity of the Hydronium Ion, H3O+
The acidity of H3O+ is calculated from the defining equilibrium:
H3O+
Ka =
+
H3O+
H2O
[H3O+] [H2O]
[H3O+]
= [H2O]
+
H2O
= 55.5 M
and the pKa of H3O+ = -Log (55.5) = -(1.74) = -1.74
Acidity Order:
pKa
H3O+ >> H2O
-1.74
15.7
Predicting Base Strengths
According to the Brønsted-Lowry definitions of acids and bases,
B: H + A:
H-A + B:
acid
conjugate
acid
base
and the conjugate pairs are:
conjugate
base
H-A / A: and B: / B:H
An important relationship between
acids and their conjugate bases is:
Ka x Kb = Kw = 1.0 x 10-14
pKa + pKb = 14
where Ka is the acidity constant of an acid, and Kb is the basicity
constant of its conjugate base defined as
Kb
A:H
+
HO
A: + H2O
Because of this intrinsic relationship between acids and their conjugate
bases, the relative order of acid strengths (pKa values) automatically
provides a relative order of the strengths of the conjugate bases.
Predicting Base Strengths
acids
pKa
CH3NH3+
C6H5NH3+
methylaminium ion
anilinium ion
10.6
4.6
acid strengths
conjugate
bases
CH3NH2
methylamine
C6H5NH2
aniline
predicted base strengths
Predicting the Strengths of Bases
• The stronger the acid, the weaker its conjugate
base will be
• An acid with a low pKa will have a weak conjugate
base
• Chloride is a very weak base because its conjugate
acid HCl is a very strong acid
45
• Methylamine is a stronger base than ammonia
• The conjugate acid of methylamine is weaker than
the conjugate acid of ammonia
46
Quiz Chapter 3 Section 5
:
=
:
= =
Based on the information provided, which organic compound below is the
stronger acid?
O
O
CH3SNH2
CH3CNH2
O
acetamide
methanesulfonamide
pKa
~15
~10
Which anion is the stronger base?
O CH3SNH
O
: :
= =
: :
=
O CH3CNH
Predicting the Outcome of Acid-Base Reactions
• Acid-base reaction always favor the formation
of the weaker acid/weaker base pair
• The weaker acid/weaker base are always on the
same side of the equation
• Example
• Acetic acid reacts with sodium hydroxide to greatly
favor products
48
Another Example: Amines in Aqueous HCl Solution
Amines (RNH2) dissolve in hydrochloric acid solution because of a
highly favorable equilibrium:
stronger base
H
+
R-N-H ClH
pKa ~
~ 9-10
stronger acid
weaker acid
:
:
R-NH2
+
H-O-H Cl+
H
pKa = -1.74
+ H2O
weaker base
products are more stable
than reactants
This equilibrium highly favors the products on the right, which
means that water-insoluble amines may be dissolved in hydrochloric
acid solution.
Quiz Chapter 3 Sections 5 and 6
Based on the information provided below, determine if each reaction below
will proceed as written.
O
N H
pKa
CH3COOH
NH4+ Cl-
O
phthalimide
8.3
O
O
N K+ + NH4+ Cl-
NO
N H +
WA
SB
SA
O
O
N K+ + CH3COOH
SB O
NH3 + KCl
O
O
WB
acetic acid
4.7
ammonium chloride
9.2
SA
YES
N H + CH3COO- K+
WA
O
WB
• Water Solubility as a Result of Salt Formation
• Organic compounds which are water insoluble can
sometimes be made soluble by turning them into
salts
• Water insoluble carboxylic acids can become
soluble in aqueous sodium hydroxide
• Water insoluble amines can become soluble in
aqueous hydrogen chloride
51
Acid Strength: Structure and Reactivity Relationships
There are several important relationships between acid
strength (Ka) and structure of compounds.
Ka
+
H-A
H2O
H3O+
+
A-
(1) Bond Strength
The acidity of H-A decreases as the bond strength increases. An
example is the order of acidity of the hydrogen halides (H-X).
pKa
3.2
H-F
H-Cl -7
H-Br -9
H-I
-10
increasing
acid
strength
increasing
H-X bond
strength
Order of Base Strength
F- > Cl - > Br- > I-
The Relationship Between Structure and Acidity
• Acidity increases going down a row of the periodic
table
• Bond strength to hydrogen decreases going down
the row and therefore acidity increases
53
• Acidity increases from left to right in a row of the
periodic table
• Increasingly electronegative atoms polarize the
bond to hydrogen and also stabilize the conjugate
base better
54
• Overview of Acidity Trends
55
Influences of Electronegativity on Acidity
Increasing electronegativity of the central atom enhances
acidity in two ways:
(1) The polarity of the covalent A-H bond increases with increasing
electronegativity of atom A, resulting in a more electropositive H.
As positive charge density on H increases, its reactivity towards a
base increases.
reactivity
towards B:-
 

H2N-H
HO-H
slower
faster
(2) As the electronegativity of A increases, the stability of the anion, A-,
increases, resulting in a larger Ka for the equilibrium:
A-H + H2O
H3O+ + A-
The Effect of Orbital Hybridization
The acidity of hydrocarbons varies considerably according to the type
of hybrid orbital projected by the carbon atom.
hydrocarbon
HH
C
H
C
HH
H
ethane
pKa
H
>50
H
C
C
H
H C
C H
H
ethene
44
ethyne
25
increasing acidity
hybrid orbital
projected by C
sp3
sp2
sp
The acidity of the hydrocarbon increases with the amount of scharacter in the hybrid orbital projected by the carbon atom.
Hybrid Orbital Electronegativities
Electrons in the 2s atomic orbital are more stable than electrons in a 2p
atomic orbital. S-type orbitals are centered on the nucleus which
enhances interaction between the positively charged nucleus and the
orbital electrons. P-type orbitals are projected out from the nucleus so
there is less stabilization of the orbital electrons by the positive nuclear
charge.
The shape and directionality of a hybrid orbital reflects the mix
of atomic orbitals. The greater the degree of s-character, the
shorter the hybrid orbital (less directionality), and the greater
the stabilization of the orbital electrons by the positive nucleus.
This influence can be called hybrid orbital electronegativity.
H
C
sp3
C
sp2
H
C
sp
increasing hybrid orbital electronegativity
H
The Effect of Hybridization on Acidity
• Hydrogens connected to orbitals with more s
character will be more acidic
• s orbitals are smaller and closer to the nucleus than
p orbitals
• Anions in hybrid orbitals with more s character will
be held more closely to the nucleus and be more
stabilized
59
Inductive Effects
Alkanes are nonpolar compounds. For example, the C-C
bond in ethane has no net polarization of charge .
H
H
H C C H
H
H
Because of symmetry,
there is no polarization
of charge within ethane.
ethane
(nonpolar)
If an electronegative group is introduced into ethane, the sigma
bond electrons are attracted (polarized) towards the electronegative
group, as shown in ethyl fluoride.
H
F 
H C C H
H  H
ethyl fluoride
The highly electronegative
fluorine atom polarizes the
sigma electrons in the C-F bond.
Charge Polarization through Sigma Bonds
H H
 
H C C F
H H
 
The strongest polarization occurs in
the C-F bond resulting in a net
positive charge on the -carbon.
In turn, the positive charge on the carbon polarizes electrons in the C-C
bond leading to a small positive charge
on the -carbon.
The polarization of charge through sigma bonds due to
electronegativity differences is called the inductive effect.
Inductive effects weaken as the
distance from the substituent
increases.


 
C
C
C
X
Inductive Effects
• Electronic effects that are transmitted through
space and through the bonds of a molecule
• In ethyl fluoride the electronegative fluorine is
drawing electron density away from the carbons
– Fluorine is an electron withdrawing group (EWG)
– The effect gets weaker with increasing distance
62
Energy Changes
The physical world is described in terms of matter and energy.
Matter is the physical stuff around us. It occupies space, has mass,
and can be measured by various methods. Energy is not always
observable.
Types of Energy
Energy is the capacity to do work.
The two general categories are kinetic and potential energy.
Kinetic energy is the capacity to do work by a moving object:
Kinetic Energy = (1/2) mv2 where m is the mass of the object and v is its
velocity. The moving object may be as small as an electron.
Potential energy is stored energy such as the energy of chemical bonds.
Kinetic and potential energy are interchangeable.
An Example: two balls attached to the ends of a spring
We can define the relaxed
position of the spring as
zero kinetic energy and zero
potential energy.
relaxed position
KE = 0 PE = 0
As the spring is stretched,
work is being done on it, and energy is
expanding spring
transferred into the spring. The expanding
KE = + PE = +
spring has kinetic energy (moving mass)
which is being converted into potential energy.
When the spring is at rest in a
stretched position, the KE is
again zero, but the potential
energy is at some finite value.
When the stretched spring is
released, a restoring force
compresses the spring. As the spring
compresses, stored potential energy is
converted into kinetic energy.
stretched spring
KE = 0 PE = ++
compressing spring
KE = + PE = +
Energy Changes in Reactions
• Kinetic energy is the energy an object has because
of its motion
• Potential energy is stored energy
– The higher the potential energy of an object the less
stable it is
• Potential energy can be converted to kinetic energy
(e.g. energy of motion)
65
Chemical Energy
Chemical energy is potential energy. Large increments of potential
energy are in the electronic structures of atoms and molecules, the
chemical bonds, and intermolecular interactions.
Relative Potential Energy and Relative Stabilities
in Chemical Systems
According to the potential energy
diagram, A is less stable than B
by the energy difference shown.
A
Relative Potential Energy
While it is difficult to describe the
absolute amount of energy in a
chemical system, it is possible and
useful to examine the relative
amount of potential energy in
different systems .
higher PE
less stable
E
B
lower PE
more stable
• Potential Energy and Covalent Bonds
• Potential energy in molecules is stored in the form
of chemical bond energy
• Enthalpy Ho is a measure of the change in bond
energies in a reaction
• Exothermic reactions
– Ho is negative and heat is evolved
– Potential energy in the bonds of reactants is more than
that of products
• Endothermic reactions
– Ho is positive and heat is absorbed
– Potential energy in the bonds of reactants is less than
that of products
67
Standard State
The relative potential energies of reactants and products are given as
relative enthalpies or heat content, H. The change in enthalpy in going
from reactants to products is H .
When the change in enthapy during the reaction is given as Ho, the
superscipt o means the reaction was carried out under standard
conditions.
Standard Conditions (25oC)
a gas at 1 atm pressure
a solute in a 1 M solution
a liquid or solid as pure material
Although reactions are usually not carried out under standard
conditions, the calculation of the Ho does give useful information.
• Example : Formation of H2 from H atoms
• Formation of bonds from atoms is always
exothermic
• The hydrogen molecule is more stable than
hydrogen atoms
69
The Equilibrium Constant and Free Energy Changes
Gibbs Free Energy
Late in the 19th century, Josiah Williard Gibbs (1839-1903) proposed
a new function of chemical states that describes the spontaneity of a
chemical reaction. Today, this state function is called the Gibbs Free
Energy or simply the Free Energy, G.
The change in free energy, G, for a chemical reaction indicates if it
is spontaneous, if it proceeds without the input of work from the
surroundings.
• If G is negative, the reaction in the forward direction
proceeds spontaneously.
• If G is zero, the reaction is at equilibrium, there is
no net driving force in the forward or reverse direction.
• If G is positive, the forward reaction will not proceed
without input of work from the surroundings.
• Go encompasses both enthalpy changes (Ho) and entropy
changes (So )
• Ho is associated with changes in bonding energy
– If Ho is negative (exothermic) this makes a negative contribution to
Go (products favored)
• So is associated with the relative order of a system
– More disorder means greater entropy
– A positive So means a system which is going from more ordered to
less ordered
– A positive So makes a negative contribution to Go (products
favored)
• In many cases So is small and Go is approximately equal to
Ho
71
Standard Free Energy Change, Go
Because free energy is a state function, like enthalpy and entropy, it is
possible to calculate the free energy of substances under standard
conditions, Go.
Standard Conditions of Substances
state of matter
standard state
solid
pure solid
liquid
pure liquid
gas
solution
elements
1 atm
1 M concentration
The standard free energy
of formation of an element
is defined as zero.
Standard Free Energy of Formation, Gof
The standard free energy of formation, Gof , of a compound is the
standard free energy change associated with the formation of one
mole of that substance from the constituent elements, with all
reactants and products in their standard states.
The values of Gof for many substances have been tabulated and can be
used to calculate the standard free energy change for reactions.
The standard free energy change for a chemical reaction is
defined as the difference between the standard free energies
of formation of the products and reactants:
Go =
 n Gof
(products) -
 m Gof
(reactants)
where n and m are the mole quantities of products and reactants.
An Example: The Hydrogenation of Ethene to Ethane
CH2=CH2(g) +
Gof
(kJ/mol)
68.2
H2(g)
CH3CH3(g)
0
-32.6
These values are from tabulated data or by definition.
Go =
 n Gof
(products) -
 m Gof
(reactants)
Go = (-32.6 kJ/mol) - (68.2 kJ/mole + 0 kJ/mol)
Go =
-100.8 kJ/mol
Since the standard free energy change is negative, this hydrogenation
reaction will spontaneously occur. If Go is more negative than about
-3 kcal/mol, the reaction is described as going to completion, meaning
more than 99% of the reactants proceed to the products at
equilibrium.
The Acidity of Carboxylic Acids
• Carboxylic acids are much more acidic than alcohols
• Deprotonation is unfavorable in both cases but much less
favorable for ethanol
75
The Relationship Between the Equilibrium Constant and Go
• Go is the standard free energy change in a reaction
• This is the overall energy change of a reaction
• It is directly related to the equilibrium constant of a reaction
– R is the gas constant (8.314 J K-1 mol-1) and T is measured in kelvin (K)
• If Go is negative, products are favored at equilibrium
(Keq >1)
• If Go is positive, reactants are favored at equilibrium
(Keq<1)
• If Go is zero, products and reactants are equally favored
(Keq = 0)
76
Analysis of the Gibbs Equation
Consider the equilibrium,
A +
B
C +
D
where the Gibbs free energy change is given by
G
=
Go
+
2.303RTLog [C][D]
[A][B]
At the start of the reaction, when the reactants A and B are mixed,
the concentrations of C and D are very low, so Q << 1. During these
early stages of the reaction, the reaction quotient makes a favorable
(2.3RTLogQ is negative) contribution to G. As long as G is negative
in value, there is a chemical driving force pushing the reaction in a
forward direction.
At some point, the concentrations of products become greater than the
concentrations of the reactants, and Q > 1. The term 2.3RTLogQ
becomes positive, and no longer contributes favorably to the overall
Gibbs free energy, G. When G = 0, equilibrium is reached and
there is no net chemical driving force in either direction.
At Equilibrium
When G = 0, equilibrium is reached and
G = 0
= Go
+
2.303RTLog Q
Therefore,
Go =
where Keq =
From above,
- 2.303RTLog Q = - 2.303RT Log Keq
[C][D]
, the equilibrium constant for the reaction.
[A][B]
Keq = 10
- Go/2.3RT
Go/RT
= e
Keq =
o/RT
G
-
e
The values of Keq in the table show that over a range of only 6
kcal/mol in free energy, reactions switch from highly favorable
to highly unfavorable at ordinary temperatures.
at 298 K
Keq
Go
-3000 cal 161
-1000
5.4
1000
0.18
3000
0.0062
Quiz Chapter 3 Section 9
The Gibbs free energy change (G) for a reaction
A +
B
C + D
is given as
G = G o +
2.3RTLog
[C] [D]
[A] [B]
What is Go?
This is the Gibbs standard free energy change for the reaction, which is
the free energy change when all reactants and products are in their
standard states.
How can it be calculated?
The Gibbs standard free energy change is the sum of the standard free
energies of formation of the products less the sum of the standard free
energies of formation of the reactants:
G o =
 nG fo (products)
-
 nG of (reactants)
What is the significance of the sign of G?
If it is negative, there is a net chemical driving force and the reaction
will spontaneously proceed in the forward direction.
If it is positive, the net chemical driving force is in the reverse direction
and the reaction will not proceed in the forward direction without the
input of energy from the surroundings.
If it is zero, there is no net chemical driving force in either direction,
the reaction is at equilibrium.
Enthalpy and Entropy
Gibbs defined his function of state in terms of enthalpy or heat (H)
and entropy (S), two thermodynamic state functions.
State functions define the properties of a thermodynamic state.
In a change between two thermodynamic states, the change in value of
the state function is given by the symbol  .
The standard free energy change (Go) is given as
Go =
Ho - TSo
where Ho is the standard enthalpy change and is equal to the
difference in the standard enthalpies of formation (Hof) between
the products and the reactants. This state function is associated with
changes in bonding between reactants and products. Changes in
enthalpy during reactions are measured by calorimetry experiments.
Standard Entropy Change, So
The standard entropy change, So, is the difference in standard
entropies between reactants and products. Entropy is a measure of
the degree of order in a chemical system due to bond rotations, other
molecular motions, and aggregation. The more random a system
(disorder), the greater the entropy. The larger a structure, the more
degrees of freedom it has, and the greater its entropy. The units of
entropy are cal/degree.
The standard entropies of materials (including elements), So, are for
one mole of pure substance at 1 atm pressure usually at 25oC. These
values are measured relative to the reference point which is the
entropy of a perfectly ordered crystal at T = 0 K (absolute zero,
where So is zero.
Note the contribution of entropy to the Gibbs free energy. A positive
value for the change in standard entropy (So ) during a reaction
makes a favorable contribution to Go because of the negative sign in
the Gibbs equation:
Go = Ho - TSo
The Hydrogenation of Ethene: Another Perspective
CH2=CH2(g) +
H2(g)
CH3CH3(g)
The following values are from the Handbook of
Chemistry and Physics.
Gof
Hof
So
(kJ/mol)
(kJ/mol)
(J/K mol)
C2H4(g)
68.1
52.2
219.2
H2(g)
0
0
130.5
C2H6(g)
-32.8
-84.5
229.3
Calculation of Standard Enthalpy Change, Ho
Ho =
Ho
=
 n Hof (products)
-

m Hof (reactants)
(-84,500 J/mol) - (52,200 J/mol) = -136,700 J/mol
Calculation of Standard Entropy Change, So
 n Sof (products)

m Sof (reactants)
So
=
So
= (229.3 J/K mol) - (219.2 J/K mol + 130.5 J/K mol)
So
= (229.3 J/K mol) - (349.7 J/K mol) = -120.4 J/K mol
-
Calculation of Gibbs Free Energy Change, Go
Go = Ho - TSo Using the values above,
Go
Go
= (-136,700 J/mol) - (298) (- 120.4 J/K mol)
= (-136,700 J/mol) - (-35,879 J/mol) = -100,821 J/mol
Go
= -100,821 J/mol, which is in agreement with earlier calculations.
This calculation shows that the favorable Gibbs free energy change in
the hydrogenation of ethene to ethane is driven by a very favorable
Ho , changes in bonding.
The So term actually makes an unfavorable contribution to the
spontaneity of the reaction. Bonding changes in chemical reactions
will often be used to evaluate Ho, and determine the feasibility of a
reaction.
A Second Example: The Hydrogenation of Ethyne to Ethane
HC
CH (g)
+
2H 2(g)
CH 3CH 3(g)
Gof
Hof
So
(kJ/mol)
(kJ/mol)
(J/K mol)
226.6
200.7
C2H2(g)
209
H2(g)
0
C2H6(g)
-32.8
0
-84.5
130.5
229.3
from the Handbook of Chemistry and Physics
Calculations
Ho
= (-84,500 J/mol) - (2 x 0 J/mol + 226,600 J/mol)
Ho
= (-84,500 J/mol) - (226,600 J/mol) = -311,100 J/mol
So
=
(229.3 J/K mol) - (2 x 130.5 J/K mol + 200.7 J/K mol)
So
=
(229.3 J/K mol) - (461.7 J/K mol) = -232.4 J/K mol
Calculation of the Gibbs Free Energy Change
Go = Ho - TSo = (-311,100 J/mole) - (298) (-232.4 J/K mol)
Go
= -311,100 J/mol + 69,255 J/mol) = -241,845 J/mol
Go
=
-241.8 kJ/mole
Again, the spontaneity of this hydrogenation reaction is due to the very
favorable standard enthalpy change (-311.1 kJ/mol) that reflects
differences in the bond strengths of the bonds lost and bonds made
during the reaction.
HC
CH (g)
+
2H 2(g)
tw o relatively w eak tw o strong bonds
are lost
 bonds are lost
CH 3CH 3(g)
four strong bonds
are made
• Explanation based on resonance effects
• Both acetic acid and acetate are stabilized by resonance
– Acetate is more stabilized by resonance than acetic acid
– This decreases Go for the deprotonation
87
The Acidity of Carboxylic Acids
Carboxylic acids are more acidic than alcohols.
=
O
CH3COH
CH3CH2OH
acetic acid
ethanol
Ka = 1.78 x 10-5
pKa = 4.75
Ka = 1.0 x 10-16
pKa = 16
Go = 27.2 kJ/mol
Go = 90.9 kJ/mol
The Go values are calculated from the Ka values that describe
the equilibrium:
HA + H2O
H3O+ + A-
Free Energy Diagrams
The Go values indicate that the ionization of acetic acid is much more
energetically favorable than the ionization of ethanol. In both
reactions, the ion product states are higher in free energy (both Go
values are positive) than the reactant states. These features are shown
in the Free Energy Diagrams below.
change in free energy
CH3CH2O- + H3O+
-
+
CH3CO2 + H3O
o
G = 90.9 kJ/mol
o
G = 27.2 kJ/mol
CH3COOH + H2O
The Go values set
the levels of the
product states.
CH3CH2OH + H2O
In the diagrams, the
reactant states are set
at the same level for
easy comparison.
The ionic state produced from ethanol is much higher in energy than the
ionic state produced from acetic acid. This difference in stability may be
explained by both resonance theory and the inductive effect.
Explanations of the Relative Acidities of Ethanol and Acetic Acid
Resonance
It is generally accepted that a major factor contributing to the acidity of
acetic acid and other carboxylic acids is resonance. Although resonance
contributes to the stability of both the reactant and product states of
acetic acid, the stabilizing influence of resonance is much more
important in the product (ion) state..
:
: O:
+
+
H
O
3
CH3C=O
: :
: :
=
:O:
CH3C-O:
Go
: O:
:
+
CH3C=OH
minor contribution
:
: :
:O:
CH3C-OH
=
Relative Free Energy
major contribution
+ H2O
Resonance reduces the
energy of the
carboxylate ion.
• Explanation based on resonance effects
• Neither ethanol nor its anion is stabilized by resonance
– There is no decrease in Go for the deprotonation
91
• Explanation based on inductive effect
• In acetic acid the highly polarized carbonyl group
draws electron density away from the acidic
hydrogen
• Also the conjugate base of acetic acid is more
stabilized by the carbonyl group
92
Inductive Effects of Other Groups
• The electron withdrawing chloro group makes
chloroacetic acid more acidic than acetic acid
– The hydroxyl proton is more polarized and more acidic
– The conjugate base is more stabilized
93
A Charge Dispersal Mechanism
The greater acidity of chloroacetic acid is attributed to the
electronegativity of the chlorine atom (compared with hydrogen). In
addition to increasing the electropositive character of the protic
hydrogen in the acid, the electronegative chlorine atom helps disperse
the negative charge in the carboxylate anion.
Cl CH2
O
C O H + H2O
O


Cl CH2 C O H + H3O+
Charge dispersal stabilizes ions through either
resonance or the inductive effect.
Quiz Chapter 3 Section 10
Select the more acidic carboxylic acid in each pair below, and explain
your choice.
Cl
CH3CHCOOH
ClCH2CH2COOH
The inductive effect
decreases rapidly with
distance.
Cl
CH3CHCOOH
F
CH3CHCOOH
The inductive effect
increases with
electronegativity.
The Effect of Solvent on Acidity
• 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
– This solvation allows acetic acid to be much more acidic in
water than in the gas phase
96
Effect of Solvent on Acidity
Most acids are much weaker in the gas phase where there is no
solvent to stabilize the ions produced in the product state.
Stabilization of the ions through ion-dipole interactions (a charge
dispersal mechanism) is worth hundreds of kilocalories per mole.
When acetic acid ionizes in water, both the carboxylate anion
and the hydronium ion are stabilized by solvation:
H
O
H
H
H
=
O
CH3-C-O-H + H2O
O
O
H
C
CH3
O
H
O
H
O
H
H
O
H
+
H
O
H
O H
H
O
H
H O
O
H
H
H
H
O
H
H
O H
H H
Solvation is a combination of hydrogen
bonding and charge-dipole interactions.
Gas Phase Acidities
The acidities of many compounds in the gas phase, in the absence of
solvation, have been measured.
H+ + AH-A
The relative acidity order in the gas phase is surprisingly
different from the acidity order in water (magnitude of Ka).
Gas Phase Acidities
OH
NH2
> RS-H >
> H-F > R C C H > CH3CH2-OH > H2O > CH4
most acidic
Because the creation and separation of charge is a very high energy
process (hundreds of kJ/mol), any factors that stabilize charge will
reduce the energy requirement. In the absence of solvation, internal
structural and electronic features determine the reactivity order.
Polarizable atoms and groups enhance acidity because they are able to
disperse charge more effectively than atoms and groups of low
polarizability.
Acidity in Water: A Closer Look
Water is a protic solvent, a solvent capable of hydrogen bonding to
solutes. When acetic acid dissolves in water, water molecules associate
with ("solvate") all solutes: acetic acid, acetate anion, hydronium ion.
O
CH3 C
OH
solvated acetic acid
H
+ H2O
O
H
H
+
CH3
O
C
O
stronger solvation of ions
The strong and ordered interaction between solvent and solute molecules
decreases randomness and creates order, decreasing entropy. When
acetic acid ionizes in water, the So is negative. Therefore, the change in
entropy makes an unfavorable contribution to the Gibbs free energy for
the ionization of acetic acid.
Thermodynamic Parameters
The change in Gibbs standard free energy for the reaction
CH3COOH + H2O
+
H3O + CH3CO-2
o
G = 27.2 kJ/mol
contains both enthalpic and entropic contributions. A negative
o
change in standard entropy ( S ) makes an unfavorable
energetic contribution ( -TS o ) to the Gibbs free energy.
o
G
o
o
= H - TS
The overall spontaneity of a chemical reaction (Go) depends on
enthalpic changes arising from bonding, solvation factors, as well
as entropic changes that measure the change in degree of order
during the reaction.
An Analysis of Acetic Acid and Chloroacetic Acid
The greater acidity of chloroacetic acid is generally explained by the
inductive effect of the electronegative chlorine. In the table below are
the results of a detailed study showing the enthalpic and entropic
contributions to the overall spontaneity of the ionization process for
each acid.
Acid
pKa
o
G
H
S
(kJ/mol)
(kJ/mol)
(J/K mol)
(kJ/mol)
(at 298 K)
o
o
o
- TS
CH3COOH
4.75
+27.2
-0.4
-92.5
+27.6
ClCH2COOH
2.86
+16.3
-4.6
-70.3
+20.9
The enhanced acidity of chloroacetic acid results from a more favorable
Ho (4.2 kJ/mol), and an even more favorable entropy contribution
(-TSo is less positive by 6.7 kJ/mol).
The thermodynamic parameters suggest that while the electronegative chlorine
disperses negative charge and stabilizes the anion (more favorable Ho), the
interaction between the chloroacetate ion and water is not as strong. The weaker
solute-solvent interaction is associated with less ordering of the solvent, and a less
negative standard entropy change (-70.3 vs -92.5).
Organic Compounds as Bases
• Any organic compound containing an atom with a
lone pair (O,N) can act as a base
102
Organic Compounds as Bases
All organic compounds with an unshared electron pair, or
a-bond are potential bases.
+ H-Cl
:
methanol
(base)
: :
:
CH3-O:
H
+
CH3-OH2
+
Cl-
(acid)
When HCl gas is dissolved in methanol, it dissociates by protonation
of the oxygen in the alcohol. This is a general reaction between
alcohols and strong acids (HX, H2SO4) analogous to the ionization of
strong acids in water.
:
R-O:
H
+ H-A
strong acid
+
ROH2 + :A-
Ethers react in a similar way:
+ H-Cl
:
: :
:
CH3-O:
R'
+
R O H +
R'
A
(acid)
Carbonyl compounds are involved in an equilibrium with strong
acids that produces a low concentration of the conjugate acids:
R
+ H-A
R'
weak base
strong acid
+ H
:O
C
R
R'
=
=
:O:
C
+ :A-
conjugate acid
(a very strong acid)
The table that follows compares the base strength of carbonyl compounds,
alcohols, ethers and water. The latter three have about the same base strength
and are comparably protonated by strong acids. Carbonyl compounds are
considerably weaker bases (factor of 104 to 105).
Some Representative pKB and pKBH+ Values
pKB
R-O :
-7
17.5
+
R-O H
R
-3.5
: :
-
H-O :
hydroxide ion
+
H-O H
H
:
16
15.74
-2
-1.74
:
:
:
H-O
H
water
+
RCH2O H
H
:
:
:
:
21
R
ether
RCH2O
H
alcohol
pKBH+
+
R-C=O H
R
:
R-C=O
R
ketone
Conjugate Acid
:
: :
Base
-2
:
H-O
H
water
16
•  Electrons can also act as bases
–  Electrons are loosely held and available for reaction
with strong acids
107
A Mechanism for an Organic Reaction
• The Substitution Reaction of tert-Butyl Alcohol
• All steps are acid-base reactions
– Step 1 is a Brønsted acid-base reaction
– Step 2 is a Lewis acid-base reaction in reverse with
heterolytic cleavage of a bond
– Step 3 is a Lewis acid-base reaction with chloride acting as a
Lewis base and the carbocation acting as Lewis acid
108
109
Acids and Bases in Nonaqueous Solutions
• Water has a leveling effect on strong acids and
bases
• 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
111
Nonaqueous Solvents
Reactions involving strong bases such as sodium amide are run in
nonaqueous solvents such as ethers, hydrocarbons, or liquid ammonia
(NH3, BP -33oC).
RC
terminal alkyne
pKa = 25
+
:
CH
+ -
Na : NH2
sodium amide
liquid NH3
-33oC
RC
-
+
C : Na + :NH3
The solvent for a reaction must be compatible with the acid
and base strengths of the reactants and products.
pKa = 38
• Alkyl lithium reagents in hexane are very strong
bases
– The alkyl lithium is made from the alkyl bromide and
lithium metal
113
Synthesis of Deuterium- and Tritium-Labeled Compounds
• Deuterium (2H) and tritium (3H) are isotopes of hydrogen
• They are used for labeling organic compounds to be able
to track where these compounds go (e.g. in biological
systems)
• An alkyne can be labeled by deprotonating with a suitable
base and then titrating with T2O
114
Introduction of Deuterium and Tritium Labels
by Acid-Base Reactions
Deuterium (2H) and tritium (3H) are used as isotopic "labels" in
mechanistic and other studies. These labels may be introduced into
organic structures by acid-base chemistry using deuterium oxide (D2O)
or tritium oxide (T2O).
D
Li
+ D2O
very fast
hexane
phenyllithium
(stronger base)
+ DO-Li+
deuteriobenzene
(stronger acid)
(weaker acid) (weaker base)