Structural effect on stability and reactivity.pdf

Structural Effects on Stability

and Reactivity

Introduction

The concepts of stability and reactivity are fundamental to understanding chemistry.

In this chapter we consider first the thermodynamic definition of chemical stability.

We then consider chemical kinetics (Section 3.2) and how it can provide information

about reactivity. We also explore how structure influences stability and reactivity. We

want to learn how to make predictions about reactivity based on the structure of the

reactants and intermediates. We begin by reviewing the principles of thermodynamics

and kinetics, which provide the basis for understanding the relationship of structure to

stability and reactivity.

Reactions are usefully described in terms of potential energy diagrams such as

shown in Figure 3.1, which identify the potential energy changes associated with the

reacting molecules as they proceed to products. The diagram plots the free energy

of the system as a function of the progress of the reaction. For each individual step

in the reaction there is a transition state representing the highest energy arrangement

of the molecules for that step. The successive intermediates are the molecules that

are formed and then react further in the course of the overall reaction. The energies

of the transition states relative to the reactants determine the rate of reaction. The

energy difference between the reactants and products is G, the free-energy change

associated with the reaction. The free energy of a chemical reaction is defined by the

equation

−

(3.1)

where H is the enthalpy change and S is the entropy change for the reaction. The

enthalpy term is a measure of the stability of the molecule and is determined by the

strength of the chemical bonds in the structure. The entropy term specifies the change

253

254

CHAPTER 3

Structural Effects on

Stability and Reactivity

Δ G ‡

A + B

D + E

Reaction Progress

Fig. 3.1. Reaction potential energy profile showing transition

states and intermediate for a reaction A + B → C → D + E.

in the order (probability) associated with the reaction. The free energy of the reaction,

G, determines the position of the equilibrium for the reaction:

=−

RT ln K

(3.2)

where K is the equilibrium constant for the reaction:

K =

The energy required to proceed from reactants to products is G ‡ , the free energy

of activation , which is the energy at the transition state relative to the reactants. We

develop the theoretical foundation for these ideas about reaction rates in Section 3.2.

We first focus attention on the methods for evaluating the inherent thermodynamic

stability of representative molecules. In Section 3.3, we consider general concepts

that interrelate the thermodynamic and kinetic aspects of reactivity. In Section 3.4,

we consider how substituents affect the stability of important intermediates, such as

carbocations, carbanions, radicals, and carbonyl addition (tetrahedral) intermediates.

In Section 3.5, we examine quantitative treatments of substituent effects. In the final

sections of the chapter we consider catalysis and the effect of the solvent medium on

reaction rates and mechanisms.

3.1. Thermodynamic Stability

Thermodynamic data provide an unambiguous measure of the stability of a

particular compound under specified conditions. The thermodynamic measure of

molecular stability is H f , the standard enthalpy of formation , which gives the

enthalpy of the compound relative to the reference state of its constituent elements

under standard conditions of 1 atm and 298 K. For each element a particular form

is assigned an enthalpy (potential energy content) of 0. For example, for hydrogen,

nitrogen, oxygen, and fluorine, the gaseous diatomic molecules are the reference states.

For carbon, 0 energy is assigned to graphite C g , which consists entirely of sp 2

carbon atoms. The H f of compounds can be measured directly or indirectly. The

H f in kcal/mol of HF, H 2 O, NH 3 , and CH 4 are found to be as follows:

255

SECTION 3.1

Thermodynamic Stability

H ° f

1/2 H 2 + 1/2 F 2

– 136.9

H 2 + 1/2 O 2

H 2 O

– 68.3

3/2 H 2 + 1/2 N 2

NH 3

– 11.0

– 17.8

2H 2 + C ( g )

CH 4

The H f of a given compound is a physical constant and is independent of the

process by which the compound is formed. Therefore, H f values are additive and

can be calculated precisely for balanced chemical equations if all the necessary data

are available. For example, it might be experimentally impossible to measure the H f

of methane directly by calorimetry, but it can be calculated as the sum of the enthalpy

for an equivalent reaction sequence, e.g:

Δ H ° (kcal/mol)

C (g) + 1/2 O 2

– :C O: +

– 26.4

– :C O: + + 3H 2

CH 4 + H 2 O

– 59.7

H 2 O

H 2 + 1/2 O 2

+ 68.3

C (g) + 2H 2

CH 4

– 17.8

The H f for many compounds has been determined experimentally and the data

tabulated. 1 In the sections that follow, we discuss approaches to computing H f when

the experimental data are not available. It is important to note that direct comparison of

the H f values for nonisomeric compounds is not meaningful. The H f for methane

through hydrogen fluoride, for example, gives us no comparative information on

stability, because the reference points are the individual elements. Other information is

needed to examine relative stability. For example, as is discussed in the next section, it

is possible to assign bond energies to the bonds in CH 4 ,NH 3 ,H 2 O, and HF. This is the

energy required to break a C

−

H, N

−

H, O

−

H, or H

−

H bonds become stronger as we go from C to F in the

second-row compounds with hydrogen.

−

Compound

X − H bond energy

(kcal/mol)

CH 4

105.0

NH 3

108.2

H 2 O

119.3

136.4

J. B. Pedley, R. D. Naylor, and S. P. Kirby, Thermochemical Data of Organic Compounds , 2nd Edition,

Chapman and Hall, London, 1986; H. Y. Afeefy, J. F. Liebman, and S. E. Stein, in NIST Chemistry

Webbook , NIST Standard Reference Database Number 69, P. J. Linstrom and W. G. Mallard, eds., 2001

(http://webbook.nist.gov).

F bond. These numbers do begin

to provide some basis for comparison of the properties of nonisomeric compounds,

as we now see that the X

256

3.1.1. Relationship between Structure and Thermodynamic Stability

for Hydrocarbons

CHAPTER 3

Structural Effects on

Stability and Reactivity

C 6 alkenes,

and several general relationships become apparent. One is that chain branching

increases the stability of alkanes . This relationship is clear, for example, in the data

for the C 6 alkanes, with a total enthalpy difference of nearly 4 kcal/mol between

the straight-chain hexane and the tetra-substituted 2,2-dimethylbutane. There is a

similar range of 4.5 kcal/mol between the least stable (octane) and most stable

C 6 and C 8 alkanes and the C 4 −

Table 3.1. Standard Enthalpy of Formation of Some Hydrocarbons

(in kcal/mol) a

Alkanes (liquid)

Butane

− 350

Octane

− 598

2-Methylpropane

− 367

2-Methylheptane

− 609

3-Methylheptane

−

603

C 5

4-Methylheptane

−

601

Pentane

−

415

2,2-Dimethylhexane

−

626

2-Methylbutane

−

427

2,3-Dimethylhexane

−

604

2,2-Dimethylpropane

− 455

2,4-Dimethylhexane

− 614

2,5-Dimethylhexane

− 622

C 6

3,4-Dimethylhexane

− 602

Hexane

− 475

3,3-Dimethyhexane

− 615

2-Methylpentane

− 489

2,2,3-Trimethylpentane

− 614

3-Methylpentane

− 484

2,2,4-Trimethylpentane

− 620

2,3-Dimethylbutane

− 496

2,3,4-Trimethylpentane

− 609

2,2-Dimethylbutane

− 511

2,3,3-Trimethylpentane

− 606

3-Ethyl-2-methylpentane

− 597

3-Ethyl-3-methylpentane

− 604

2,2,3,3-Tetramethylbutane

− 643

B. Alkenes (liquid)

C 4

C 6

1-Butene

− 490

1-Hexene

− 177

E-2-Butene

− 789

E-2-Hexene

− 204

Z-2-Butene

− 710

Z-2-Hexene

− 201

2-Methylpropene

− 896

E-3-Hexene

− 206

Z-3-Hexene

− 189

C 5

2-Methyl-1-pentene

− 215

1-Pentene

− 112

3-Methyl-1-pentene

− 187

E-2-Pentene

− 139

4-Methyl-1-pentene

− 191

Z-2-Pentene

− 128

2-Methyl-2-pentene

− 235

2-Methyl-1-butene

− 123

3-Methyl-2-pentene

− 226

3-Methyl-1-butene

−

164

E-3-Methyl-2-pentene

−

226

Z-3-Methyl-2-pentene

−

226

E-4-Methyl-2-pentene

−

219

Z-4-Methyl-2-pentene

−

208

2,3-Dimethyl-1-butene

− 223

3,3-Dimethyl-1-butene

− 209

2-Ethyl-1-butene

− 208

3,3-Dimethyl-1-butene

− 209

2,3-Dimethyl-2-butene

− 243

a. From Thermochemical Data of Organic Compounds , 2nd Edition, J. B. Pedley, R. O. Naylor,

and S. P. Kirby, Chapman and Hall, London, 1986.

Extensive thermodynamic data are available for the major classes of hydro-

carbons. Table 3.1 gives data for the C 4 −

C 4

C 8

(2,2,3,3-tetramethylbutane) of the C 8 isomers. For alkenes, substitution on the double

bond is stabilizing. There is a range of nearly 7 kcal/mol for the C 6 H 12 isomers. The

data for C 6 alkenes, for example show:

257

SECTION 3.1

Thermodynamic Stability

1-hexene

E -3-hexene

2-methyl-2-pentene

2,3-dimethyl-2-butene

H f ° = –17.7 kcal/mol

H f ° = –20.6 kcal/mol

H f ° = –23.5 kcal/mol

H f ° = –24.3 kcal/mol

These relationships are a result of the stabilizing effect of branching and double-bond

substitution and hold quite generally, except when branching or substitution results in

van der Waals repulsions (see Section 2.3), which have a destabilizing effect.

3.1.2. Calculation of Enthalpy of Formation and Enthalpy of Reaction

In Chapter 1, we introduced various concepts of structure and the idea that the

properties of molecules are derived from the combination of the properties of the atoms.

One of the qualitative conclusions from these considerations is that the properties

of CH 3 ,CH 2 , CH, and C groups in hydrocarbons are expected to be similar from

molecule to molecule, as long as they are not perturbed by a nearby functional group.

Several methods for the calculation of thermodynamic data based on summation of

group properties have been developed and are discussed in the next two sections.

3.1.2.1. Calculations of Enthalpy of Reaction Based on Summation of Bond Energies.

The computation of molecular energy by MO or DFT methods gives the total binding

energy of a molecule. This is a very large number, since it includes all the electron-

nuclei forces in the atoms, not just the additional attractive forces of the bonding

electrons. The total energy can be converted to an energy representing all bonding

between atoms by subtracting the energy of the individual atoms. This difference in

energy is called the energy of atomization . This quantity still represents an energy that

is far larger then the change involved in chemical reactions, which is of primary interest

to chemists. The focus of chemical reactivity is on the bonds that are being formed

and broken in the reaction. Useful relationships between structure and reactivity can

be developed by focusing on bond dissociation energies (BDE). The most completely

developed information pertains to homolytic bond dissociation , 2 which is the energy

required to break a specific bond in a molecule with one electron going to each of the

atoms. From the general bond energies in Part A of Table 3.2 we can discern several

trends. One is that C

−

N bonds). We can also

note that the bonds in the dihalogens are relatively weak, with a somewhat irregular

trend with respect to position in the periodic table: F 2 < Cl 2 > Br 2 > I 2 . The bonds to

hydrogen are also slightly irregular: N < C < O < F. For the hydrogen halides, there

is a sharp drop going down the periodic table.

It is known that the immediate molecular environment significantly affects the

bond energy, as is illustrated by the data in Part B of Table 3.2. For hydrocarbons the

−

O and N

−

H bond dissociation energy depends on the degree of substitution and hybridization

2 For a discussion of the measurement and application of bond dissociation energies, sec S. J. Blanksby

and G. B. Ellison, Acc. Chem. Res. , 36 , 255 (2003).

C bonds are considerably stronger than the other homonuclear

bonds for the second-row elements (compare with O

Plik z chomika:

Inne pliki z tego folderu:

Inne foldery tego chomika: