Organic Acid and Bases

1. • An organic acid is an organic compound with acidic properties. The most common
organic acids are the carboxylic acids, whose acidity is associated with their
carboxyl group –COOH. ... The relative stability of the conjugate base of the acid
determines its acidity.
• An organic base is an organic compound which acts as a base. Organic bases are
usually, but not always, proton acceptors. They usually contain nitrogen atoms,
which can easily be protonated, for example amines have a lone pair of electrons
on the nitrogen atom and can thus act as proton acceptors.
• According to Arrhenius acids are the substances that produce hydrogen ions (H+)
when they dissolve in water .Because acids produce ions in water ,they are also
electrolyte. Bases are the substances that produce (OH-) ions when they dissolve in
water.
ORGANIC ACID AND BASES

2. HCl(g) + H2O H(aq)
+ + Cl-
(aq)
NaOH (S) + H2O Na(aq)
+ + OH-

3. BRONSTED –LOWRY ACIDS AND BASES
• Bronsted –lowry acid can donate a hydrogen ion ,(H+) and Bronsted – lowry
base can accept a hydrogen ion.
• CONJUGATE ACID –BASE PAIR : A conjugate acid – base pair consists
of molecules or ions related by loss of one H+ by an acid and gain of one H+
by a base .Every acid – base reaction contains 2 conjugate acid- base pairs
because an H+ is transferred in both the forward and reverse directions.
• HF(aq) + H2O(l) F(aq) + H3O (aq)

4. STRENGTHS OF ACID AND BASES
• In the process called as dissociation , an acid or a base separates into ions in water.
The strength of an acid is determined by moles oh hydronium ion produced for
each mole of acid that dissociates .The strength of base is determined by moles of
OH- that are produced for each mole of base that dissolves .
• STRONG AND WEAK ACIDS : Strong acids are examples for strong electrolytes
because they donate H + easily that their dissociation in water is essentially
complete.( EX; HCl).
• Weak acids are weak electrolytes because they dissociate slightly in water forming
only a small amount of hydronium ions .A weak acid has a strong conjugate base
,which is why the reverse reaction is more prevalent.(Acetic acid)
• CH3COOH(Aq) + H2O(l) C2H3O2
-
(aq) + H3O+
(aq)
• (Acetic acid ) (acetate ion)

5. STRONG AND WEAK BASE
• Strong base dissociates completely in water because these strong bases
asre ionic compounds ,they dissociate in water to give an aqueous solution
of metal ions and hydroxide ions.
• KOH(S) + H2O K+
(aq) + OH-
(aq)
• Weak base are weak electrolytes that are poor acceptors of hydrogen ions
and produce very few ions in solution .A typical example for weak base ,
ammonia NH3 , in aqueous solution , only a few ammonia molecules
accept hydrogen ions to form NH4
+ and OH- .
• NH3 (g) + H2O(l) NH4
+
(aq) + OH-
(aq)
• (Ammonia ) ( Ammonium)

6. DISSOCIATION CONSTANTS FOR ACIDS AND BASES
• Strong acids dissociate completely in water and the reaction is not considered to
be an equilibrium situation ,whereas weak base partially dissociates in water and
the ionic product reach equilibrium with the undissociated weak acid molecules.
• HCHO2(aq) + H2O(aq) H3O +
(aq) + CHO2
-
(aq)
• (formic acid ) (formate ion )
• Ka = [H3O+] [CHO-
2 ]
• [HCHO2]
• = 1.8 × 10-4
• The dissociation constant for formic acid is small, which confirms that the
equilibrium mixture of formic acid contains only small amounts of the products .

7. EFFECT OF SOLVENT ON ACID AND BASE STRENGTH
• LEVELING EFFECT :Leveling effect or solvent leveling refers to the effect
of solvent on the properties of acids and bases. The strength of a strong acid is
limited ("leveled") by the basicity of the solvent. Similarly the strength of a strong
base is leveled by the acidity of the solvent. When a strong acid is dissolved in
water, it reacts with it to form hydronium ion (H3O+).[2]
• Any acid that is stronger than H3O+ reacts with H2O to form H3O+. Therefore, no
acid stronger than H3O+ exists in H2O.The strongest base in water is OH.
• A weakly basic solvent has less tendency than a strongly basic one to accept
a proton. Similarly a weak acid has less tendency to donate protons than a strong
acid. As a result, a strong acid such as perchloric acid exhibits more strongly acidic
properties than a weak acid such as acetic acid when dissolved in a weakly basic
solvent. On the other hand, all acids tend to become indistinguishable in strength
when dissolved in strongly basic solvents owing to the greater affinity of strong
bases for protons.

8. HCl(g) + H2O H 3 O(aq)
+ + Cl-
(aq) (1)
CH3COOH + H2 O H3 O+ + CH3COO- (2)
CH3COOH + NH3 NH4 + CH3COO- (3)
• In the first reaction the HCl is strong acid and water is a weak base and
HCl readily gives proton to the leading to formation of hydronium ion and
here the dissociation is said to be complete . Hydronium ion formed levels
the acidic strength of HCl and water is said to be leveling solvent .
• In the second reaction Acetic acid being a weak acid has less tendency to
donate a proton and water being a weak base has tendency to accept the
proton so partial dissociation takes place ,hence water here is called as
Differentiating solvent .

9. In third reaction
• Acetic acid being weak acid has less tendency to donate a proton whereas
ammonia pulls the protons from the acetic acid and hence called leveling
solvent.
• Due to this acetic acid is a strong acid in liq . Ammonia and a weak acid in water.
• HF in water partially dissociates so water here is called as differentiating solvent
,whereas HF completely dissociates in liq. Ammonia so here ammonia is calles as
levelling solvent.
• Acetic acid acts as leveling solvents for bases and differentiating solvents for for
mineral acid.

10. Types of solvent on the basis of proton interaction
• (i) Protophilic solvents: Solvents which have greater tendency to accept
protons, i.e., water, alcohol, liquid ammonia, etc.
• (ii) Protogenic solvents: Solvents which have the tendency to produce
protons, i.e., water, liquid hydrogen chloride, glacial acetic acid, etc.
• (iii) Amphiprotic solvents: Solvents which act both as protophilic or
protogenic, e.g., water, ammonia, ethyl alcohol, etc.
• (iv) Aprotic solvents: Solvents which neither donate nor accept protons,
e.g., benzene, carbon tetrachloride, carbon disulphide, etc.

11. EFFECT OF STRUCTURE OF ORGANIC COMPOUND ON
ACID AND BASE STRENGTH
• Bond Strengths
• In general, the stronger the A–H or B–H+ bond, the less likely the bond is
to break to form H+ ions and thus the less acidic the substance. This effect
can be illustrated using the hydrogen halides:
• The trend in bond energies is due to a steady decrease in overlap between
the 1s orbital of hydrogen and the valence orbital of the halogen atom as
the size of the halogen increases. The larger the atom to which H is
bonded, the weaker the bond.

12. Stability of the Conjugate Base
• Whether we write an acid–base reaction as AH ⇌ A- + H+ or as BH ⇌B +
H+, the conjugate base (A- or B) contains one more lone pair of electrons
than the parent acid (AH or BH+). Any factor that stabilizes the lone pair on
the conjugate base favors dissociation of H+ and makes the parent acid a
stronger acid
• Inductive Effects
• Atoms or groups of atoms in a molecule other than those to
which H is bonded can induce a change in the distribution of
electrons within the molecule. This is called an inductive
effect, and, much like the coordination of water to a metal
ion, it can have a major effect on the acidity or basicity of the
molecule. For example, the hypohalous acids (general formula
HOX, with X representing a halogen) all have a hydrogen atom
bonded to an oxygen atom.

13. HOX(aq) H(aq)
+ + OX-
(aq)
• As the electronegativity of X increases, the distribution of electron
density within the molecule changes: the electrons are drawn more
strongly toward the halogen atom and, in turn, away from the H in the O–
H bond, thus weakening the O–H bond and allowing dissociation of
hydrogen as H+ .
• Acidity increases with increase in oxygen atoms. Because oxygen is the
second most electronegative element, adding terminal oxygen atoms
causes electrons to be drawn away from the O–H bond, making it weaker
and thereby increasing the strength of the acid.

14. pKa for acids
• The pKa value is one method used to indicate the strength of an acid. pKa is the
negative log of the acid dissociation constant or Ka value. A lower pKa value
indicates a stronger acid. That is, the lower value indicates the acid more fully
dissociates in water.
• pKb for base : pKb is the negative base-10 logarithm of the base dissociation
constant (Kb) of a solution. It is used to determine the strength of a base or alkaline
solution. pKb = -log10Kb.
• A large Kb value indicates the high level of dissociation of a strong base. A lower
pKb value indicates a stronger base. pKa and pKb are related by the simple
relation: pKa + pKb = 14.

15. Reference : Timberlake ,General organic And Biological Chemistry.
Swathi S Rao
198229