Chemical Reactivity I

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 reactions are commonly
written as equations:
 

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 of such functions 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.


      2. Classification by Functional Group
Functional groups are atoms or small groups of atoms (usually two to four) that exhibit a characteristic reactivity when treated with certain reagents. To view a table of the common functional groups and their class names Click Here. A particular functional group will almost always display its characteristic chemical behavior when it is present in a compound. Because of this, the discussion of organic reactions is often organized according to functional groups. The following table summarizes the general chemical behavior of the common functional groups. For reference, the alkanes provide a background of behavior in the absence of more localized functional groups.

Functional ClassFormulaCharacteristic Reactions
AlkanesC–C,   C–HSubstitution (of H, commonly by Cl or Br)
Combustion (conversion to CO2 & H2O)
AlkenesC=C–C–HAddition
Substitution (of H)
AlkynesC≡C–HAddition
Substitution (of H)
Alkyl HalidesH–C–C–XSubstitution (of X)
Elimination (of HX)
AlcoholsH–C–C–O–HSubstitution (of H); Substitution (of OH)
Elimination (of HOH); Oxidation (elimination of 2H)
Ethers(α)C–O–R Substitution (of OR); Substitution (of α–H)
AminesC–NRHSubstitution (of H);
Addition (to N); Oxidation (of N)
Benzene RingC6H6Substitution (of H)
Aldehydes(α)C–CH=OAddition
Substitution (of H or α–H)
Ketones(α)C–CR=OAddition
Substitution (of α–H)
Carboxylic Acids(α)C–CO2HSubstitution (of H); Substitution (of OH)
Substitution (of α–H); Addition (to C=O)
Carboxylic Derivatives(α)C–CZ=O
(Z = OR, Cl, NHR, etc.)
Substitution (of Z); Substitution (of α–H)
Addition (to C=O)

This table does not include any reference to rearrangement, due to the fact that such reactions are found in all functional classes, and are highly dependent on the structure of the reactant. Furthermore, a review of the overall reaction patterns presented in this table discloses only a broad and rather non-specific set of reactivity trends. This is not surprising, since the three remaining categories provide only a coarse discrimination (comparable to identifying an object as animal, vegetable or mineral). Consequently, apparent similarities may fail to reflect important differences. For example, addition reactions to C=C are significantly different from additions to C=O, and substitution reactions of C-X proceed in very different ways, depending on the hybridization state of carbon.


      3. Other Classifications
As we learn more about the reactions of organic compounds, other types of classification will become useful. For example, oxidations and reductions are common transformations that can be effected by a variety of reagents. Although the former are often eliminations and the latter additions, the correlation is far from exact, making oxidation & reduction useful categories for operational consideration. This subject will be developed in a later chapter.


Mechanisms of Organic Reactions

A detailed description of the changes in structure and bonding that take place in the course of a reaction, and the sequence of such events is called the reaction mechanism. A reaction mechanism should include a representation of plausible electron reorganization, the preferred spatial orientation of reactants, as well as the identification of any intermediate species that may be formed as the reaction progresses. These features are elaborated in the following sections.

      1. The Arrow Notation in Mechanisms
Since chemical reactions involve the breaking and making of bonds, a consideration of the movement of bonding ( and non-bonding ) valence shell electrons is essential to this understanding. It is now common practice to show the movement of electrons with curved arrows, and a sequence of equations depicting the consequences of such electron shifts is termed a mechanism. In general, two kinds of curved arrows are used in drawing mechanisms:

A full head on the arrow indicates the movement or shift of an electron pair:
A partial head (fishhook) on the arrow indicates the shift of a single electron:

The use of these symbols in bond-breaking and bond-making reactions is illustrated below. If a covalent single bond is broken so that one electron of the shared pair remains with each fragment, as in the first example, this bond-breaking is called homolysis. If the bond breaks with both electrons of the shared pair remaining with one fragment, as in the second and third examples, this is called heterolysis. Heterolysis produces ionic species.
Bond-Breaking   Bond-Making
For further information about the use of curved arrows in reaction mechanisms Click Here.

Other Arrow Symbols

Chemists also use arrow symbols for other purposes, and it is essential that they be used correctly.

The Reaction Arrow

The Equilibrium Arrow

The Resonance Arrow

The following equations illustrate the proper use of these symbols:


      2. Reactive Intermediates
The products of bond breaking, shown above, are not stable in the usual sense, and cannot be isolated for prolonged study. Such species are referred to as reactive intermediates, and are believed to be transient intermediates in many reactions. The general structures and names of four such intermediates are given below.

A pair of widely used terms, related to Lewis acid-base notation, should also be introduced here.

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).

Using these definitions, it is clear that carbocations ( called carbonium ions in the older literature ) are electrophiles and carbanions are nucleophiles. Carbenes have only a valence shell sextet of electrons and are therefore electron deficient. In this sense they are electrophiles, but the non-bonding electron pair also gives carbenes nucleophilic character. As a rule, the electrophilic character dominates carbene reactivity. Carbon radicals have only seven valence electrons, and may be considered electron deficient; however, they do not in general bond to nucleophilic electron pairs, so their chemistry exhibits unique differences from that of conventional electrophiles. Radical intermediates are often called free radicals.
The importance of electrophile / nucleophile terminology comes from the fact that many organic reactions involve at some stage the bonding of a nucleophile to an electrophile, a process that generally leads to a stable intermediate or product. Reactions of this kind are sometimes called ionic reactions, since ionic reactants or products are often involved. Some common examples of ionic reactions and their mechanisms are shown in the next section

The shapes ideally assumed by these intermediates becomes important when considering the stereochemistry of reactions in which they play a role. A simple tetravalent compound like methane, CH4, has a tetrahedral configuration. Carbocations have only three bonds to the charge bearing carbon, so it adopts a planar trigonal configuration. Carbanions are pyramidal in shape ( tetrahedral if the electron pair is viewed as a substituent ), but these species invert rapidly at room temperature, passing through a higher energy planar form in which the electron pair occupies a p-orbital. Radicals are intermediate in configuration, the energy difference between pyramidal and planar forms being very small. Since three points determine a plane, the shape of carbenes must be planar; however, the valence electron distribution varies.

      3. Acid-Base Reactions
A few common reaction types are encountered repeatedly as the chemical behavior of different compounds is examined. For example, proton transfer equilibria (acid-base reactions) play a key role in most of the functional group reactions listed above.
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, whereas 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 overall complex, our understanding of how they occur often begins with a simple acid-base reaction.

Two acid-base theories are commonly used:
the Brønsted theory and the Lewis theory

Brønsted Acid-Base Equilibria

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.  


Some useful principles of acid-base reactions are:
            The stronger the acid the weaker its conjugate base.
            The stronger the base the weaker its conjugate acid.
            Acid-base equilibria always favor the weakest acid and the weakest base.


Examples of Brønsted Acid-Base Equilibria

Acid-Base Reaction Conjugate
Acids
Conjugate
Bases
KapKa
HBr   +   H2OH3O(+)   +   Br(–)HBr
H3O(+)
Br(–)
H2O
105-5
CH3CO2H   +   H2O H3O(+)   +   CH3CO2(–) CH3CO2H
H3O(+)
CH3CO2(–)
H2O
1.77*10-54.75
C2H5OH   +   H2O H3O(+)   +   C2H5O(–)C2H5OH
H3O(+)
C2H5O(–)
H2O
10-1616
NH3   +   H2OH3O(+)   +   NH2(–)NH3
H3O(+)
NH2(–)
H2O
10-3434

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.

For further discussion of how acidity is influenced by molecular structure Click Here.


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.
The distinction between the pKb of a base and the pKa of its conjugate acid can be confusing because of their reciprocal relationship. We see this relationship in the following two equilibria. In the upper equation, the base ammonia accepts a proton from water, a Brønsted acid. In the lower equation, the ammonium cation, a Brønsted acid, transfers a proton to water, acting as a base.

Acid-Base Reaction Conjugate
Acids
Conjugate
Bases
KpK
NH3   +   H2ONH4(+)   +   OH(–)NH4(+)
H2O
NH3
OH(–)
Kb = 1.8*10-5pKb = 4.74
NH4(+)   +   H2OH3O(+)   +   NH3 NH4(+)
H3O(+)
NH3
H2O
Ka = 5.5*10-10pKa = 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.

Lewis Acid-Base Reactions

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. Two examples of Lewis acid-base equilibria are shown in equations 1 & 2 below. 

In the first example, an electron deficient aluminum atom bonds to a covalent chlorine atom be sharing one of its non-bonding valence electron pairs, and thus achieves an argon-like valence shell 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.

As noted earlier, a related terminology is commonly 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).


      4. Illustrative Examples
The previous topics may now be illustrated by the following examples. Reactions such as these are called ionic or polar reactions, because they often involve charged species and the bonding together of electrophiles and nucleophiles. Ionic reactions normally take place in liquid solutions, where solvent molecules assist the formation of charged intermediates.

Ionic Reactions

 

The substitution reaction shown on the left can be viewed as taking place in three steps. The first is an acid-base equilibrium, in which HCl protonates the oxygen atom of the alcohol. The resulting conjugate acid then loses water in a second step to give a carbocation intermediate. Finally, this electrophile combines with the chloride anion nucleophile to give the final product.

The addition reaction shown on the left can be viewed as taking place in two steps. The first step can again be considered an acid-base equilibrium, with the pi-electrons of the carbon-carbon double bond functioning as a base. The resulting conjugate acid is a carbocation, and this electrophile combines with the nucleophilic bromide anion.

The elimination reaction shown on the left takes place in one step. The bond breaking and making operations that take place in this step are described by the curved arrows. The initial stage may also be viewed as an acid-base interaction, with hydroxide ion serving as the base and a hydrogen atom component of the alkyl chloride as an acid.

There are many kinds of molecular rearrangements. The examples shown on the left are from an important class called tautomerization or, more specifically, keto-enol tautomerization. Tautomers are rapidly interconverted constitutional isomers, usually distinguished by a different bonding location for a labile hydrogen atom (colored red here) and a differently located double bond. The equilibrium between tautomers is not only rapid under normal conditions, but it often strongly favors one of the isomers (acetone, for example, is 99.999% keto tautomer). Even in such one-sided equilibria, evidence for the presence of the minor tautomer comes from the chemical behavior of the compound. Tautomeric equilibria are catalyzed by traces of acids or bases that are generally present in most chemical samples.

Since many ionic reactions proceed by bonding interactions between electrophiles and nucleophiles, it is important to understand how these qualities vary from compound to compound, and how they may be enhanced by acid or base catalysts. These subjects will be further developed in future chapters.

Those wishing to further explore the relationship of nucleophilicity and basicity may do so by Clicking Here.


Radical Reactions

If methane gas is mixed with chlorine gas and exposed to sunlight an explosive reaction takes place in which chlorinated methane products are produced along with hydrogen chloride. An unbalanced equation illustrating this reaction is shown below; the relative amounts of the various products depends on the proportion of the two reactants that are used.

CH4   +   Cl2   +   energy CH3Cl   +   CH2Cl2   +   CHCl3   +   CCl4   +   HCl

How does this reaction take place? Gas phase reactions, such as the chlorination of methane, do not normally proceed via ionic intermediates. Strong evidence indicates that neutral radical intermediates, sometimes called free radicals, play a role in this and many other similar transformations. A radical is an atomic or molecular species having an unpaired, or odd, electron. Some radicals, such as nitrogen dioxide (NO2) and nitric oxide (NO) are relatively stable, but most are so reactive that isolation and long-term study under normal conditions is not possible.
A set of radical reactions called a chain reaction can account for all the facts observed for this process.

The reaction is initiated by the input of energy (heat or light). The weak chlorine-chlorine bond is broken homolytically to give chlorine atoms.

In these two reactions radical intermediates abstract an atom from one of the reactant molecules. If a chlorine atom abstracts a hydrogen from methane in the first step, the resulting methyl radical abstracts a chlorine atom from chlorine in the second step, regenerating a chlorine atom. This is therefore a chain reaction.

In principle, a chain reaction should continue until one or both of the reactants are consumed. In practice, however, such reactions stop before completion and have to be reinitiated. This happens whenever two radical intermediates meet and combine to give a stable molecule, thus terminating the chain of reactions. Since radical intermediates are extremely reactive and are present in very low concentration, the probability that two such intermediates will collide is small. Consequently, the chain reaction will proceed through many cycles before termination occurs.

Reaction Energetics

The potential energy of a reacting system changes as the reaction progresses. The overall change may be exothermic (energy is released as heat) or endothermic (energy must be added, often as heat), and there is usually an activation energy requirement as well. By convention, the heat released in an exothermic reaction is reported as a negative number, and the heat required to effect an endothermic change as a positive number. Discovering the energy requirements and changes that characterize an organic reaction is an essential part of achieving a full understanding of its mechanism.
Chemical reactions involve breaking and making some (or even all) of the bonds that hold together the atoms of reactant and product molecules. Energy is required to break bonds, and since the strengths of different kinds of bonds differ, there is often a significant overall energy change in the course of a reaction. In the combustion of methane, for example, all six bonds in the reactant molecules are broken, and six new bonds are formed in the product molecules (equation 1).

Reactants         Products

As a rule, compounds constructed of strong covalent bonds are more stable than compounds incorporating one or more relatively weak bonds. In this case the sum of the product bond strengths is greater than the sum of reactant bond strengths; consequently, the products are energetically (or thermodynamically) more stable than the reactants, and energy is released in the form of heat. Such reactions are called exothermic. It is helpful to think of exothermic reactions as proceeding from a higher energy (less stable) reactant state to a lower energy (more stable) product state, as shown in the diagram on the right. Reactions in which the products are higher energy than the reactants require an energy input to occur, and are called endothermic. Photosynthesis (equation 2) is an important example of an endothermic process. Energy in the form of photons (sunlight) drives the reaction, which requires chlorophyll as a catalyst.

Reactants Products            

Since reactions of organic compounds involve the making and breaking of bonds, the strength of bonds, or their resistance to breaking, becomes an important consideration. For example, the chlorination of methane, discussed earlier, was induced by breaking a relatively weak Cl-Cl covalent bond. If that bond fails to break, no reactive chlorine atoms are formed, and no reaction takes place.

      1. Bond Energy
The millions of molecules that exist in this world are structurally stable relative to their component atoms, otherwise they would break apart spontaneously. The forces that hold the atoms together in unique constitutions and configurations are called bonds. Bond energy is the energy required to break a covalent bond homolytically (into neutral fragments or atoms). Bond energies are commonly given in units of kcal/mol or kJ/mol, and are generally called bond dissociation energies when given for specific bonds, or average bond energies when summarized for a given type of bond over many kinds of compounds. Tables of bond energies may be found in most text books and handbooks. The following table is a collection of average bond energies for a variety of common bonds. Such average values are often referred to as standard bond energies, and are given here in units of kcal/mole.

Standard Bond Energies

Single Bonds

ΔH*

 

Single Bonds

ΔH*

 

Multiple Bonds

ΔH*

H–H

104.2

B–F

150

C=C

146

C–C

83

B–O

125

N=N

109

N–N

38.4

C–N

73

O=O

119

O–O

35

N–CO

86

C=N

147

F–F

36.6

C–O

85.5

C=O (CO2)

192

Si–Si

52

O–CO

110

C=O (aldehyde)

177

P–P

50

C–S

65

C=O (ketone)

178

S–S

54

C–F

116

C=O (ester)

179

Cl–Cl

58

C–Cl

81

C=O (amide)

179

Br–Br

46

C–Br

68

C=O (halide)

177

I–I

36.

C–I

51

C=S (CS2)

138

H–C

99

C–B

90

N=O (HONO)

143

H–N

93

C–Si

76

P=O (POCl3)

110

H–O

111

C–P

70

P=S (PSCl3)

70

H–F

135

N–O

55

S=O (SO2)

128

H–Cl

103

S–O

87

S=O (DMSO)

93

H–Br

87.5

Si–F

135

P=P

84

H–I

71

Si–Cl

90

P≡P

117

H–B

90

Si–O

110

C≡O

258

H–S

81

P–Cl

79

C≡C

200

H–Si

75

P–Br

65

N≡N

226

H–P

77

P–O

90

C≡N

213

* Average Bond Dissociation Enthalpies in kcal mole-1
  The SI unit of energy is the joule, symbol J. To convert kilocalories into kilojoules multiply by 4.184.
  A useful site for unit conversions may be reached by Clicking Here.

Some useful and interesting conclusions may be drawn from this table. First, a single bond between two given atoms is weaker than a double bond, which in turn is weaker than a triple bond. Second, hydrogen forms relatively strong bonds (90 to 110 kcal) to the common elements found in organic compounds (C, N & O). Third, with the exception of carbon and hydrogen, single bonds between atoms of the same element are relatively weak (35 to 64 kcal). Indeed, the fact that carbon forms relatively strong bonds to itself as well as to nitrogen, oxygen and hydrogen is a primary factor accounting for the very large number of stable organic compounds.

Bond energy is energy that must be introduced to break a bond, it is not a component of a molecule's potential energy.


      2. Isomer Stability
Common sense suggests that molecules in which the bonds are all strong will be more stable than molecules having weaker bonds. Previously we defined bond strengths as the energy required to break a bond into neutral fragments (radicals or atoms). The sum of all the bond energies of a molecule can therefore be considered its atomization energy, i.e. the energy required to break the molecule completely into its component atoms. Such a transformation would be highly endothermic. The reverse process, forming a particular molecule from its component atoms, would be exothermic and the energy released is called the heat of formation. If this concept is applied to a group of isomers, it should be clear that all the isomers will have a common atomic state, and that the total bond energy of each isomer is related to that isomer's potential energy. Thus, that isomer having the greatest total bond energy has the lowest potential energy and is thermodynamically most stable.
The three C6H12 isomers shown below illustrate this relationship. Cyclohexane is made up of six C-C sigma bonds and twelve C-H sigma bonds configured in a strain-free six-membered ring. The isomer having a double bond, 1-hexene, on the other hand, has four C-C single bonds (all sigma) and one C-C double bond (one sigma plus one pi bond). Since the pi bond is weaker than a sigma bond, cyclohexane has a larger total bond energy (by nearly 20 kcal/mol) and is thermodynamically more stable than 1-hexene. The four-membered ring compound, ethylcyclobutane, has the same kinds of bonds as cyclohexane, but they are weakened by ring strain to such a degree that this isomer is even less stable (in the thermodynamic sense) than 1-hexene. This relationship is depicted by the potential energy diagram on the right. Here the numbers are heats of formation, and are negative to reflect the exothermic nature of bond formation. They are proportional to potential energy, with a lower (more negative) value indicating increased thermodynamic stability.

The Nature of Energy
The term energy is used here in a general and rather imprecise way.
For a more comprehensive treatment Click Here.



Practice Problems

The following problems include acid base relationships, recognition of different functional groups, recognition of nucleophiles and electrophiles, classification of reactions by structural change, and the use of curved arrow notation.



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