Ch 201 Problem 8.66

Using bond enthalpies (Table 8.4), estimate ΔH for the following gas-phase reactions:

(a)

OK, here's the idea: ΔH is just the amount of energy absorbed or released in a reaction. It has the usual units of energy (kJ) - don't worry about the mol units right now. We can get a pretty good estimate of ΔH by looking at the amount of energy required to break individual bonds (it always requires energy to break a bond) and the amount of energy required (or released) when a bond is formed.

To estimate ΔH, we state that ΔH = Σ(enthalpies of bonds broken) - Σ(enthalpies of bonds formed)

In the above reaction, one C - H and one Cl - Cl bond are broken, and a C-Cl bond and an H-Cl bond are formed. From Table 8.4,

now,

so 104 kJ of energy are released - this means that the products are more stable than the reactants. A negative ΔH means that heat is released!

(b)

Here, we're breaking 2 C - S bonds and 2 H - Br bonds. We're making 2 C - Br bonds and 2 H - S bonds. Notice that bond enthalpies depend on amount of material, so if we break 2 C-S bonds, we multiply the bond enthalpy of a C-S bond x 2, and if we make 2 C-Br bonds, we multiply by 2 also. From Table 8.4,

A positive ΔH means that heat is required to convert reactants to products.

(c) Look at the Lewis structures given: one N-N bond and one Cl-Cl bond are broken, and 2 N-Cl bonds are formed:

heat is required to get the reaction going!

DAC 7/16/08