Chemistry · Chemical Bonding · NEET
When a molecule has more than one bond of the same type (like the 4 C-H bonds in CH4), we cannot give one single 'bond dissociation enthalpy' for each bond because breaking them one at a time gives different numbers. So we add the energy to break ALL of them and divide by how many there are. That average value is the mean bond enthalpy. For CH4: total energy to turn it fully into atoms is 1665 kJ/mol, and there are 4 C-H bonds, so mean C-H bond enthalpy = 1665/4 = 416 kJ/mol.
In the intact CH4 molecule, all four C-H bonds ARE identical in length and energy. The difference shows up only when you break them one by one. After you remove the first H, the leftover fragment (CH3, then CH2, then CH) is a different species with different electron arrangement, so the next bond needs a different amount of energy. NCERT gives them as 427, 439, 452 and 347 kJ/mol. Because the molecule changes after each break, we report the AVERAGE instead of four separate values.
Bond dissociation enthalpy is the energy to break ONE specific bond in a specific molecule. For a diatomic molecule like H2 or Cl2 there is only one bond, so this single value is exact (H-H = 435 kJ/mol). Mean bond enthalpy is used for POLYATOMIC molecules with many identical bonds; it is the AVERAGE over all those bonds. NCERT note: the same symbol (Δ_bondH°) is used for both, so read the context. Rule: diatomic = bond dissociation enthalpy; polyatomic same-type bonds = mean bond enthalpy.
Step 1: Write the atomization reaction, where the whole molecule breaks into free gaseous atoms, e.g. CH4(g) -> C(g) + 4H(g). Step 2: Find the enthalpy of atomization (Δ_aH°) for this = 1665 kJ/mol. Step 3: Count how many bonds of that type break = 4 C-H bonds. Step 4: Divide. Mean C-H bond enthalpy = 1665/4 = 416 kJ/mol. For water H2O(g) -> 2H(g) + O(g), you would divide the total by 2 to get the mean O-H bond enthalpy.
Only for diatomic molecules. For H2, breaking the single bond gives 2 free H atoms, so enthalpy of atomization = bond dissociation enthalpy = 435 kJ/mol. For polyatomic molecules like CH4 the enthalpy of atomization (1665 kJ/mol) is the TOTAL to make all atoms free, which equals the SUM of all the successive bond enthalpies. Dividing that total by the number of bonds gives the mean bond enthalpy.
No, it changes a little from molecule to molecule. NCERT says the mean C-H bond enthalpy is 416 kJ/mol in CH4 but differs slightly in compounds like CH3CH2Cl or CH3NO2, because the neighbouring atoms change the environment of the bond. The change is small, so tabulated mean bond enthalpies are still useful for estimating reaction enthalpies using Hess's law.
Identify the correct orders against the property mentioned: A. H2O > NH3 > CHCl3 (dipole moment); B. XeF4 > XeO3 > XeF2 (lone pairs on central atom); C. O-H > C-H > N-O (bond length); D. N2 > O2 > H2 (bond enthalpy). Choose the correct answer.
Try the real previous-year questions from this chapter — each with the answer and a full solution.
416 kJ/mol. It comes from the enthalpy of atomization of CH4 (1665 kJ/mol) divided by the 4 C-H bonds: 1665/4 = 416 kJ/mol.
Because a polyatomic molecule has several identical bonds that break with slightly different energies in successive steps. Averaging them gives one useful value. A diatomic molecule has only one bond, so its exact bond dissociation enthalpy is used directly.
In the whole CH4 molecule they are equal in length and energy. They only appear unequal when broken one at a time, because each break leaves a different fragment (CH3, CH2, CH).
kJ/mol (kilojoules per mole). It is always positive because breaking bonds needs energy.
Yes. NCERT states that using tabulated mean bond enthalpies with Hess's law lets you estimate the enthalpy change of a reaction from bonds broken and bonds formed.
