# I am unable to figure out the reason for first bond enthalpy being higher than second one in water

In $$\ce{H2O}$$, I am able to understand that the enthalpy needed to break two $$\ce{OH}$$ bonds are not the same. But I don't understand why is the bond energy for breaking the first $$\ce{O-H}$$ bond more than the second one.

My logic for second bond energy more than the first one is, when a $$\ce{H}$$ is removed from $$\ce{H2O}$$, The remaining $$\ce{O}$$ gets a negative charge and that ensures strong bonding with other $$\ce{H}$$ which will imply more energy for breaking the second bond.

First bond enthalpy is 502 kJ/mol and second bond enthalpy is 427 kJ/mol.

I would appreciate someone who could correct me

• After dissociation of the first bond, the remaining hydroxyl radical does not get a negative charge. Nov 7 at 9:25
• How can that happen? Nov 7 at 9:40
• You seem to think that water is made of ions, because it dissociates to ions (which it does). This is wrong. Nov 7 at 10:00
• Oh OK. But what's the reason for such trend, first bond enthalpy being more than second one Nov 7 at 10:02
• There is no reason to expect anything. O in OH is in a pathological state, never seen in a stable molecule. Nov 7 at 10:44