Enthalpy Practice Problems

In the iodination of ethane shown below, tetraiodomethane (CI₄) is used as the source of iodine in the presence of a free-radical initiator such as hydrogen peroxide (H₂O₂).
Propose a mechanism for the reaction, and determine the value of ∆H for each of the steps in your mechanism using the appropriate bond-dissociation energies.

H₃C−CH₃ + CI₄ → H₃C−CH₂I + HCI₃

For the reaction-energy diagram shown below:

(i) Give the appropriate label for the reactants, products, and activation energy for the first and the second step of the reaction.

(ii) Determine the sign of ∆Hº for the overall reaction, and identify whether the overall reaction is exothermic or endothermic.

(iii) Identify parts of the reaction-energy diagram that correspond to intermediates and transition states.

(iv) Identify part of the diagram that corresponds to the transition state of the rate-limiting step. Is the structure of this transition state similar to the reactants, the products, or an intermediate?

Use the appropriate bond-dissociation enthalpies to compute ∆Hº value for the reaction below.
CH₃−CH₃ + Br₂ → CH₃CH₂Br + HBr

Compute the overall ∆Hº for the iodination of methane which proceeds through the following steps:
1. I₂ + 2I•
∆Hº (per mol) = +149 kJ (+35.6 kcal)
2. CH₄ + I• → •CH₃ + HI
∆Hº (per mol) = +141 kJ (+33.7)
3. •CH₃ + I₂ → CH₃I + I•
∆Hº (per mol) = −92 kJ (−22.0 kcal) 

(i) Propose a mechanism for the reaction shown below:

(ii) Use the appropriate bond-dissociation enthalpies to find the value of ∆Hº for each step shown in the mechanism you proposed, and calculate the value of ∆Hº for the overall reaction if the bond-dissociation enthalpy for PhCH₂−Br bond is approximately 280 kJ/mol. Are these values appropriate for a rapid free-radical chain reaction?

(i) Calculate the heat of reaction for each step in the free-radical chlorination of methane using bond-dissociation enthalpies from the table below.

(ii) Calculate the overall heat of the reaction.

Refer to the information given in the table below. Provide the ∆H°_hydr for each step involved in the reduction of allylbenzene. 

∆H₁ + ∆H₂ + ∆H₃ + ∆H₄ = ∆H_tot = –78.2 kcal/mol (–328 kJ/mol)

Utilizing bond dissociation energies (BDE), determine the change in enthalpy (∆H°) for the reaction given below.

The reaction below is said to be favorable. Determine the sign of ΔH and explain why.

Provide the bond dissociation energy for the indicated bonds in the compound below:

Use the following BDE values:

3° C—H = 381 kJ/mol

2° C—H = 397 kJ/mol

RO—H = 426 kJ/mol

The following reaction illustrates an alkene addition reaction:

Provide the value of its ΔH° and predict the sign of ΔS°.

Use the following BDE values:

C—C π bond = 65 kcal/mol

I—I bond = 36 kcal/mol

2° C—I bond = 53 kcal/mol

The reaction (i) is disfavored overall (ΔG° > 0) yet favored by ΔH°:

Meanwhile, reaction (ii) is favored (ΔG° < 0):

Determine the bonds formed and broken for reactions (i) and (ii), and explain why reaction (ii) is favored.

BDE values for reference: 

C—O π bond = 356 kJ/mol, RO—H bond = 426 kJ/mol, 1° C—OH bond = 381 kJ/mol

Based on the bond dissociation energies of the indicated bonds below, which will be easier to break?

How can this explain the following mechanism?

Use the following BDE: 

I—I bond = 36 kcal/mol

H—I bond = 71 kcal/mol

Determine the enthalpy change (∆H°) for the given equilibrium process.

For the following reaction, the value of ∆H° is −42 kcal/mol. Calculate the bond dissociation enthalpy of the C—H bond when the bond dissociation enthalpies of O—Cl, C—Cl, and O—H bonds are 47, 85, and 105 Kcal/mol, respectively.

A. 101 Kcal/mol

B. 120 Kcal/mol

C. 50 Kcal/mol

D. 320 Kcal/mol

Consider the following reverse reaction:

•CH(CH₃)₂ + HF → (CH ₃)₂CH₂ + F•

The activation energy and the heat of reaction (ΔH°) of its forward reaction are 26.1 kJ/mol (6.24 kcal/mol) and −157 kJ/mol (−37.5 kcal/mol), respectively. Determine the activation energy and the heat of reaction (ΔH°) for the reverse reaction.

Calculate the value of the standard enthalpy change (ΔH°) for each step in the chlorination of ethane. Determine the overall value of ΔH°. 

Use the following BDE values: 

CH₃CH₂—Cl bond = 355 kJ/mol

CH₃CH₂—H bond = 423 kJ/mol

Cl—Cl bond = 240 kJ/mol

H—Cl bond = 432 kJ/mol

Provide the value of the standard enthalpy change (ΔH°) for the following reaction:

CH₃Br + HCl → CH₃Cl + HBr

BDE values:

CH₃—Br bond = 302 kJ/mol

H—Cl bond = 432 kJ/mol

CH₃—Cl bond = 350 kJ/mol

H—Br bond = 366 kJ/mol

Estimate the value of the standard enthalpy change, ΔH°, for the following reaction:

(CH₃)₂CHOH + HI →  (CH₃)₂CHI + H₂O

BDE values:

(CH₃)₂CH—OH bond = 95.0 kcal/mol

H—I bond = 71.0 kcal/mol

(CH₃)₂CH—I bond = 57.0 kcal/mol

HO—H bond = 119 kcal/mol

Determine the value of the standard enthalpy change, ΔH°, for the following reaction:

CH₃CH₂CH₂CH₃ + H₂ → 2 CH₃CH₃

BDE values:

CH₃CH₂—CH₂CH₃ bond = 88.0 kcal/mol

H—H bond = 104 kcal/mol

CH₃CH₂—H bond = 101 kcal/mol