Bond Enthalpies of Diatomic Molecules

[Jackson Crist 1G]

I thought we said in class that diatomic molecules all have zero bond enthalpy... like O2 or H2 has bond enthalpy 0. So why is there a table in the textbook (4E.2) that is labeled "Bond Enthalpies of Diatomic Molecules" and has deltaH values for those molecules when they should just be zero? (the values are in the hundreds for each one).

[Jackson Crist 1G]

Also, this table contains molecules CO, HF, HCl, HBR, HI... I'm confused. Since when are these diatomic?

[Layla Shapouri 2L]

The table in 4E.2 show the enthalpy required to break the bond of the diatomic molecules. For examples, the enthalpy required to break H2(g) into 2H(g) (H-H) is 436 kJ/mol. The other molecules (CO, HF, etc) are diatomic molecules, not elements.

[Audrey2B]

Do they have 0 bond enthalpies because they are in their standard state?

[Ivy Zhang 2A]

Yes, diatomic molecules such as O2 have 0 bond enthalpy since they are measured in their standard state which acts as a reference point to measure changes in energy. Since diatomic elements are in their most stable state, the reactants and products remain the same and thus no enthalpy change takes place.

[Srothery 2k]

There are multiple meanings that can be derived from diatomic molecule, the traditional sense only includes molecules that are single element and naturally found together (like H2, O2, etc), and the most literal being a molecule with two atoms (hence why it includes other molecules). Because the traditional Diatomics are naturally stable together, the enthalpy of formation is 0, but this table is for the breaking, not forming, of the molecules so the values will be non zero.