Why is formamide planar

Connect and share knowledge within a single location that is structured and easy to search. However, analysis shows that the molecule is actually very nearly planar with bond angles close to degrees. EDIT: as suggested by Martin and another poster, hybridization is a rough concept. This however would still necessitate planarity, correct? Pi bonds are formed through the above and below electron-pairing in p-orbitals; effective bonding is achieved when these p-orbitals are parallel with respect to each other.

I'm thinking this has to do with the partial double bond character in the molecule also appears to be some ionic character to the molecule - likely due to electron-withdrawing effects of the nitrogen and the oxygen. This is the standard answer. However, would intramolecular hydrogen bonding also play a role? Couldn't there be a hydrogen bond between the peripheral hydrogen on the nitrogen and the oxygen?

Couldn't this also help with achieving the degree bond angles? Most of the amides are planar due to steric reasons the restriction may be lifted and so is also formamide. However, the inversion barrier for these molecules is depending on the substituents very low. You can now stabilise the intermediated structure with conjugation, and that is exactley the case here. In your chart, that refers to entry 2. All those resonance structures are only descriptions of extreme states, the truth lies between them.

The following scheme considers the most common ones and adds a third, that might explain delocalisation in a non-traditional Lewis way up to a certain visual point. In molecular orbital theory you can form 3-centre orbitals from all molecules perpendicular to the molecular plain. The HOMO is an in-plane lone pair of oxygen.

But I will not go into detail about that, because it would involve breaking away from the very handy hybridisation concept. Sign up to join this community. The best answers are voted up and rise to the top. Stack Overflow for Teams — Collaborate and share knowledge with a private group.

Create a free Team What is Teams? Learn more. The positions of H atoms are the consequence of a balance between intra-molecular forces and inter-molecular forces, so that relying on geometric considerations alone without insights into the true balance between these two effects is unlikely to be reliable. Thus, another aim of this work is to provide quantitative insights on the extent to which the planarity of amide groups may be deformed under the influence of intermolecular hydrogen bonding, in order to assist the interpretation of the geometries of such groups when carrying out structure determination from powder XRD data.

For this study, we focus on formamide and urea as model amide molecules, recognizing as stated above that the degree of planarity of the amide group in the isolated molecule is significantly different in the case of formamide and urea. Furthermore, we also consider glycyl glycine as an example of an amide group in a model peptide. We consider complexes comprising these molecules together with one or more hydrogen-bond acceptor or hydrogen-bond donor molecules, focusing on hydrogen cyanide NCH as a model hydrogen-bond acceptor and hydrogen fluoride HF as a model hydrogen-bond donor.

Rather than searching for minima on the potential energy surface, we take the approach of fixing the position of the hydrogen-bond donor or acceptor to explore the response of the amide in a systematic and controlled manner.

In each case, the linear geometry of these molecules ensures that the results are not influenced by the formation of secondary interactions. In the case of glycyl glycine, water is also considered as a hydrogen-bond donor.

In this study, we use E bind to denote the binding energy of the complex, representing the total energy associated with formation of the geometry-optimized complex starting from the individual isolated molecules in their ground-state geometries. The binding energy E bind can be broken down into two separate contributions denoted E def deformation energy and E int interaction energy.

The most striking geometric feature is the non-planarity of the NH 2 group, which adopts a slightly asymmetric pyramidal form, evident from the dihedral angles involving the H atoms in the syn and anti positions. The degree of non-planarity can be further quantified by the sum of the angles around the N atom, which equals The formation of the hydrogen bond with HF also leads to a decrease in r C—O and an increase in r C—N , suggesting that the delocalization across the amide bond is disrupted.

While B3LYP performs well for hydrogen bonding, we felt it necessary to examine the dependence of these properties on the method used. The components of this binding energy give further insights into the nature of this interaction: the deformation energy E def is 1.

Thus, it is clear that the energy required to deform the individual molecules into the geometries adopted in the hydrogen-bonded complex is more than compensated by the stabilising energy of the hydrogen bond. Possible underlying reasons are explored in more detail below. In addition to acting as a hydrogen-bond acceptor through the N atom, formamide can also act as a hydrogen-bond donor involving one or both of the N—H bonds, and we have studied this type of interaction using hydrogen cyanide NCH as a model acceptor.

With a single NCH molecule, the H atom involved in the hydrogen bond moves out of the plane rather more than the non-interacting H atom, which also moves out of the plane even though it is not close to the NCH molecule.

The configuration in which both NCH molecules are in the same plane as formamide is confirmed as a true energy minimum by harmonic frequency calculation. Geometric and energetic aspects of this complex are given in Table 2. As observed for formamide , the r C—O and r C—N distances also change substantially; in particular, r C—O decreases and r C—N increases, which may indicate that the hydrogen bonding leads to a reduction in the delocalization over the OCNH unit. Interestingly, the NH 2 group that is not involved in hydrogen bonding is perturbed by the hydrogen-bonding interaction of the other NH 2 group with HF.

Complexes of urea with NCH as a hydrogen-bond acceptor were also studied, and the results are reported in Table 2.

As observed for formamide , the geometry of the NH 2 group of urea is sensitive to the position of the NCH molecule. Table 2 also reports results for a twisted complex, with one NCH molecule above the plane of the non-H atoms of urea and one NCH molecule below this plane. Again, the NH 2 group deforms in order to establish hydrogen bonding, resulting in one negative dihedral angle and one positive dihedral angle.

Thus, urea is a better hydrogen-bond acceptor but a worse hydrogen-bond donor than formamide. The deformation energies for urea are generally slightly smaller than for formamide. The overall value of 0. Several geometric parameters reported in Tables 1 and 2 suggest that hydrogen bonding affects the electronic delocalization across the amide bond. In the isolated formamide molecule, the C—O and C—N bond orders support the delocalized picture of bonding.

For the isolated urea molecule, the charges on the O and N atoms are more negative, and the C—O and C—N bond orders are lower, than in formamide. We propose that these differences are responsible for the contrasting degrees of planarity of the isolated formamide planar and urea non-planar molecules; thus, each N atom in urea donates some electron density to the C O group, but the capacity of the C O group to accept this electron density is limited.

These changes may be interpreted as evidence that hydrogen bonding to one N atom allows the other N atom to participate in greater delocalization with C O, in general agreement with the geometric changes noted in Table 2.

The reason underlying this difference is not immediately clear, but may indicate that the H atoms in an amide group are relatively free to move to maximise their strength as hydrogen bond donors without overall disruption of electronic structure, whereas for the N atom to act as a hydrogen bond acceptor inevitably requires substantial electronic rearrangement throughout the group.

Thus, the zero-point energy is insufficient to overcome the barrier, but at sufficiently high temperature, thermal motion may be expected to lead to the averaged structure being observed. Hydrogen bonding to HF alters this picture drastically. Thus, even at elevated temperatures, we predict that hydrogen bonding should promote the non-planar form of urea.

As discussed above, the crystal structure of urea has a perfectly planar C 2v structure of the urea molecule, with H atom positions determined accurately from neutron diffraction data. The clusters are constructed to introduce the local crystal environment for a central urea molecule in a systematic manner. The clusters were constructed to introduce the local crystal environment for a central urea molecule in a systematic manner.

In both cases, the other urea molecules that represent the crystal environment were held fixed at their geometries in the crystal structure at 12 K. These data show that constraining non-H atoms to their positions in the crystal structure affects even the planarity of a single urea molecule, with the fully optimised molecule exhibiting slightly larger dihedral angles.

The presence of a single neighbour is sufficient to reduce the non-planarity of the urea molecule. In the free optimisation of the dimer, the free urea molecule moves out of the plane of the dimer and the hydrogen bonds between the molecules are elongated relative to the crystal structure see Fig.

The constrained dimer retains the overall pattern of hydrogen bonding observed in the crystal structure, and exhibits larger changes towards planarity in the dihedral angles for both the syn and anti H atoms. In both optimisations, the syn H atoms are slightly further out-of-plane than in the monomer or dimer, whereas the anti H atoms lie slightly more in the plane than in the dimer.

This trend is continued in the pentamer, which is the only structure considered that includes hydrogen bonds to both sets of H atoms of the central molecule.