Вот мне всегда было интересно, почему P=O связь такая стабильная? По сути разница радиусов/размеров атомов фосфора и кислорода достаточно большая, потому перекрывание их орбиталей будет плохим, вдобавок разница электроотрицательности всего лишь 1,4 что по мне достаточно слабо чтобы иметь большой ионный характер и электрическое притяжение. Да даже, оксид фосфора P4O6 не содержит двойных связей.
@eudaimon provide a detailed explanation of stability of P=O double bond, include different points of view and theories
Indeed, the stability of the phosphorus-oxygen double bond (in P=O) is an interesting subject of inquiry. To understand this stability, one must consider several factors, including molecular orbital theory, resonance, and the nature of the atoms involved.
Firstly, let’s consider the molecular orbital theory. Phosphorus and oxygen can form a double bond consisting of one sigma (σ) bond and one pi (π) bond. The σ bond is formed by the overlap of sp^3 hybrid orbitals from phosphorus and sp^2 or p orbital from oxygen. The π bond is formed by the side-by-side overlap of p orbitals.
The difference in electronegativity between phosphorus (2.19) and oxygen (3.44) is significant enough to create a polar bond, which increases the bond strength due to the added ionic character. Although the size mismatch between the phosphorus and oxygen atoms might suggest less effective orbital overlap, in practice, the π bond is sufficiently strong, possibly due to effective p-p orbital overlap and the ability of phosphorus to expand its octet and accommodate additional electron density.
Furthermore, resonance stabilization plays a crucial role. Phosphorus compounds can exhibit resonance with structures that show a positive charge on phosphorus and a negative charge on oxygen (for example, in phosphates). This delocalization of electrons adds stability to the P=O bond.
Lastly, in the context of thermodynamics, the P=O bond is strong, with a bond dissociation energy typically higher than that of a P-C or P-H bond. This suggests that once formed, the P=O bond is relatively stable and not easily broken.
Regarding the comparison with P_4O_6, it’s important to note that in this molecule, the phosphorus atoms are in a different oxidation state and are bonded to each other, forming a more complex structure with single P-O bonds. The lack of P=O double bonds in P_4O_6 does not directly relate to the stability of P=O bonds in other compounds, such as phosphine oxides or phosphates, where the bonding environment is different.
Thus, the stability of the P=O bond is a result of several contributing factors, including covalent bonding with some ionic character, possible resonance structures, and molecular orbital considerations.
1 лайк
I do not think you should have included that. I asked about why is it stable, not whether it is stable or not, so your argument is irrelevant as the fact that this bond is stable is already implied
(простите за потраченные центы)
Apologies for any confusion. Let us delve deeper into the essence of the stability of the P=O bond, focusing on the theoretical underpinnings.
The P=O bond possesses stability for several reasons:
-
Orbital Overlap: Despite the size difference between phosphorus and oxygen, the orbitals can still overlap effectively. Phosphorus can utilize d-orbitals to expand its valence shell and form stronger pi bonds with oxygen.
-
Electronegativity and Polarity: The electronegativity difference leads to a polar bond, where the electron density is drawn towards the more electronegative oxygen. This polarization can enhance the stability of the bond through increased electrostatic attraction between the positively charged phosphorus and the negatively charged oxygen.
-
Resonance and Delocalization: In many phosphorus-oxygen compounds, the presence of resonance structures can delocalize the electron density over the molecule, thereby stabilizing the bond. For example, in phosphate ions, the P=O bond is one of several resonance structures that delocalize charge and contribute to bond stability.
-
Bond Energy: From a thermodynamic perspective, the bond dissociation energy of the P=O bond is quite high, indicating a stable bond. This is a result of the factors above, as well as the intrinsic strength of the double bond formed between the two atoms.
-
Electron Donation and Back-donation: Phosphorus, being less electronegative, can act as an electron donor, while the oxygen atom can act as an electron acceptor. There can also be back-donation from the oxygen’s lone pairs to the empty d-orbitals of phosphorus, which can strengthen the bond.
-
Comparison with P=O in Oxides: In molecular species like P_4O_6, the absence of P=O double bonds is due to the specific bonding requirements and the structure of the molecule, which prefers P-O single bonds to satisfy its stability. This does not undermine the stability of P=O double bonds in other molecular environments where such double bonds are favorable.
The stability of the P=O bond, therefore, arises from a combination of quantum mechanical effects, including orbital interactions and electron delocalization, as well as the inherent properties of the atoms involved.
1 лайк
@eudaimon do you happen to know any specific insights on the nature of P=O bond from the electronic structure theory? Maybe there are some contributions from d orbitals? Maybe there are other quantum effects? When describing those be very specific and precise, do not make overly general statements. Respond as if you’re talking to someone who is doing PhD in electronic structure theory.
1 лайк
Certainly, Anton. When considering the nature of the P=O bond from an electronic structure theory perspective, we can dive into a more rigorous quantum mechanical description.
-
Orbital Contributions: Traditional descriptions often invoke d-orbital participation on the part of phosphorus to rationalize the formation of the P=O π bond. However, contemporary computational studies using methods like Natural Bond Orbital (NBO) analysis suggest that the role of d-orbitals in main group chemistry is often overstated. In the case of the P=O bond, the bonding can largely be described by the overlap of p-orbitals on oxygen with the sp^3 hybridized orbitals on phosphorus for the σ bond and p-orbitals for the π bond, without significant d-orbital contribution. The effective overlap for the π bond is still a matter of debate, but it seems that the p-p π bonding is sufficient to account for the observed stability without invoking d-orbitals.
-
Electron Delocalization: Through methods like Natural Resonance Theory (NRT) analysis, one can observe the degree of electron delocalization. In phosphorus oxides and oxyacids, there is a delocalization of electron density from the lone pairs on oxygen into the antibonding orbitals of the P-O σ bonds, which contributes to the double bond character and overall stability of the P=O bond. This delocalization is a stabilizing factor and is well-represented in resonance structures in valence bond theory.
-
Electronegativity and Charge Transfer: More electronegative oxygen atoms attract more electron density in the bond, which can be quantified using the concept of partial charges derived from Mulliken or Löwdin population analyses. This results in a polar covalent bond, where the polarization itself can stabilize the bond due to the partial ionic character.
-
Quantum Chemical Calculations: Advanced ab initio or density functional theory (DFT) calculations, which include electron correlation effects, can provide insights into the nature of the bonding. Methods such as coupled-cluster with single, double, and perturbative triple excitations [CCSD(T)] can yield accurate bond energies and electron density distributions, which confirm the stability and nature of the P=O bond.
-
Bonding Energy Considerations: The bond dissociation energies calculated from quantum chemical methods can be compared with experimental values to confirm the strength and thus the stability of the P=O bond. These calculations can be refined by including zero-point vibrational energy corrections and taking into account basis set superposition error (BSSE).
In summary, from an electronic structure theory perspective, the stability of the P=O bond can be attributed to effective p-p orbital overlap for both σ and π components of the bond, significant polarization due to electronegativity differences, and electron delocalization that is consistent with resonance structures. These factors, validated by quantum chemical calculations, support the observed stability of the P=O double bond without necessitating a significant role for d-orbital participation.
2 лайка