Chemical Reactivity

Organic chemistry encompasses a very large number of compounds ( many
millions ), and our previous discussion and illustrations have focused
on their structural characteristics. Now that we can recognize these
actors ( compounds ), we turn to the roles they are inclined to play in
the scientific drama staged by the multitude of chemical reactions that
define organic chemistry.
We begin by defining some basic terms that will be used frequently as this subject is elaborated.

Chemical Reaction: A transformation resulting in a change of
composition, constitution and/or configuration of a compound ( referred
to as the reactant or substrate ).
Reactant or Substrate: The organic compound undergoing change in a
chemical reaction. Other compounds may also be involved, and common
reactive partners ( reagents ) may be identified. The reactant is often
( but not always ) the larger and more complex molecule in the reacting
system. Most ( or all ) of the reactant molecule is normally
incorporated as part of the product molecule.
Reagent: A common partner of the reactant in many chemical reactions.
It may be organic or inorganic; small or large; gas, liquid or solid.
The portion of a reagent that ends up being incorporated in the product
may range from all to very little or none.
Product(s) The final form taken by the major reactant(s) of a reaction.
Reaction Conditions The environmental conditions, such as temperature,
pressure, catalysts & solvent, under which a reaction progresses
optimally. Catalysts are substances that accelerate the rate ( velocity
) of a chemical reaction without themselves being consumed or appearing
as part of the reaction product. Catalysts do not change equilibria
positions.

Classifying Organic Chemical Reactions

If you scan any organic textbook you will encounter what appears to be
a very large, often intimidating, number of reactions. These are the
"tools" of a chemist, and to use these tools effectively, we must
organize them in a sensible manner and look for patterns of reactivity
that permit us make plausible predictions. Most of these reactions
occur at special sites of reactivity known as functional groups, and
these constitute one organizational scheme that helps us catalog and
remember reactions.
Ultimately, the best way to achieve proficiency in organic chemistry is
to understand how reactions take place, and to recognize the various
factors that influence their course.
This is best accomplished by perceiving the reaction pathway or mechanism of a reaction.

1. Classification by Structural Change

First, we identify four broad classes of reactions based solely on the
structural change occurring in the reactant molecules. This
classification does not require knowledge or speculation concerning
reaction paths or mechanisms.
The letter R in the following illustrations is widely used as a symbol
for a generic group. It may stand for simple substituents such as H– or
CH3–, or for complex groups composed of many atoms of carbon and other
elements.

Four Reaction Classes
Addition

Elimination


Substitution

Rearrangement


In an addition reaction the number of σ-bonds in the substrate molecule
increases, usually at the expense of one or more π-bonds. The reverse
is true of elimination reactions, i.e.the number of σ-bonds in the
substrate decreases, and new π-bonds are often formed. Substitution
reactions, as the name implies, are characterized by replacement of an
atom or group (Y) by another atom or group (Z). Aside from these
groups, the number of bonds does not change. A rearrangement reaction
generates an isomer, and again the number of bonds normally does not
change.
The examples illustrated above involve simple alkyl and alkene systems,
but these reaction types are general for most functional groups,
including those incorporating carbon-oxygen double bonds and
carbon-nitrogen double and triple bonds. Some common reactions may
actually be a combination of reaction types. The reaction of an ester
with ammonia to give an amide, as shown below, appears to be a
substitution reaction ( Y = CH3O & Z = NH2 ); however, it is
actually two reactions, an addition followed by an elimination.

The addition of water to a nitrile does not seem to fit any of the
above reaction types, but it is simply a slow addition reaction
followed by a rapid rearrangement, as shown in the following equation.
Rapid rearrangements of this kind are called tautomerizations.


Additional examples illustrating these classes of reaction may be examined by Clicking Here

2. Classification by Reaction Type

At the beginning, it is helpful to identify some common reaction types
that will surface repeatedly as the chemical behavior of different
compounds is examined. This is not intended to be a complete and
comprehensive list, but should set the stage for future elaborations.
Acidity and Basicity

It is useful to begin a discussion of organic chemical reactions with a
review of acid-base chemistry and terminology for several reasons.
First, acid-base reactions are among the simplest to recognize and
understand. Second, some classes of organic compounds have distinctly
acidic properties, and some other classes behave as bases, so we need
to identify these aspects of their chemistry. Finally, many organic
reactions are catalyzed by acids and/or bases, and although such
transformations may seem complex, our understanding of how they occur
often begins with the functioning of the catalyst.
Organic chemists use two acid-base theories for interpreting and planning their work: the Brønsted theory and the Lewis theory.
Brønsted Theory

According to the Brønsted theory, an acid is a proton donor, and a
base is a proton acceptor. In an acid-base reaction, each side of the
equilibrium has an acid and a base reactant or product, and these may
be neutral species or ions.

H-A + B–) A–) + B-H
(acid1) (base1) (base2) (acid2)

Structurally related acid-base pairs, such as {H-A and A–)} or {B–)
and B-H} are called conjugate pairs. Substances that can serve as both
acids and bases, such as water, are termed amphoteric.

H-Cl + H2O Cl–) + H3O(+)
(acid) (base) (base) (acid)

H3N: + H2O NH4(+) + HO(–)
(base) (acid) (acid) (base)

The relative strength of a group of acids (or bases) may be
evaluated by measuring the extent of reaction that each group member
undergoes with a common base (or acid). Water serves nicely as the
common base or acid for such determinations. Thus, for an acid H-A, its
strength is proportional to the extent of its reaction with the base
water, which is given by the equilibrium constant Keq.


H-A + H2O

H3O(+) + A–)

Since these studies are generally extrapolated to high dilution,
the molar concentration of water (55.5) is constant and may be
eliminated from the denominator. The resulting K value is called the
acidity constant, Ka. Clearly, strong acids have larger Ka's than do
weaker acids. Because of the very large range of acid strengths
(greater than 1040), a logarithmic scale of acidity (pKa) is normally
employed. Stronger acids have smaller or more negative pKa values than
do weaker acids.



Examples of Brønsted Acid-Base Equilibria

Acid-Base Reaction Conjugate
Acids Conjugate
Bases Ka pKa
HBr + H2O H3O(+) + Br(–) HBr
H3O(+) Br(–)
H2O 105 -5
CH3CO2H + H2O H3O(+) + CH3CO2(–) CH3CO2H
H3O(+) CH3CO2(–)
H2O 1.77*10-5 4.75
C2H5OH + H2O H3O(+) + C2H5O(–) C2H5OH
H3O(+) C2H5O(–)
H2O 10-16 16
NH3 + H2O H3O(+) + NH2(–) NH3
H3O(+) NH2(–)
H2O 10-34 34

In all the above examples water acts as a common base. The last example
( NH3 ) cannot be measured directly in water, since the strongest base
that can exist in this solvent is hydroxide ion. Consequently, the
value reported here is extrapolated from measurements in much less
acidic solvents, such as acetonitrile.

Since many organic reactions either take place in aqueous environments
( living cells ), or are quenched or worked-up in water, it is
important to consider how a conjugate acid-base equilibrium mixture
changes with pH. A simple relationship known as the
Henderson-Hasselbach equation provides this information.

When the pH of an aqueous solution or mixture is equal to the pKa of an
acidic component, the concentrations of the acid and base conjugate
forms must be equal ( the log of 1 is 0 ). If the pH is lowered by two
or more units relative to the pKa, the acid concentration will be
greater than 99%. On the other hand, if the pH ( relative to pKa ) is
raised by two or more units the conjugate base concentration will be
over 99%. Consequently, mixtures of acidic and non-acidic compounds are
easily separated by adjusting the pH of the water component in a two
phase solvent extraction.
For example, if a solution of benzoic acid ( pKa = 4.2 ) in benzyl
alcohol ( pKa = 15 ) is dissolved in ether and shaken with an excess of
0.1 N sodium hydroxide ( pH = 13 ), the acid is completely converted to
its water soluble ( ether insoluble ) sodium salt, while the alcohol is
unaffected. The ether solution of the alcohol may then be separated
from the water layer, and pure alcohol recovered by distillation of the
volatile ether solvent. The pH of the water solution of sodium benzoate
may then be lowered to 1.0 by addition of hydrochloric acid, at which
point pure benzoic acid crystallizes, and may be isolated by filtration.

Basicity

The basicity of oxygen, nitrogen, sulfur and phosphorus compounds
or ions may be treated in an analogous fashion. Thus, we may write
base-acid equilibria, which define a Kb and a corresponding pKb.
However, a more common procedure is to report the acidities of the
conjugate acids of the bases ( these conjugate acids are often "onium"
cations ). The pKa's reported for bases in this system are proportional
to the base strength of the base. A useful rule here is: pKa + pKb = 14.
We see this relationship in the following two equilibria:

Acid-Base Reaction Conjugate
Acids Conjugate
Bases K pK
NH3 + H2O NH4(+) + OH(–) NH4(+)
H2O NH3
OH(–) Kb = 1.8*10-5 pKb = 4.74
NH4(+) + H2O H3O(+) + NH3 NH4(+)
H3O(+) NH3
H2O Ka = 5.5*10-10 pKa = 9.25

Tables of pKa values for inorganic and organic acids ( and bases)
are available in many reference books, and may be examined here by
clicking on the appropriate link:
Inorganic Acidity Constants
Organic Acidity Constants
Basicity Constants

Although it is convenient and informative to express pKa values
for a common solvent system (usually water), there are serious
limitations for very strong and very weak acids. Thus acids that are
stronger than the hydronium cation, H3O(+), and weak acids having
conjugate bases stronger than hydroxide anion, OH(–), cannot be
measured directly in water solution. Solvents such as acetic acid,
acetonitrile and nitromethane are often used for studying very strong
acids. Relative acidity measurements in these solvents may be
extrapolated to water. Likewise, very weakly acidic solvents such as
DMSO, acetonitrile, toluene, amines and ammonia may be used to study
the acidities of very weak acids. For both these groups, the reported
pKa values extrapolated to water are approximate, and many have large
uncertainties. A useful table of pKa values in DMSO solution has been
compiled from the work of F.G. Bordwell, and may be reached by Clicking
Here.
Lewis Theory

According to the Lewis theory, an acid is an electron pair
acceptor, and a base is an electron pair donor. Lewis bases are also
Brønsted bases; however, many Lewis acids, such as BF3, AlCl3 and Mg2+,
are not Brønsted acids. The product of a Lewis acid-base reaction, is a
neutral, dipolar or charged complex, which may be a stable covalent
molecule. As shown at the top of the following drawing, coordinate
covalent bonding of a phosphorous Lewis base to a boron Lewis acid
creates a complex in which the formal charge of boron is negative and
that of phosphorous is positive. In this complex, boron acquires a neon
valence ****l configuration and phosphorous an argon configuration. If
the substituents (R) on these atoms are not large, the complex will be
favored at equilibrium. However, steric hindrance of bulky substituents
may prohibit complex formation. The resulting mixture of non-bonded
Lewis acid/base pairs has been termed "frustrated", and exhibits
unusual chemical behavior.
Two examples of Lewis acid-base equilibria that play a role in chemical reactions are shown in equations 1 & 2 below.

In the first example, an electron deficient aluminum atom bonds to
a covalent chlorine atom by sharing one of its non-bonding valence
electron pairs, and thus achieves an argon-like valence ****l octet.
Because this sharing is unilateral (chlorine contributes both
electrons), both the aluminum and the chlorine have formal charges, as
shown. If the carbon chlorine bond in this complex breaks with both the
bonding electrons remaining with the more electronegative atom
(chlorine), the carbon assumes a positive charge. We refer to such
carbon species as carbocations. Carbocations are also Lewis acids, as
the reverse reaction demonstrates.
Many carbocations (but not all) may also function as Brønsted acids.
Equation 3 illustrates this dual behavior; the Lewis acidic site is
colored red and three of the nine acidic hydrogen atoms are colored
orange. In its Brønsted acid role the carbocation donates a proton to
the base (hydroxide anion), and is converted to a stable neutral
molecule having a carbon-carbon double bond.

A terminology related to the Lewis acid-base nomenclature is often
used by organic chemists. Here the term electrophile corresponds to a
Lewis acid, and nucleophile corresponds to a Lewis base.
Electrophile: An electron deficient atom, ion or molecule that has an
affinity for an electron pair, and will bond to a base or nucleophile.
Nucleophile: An atom, ion or molecule that has an electron pair that
may be donated in bonding to an electrophile (or Lewis acid).