Chapter 6: Thermodynamics and Mechanisms of Organic Reactions

Overview of Organic Reactions

Kinds of Organic Reactions

Organic reactions can be classified into several fundamental types, each defined by the changes in molecular structure and functional groups:

• Addition Reactions: Two molecules combine to form a single product. Common in alkenes and alkynes.

• Elimination Reactions: A single molecule splits into two products, often forming a double bond.

• Substitution Reactions: An atom or group in a molecule is replaced by another atom or group.

• Rearrangement Reactions: The structure of a molecule is reorganized to form an isomer.

Example: The addition of HBr to ethylene is an example of an addition reaction.

Reaction Mechanisms

Bond Breaking and Bond Making

Organic reaction mechanisms describe the stepwise process by which reactants are converted to products. Mechanisms involve:

• Bond Breaking: Can occur via homolytic (each atom takes one electron) or heterolytic (one atom takes both electrons) cleavage.

• Bond Making: Formation of new bonds through electron pair donation.

Additional info: Homolytic cleavage produces radicals, while heterolytic cleavage produces ions.

Polar Reactions

Polar reactions involve the movement of electrons from electron-rich (nucleophilic) sites to electron-poor (electrophilic) sites. The polarity of molecules is influenced by their functional groups and polarizability.

• Nucleophile: Electron-rich species, often negatively charged or with lone pairs.

• Electrophile: Electron-poor species, often positively charged or with electron-deficient atoms.

Example: Chloromethane (CH3Cl) and methyllithium (CH3Li) demonstrate polar reactions, with methyllithium acting as a nucleophile.

Identifying Nucleophiles and Electrophiles

Electrostatic potential maps help identify nucleophilic (negative) and electrophilic (positive) atoms. Consider the following species:

Species	Type	Reason
NO2+	Electrophile	Positively charged
CN–	Nucleophile	Negatively charged
CH3NH2	Both	Lone pair on N (nucleophile); N–H hydrogens (electrophile)
(CH3)3S+	Electrophile	Positively charged

Curved Arrows in Polar Reaction Mechanisms

Rules for Using Curved Arrows

Curved arrows are used to show the movement of electron pairs in reaction mechanisms:

• Rule 1: Arrows start at electron-rich sites (lone pairs or bonds) and point to electron-poor sites.

• Rule 2: Arrows must not violate the octet rule.

• Rule 3: Each arrow represents the movement of a pair of electrons.

• Rule 4: Multiple arrows may be needed for multistep processes.

Example: In the reaction of a nucleophilic carbon with CH3Br, a curved arrow shows electron donation from the nucleophile to the electrophile, and another arrow shows the breaking of the C–Br bond.

Thermodynamics of Organic Reactions

Equilibrium Constant (Keq)

The equilibrium constant describes the ratio of products to reactants at equilibrium:

• If $K_{eq} \gg 1$, the forward reaction is favored.

• If $K_{eq} \ll 1$, the backward reaction is favored.

• If $K_{eq} = 1$, neither direction is favored.

Significance: Keq indicates the theoretical yield but not the rate of reaction.

Limitations of Keq

Keq only tells the position of equilibrium, not how fast equilibrium is reached. Reaction rate and equilibrium are separate criteria for reaction favorability.

Gibbs Free-Energy Change (ΔG°) and Equilibrium

The relationship between Gibbs free energy and equilibrium constant is given by:

$\Delta G^\circ = -RT \ln K_{eq}$

Where R is the gas constant and T is temperature in Kelvin.

Thermodynamic quantities are related as:

$\Delta G^\circ = \Delta H^\circ - T\Delta S^\circ$

• ΔG°: Standard free-energy change

• ΔH°: Standard enthalpy change

• ΔS°: Standard entropy change

Thermodynamically Favorable Reaction:

• Products more stable than reactants

• $K_{eq} \gg 1$

• $\Delta G^\circ$ is negative

• $\Delta S^\circ$ is positive

Bond Dissociation Energy

Definition and Significance

Bond dissociation energy (D) is the energy required to break a bond and produce two radical fragments in the gas phase at 25°C.

• Heat is released ($\Delta H^\circ$ negative) when a bond is formed.

• Heat is absorbed ($\Delta H^\circ$ positive) when a bond is broken.

Stronger bonds have higher dissociation energies; weaker bonds have lower dissociation energies.

Energy Diagrams and Transition States

Energy Diagrams

Energy diagrams illustrate the energy changes during a reaction:

• Reactants must overcome an energy barrier (activation energy, $\Delta G^\ddagger$) to form products.

• Transition State: The highest energy point along the reaction path; unstable and cannot be isolated.

• Reaction Intermediate: A short-lived species formed during multistep reactions; can sometimes be isolated.

Example: The addition of HBr to ethylene involves two steps, each with its own transition state and intermediate.

Energy Diagram Features

• $\Delta G^\ddagger$ (activation energy) determines reaction rate.

• $\Delta G^\circ$ (free-energy change) determines equilibrium position.

Example: A fast exergonic reaction has low $\Delta G^\ddagger$ and negative $\Delta G^\circ$.

Worked Examples: Drawing Energy Diagrams

• One-step fast exergonic reaction: Small $\Delta G^\ddagger$, large negative $\Delta G^\circ$.

• Two-step exergonic reaction: Two transition states, intermediate between them; overall $\Delta G^\ddagger$ is the difference between reactants and the highest transition state.

Comparison of Biological and Laboratory Reactions

Biological vs. Laboratory Reactions

Biological reactions (enzyme-catalyzed) typically involve many steps, each with small activation energies, leading to efficient and controlled processes. Laboratory reactions may have fewer steps but higher activation energies.

Feature	Biological Reaction	Laboratory Reaction
Number of Steps	Many	Few
Activation Energy	Small for each step	Large
Control	High (enzymes)	Lower
End Result	Same as lab reaction	Same as biological reaction

Additional info: Enzymes lower activation energy, increasing reaction rates without altering equilibrium.