Question

Consider the following reaction: CH4(g) → C(g)+4H (g); ∆aH° = 1665 kJ mol-¹ Which of the statements is FALSE?

• A. The energy required to break individual C-H bonds in successive steps is different.
• B. ∆aH° is the mean bond enthalpy of a C-H bond.
• C. All four C-H bonds in CH4 are identical in bond length and energy.
• D. Mean C-H bond enthalpies differ slightly from compound to compound.

Answer 1

Answer: (B)

∆aH° is the mean bond enthalpy of a C-H bond.

Answer 2

The given reaction is the atomization of methane:

CH4(g) → C(g) + 4H(g); ∆aH° = 1665 kJ mol⁻¹

∆aH° is the atomization enthalpy, which is the total energy required to break all the bonds in one mole of gaseous methane to form gaseous atoms. In methane, there are four C-H bonds.

• Statement A is TRUE: The energy required to break individual C-H bonds in successive steps is different because the chemical environment changes after each hydrogen atom is removed.

• Statement B is FALSE: ∆aH° is the total energy required to break all four C-H bonds, not the mean bond enthalpy of a C-H bond. The mean bond enthalpy would be ∆aH° / 4.

• Statement C is TRUE: All four C-H bonds in CH4 are identical in bond length and energy due to the tetrahedral structure of methane.

• Statement D is TRUE: Mean C-H bond enthalpies differ slightly from compound to compound because the electronic environment around the C-H bond is influenced by the rest of the molecule.

Therefore, the false statement is B.