Primary amines (2 Viewers)

[=)(=]

Would they have 2 H-bonds per molecule cause of the two hydrogens on nitrogen or is it only 1 H bond cause nitrogen only has 1 lone pair?

 

[SadCeliac]

Would they have 2 H-bonds per molecule cause of the two hydrogens on nitrogen or is it only 1 H bond cause nitrogen only has 1 lone pair?

jeez that is a huge atar aim - how are you tracking for that

I always thought it was 2 (1 per hydrogen):
H bonds form between an H (attached to N, O, F --electroneg difference needs to be great enough) and so the H's of an amine can form a hydrogen bond each...?

update just found this diagram if it helps???

 

[its_ace21]

jeez that is a huge atar aim - how are you tracking for that

says the one with a 98 aim and getting 69 in math papers

 

[=)(=]

jeez that is a huge atar aim - how are you tracking for that

I always thought it was 2 (1 per hydrogen):
H bonds form between an H (attached to N, O, F --electroneg difference needs to be great enough) and so the H's of an amine can form a hydrogen bond each...?

update just found this diagram if it helps???
View attachment 39640

At the moment its possible just gotta grind rlly hard, lol

Ahh interesting so it seems only 1 H bond, cheers

 

[wizzkids]

This is not a valid question. Perhaps you have forgotten the fundamental nature of hydrogen bonding. Hydrogen bonds only last for between one to ten picoseconds. They are highly transitory and subject to statistical distribution and therefore cannot be counted as such. All we can say is there is a certain probability of an average number of hydrogen bonds per molecule, and this gives rise to increased attractive forces between molecules on average. We know how many charge centres are located on the molecule, but we can't say for sure how many hydrogen bonds a molecule will have at any given instant.
Experiments and calculations done on liquid water showed that although each water molecule is capable of forming four hydrogen bonds, in fact this is highly dependent on time and temperature, and at 25^oC water forms about 3.5 hydrogen bonds per molecule.
We can say for sure that primary amines are more strongly attracted than secondary amines, for the same molecular weight, but you can't count the bonds that will occur on any given molecule; this is simply an invalid concept.

 

[SadCeliac]

This is not a valid question. Perhaps you have forgotten the fundamental nature of hydrogen bonding. Hydrogen bonds only last for between one to ten picoseconds. They are highly transitory and subject to statistical distribution and therefore cannot be counted as such. All we can say is there is a certain probability of an average number of hydrogen bonds per molecule, and this gives rise to increased attractive forces between molecules on average. We know how many charge centres are located on the molecule, but we can't say for sure how many hydrogen bonds a molecule will have at any given instant.
Experiments and calculations done on liquid water showed that although each water molecule is capable of forming four hydrogen bonds, in fact this is highly dependent on time and temperature, and at 25^oC water forms about 3.5 hydrogen bonds per molecule.
We can say for sure that primary amines are more strongly attracted than secondary amines, for the same molecular weight, but you can't count the bonds that will occur on any given molecule; this is simply an invalid concept.

but... whilst it may be an invalid concept it still makes it much easier to deduce what will be more attracted to what - so what's the issue?

 

[=)(=]

ohh okie but like how would you describe the differences between amide and carboxylic acids if you cant talk about number of H bonds tho?

 

[wizzkids]

All hydrogen bonds are not the same. We can assign a bond enthalpy to the hydrogen bonds, and typical enthalpies can range from -1 or -2 kJ/mol up to as much as -25 kJ/mol. You can talk about the potential for more hydrogen bonding between molecules. The difference between amide and carboxylic acid hydrogen bonding revolves around the strength of the hydrogen bonding; differing electronegativities of the oxygen and the nitrogen atoms give rise to different bond enthalpy.
Data on boiling points suggests than amides have stronger hydrogen bonding than carboxylic acids of equivalent molecular weight due to the higher hydrogen bond enthalpy, e.g. ethanamide CH₃CONH₂ b.p. 221^oC versus ethanoic acid CH₃COOH b.p. 118^oC

 

[Luukas]

Would they have 2 H-bonds per molecule cause of the two hydrogens on nitrogen or is it only 1 H bond cause nitrogen only has 1 lone pair?

As the diagram in SadCeliac's post (reproduced below) shows, for an H-bond you need both the hydrogen bond donor (the H atom bound to the very electronegative N, O, or F atom) and the hydrogen bond acceptor (the lone pair on the second very electronegative atom). So, a primary amine like methanamine can, with two donors but only one receptor, can only average one H-bond per molecule. There will be some molecules with both H's involved in an H-bond, and some with none, but any average of one per molecule is reasonable.

For this reason, an alcohol and an amine will have similar amounts of H-bonding. The alcohol will typically have the higher BP as the H-bonds in the amine are weaker (O is more electronegative than N) and because of steric factors (it is easier to get the one donor H of an alcohol near to one of the two acceptor lone pairs than it is to get one of the two donor H's of a primary amine near the one acceptor lone pair.

View attachment 39640

In the common year 11 question showing the BP trends of the group 4, 5, 6, and 7 hydrides against molar mass (where HF, H₂O, and NH₃ all have much higher BPs than any of the other species due to H-bonding), the BP(H₂O) = 100 ^oC is greater than both BP(HF) = 20 ^oC and BP(NH₃) = -33 ^oC because only water has 2 donors and 2 acceptors and is thus capable of double the H-bonding of HF (1 donor, 3 acceptors) or NH₃ (3 donors, 1 acceptor). The graph is available on the Doc Brown Chemistry notes website at https: //docbrown.info/page07/equilibria8f.htm .

 

[Luukas]

All hydrogen bonds are not the same. We can assign a bond enthalpy to the hydrogen bonds, and typical enthalpies can range from -1 or -2 kJ/mol up to as much as -25 kJ/mol. You can talk about the potential for more hydrogen bonding between molecules. The difference between amide and carboxylic acid hydrogen bonding revolves around the strength of the hydrogen bonding; differing electronegativities of the oxygen and the nitrogen atoms give rise to different bond enthalpy.
Carboxylic acids have much higher boiling points than amides of equivalent molecular weight due to the higher hydrogen bond enthalpy.

I've been taught that amides have higher BP than carboxylic acids of similar molar mass.

This is certainly true in the example in the Caringbah 2023 Chemistry Trial that was recently uploaded, question 34, which gives the following data:

• Propanoic acid, CH₃CH₂COOH, has molar mass 74.08 g mol^-1 and BP = 141.2 ^oC.
• Propanamide, CH₃CH₂CONH₂, has molar mass 73.09 g mol^-1 and BP = 213 ^oC.

Its explanation is flawed, though, as it attributes the difference to dimerisation in the amide, but this also occurs in the carboxylic acid.

Similar data for the two carbon compounds:

• Acetic acid, CH₃COOH, has molar mass 60.052 g mol^-1 and BP = 118 ^oC.
• Acetamide, CH₃CH₂CONH₂, has molar mass 59.07 g mol^-1 and BP = 222 ^oC.

And for the one carbon cases:

• BP(formamide, HCONH₂) = 210 ^o > BP(formic acid, HCOOH) = 100.8 ^oC.

 

[=)(=]

Can we just simply say amides have more extensive Hydrogen bonding networks than carboxylic acids due to the potential to form more hydrogen bonding?