INTRODUCTION TO IONIC MECHANISMS
PART I: FUNDAMENTALS OF BRONSTED-LOWRY ACID-BASE CHEMISTRY
HYDROGEN ATOMS AND PROTONS IN ORGANIC MOLECULES - A hydrogen atom that has lost its only electron is sometimes
referred to as a proton. That is because once the electron is lost, all that remains is the nucleus, which in the case of hydrogen consists of
only one proton.
The large majority of organic reactions, or transformations, involve breaking old bonds and forming new ones. If a covalent bond is broken
heterolytically, the products are ions. In the following example, the bond between carbon and oxygen in the t-butyl alcohol molecule breaks
to yield a carbocation and hydroxide ion.
H3C
CH3
H3C
OH
+
H3C
OH
CH3
H3C
Hydroxide
ion
A tertiary
carbocation
The full-headed curved arrow is being used to indicate the movement of an electron pair. In this case, the two electrons that make up the
carbon-oxygen bond move towards the oxygen. The bond breaks, leaving the carbon with a positive charge, and the oxygen with a negative
charge. In the absence of other factors, it is the difference in electronegativity between the two atoms that drives the direction of electron
movement. When pushing arrows, remember that electrons move towards electronegative atoms, or towards areas of electron deficiency
(positive, or partial positive charges). The electron pair moves towards the oxygen because it is the more electronegative of the two atoms.
If we examine the outcome of heterolytic bond cleavage between oxygen and hydrogen, we see that, once again, oxygen takes the two electrons
because it is the more electronegative atom. Hydrogen is left with only a positive charge. In other words, it becomes a proton.
H3C
H3C
H3C
H3C
O
H
H3C
O
+
H
H3C
An alkoxide
ion
Hydrogen ion,
or proton
BRONSTED-LOWRY ACIDS AND “ACIDIC PROTONS” - A hydrogen bonded to a very electronegative atom makes for a highly
polar bond. The dipole moment favors electron density around the more electronegative atom, leaving the hydrogen with a partial positive
charge. This bond is different from other bonds in the molecule because of its propensity to break into a negative ion and a positive hydrogen
ion. This propensity is driven by the tendency of the more electronegative atom to take up the electrons that make up the covalent bond. In
fact, one can write resonance structures for such molecule that show the bond in question already broken. The t-butyl alcohol molecule can
be used again to illustrate this point.
H3C
H3C
H3C
O
H3C
H
O
H
H3C
H3C
(II)
t-butyl alcohol (I)
The greatest contributor to the hybrid is obviously structure I because it is neutral. Structure II has charge separation and therefore is a minor
contributor. However, the significance of structure II is that it shows the negative character of the oxygen and the positive character of the
hydrogen, and therefore the polarity of this bond. A better representation of the hybrid could be structure III, which shows the oxygen with
partial negative character, and the hydrogen with partial positive character.
H3C
H3C
δ−
O
δ+
H
(III)
H3C
In the Bronsted-Lowry theory of acids and bases, an acid is a hydrogen ion donor, or proton donor, and a base is a hydrogen ion acceptor,
or proton acceptor. Hydrogen atoms that have a substantial degree of partial positive charge (i.e. low electron density around them) are
commonly referred to as acidic protons. In the example above, the hydrogen bonded to oxygen is considered to be acidic, and the molecule
as a whole is considered a Bronsted acid because it has a propensity to release a hydrogen ion, or proton.
ACID-BASE REACTIONS AS “PROTON” TRANSFERS - When a Bronsted acid (or simply acid) reacts with a Bronsted base
(or simply base) a proton is transferred from the acid to the base. This results in formation of another acid, called the conjugate acid, and
another base, called the conjugate base. For example, when hydroxide ion (a base) reacts with hydrogen chloride (an acid), a new acid
(water) is formed. Water is then the conjugate acid of hydroxide ion. Likewise, a new base (chloride ion) is formed. Chloride ion is then the
conjugate base of hydrogen chloride. The reaction is an equilibrium process because the new acid and the new base can react together to
revert to the original reactants. Therefore we can also say that hydroxide ion is the conjugate base of water, and that hydrogen chloride is
the conjugate acid of chloride ion. This relativity of concepts is characteristic of the Bronsted-Lowry acid-base theory.
H
+ H
O
Cl
δ+
hydroxide
ion
BASE1
H
δ−
O
H
+
water
hydrogen
chloride
ACID2
+ ACID1
Cl
chloride
ion
+
BASE2
conjugate acid-base pairs
Conjugate acid-base pairs differ only by a proton. Other examples of conjugate acid-base pairs are H2O / H3O+ and NH3 / NH4+. The above
reaction also shows the direction of electron movement. In acid-base reactions, electron movement always originates at the base and
moves towards the acidic proton in the acid. The base is the electron-rich species. We can identify bases because they usually have an
atom with unshared electron pairs (lone pairs). Sometimes this atom also carries a negative charge, but this is not a requirement. Likewise,
we can identify the acid because it is the molecule that has acidic protons (hydrogens that carry a strong partial positive charge). The following
are examples of other acids and bases. The acidic protons are shown in red, and the basic atoms in blue. Keep in mind that the concept of
acid or base is always relative in the Bronsted theory. Some molecules such as water can act as acids or as bases, depending on who they
are reacting with. We will expand on this later.
Bronsted acids
O
δ+ δ−
H
H3C
F
δ− δ+
O
δ+
H
δ−
O
H
δ+ δ−
O
H
S
H
H
N
H
N
H
H
(H2SO4)
H
H
O
O
(CH3COOH)
Bronsted bases
O
H
O
H
H3C
O
CH3
O
H
THE SCALE OF ACIDITY: pKa VALUES - Many acid-base reactions take place in water, one of the most universal solvents. Water
also has the dual capability of acting as a proton donor or as a proton acceptor. It makes sense, then, to develop a scale of acidity based on
the behavior of the substance of interest towards water. Since most acid-base reactions are equilibrium processes, the equilibrium constant
of the reaction between an acid (or base) and water forms the basis for the pKa scale. Most general and organic chemistry textbooks contain
adequate discussions of this parameter. What is of interest to us here is the relationship between pKa and acidity. The equation
pKa = -log Ka = log(1/Ka)
shows that such relationship is inverse. The stronger the acid (i.e. the higher its acidity constant Ka), the lower its pKa value, and viceversa.
Tables of pKa values usually show the acids and their conjugate bases arranged by order of decreasing (or increasing) acidity. Chemistry
students should become proficient at reading and using data presented in such tables.
PREDICTING EQUILIBRIUM IN ACID-BASE REACTIONS - Equilibrium in acid-base reactions always favors the weaker
side. In the following example the pKa values for the substances acting as acids are shown under their structures. Equilibrium favors the left
side because the substances on the left are the weaker acid and the weaker base.
CN + NH3
38
weaker
acid
HCN +
NH2
9.2
stronger
acid
We can arrive at the same conclusion looking at the bases. The strength of bases is measured by the pKa of their conjugate acids. To
understand how this works, we must remember that the relative strengths of the acid and the base in a conjugate pair hold an inverse relationship.
The stronger the acid, the weaker its conjugate base, and viceversa. Therefore, the relationship between the pKa of an acid and the strength
of its conjugate base is direct: the stronger the base, the higher the pKa value of its conjugate acid. We can look at the same reaction
again, except that now we’re focusing on the bases, and arrive at the same conclusion that equilibrium favors the left side.
CN
+ NH3
pKa of
conj. acid = 9.2
weaker base
HCN
+
NH2
pKa of
conj. acid = 38
stronger base
The weaker acid and the weaker base are always on the same side. If you arrive at a different conclusion, something is not right.
IDENTIFYING ACIDIC PROTONS - The most general principle ruling acid strength can be stated thus: strong acids have relatively
stable conjugate bases. In general, the more stable the conjugate base, the stronger the acid. An important thing to remember is that stability
and reactivity are inverse. The more stable a substance is, the less reactive it is, and viceversa. Therefore, another way of stating the rule
above is by saying that strong acids have weak conjugate bases. HCl and H3O+ are strong acids. Accordingly, the corresponding conjugate
bases, Cl- and H2O, are weak (very stable). Chloride ion is stable because the negative charge resides on a very electronegative atom. Water
molecule is one of the most stable substances known.
How do we know which proton is the most acidic in a molecule (such as acetic acid) that contains more than one type of proton? Remember
that the higher the degree of positive character on the proton, the more acidic it is. Examination of a pKa table reveals some trends for acidic
protons. The following guidelines can be used to predict acidity.
1. Hydrogens directly attached to very electronegative atoms such as oxygen, sulphur, and the halogens carry a substantial degree of acidity.
δ+ δ−
H
δ+ δ−
Cl
H
O
O
S
O
δ−
δ+
H
H
S
H
O
2. Hydrogens attached to a positively charged nitrogen, oxygen, or sulfur are acidic. The high electronegativity of these atoms makes them
uncomfortable with the positive charge. They seek to diffuse the charge among the neighboring atoms by withdrawing electron density from
them. This can be shown by drawing resonance structures as shown.
H
H
H
N
H
H
H
H
H
N
H
O
H
H
O
H
H
A
B
A
CH3 SH2
A
CH3 SH H
B
In all cases structure B reveals the positive character of hydrogen, and therefore its acidic nature.
B
H
3. As evidenced by the pKa values of alkanes and alkenes, hydrogens attached to carbon are of very low acidity. Such substances are not
normally considered acids at all. However, some hydrocarbons can be weakly acidic if their conjugate bases are stable ions. This can happen
in the following cases.
a) There is one or more electronegative atoms near the proton under consideration. The inductive effect of these electronegative atoms
leaves the hydrogens in the vicinity deprived of electron density, and therefore with partial positive character.
δ+
δ−
Cl
δ−
δ+
H δ+
Cl
CH3
−
Br δ
H3C
Cl
δ+ H
C
−
N δ
CH3
−
δ+
δ
b) A hydrogen atom bonded to a carbon which is in turn bonded to another carbon that carries a partial or a full positive charge is acidic.
CH3
CH3
H
H3C
C
H3C
H
H
CH2
H
carbocations
δ−
O
O
O
H
H3C
δ+ CH3
acetone
H3C
C
H
H3C
CH2 H
H
The acidity of the protons shown becomes apparent in elimination reactions (chapter 6) and in the chemistry of enols (chapter 22), when
the presence of a base leads to formation of alkenes or enolate ions through a step involving a proton transfer.
c) The conjugate base is resonance-stabilized. This effect is most important when there is another factor enhancing the acidity, such as
the presence of a dipole or electronegative atom (as in the nitrile functional group, –CN). Otherwise resonance stabilization alone is
not enough to dramatically increase the acidity of a hydrogen attached to carbon (as in toluene, where the pKa is only 40).
R
CH2 C
-H
N
R
CH
C
N
H2C
nitrile, pKa = 25
CH3
C
N
conjugate base
-H
(write all the resonance structures ;-)
CH2
Toluene, pKa = 40
d) The hydrogen is attached to an sp-hybridized carbon. Hybridization effects on acidity are discussed in chapter 9. The trends in
hybridization can be extended to oxygen and nitrogen besides carbon, as in the example on the right.
H
H
CH3 C
CH
Propyne, pKa = 25
H
C
N
O
more acidic than
O
H
pKa = 9.2
proton attached
2
to sp oxygen
proton attached
to sp3 oxygen
SOME ACIDS CAN ACT AS BASES AND VICEVERSA - The relative nature of Bronsted acid-base terminology becomes apparent
when we consider substances that can act as either proton donors or acceptors. When two such substances react, how can we predict which
will be the proton donor (acid) and which will be the proton acceptor (base)? The answer is that the stronger acid will force the other
substance to act as a base. In other words, the substance with the lower pKa will act as the acid, and the other as the base.
In the first example below, methanol is forced to act as a proton acceptor by the strong sulfuric acid. However, in the second example methanol
is the proton donor because it is a stronger acid than the amide ion.
CH3OH
+
CH3OH2 + HSO4
pKa = -10
pKa = 15
CH3OH
H2SO4
+
NH2
CH3O
+
NH3
PART II: IDENTIFYING BRONSTED BASES AND CENTERS OF REACTIVITY
IN ORGANIC MOLECULES. ELECTRON MOVEMENT IN IONIC REACTIONS.
FUNCTIONAL GROUPS AND REACTIVITY SITES IN ORGANIC MOLECULES - As explained before in relation to functional
groups, the very definition of this term is as a reactive center, or site, in the molecule. We can view an organic molecule as consisting of
two major structural categories: the basic carbon skeleton, and functional groups.
The basic carbon skeleton is an alkane-like structure, made up of only sp3 carbons and hydrogen atoms, and therefore only sigma (single)
bonds. It constitutes the framework that supports the reactive sites, or functional groups. The carbon skeleton is for most practical purposes
considered a nonpolar part of the molecule. It is also considered to be nonreactive, except under very harsh, extreme, or special conditions.
Alkanes contain no functional groups, but as soon as their structure is chemically modified in any way that leads to relatively stable products,
functional groups are introduced. Some of the simplest modifications that introduce functional groups are the creation of a p-bond and the
introduction of a heteroatom. Technically speaking, heteroatoms are atoms other than carbon and hydrogen. The term is most commonly
used in reference to nonmetals of the second and third rows. That is, N, O, S, P, and the halogens.
The introduction of p-bonds, as in alkenes and alkynes, does not affect the polarity of the molecule very much, but it affects the flexibility of it
by restricting free rotation around carbon-carbon bonds. The presence of a p-bond introduces a center of reactivity by increasing the electron
density at that site. At the same time, “mobile,” or “available” electrons are introduced. Remember that p-bonds are weaker than s-bonds,
and therefore p-electrons are easier to move for use in chemical reactions.
C
C
The p-bond is a high electron density region.
p-electrons can be used for chemical reactions under the right conditions.
The introduction of heteroatoms such as oxygen almost invariably introduces a dipole, since most of these elements are more electronegative
than carbon and hydrogen. By introducing a dipole, an electron imbalance is created that results in formation of a low electron density area
and a high electron density area. The low electron density area will then be reactive towards electron-rich molecules. The high electron density
area will be reactive towards electron-poor (deficient) molecules.
Identifying reactive sites in organic molecules amounts to identifying areas of low or high electron density. In the case of carboncarbon double or triple bonds (alkenes and alkynes), the region is an area of high electron density. Such molecules typically act as electron
donors, even though the p-bond must be broken before those electrons become available. If we’re not dealing with p-bonds, then we must
identify the strongest dipoles in the molecule and the areas of low or high electron density they create, such as in the examples below.
δ−
−
Oδ
δ+ C
O
H3C
H
δ+
δ+
Obviously, when it is a hydrogen atom that is positioned at the d+ end of the dipole, we have an acidic proton that can engage in reactions
with Bronsted bases. We have already learned to identify acidic protons, but how do we identify basic sites in molecules?
IDENTIFYING BRONSTED BASES - Basic sites in organic molecules are areas of high electron density. We’ve already talked
about p-bonds as examples of such areas. However, the large majority of Bronsted bases comprises molecules that have atoms with
nonbonding electrons (lone pairs). The following trends can be observed.
1. Nonbonding electrons are the most readily available for reaction with acids because no energy has to be invested in breaking a bond before
they can be used. The strongest Bronsted bases contain atoms with unshared electrons which are localized. Examples are:
H
O
CH3O
H
NH3
Br
2. In a conjugate acid-base pair, the more negatively charged (or less positively charged) partner is the most basic. Examples of relative
basicities in such pairs are:
H
OH
CH3O
more basic than
H
more basic than
O
H
CH3OH
H
O
H
NH2
more basic than
H
more basic than
O
NH3
H
3. The periodic trend in basicity of atoms is the reverse of the acidity trend. This is particularly useful when comparing conjugate bases. Other
factors being comparable, the basicity of the atom increases from bottom to top and from right to left. Examples are:
F
more basic than
CH3CH2CH2CH2
Cl
more basic than
CH3CH2CH2N
H
H
4. The above can also be applied to resonance structures. This concept can be used to predict the outcome of an acid-base reaction when
there are several possible reactive sites.
O
O
more basic than
5. Electron delocalization decreases the basic character of the atom. Delocalization stabilizes the molecule and makes it less reactive.
Delocalized electrons are not as available as localized electrons because they are diffused over a larger area.
O
more basic than
O
6. Resonance structures involving p-bonds are the most commonly seen type. However, resonance structures involving s-bonds can be written
when the bond is highly polarized, as in the case of a C-Li bond. Such structures often reveal reactivity sites not obvious from the main
contributor. In the example below, structure II reveals the basic character of the carbon atom.
CH3
CH2
CH2
I
CH 2
Li
CH3
CH2
CH2
II
CH 2
Li
7. Alkenes are considered weak bases, since the p-bond must be broken before its p-electrons can be used for reaction with an acid.
8. Finally, a pKa table can provide the best measure of basicity when properly used.
ELECTRON MOVEMENT IN IONIC MECHANISMS - Having identified basic sites as areas of high electron density and acidic
protons as hydrogen atoms of low electron density, we can now establish how curved arrows are used to indicate the movement of electrons
in acid-base reactions. The rules are basically the same as for resonance structures.
(a) Full headed arrows indicate electron pair movement.
(b) The arrow always originates at the basic site, i.e. the source of electrons (typically nonbonding electrons, but can also be a p-bond).
(c) The arrow indicates a newly formed s-bond between the basic site of one molecule and the acidic proton of another molecule. As this new
bond is being formed, the atom (or group of atoms) to which the acidic proton is attached leaves as the conjugate base. Sometimes this atom
or group of atoms is referred to as a leaving group.
SINCE STRONG ACIDS HAVE WEAK CONJUGATE BASES, THE BEST LEAVING GROUPS ARE WEAK BASES.
IN OTHER WORDS, EQUILIBRIUM FAVORS DISPLACEMENT OF THE WEAKER BASE BY THE STRONGER BASE.
The red arrow originates at the electron source and moves towards the acidic proton.
It indicates a new bond that forms between the oxygen and the acidic proton in HCl
H
O
+
H
Cl
δ+
δ−
H
O
H
+
Cl
As oxygen bonds to the acidic proton, chlorine leaves with the electrons as chloride ion,
the conjugate base of HCl. It can therefore be referred to as the leaving group.
H
CH3CH2
OH +
H
O
H
CH3CH2
H
+
O
H
H
O
H
In this case, water acts
as the leaving group.
USING RESONANCE STRUCTURES TO PREDICT RELATIVE REACTIVITIES OF BASIC SITES
Sometimes a molecule can have several atoms which could be potentially basic. The question then is, is one more basic than the other, or
are they of equal basicity? Sometimes resonance structures can reveal relative reactivities that are not apparent from the main, or most stable
contributor. An example is acetic acid, which contains two oxygen atoms that could in principle become protonated by acid. The problem can
be approached in two ways. One is to look at the relative stabilities of the conjugate bases. It is shown below that the conjugate base that
results from protonation of oxygen (a) is resonance stabilized, whereas the one that results from protonation of oxygen (b) is not. Based on
this criterion, one predicts that oxygen (a) is more basic than oxygen (b), and that protonation will occur predominantly at oxygen (a).
Oa
a
H+
H3C
O
H3C
O
H
Oa
H
O
H
H3C
H
O
The conjugate base of protonation
at a is resonance stabilized.
H
H+
b
O
Oxygens a and b are
potential protonation sites
H3C
The conjugate base of protonation
at b is NOT resonance stabilized.
O
b
O
H
H3C
H
b
+2
O
H
An attempt to write resonance structures
would result in an oxygen with too many bonds
and a structure with too much charge separation.
H
One can arrive at the same conclusion by examining the resonance structures of acetic acid, which reveal that oxygen (a) carries a partial
negative charge and that oxygen (b) carries a partial positive charge. The more negative character of (a) makes it the more basic of the two.
a
a
O
H3C
O
O
b
H
H3C
Try moving electrons the other way around
to reveal the unfeasibility of the resulting structures.
O
b
H
Resonance structures can be used in conjunction with other information, such as periodic trends, to make a prediction. In the example below,
the two resonance structures of the cyclohexanone enolate ion show carbon and oxygen sharing the negative charge. However, basicity trends
in the periodic table predict that in comparable situations, carbon is more basic than oxygen. Remember that we’re not looking at electronegativity
trends here, but basicity, which runs opposite to electronegativity across the same row. The molecule will react with acids to favor protonation
of the carbon over the oxygen atom.
O
O
Can you push electrons to show how to
arrive from one structure to the other?
Cyclohexanone
enolate ion
Writing resonance structures involving s-bonds is not as common as doing it with p-bonds. However it’s perfectly acceptable to do it in the
case of highly polarized s-bonds, such as might be the case in H-Cl.
δ+
H
δ−
Cl
H+
Cl-
Although fairly obvious to a chemistry student, the structure on the right makes the acidic nature of the hydrogen atom apparent. The same
can be done with bases. One way to have a highly polarized s-bond involving carbon is by bonding it to a metal. The higher electronegativity
of carbon biases the electron density distribution in its favor. In this case, resonance structures reveal the negative character of such carbon,
and therefore its basic character.
CH3 CH2 CH2 CH2
Li
CH3 CH2 CH2 CH2
n-butyllithium is a strong base (pKa ~ 48)
MgBr
MgBr
Another category of strong bases is the Grignard reagents,
which specifically contain carbon-magnesium bonds.
Li
ALKENES AS WEAK BASES. OUTCOMES OF PROTONATION AT p-BONDS - As stated before, the double bond in alkenes
is a source of electrons. Alkenes are weak bases because the p-electrons are only available after breaking the p-bond first. Nonetheless,
alkenes are capable of becoming protonated by strong acids. The most important principle to keep in mind is that with unsymmetrical alkenes,
formation of the most stable carbocation is the preferred outcome.
H3C
CH3
H
H3C
H3C
CH3
+
H3C
H
+
F
Major outcome
H
F
H
H
H3C
CH3
+
H3C
F
Minor outcome
H
Protonation of the unsymmetrical alkene favors formation of the most stable
carbocation. In this example the tertiary cation is favored over the secondary one.
CH3
CH3
Major outcome
+ HF
CH3
Minor outcome
Can you show the electron movement taking place here and the
exact location of the new C-H bond in the carbocation that forms?