Drago 39;s Rule Pdf

[In Libman]

I could not find much information regarding Drago's rule on the internet. Not even on Wikipedia. One of my acquaintances (who is a chemist) said that no such rule exists. But I didn't get much time to discuss the matter with him.

penta- and hexacoordination; tetracoordination with additioal lone pairs and related: attempt to form as many normal bonds with p orbitals as possible; keep one lone pair in an $\mathrms$ orbital if possible. Use remaining $\mathrmp$ orbitals to construct four-electron-3-centre bonds to the remaining atoms. (Predicted bond angles: diverse. $90^\circ$ going from 2e2c bonds to 4e3c bonds; $180^\circ$ between a pair of coordinating atoms contributing to the same 4e3c bond.)

Using the aforementioned 4e3c bonds nicely explains why electronegative atoms prefer to occupy these positions. You should always remember that 4e3c bonds can be described by the following resonance structures (drawn for $\ceClF3$):$$\ceF-Cl^+\bond...F- F^-\bond...Cl^+-F$$In this resonance, each outer atom taking part in a 4e3c bond has a formal average charge of $-0.5$; it is generally better to assign charges to electronegative atoms.

The reason for the special behaviour of nitrogen and oxygen is not their high electronegativity but their small size. Chlorine, which has an electronegativity similar to nitrogen, behaves as predicted for large atoms.

These rules go back to a publication by Russel S. Drago criticizing the use of the VSEPR model in Introductory Chemistry. As an alternative, he suggests a set of heuristics. First, if there are no lone pairs, use essentially the same rules as VSEPR. How to treat lone pairs is described below:

Here we shall consider central atoms with eight orless electrons about it. When there are lone pairs on thecentral atom, each lone pair counts as a group in theabove scheme if the central atom is a second row element.If the central element is a third, fourth, etc., row atomand if the groups attached to the central atom are oxygenor a halogen, the lone pair also counts as a group. If thegroups attached are less electronegative than bromine(that is, if almost anything other than a halogen or oxygenis attached), the lone pair occupies an unhybridized orbitaland does not count as a group. When the s orbital isused to accommodate a lone pair, only p orbitals remainfor bonding, and the geometry involves an arrangement of the groups at about 90 angles to each other. Some examplesare listed in Table 2.

He goes on to say that chemists, when they encounter a molecule they don't know, would draw a Lewis structure and compare it to those of molecules they know rather than trying to predict a structure on paper from first principles. Another quote:

Drago's rule is a principle in ecology that states that the larger species in a community will have a greater impact on the environment than smaller species. It is important in science because it helps us understand the dynamics of species interactions and their effects on ecosystems.

Drago's rule is related to competition because it suggests that larger species are better able to outcompete smaller species for resources in a given environment. This can lead to changes in species composition and diversity within a community.

No, Drago's rule is not applicable to all ecosystems. It is most commonly observed in terrestrial environments, but it can also be seen in aquatic ecosystems. However, there are exceptions to the rule and it may not hold true in all cases.

Drago's rule can provide a general framework for understanding the outcome of species interactions, but it is not a definitive predictor. Other factors such as environmental conditions and species adaptations also play a role in determining the outcome of interactions.

Drago's rule can be applied in conservation efforts by helping us identify which species may have a greater impact on the environment and which species may be more vulnerable to competition. This information can be used to inform management and conservation strategies for protecting species and maintaining healthy ecosystems.

The same logic holds true for all the other elements too. Now that we have established that the lone pair is NOT in any hybrid orbital but in a s- orbital, we can explain all the observations we listed in the earlier post.

As the lone pair is in a s- orbital, it is closer to the nucleus and so very tightly bound to it. Thus, this lone pair of electrons is NOT available for sharing or donating. So, it takes concentrated acid to react with phosphine. Also, as these lone pairs cannot be shared, these compounds do not form coordination compounds. Ammonia on the other hand can lend its lone pair to form coordinate covalent bonds and so can form coordination compounds.

As opposed to this, hybridization takes place in the elements of the 2nd period. Thus, the lone pairs occupy the hybrid orbitals, which have less %s character. Naturally, they are less tightly bound to the nucleus and are available for sharing or bonding.

Hi Bruno.. I have calculated the %s- character for all the three bonds in phosphine using the Draco rule. The value of s- character in each bond comes to 6%. Thus, for 3 P-H bonds the value will be 3*6 = 18 %.

The s- character in 3s2 orbital is much more than 50%.
So , if the s-character is more , we can safely conclude that the orbital cannot be a hybrid one.
In the case of phosphine, the lone pairs are in an orbital with 82% s- character. Thus, the orbital holding these lone pair of electrons cannot be a hybrid orbital. It HAS TO BE a regular pure s- orbital.

According to drago,s rule when the following conditions are satisfied, then the energy difference between the participating atomic orbitals will be very high and thus no mixing of orbitals or hybridization takes place.

In chemistry, Bent's rule describes and explains the relationship between the orbital hybridization and the electronegativities of substituents.[1][2] The rule was stated by Henry A. Bent as follows:[2]

Valence bond theory gives a good approximation of molecular structure. Bent's rule addresses disparities between the observed and idealized geometries.[3] According to Bent's rule, a central atom bonded to multiple groups will rehybridize so that orbitals with more s character are directed towards electropositive groups, and orbitals with more p character will be directed towards groups that are more electronegative. By removing the assumption that all hybrid orbitals are equivalent, Bent's rule leads to improved predictions of molecular geometry and bond strengths.[4][5] Bent's rule can be justified through the relative energy levels of s and p orbitals. Bent's rule represents a modification of VSEPR theory for molecules of lower than ideal symmetry.[6] For bonds with the larger atoms from the lower periods, trends in orbital hybridization depend strongly on both electronegativity and orbital size.

In the early 1930s, shortly after much of the initial development of quantum mechanics, those theories began to be applied towards molecular structure by Pauling,[7] Slater,[8] Coulson,[9] and others. In particular, Pauling introduced the concept of hybridisation, where atomic s and p orbitals are combined to give hybrid sp, sp2, and sp3 orbitals. Hybrid orbitals proved powerful in explaining the molecular geometries of simple molecules like methane, which is tetrahedral with an sp3 carbon atom and bond angles of 109.5 between the four equivalent C-H bonds. However, slight deviations from these ideal geometries became apparent in the 1940s.[10] A particularly well known example is water, where the angle between the two O-H bonds is only 104.5. To explain such discrepancies, it was proposed that hybridisation can result in orbitals with unequal s and p character. A. D. Walsh described in 1947[10] a relationship between the electronegativity of groups bonded to carbon and the hybridisation of said carbon atom. Finally, in 1961, Bent published a major review of the literature that related molecular structure, central atom hybridisation, and substituent electronegativities [2] and it is for this work that Bent's rule takes its name.

Bent's original paper considers the group electronegativity of the methyl group to be less than that of the hydrogen atom because methyl substitution reduces the acid dissociation constants of formic acid and of acetic acid.[2]

Bent's rule can be extended to rationalize the hybridization of nonbonding orbitals as well. On the one hand, a lone pair (an occupied nonbonding orbital) can be thought of as the limiting case of an electropositive substituent, with electron density completely polarized towards the central atom. Bent's rule predicts that, in order to stabilize the unshared, closely held nonbonding electrons, lone pair orbitals should take on high s character. On the other hand, an unoccupied (empty) nonbonding orbital can be thought of as the limiting case of an electronegative substituent, with electron density completely polarized towards the ligand and away from the central atom. Bent's rule predicts that, in order to leave as much s character as possible for the remaining occupied orbitals, unoccupied nonbonding orbitals should maximize p character.

Experimentally, the first conclusion is in line with the reduced bond angles of molecules with lone pairs like water or ammonia compared to methane, while the second conclusion accords with the planar structure of molecules with unoccupied nonbonding orbitals, like monomeric borane and carbenium ions.

Bent's rule can be used to explain trends in both molecular structure and reactivity. After determining how the hybridisation of the central atom should affect a particular property, the electronegativity of substituents can be examined to see if Bent's rule holds.

Valence bond theory predicts that methane is tetrahedral and that ethylene is planar. In water and ammonia, the situation is more complicated because the bond angles are 104.5 and 107 respectively, which are less than the expected tetrahedral angle of 109.5. One rationale for those deviations is VSEPR theory, where valence electrons are assumed to lie in localized regions and lone pairs are assumed to repel each other to a greater extent than bonding pairs. Bent's rule provides an alternative explanation.

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