Chemical Reactivity and Mechanisms – Enthalpy, Bond Dissociation, and Energy Diagrams

Chemical Reactivity and Mechanisms

Enthalpy (ΔH) – Definition and Bond Breaking

Enthalpy (ΔH or q) is a measure of the heat energy exchanged between a chemical reaction and its surroundings. Understanding enthalpy is essential for predicting whether a reaction will absorb or release energy.

• Bond breaking requires the system to absorb energy, as electrons must overcome the stability of the bond.

• When a bond is broken, electrons absorb kinetic energy to overcome the bond's stability.

• Potential energy diagrams illustrate the energy changes during bond formation and cleavage.

Bond Cleavage: Homolytic vs. Heterolytic

Bonds can break in two distinct ways, each producing different species:

• Homolytic cleavage: Each atom retains one electron from the bond, forming radicals.

• Heterolytic cleavage: One atom retains both electrons, forming ions.

• Bond dissociation energy (BDE) or ΔH for bond breaking typically refers to homolytic bond cleavage.

Bond Dissociation Energy (BDE) – Table of Common Bonds to Hydrogen

BDE values indicate the energy required to break specific bonds. These values are crucial for calculating reaction enthalpy.

Bonds to H	kJ/mol	kcal/mol
H—H	435	104
H—CH3	435	104
H—CH2CH3	410	98
H—F	569	136
H—Cl	431	103
H—Br	368	88
H—I	297	71
H—OH	498	119

Bond Dissociation Energy – More Examples

Additional BDE values for C—H and C—X bonds are useful for organic reaction calculations.

C—H bonds	kJ/mol	kcal/mol
CH3—CH3	368	88
CH2CH3—CH3	356	85
CH2CH3—CH3	351	84
CH3—Br	293	70
CH3—I	234	56
CH3—OH	381	91
C—X bonds	kJ/mol	kcal/mol
C—F	159	38
C—Cl	243	58

Exothermic and Endothermic Reactions

Chemical reactions often involve multiple bonds breaking and forming. The overall energy change determines whether the reaction is exothermic or endothermic.

• Exothermic reaction: The energy gained by bonds formed exceeds the energy needed for bonds broken. Products are more stable than reactants.

• Endothermic reaction: The energy needed for bonds broken exceeds the stability gained by the bonds formed. Products are less stable than reactants.

Energy Diagrams for Exothermic vs. Endothermic Reactions

Energy diagrams visually represent the energy changes during a reaction.

• Exothermic reactions:

  • Products are lower in energy than reactants.

  • Energy is released as heat (potential energy converted to kinetic energy).

  • $\Delta H$ is negative.

  • Temperature of surroundings increases.

• Endothermic reactions:

  • Products are higher in energy than reactants.

  • Energy is consumed (kinetic energy converted to potential energy).

  • $\Delta H$ is positive.

  • Temperature of surroundings decreases.

Sign of ΔH and Energy Diagrams

The sign of ΔH indicates whether a reaction is exothermic (negative ΔH) or endothermic (positive ΔH). Energy diagrams are used to describe both the kinetics and thermodynamics of a chemical reaction.

• Potential energy (PE) is shown on the y-axis.

• Reaction progress is shown on the x-axis (reaction coordinate).

• Practice predicting ΔH using bond dissociation energies and energy diagrams.

Key Formula: Calculating ΔH Using BDEs

• The enthalpy change for a reaction can be calculated using bond dissociation energies:

$\Delta H = \sum \text{BDE (bonds broken)} - \sum \text{BDE (bonds formed)}$

• If ΔH is negative, the reaction is exothermic.

• If ΔH is positive, the reaction is endothermic.

Example: Calculating ΔH

• Bonds broken: C—H (397 kJ/mol) and Br—Br (193 kJ/mol)

• Bonds formed: C—Br (285 kJ/mol) and H—Br (368 kJ/mol)

$\Delta H = (397 + 193) - (285 + 368) = -63 \text{ kJ/mol}$

• Since ΔH is negative, the reaction is exothermic.