ACIDS AND BASES IN

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CHAPTER 3 ACIDS AND BASES IN
ORGANIC CHEMISTRY
3.1 INRODUCTION

• The important of acid-base reaction
• is a simple, fundamental reaction.
• enable you to see the mechanism of the reaction
• illustrate the process of bond breaking and bond
  making
• examine important ideas about the relationship
  between the
• structure of molecules and their
  reactivity.
• 5. illustrate the important role solvents play
  in chemical reactions
• 6. find something familiar in General
  chemistry.

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3.1A THE BRNSTED-LOWRY DEFINATION
OF ACIDS AND BASES
According to the brnsted-Lowry theory, an acid is
a substance that can donate a proton, and a base
is a substance that can accept a proton. For
example
Hydrogen chloride, a very strong acid, transfers
its proton to water. Water acts as a base and
accepts the proton.
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The molecule or ion forms when an acid loses
its proton is called the conjugate base of that
acid. Such as the chloride ion. The molecule
or ion that forms when a base accepts a proton is
called the conjugate acid of that base. Such as
the hydronium ion.
Other strong acids that completely transfer a
proton when dissolved in water are hydrogen
iodide, hydrogen bromide, and sulfuric acid.
The proton transfer is stepwise in sulfuric
acid, the first proton Transfer completely, the
second only to the extent of 10.
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When an aqueous solution of sodium hydroxide is
mixed with an aqueous solution of hydrogen
chloride, the reaction that occurs is between
hydronium and hydroxide ions.
The net reaction is simply
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3.1B THE LEWS DEFINITION OF ACIDS AND BASES
Lewis proposed that acids be defined as
electron-pair acceptors and bases be defined as
electro-pair donors. For example
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The Lewis theory , by virtue of its broader
definition of acids, allows acid-base theory to
include all of the Brnsted-Lowry reactions aqnd ,
as we shall see, a great many others.
Any electron-deficient atom can act as a Lewis
acid. Many compounds containing group IIIA
elements such as aluminium are Lewis acids
because group IIIA atoms have only a sextet of
electrons in their outer shell. Many other
compounds that have atoms with orbitals also act
as Lewis acids. For example
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3.2 THE USE OF CURVED ARROWS IN
ILLUSTRATING REACTIONS
The curved arrows is commonly used by organic
chemists to show the direction of electron flow
in a reaction. Besides it is a useful Method for
indicating which bonds form and which bonds break.
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Now, lets us illustrate some of the basic ideas
of the curved-arrow notation with simple Lewis
acid-base reactions
The following acid-base reactions gives other
examples of the use of the curved-arrow notation
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3.3 THE STRENGTH OF ACIDS AND BASES Ka AND pKa
When acetic acid dissolves in water, the
following reaction dose not proceed to completion
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3.3A THE ACIDITY CONATANT Ka
Experiments show that in a 0.1M solution of
acetic acid at 25? only about 1 of the acetic
acid molecules ionize by transferring
their protons to water. The reaction is a
equilibrium, we can describe it with an
expression for the equilibrium constant.
For dilute aqueous solutions, the concentration
of water is essentially constant, so the
equilibrium constant can be expressed with the
acidity constant (Ka).
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A large value of Ka means the acid is a strong
acid, and a small value of Ka means the acid is a
weaker acid. If the Ka is greater than 10, the
acid will be completely dissociated in water.
3.3B ACIDITY AND pKa
Chemists usually express the acidity constant,
Ka, as its negative logarithm, pKa.
pKa - logKa
For acetic acid the pKa is 4.75
Notice that there is an inverse relationship
between the magnitude of the pKa and the
strength of the acid
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The larger the value of the pKa, the weaker is
the acid.
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3.3C PREDICTING THE STRENGTH OF BASES
The principle which allows us to estimate the
strengths of bases the stronger the acid, the
weaker will be its conjugated base.
So relate the strength of a base to the pKa of
its conjugate acid. The larger the pKa of the
conjugate acid, the stronger is the base.
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Amines are like ammonia in that they are weak
bases. Dissolving in water bring about the
following equilibrium.
Dissolving methylamine in water causes the
establishment of a similar equilibrium.
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3.4 PREDICTING THE OUTCOME OF ACID-BASE
REACTIONS
There is a general order of a acidity and basity
for some of the common acids and base. For
example acetic acid has a pKa 4.76 and
carboxylic acids generally have pKa values near
35. The pKa of ethyl alcohol is 16, and alcohols
generally have pKa values near 1518, and so on.
But there are exceptions.
The general principle of predicting whether or
not an acid-base will occur acid-base reaction
always favor the formation of the weaker acid and
the weaker base.
Using the principle, we can predict that a
carboxylic acid will react with aqueous NaOH in
the following way because the reaction will lead
to the formation of the weaker acid and the
weaker base.
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Because of the acidity water-insoluble carboxylic
acids dissolve in aqueous sodium hydroxide they
do so by reacting to form water- soluble salts.
The water-insoluble amines dissolve readily in
hydrochloric acid because the acid-base reaction
convert them to soluble salts.
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3.5 THE RELATIONSHIP BETWEEN STRUCTURE
AND ACIDIY
The strength of the bond to the proton is the
dominating effect if we compare compounds in a
vertical column of the periodic table.
HF HCl HBr
HI pKa 3.2 pKa -7 pKa -9 pKa
-10
The important factor is the strength of the H-X
bond, the stronger the bond, the weaker the acid.
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We see the same trend of acidities and basicities
in other vertical columns of the periodic table.
For example
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When we compare compounds in the same horizontal
row of the periodic table, the dominant factor
becomes the electronegativity of the atom bond
to the hydrogen.
Lets see an example of CH4, NH3, H2O and HF.
These compounds are all hydrides of first-row
elements and electronegaticity increases across a
row of the periodic table from left to right.
Because fluorine is the most electronegative, the
bond in H-F is Most polarized, and the proton in
H-F is the most positive, so H-F is the most
acidic.
CH3H NH2H HOH FH pKa 48
pKa 38 pKa 15.7 pKa 3.2
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Because H-F is the strongest acid, its conjugate
base, the fluoride ion will be the weakest base,
fluorine is the most electronegative atom and it
accommodates the negative charge most readily.
3.5A THE EFFECT OF HYBRIDIZATION
The protons of ethyne are more acidic than
those of ethene, which in turn, are more acidic
than of ethane.
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We can explain this order of acidities on the
basis of the hybridiza- tion state of carbon in
each compound. Electrons of 2s orbitals have
Lower energy than those of 2p orbitals because
in 2s orbitals tend to be much closer to the
nucleus than electrons in 2p orbitals.
With hydrid, therefore, having more s charater
means that the electrons of the anion will, on
the average, be lower in enegy,and the anion will
be more state.
Now we can see how the order of relative
acidities of ethyne, ethene, and ethane parallels
the effective electronegativity of the carbon
atom in each compound
Relative acidity
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The more the electronegative, the most positive
the hydrogens. So ethyne donates a proton to a
base more readily. And in the same way, the
ethynide ion id the weaker base because the more
electro- negative carbon of ethyne is best able
to stabilize the negative charge.
Relative basicity
3.5B INDUCTIVE EFFECTS
The carbon-carbon bond of ethane is completely
nonpolar because at each end of the bond there
are two equivalent methyl groups.
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This is not the case with ethyl fluoride, however
Inductive effect the effect of the polarization
of the carbon-carbon bond results from an
intrinsic electron-at-tracting ability of the
fluorine that is transmitted through space and
through the bonds of the molecule.
Inductive effect weaken steadily as the distance
from the substituent increase. In this instance,
the positive charge that the fluorine imparts to
C-1 is greater than that imparted to C-2 because
the fluorine is closer to C-1.
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Inductive effects help us to understand why
carboxylic acids are much more acidic than
alcohols.
The key to the much greater acidity of acetic
acid is the power electro-attracting inductive
effect of its carbonyl group(CO) when compare to
the CH2 group in the corresponding position of
ethyl alcohol.
The carbonyl group of acetic acid, because its
bears a large positive charge, adds its
electron-attracting effect to that of the oxygen
of the hydroxyl group attached to it this makes
the hydroxyl proton much more positive than the
proton of the alcohol.
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This greater positive charge on the proton of the
acid means that the proton separates more
readily. The electron-attracting effect of the
carbonyl group also stabilizes the anion that
forms from the carboxylic acid, and, therefore,
the carboxylate ion is weaker base than the
ethoxide ion.
The acid-strengthen effect of electron-attracting
group can also be shown by comparing the
acidities of acetic acid and chloroacetic acid.
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The greater acidity of chloroacetic acid can be
attributed, in part, to the extra
electron-attracting inductive effect of the
electronegative chlorine atom.
So, dispersal of charge always makes a species
more stable, and any factor that stabilizes the
conjugate base of an acid will increase the
strength of the acid.
3.6 THE RELATIONSHIP BETWEEN THE EQUILIBRIUM
CONSTANT AND THE STANDARD FREE-ENERGY
CHANGE, ?Gº
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An important relationship exists between the
equilibrium constant and the standard
free-energy change that accompanies the reaction.
?Gº -2.303RTlogKeq
R is the gas constant and equals 1.987 cal
/(Kmol) T is the absolute temperature in
kelvins (K).
A negative value of ?Gº is associated with
reactions that favor the formation of products
when equilibrium is reached, and a positive value
of ?Gº is associated with reactions for which the
formation of the products at equilibrium is
unfavorable.
The free-energy change(?Gº) has two components,
the enthalpy change(?Hº) and the entropy
chang(?Sº). The relationship between these
thermodynamic quantities is
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?Gº ?Hº - T?Sº
?Hº is associated with changes in bonding that
occur in a reaction. If, collectively, stronger
bonds are formed in the products than exited in
the starting materials, then ?Hº will be
negative,vice versa. A negative value for ?Hº ,
therefore, will contribute to making ?Gº
negative, and will, consequently favor the
formation of products.
The more random a system is, the greater is its
entropy. A positive entropy change ( ?Sº ) is
always associated with a change from a more
ordered system to a cess. We can see from the
above relationship that a positive entropy change
makes a negative contri- bution to ?Gº and is
energetically favorable for the formation
of products.
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3.7 THE EFFECT OF THE SOLVENT ON A ?Hº CIDITY
In the absence of a solvent, most acids are far
weaker than they are in solution. For example
in gas phase acetic acid is estimated to have a
pKa of about 130! The reason when an acetic acid
molecule donates a proton to a water molecule in
the gas phase, the ions that are formed are
oppositely charged particles and these particles
must become separated.
A protic solvent solvent that has a hydrogen
atom attached to a strongly electronegative
element such an oxygen or nitrogen.
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Molecules of a protic solvent can form hydrogen
bonds to the un- shared electron pairs of oxygen
atoms of an acid and its conjugate base, but
they may not stabilize both equally. For example
if acetic acid in aqueous solution, hydrogen
bonding to CH3COO is much stronger than to
CH3COOH because the water molecules are more
attracted by the negative charge.
This differential solvation, moreover, has
important consequences for the entropy charge
that accompanies the ionization. Solvation of any
species decreases the entropy of the solvent
because the solvent mole- cules become much more
ordered as they surround molecules of the solute.
The following table list the thermodynamic values
for the dissociation of acetic and chloroacetic
acids in H2O at 25?
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3.8 ACID AND BASES IN NONAQUEOUS AOLUTIONS
The reaction of amide ion in aqueous solution as
the following
This example illustrate what is called the
leveling effect of the solvent. The solvent
here, water, converts any base stronger than a
hydroxide ion to a hydroxide ion by donating.
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But we can convert ethyne to its conjugate base
by treating it with sodium amide in liquid
ammonia.
Most alkynes with a proton attached to a triply
bonded carbon have pKa values of about 25,
therefore, all react with sodium amide in liquid
ammonia in the same way that ethyne dose
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Alcohols are often used as solvents for organic
reactions because being somewhat less polar than
water, they dissolve less polar organic compound.
Another example
Alkyllithiums react as though they contained
alkanide ions, and being the conjugate bases of
alkanes, alkanide ions are the strongest bases
that we shall encounter.
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Alkyllithiums can be easily prepared by allowing
an alkyl bromide to react with lithium metal in
an ether solvent.
General reaction
Specific reaction
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3.9 ACID-BASE REACTIONS AND THE SYNTHESIS OF
DEUTERIUM AND TRITIUM-LABELED COMPOUNDS
Chemists often use compounds in which
deuterium(²H) or tritium(³H) atoms have replaced
one or more hydrogen atoms of the compound as a
method of labeling or identifying particular
hydrogen atoms.
The extra mass and additional neutrons associated
with a deuterium or tritium atom often makes its
position in a molecule easy to locate by certain
spectroscopic.
One way to introduce a deuterium or tritium atom
into a specific location in a molecule is
through the acid-base reaction that takes place
when a very strong base is treated with D2O or
T2O. For example
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3.10 SOME IMPORTANT TERMS AND CONCEPTS
A Brnsted-Loery acid a substance that donate a
proton A Brnsted-Lowry base a substance that
can accept a proton. A Lewis acid an
electron-pair acceptor, a Lewis base is an
electron-pairdoner. Curved
arrows used to show the direction of electron
flow when mechanisms
are written. The strength of an acid can be
expressed by its acidity constant, Ka.
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Or by its pKa

pKa -logKa
The strength of a base inversely related to the
strength of its conju-
gated acid the weaker the
conjugated acid,
the stronger the base. The outcome of the
acid-base reactionpredicted on the basis of the

principle that acid-base reactions proceed
toward
equilibrium so as to favor formation
of the weaker acid
and the weaker base. An inductive effectreflects
the ability of a substituent to attract or
release electrons
because of its electronegativity.
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Dispersal of electrical charge always makes a
chemical entity
more stable. The relationship
between Keq and the standard free-energy
change(?Gº )
?Gº -2.303RTlogKeq The relationship of ?Gº ,
?Hº and ?Sº
?Gº ?Hº - T?Sº A protic solvent one that has
a hydrogen atom attached to a
strongly electronegative atom.