First Row BDE's: Why is H-NH2 lower than H-CH3?

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Discussion Overview

The discussion centers on the bond dissociation energies (BDEs) of first-row hydrides, specifically comparing H-NH2 and H-CH3. Participants explore the apparent contradiction in the expected ionic character and bond strength, as well as the implications of measurement errors in reported BDE values.

Discussion Character

  • Debate/contested
  • Technical explanation
  • Mathematical reasoning

Main Points Raised

  • One participant presents BDE values for H-CH3 (105 kcal/mol) and H-NH2 (103 kcal/mol), suggesting that the H-NH2 bond should be stronger due to higher ionic character.
  • Another participant points out that bond dissociation energies are typically reported for homolytic dissociation and questions the source of the initial values, noting discrepancies in trends from a referenced source.
  • A participant acknowledges the homolytic nature of the reported values and agrees that the BDEs of H-CH3 and H-NH2 are comparable, while questioning the larger difference in BDEs between N-H and O-H bonds compared to C-H and O-H bonds.
  • One participant emphasizes that measurement errors in BDEs are significant enough that the differences between H-CH3 and H-NH2 cannot be definitively established, suggesting that they are comparable.

Areas of Agreement / Disagreement

Participants generally agree that the BDEs of H-CH3 and H-NH2 are comparable, but there is disagreement regarding the implications of measurement errors and the relative strengths of these bonds compared to others, such as N-H and O-H.

Contextual Notes

Participants note that the measurement error in BDE values for C-H and N-H bonds is larger than the observed difference, which complicates definitive conclusions about their relative strengths.

Steven Hanna
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Here are some first row BDE's (in kcal/mol):
H-CH3 = 105 <---> [H]+ [CH3]-
H-NH2 = 103
H-OH = 119
H-F = 136 <---> [H]+ [F]-
This trend is often rationalized in terms of increasing ionic character (or with no-bond resonance). However, the H-NH2 BDE should have a higher ionic contribution than the H-CH3 BDE. Yet the H-NH2 bond is weaker. Can anyone explain this apparent contradiction?
 
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Bond dissociation energies are usually reported for homolytic bond dissociation: HF would dissociate to neutral H and F atoms, rather than H+ and F-. Where are you getting these values from? Here's one comparable source:

https://labs.chem.ucsb.edu/zakarian/armen/11---bonddissociationenergy.pdf

(with values in kJ/mol), where we note two things: First, the trend in the link provided is different (BDE of C-H is less than N-H), and second (and more importantly), the error in the measurements of CH and NH bond scission (8 kJ/mol) is larger than the difference between their BDE's (4 kJ/mol). So I think the best way to think about this is that H-CH3 and H-NH2 BDE's are comparable, meaning that the effects which stabilize the bonds in H2O and HF are of far less importance in CH4 and NH3.
 
the reported values are homolytic. I'll have to check with my prof. to see where he got them. but yes, I think you're right, the two bde's are comparable. but then why is the difference between N and O larger than that between C and O?

ps. sorry, <---> these are resonance arrows.
 
Again, referring back to my original post, whether the N-H BDE is higher or lower than the C-H BDE is a matter of who you ask (compare the numbers from the pdf I posted to the ones you posted). The point is that the measurement error is larger than the difference between them, so that you can't really say for certain which one is lower. The only thing you can conclude is that they are comparable. The takeaway from this is that the processes that drive the BDE's higher for HF and H2O are clearly not very important in CH4 and NH3 (or at least, they don't lead to substantial differences between CH4 and NH3).