The Electron-Pair Repulsuion Model for Molecular Geometry (Gillespie, 1970)

Our first “real” inorganic chemistry paper, or is it? You decide…

Thank you all for hanging in there as we moved through the early topics. We will now be moving out from the nucleus to explain why complexes have the shapes that they do. We will try to cover topics that include coordination number, coordination geometry, Werner complexes and other essential pieces of introductory inorganic chemistry.

Gillespie is one of the few remaining pre-war figures in inorganic chemistry. Still at McMaster University in Canada, his insights have been proven to be greatly useful in the field. His wikipedia page does not do him justice, so make sure you dig futher into his work.

The paper will be discussed on either Friday or next Monday.

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9 Responses to “The Electron-Pair Repulsuion Model for Molecular Geometry (Gillespie, 1970)”

  1. Hunter Burgin Says:

    Based on my secondary reading online it is clear to me that Dr Ronald James Gillespie is an incredibly important figure in the world of chemistry, specifically in the concentration of molecular geometry. I found it quite interesting to read about RJ Gillespie specifically because he is the first scientist we have read about that is still alive and making contributions to the scientific community.
    Gillespie’s 1970 paper explains the molecular geometry of a multitude of chemicals and chemical compounds as determined by their electron domains and electron pair repulsions. The paper constantly sites the valence shell electron pair repulsion theory (VSEPR, also known as the Gillespie-Nyholm theory) as the deciding factor in determining the geometric configuration of various chemical compounds ranging from anionic compounds, ionic compounds, triple bonds, double bonds, and even quadruple bonds. Although I do not doubt Dr Gillespie’s findings regarding the electron configurations found based upon his theory, I would have liked to have seen some form of experimental evidence supporting findings. Did he use X-Ray crystallography? Was there a mathematical formula developed to support his geometric findings for the various chemical structures listed in his paper?
    On a side note I also found it amazing that a chemical such as B2H6 could exist where two hydrogens could form four bonds on either side creating a bridge from one bromine to the other.

    • Adam Settimo Says:

      I agree, I would like to see what kind of experimental work was done, and the results. Along those lines, I’m a little confused by table 2. Their are two different bond angles shown for H2C=CF2 that don’t match. Why?

  2. Adam Settimo Says:

    The paper states that it is well know that electron pair arrangements and the general shape of a molecule can be predicted if the bonding and non bonding elections of the central molecules valence are know. My question is if this takes into account numbers of bonds that have multiple options for their configuration? For example, 4 and 5 bonds(to a central atom) have two different configurations, each. It has been a while since I studied this in depth, but is there a way, with the above information to determine whether a central atom with 5 bonds (paired or unpaired) for instance, would “choose” square pyramidal vs. Trigonal biparamidal? It is shown that lone pairs occupy the most amount of space and other bonds will stretch or bend away to give them more room. For this example, would a five bonded valence shell prefer square pyramidal to give the lone pairs the most amount of space?

  3. Porter Marsh Says:

    Gillespie gives the equation F=1/(r^n) to define the repulsive force between electrons. He defines r as the distance between the two orbitals but he doesn’t define what n is. For most cases, n is equal to infinity but for some cases (7 or 10 electron pairs) it is not. Gillespie claims both that “there seems to be no reliable way of estimating a reasonable value of n for any particular case” and that the value of n can decrease. If it is not a fixed value then it is not a constant but if we can’t estimate it then it’s not a calculated value either. Since he doesn’t define it and it doesn’t seem to be knowable, what was he talking about? The value n is part of his law for electron repulsion so I think it deserves more of an explanation than it’s given.

  4. Kevin Greenwood Says:

    This really cleared up some misconceptions I had about VSEPR theory. This paper particularly helped me understand that much of what dictates a molecule’s structure depends on orbital interactions as they occur near or at the surface of the central atom.
    This leads me to the shape of borane as it is depicted is this paper. I was having trouble understanding how one hydrogen with a single 1s orbital could covalently bond to two atoms. After some looking into MO theory as it applies to borane (link at bottom of post), this picture as it is drawn isn’t quite right. If each of the bridging sp3 orbitals from boron is hybridizing with the 1s orbitals from the bridging hydrogens, then according to VSEPR theory, the bond angle at (or very near) each hydrogen should be around 180°, overall giving each bridging bond a banana shape. It may be a small detail but I thought it was confusing.
    Another molecule with an interesting shape is tetraboron tetrachloride. The diagram shows the three center bonds as the faces of the tetrahedron formed by the four boron atoms. As shown, one would expect electron density to be low at the vertices of the tetrahedron as there are four bonding orbitals converging at each. Each bonding plane has a three-sided inverted pyramid with electrons for faces at it’s back, and with the atomic radius of boron being relatively small, shouldn’t each face be bulged out and away from the molecule’s center? The only thing I see that would prevent this would be the wag of the B-Cl bonds, but that seems pretty far to bend even considering the increased size of the chlorine atom relative to boron. If anyone could explain why these bonds are planar (if indeed they are), I would appreciate it.

    (http://www.chem.ucla.edu/~cantrill/30A_F05/Diborane.pdf)

  5. Josh Ellsworth Says:

    I was confused by the explanation given for the formation of square planar complexes with transition metals. Gillespie seems to postulate that interaction with the d shell causes an octahedral complex to lose its axial ligands and electron pairs, thereby assuming the square planar shape. While this seems more likely to lead to a planar molecule than the distortion of a tetrahedral arrangement by the d shell, ( which seems to be more likely to only form a squarish, kinda planar complex), I have a hard time visualizing the axial ligands simply leaving the party with an electron pair in tow. Does x-ray crystallography of square planar compounds (coordinated to the same species of ligand) show a high degree of symmetry? If so, are there any methods available to trap the proposed octahedral-square planar intermediate (there’s got to be an intermediate/ transition state)? I was also struck by the explanation of the octahedral symmetry of the Tellurium-Bromine complex. If the lone pair truly is “dropped” into an inner shell, does that give a benefit that can be exploited?, or is it just an interesting side note that illustrates how fungible electron orbits become as the elements negotiate the 5p block? Does the p shell allow for the formation of a hybridized orbital, or is it more correct to picture the 6s shell as being close enough in energy to provide the orbital?

  6. Daniel Begay Says:

    At first, I was a bit worried to read through this paper because it looked extremely dense and I thought I would be lost reading through it. I was wrong because the paper does an excellent job of explaining theories such as the VSEPR theory clearly. It was able to give me a quick refresher on coordination geometry of inorganic chemistry
    It’s also interesting to see how these theories work with cluster compounds. It shows two metal cluster different in structure. They do not necessarily have a center atom to work with, but about 6 metal atoms bonding together. By taking a look at both cluster compounds, By looking at the position, does the Mo6Cl8 have a lone pair on each metal atom that positions the ligands to be as far apart as possible? I’m curious because the paper describes the metal atoms as octohedron to each other but does not describe the position of the ligands to individual metal atoms.

  7. Antony Says:

    It is interesting to see what Inorganic is about (right? Now we’re on it?). My first question involves the axial and equatorial atoms mentioned in the “Trigonal Bipyramidal Molocules” section. There is two questions that I have but they aren’t related to each other so don’t take them as a package but individually. First, how do they assign one over the other, axial over equatorial, if it all depends what position you look at it? Second, what would put the axial atom/group further than the equatorial; does it relate in any way to the energy difference in cyclohexane chair conformation energy differences?

  8. Antony Says:

    It is interesting to see what Inorganic is about (right? Now we’re on it?). My first question involves the axial and equatorial atoms mentioned in the “Trigonal Bipyramidal Molocules” section. There is two questions that I have but they aren’t related to each other so don’t take them as a package but individually. First, how do they assign one over the other, axial over equatorial, if it all depends what position you look at it? Second, what would put the axial atom/group further than the equatorial; does it relate in any way to the energy difference in cyclohexane chair conformation energy differences?

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