The methane molecule, CH4, has the geometry shown in Figure 2.19. Imagine a hypothetical process in which the methane molecule is "expanded," by simultaneously extending all four C-H bonds to infinity. We then have the process
CH4(g) → C(g) + 4H(g)
(a) Compare this process with the reverse of the reaction that represents the standard enthalpy of formation of CH4(g).
(b) Calculate the enthalpy change in each case. Which is the more endothermic process? What accounts for the difference in ΔH° values?
(c) Suppose that 3.45 g CH4(g) reacts with 1.22 g F2(g), forming CF4(g) and HF(g) as sole products. What is the limiting reagent in this reaction? If the reaction occurs at constant pressure, what amount of heat is evolved?