Talk:Molecular geometry

message: I would like a section to be added on the quantum-mechanical view of molecular structure. Atoms (or nuclei) are not point particles and the equilibrium geometries calculated by quantum chemists (minima of potential energy surfaces) are a theoretical construct which arises because of the Born-Oppenheimer approximation. "In reality" nuclei are also delocalised, and in some situations (e.g., highly excited ro-vibrational states of molecular species such as H3+) the classical concept of molecular structure stops being relevant (nuclei are very highly delocalised). Also it is worth mentioning some apparent paradoxes: e.g., that molecules in a rotationless state have spherical symmetry; that is, if you plotted the electronic & nuclear densities you'd have a blob with no structure! Of course one can then consider the expectation values of inter-atomic distances (and also possible higher momenta of such operators) and the classical concept is then restored.
I have no time at the moment to write on such a topic but I may in the future. L0rents (talk) 15:43, 25 February 2010 (UTC)
thread_title: Quantum view of molecular structure

message: I added different 3D representations of molecular geometry.  It could go in Molecular graphics but we couldn't really talk about molecular geometry without models or some kind of picture. 
Mcpazzo (talk) - 21:08 8 Oct 2008 PDT  —Preceding undated comment was added at 05:09, 9 October 2008 (UTC)
replies: text: Good stuff.  Ribbon style would also be nice.  And another name for the (volumetric) sticks model is "licorice", I believe.  —63.249.110.34 (talk) 08:43, 18 March 2009 (UTC)
thread_title: 3D representations added

message: I didn't like wobbling for two reasons: first the statement is not true. At room temperature
molecules do not "wobble"  in a way that their geometry is not determined, and second the concept "wobble" is ill-defined. My first reaction was to change the offensive sentence, but then I thought that it would perhaps be more illuminating to explain my objections by adding a short section. --P.wormer 11:50, 11 January 2007 (UTC)
replies: text: On second thought I came to the conclusion that wobbling means rotating and that geometries are confused with spectra.  Most spectra are broadened with increasing temperatures and the errors in the   measured geometries become larger for higher temperatures, but this does not mean that the concept of molecular geometry loses its value. (My confusion wouldn't have arisen if proper scientific terminology was used instead of the silly word wobbling. Lesson 101: use proper terms). --P.wormer 13:06, 11 January 2007 (UTC)

text: * Removed "wobbling" and adapted section about temperature dependence of geometry.--P.wormer 16:30, 10 February 2007 (UTC)
thread_title:  Wobbling 

message: I noticed that all geometries listed in this article provide examples apart from octahedral geometry.
My first thought was the common example XeF6, xenon hexafluoride.
Would anyone object to the addition of this example for octahedral geometry?
Marshmellis 01:25, 21 October 2007 (UTC)
thread_title:  Octahedral geometry 

message: This was added today at the top of the article. It is clearly not appropriate there, but someone might find a use for it. --Bduke (talk) 07:20, 17 July 2008 (UTC)
thread_title: Recent edit

message: Coordination geometries in TM complexes and in solid state are not well catered for in this article and in the related articles on specific geometries. This article is heavily VSEPR in focus, educationally not necessarily a bad thing, but restricts its applicability to main group and molecular species.  I can't see the point of a new series of articles on coordination geometries (see polyhedral symbol for list) as there would be a lot of overlap- shouldn't this article be expanded in scope and ALL of the geometry articles changed appropriately- rather than start anew?   --Axiosaurus (talk) 10:12, 1 August 2008 (UTC)
thread_title:  Molecular Geometry- Coordination Geometry 

message: For a long time now, I have been reverting IP editors who come along and alter the bond angle in the column labeled "ideal bond angle" (e.g,  109.5 for bent) and replace it with the actual bond angle for, e.g. water. This is in spite of comments I added by the angles and in spite of a paragraph above the table explaining this. This is getting boring. I have mofified the table to give both the ideal and the actual angles. I hope this stops this minor vandalism. Other ideas are welcome. --Bduke  (Discussion) 20:24, 27 January 2009 (UTC)
replies: text:  I actually felt that these edits were real vandalism and not well meant edits (after all, the column title clearly states "ideal bond angles"). Having ideal angles and real angles of the example molecules in one column is very confusing. Instead, we could add the real angles in a new column after the molecules. Cacycle (talk) 23:06, 27 January 2009 (UTC)
replies: text: Well that is the first time in three years that someone has thought I was a vandal-:) The problem has been that IP editors keep ignoring the heading and changing the angles for water and ammonia. Maybe my change in the heading of that column could be improved, but I think your idea of a separate column after the molecules is a better one. I would still put money on somebody changing the 109.4 for water to 104.5 sometime soon. I have a look at it when I can but life is wearing here as it is over 40C today and will be the same everyday until Saturday. --Bduke  (Discussion) 07:51, 28 January 2009 (UTC)
replies: text:  Not your edits - I was talking about the anonymous edits :-) We can put a html comment into the source text. Cacycle (talk) 13:54, 28 January 2009 (UTC)
thread_title:  VSEPR Table 

message: Forgive me if this is not the proper way to add a discussion topic, but it should be noted that "Electron geometry" re-directs to this page for some reason, I don't know about the rest of you, but I do not believe them to be the same, nor does my University professor and the authors of my Chemistry book. I don't know if there exists a page on Electron geometry, whether there is or not, this should be fixed.  —Preceding unsigned comment added by 207.161.146.197 (talk) 23:22, 20 November 2009 (UTC)
replies: text: An electron is a point particle, and has no geometry. I believe the redirect was created in view of molecular geometry being dictated by the behaviour of electron clouds orbiting the nuclei.—Tetracube (talk) 23:32, 20 November 2009 (UTC)
thread_title: Electron geometry

message: I've added a new See also, I dont know how to create an article but if someone know to do it, please, I think it is interesting:
Foundation
We begin by assuming a Lewis structure model for chemical bonding based on valence shell electron pair sharing and the octet rule. We thus assume the nuclear structure of the atom, and we further assume the existence of a valence shell of electrons in each atom which dominates the chemical behavior of that atom. A covalent chemical bond is formed when the two bonded atoms share a pair of valence shell electrons between them. In general, atoms of Groups IV through VII bond so as to complete an octet of valence shell electrons. A number of atoms, including C, N, O, P, and S, can form double or triple bonds as needed to complete an octet. We know that double bonds are generally stronger and have shorter lengths than single bonds, and triple bonds are stronger and shorter than double bonds.
Goals
We should expect that the properties of molecules, and correspondingly the substances which they comprise, should depend on the details of the structure and bonding in these molecules. The relationship between bonding, structure, and properties is comparatively simple in diatomic molecules, which contain two atoms only, e.g. HCl or O2. A polyatomic molecule contains more than two atoms. An example of the complexities which arise with polyatomic molecules is molecular geometry: how are the atoms in the molecule arranged with respect to one another? In a diatomic molecule, only a single molecular geometry is possible since the two atoms must lie on a line. However, with a triatomic molecule (three atoms), there are two possible geometries: the atoms may lie on a line, producing a linear molecule, or not, producing a bent molecule. In molecules with more than three atoms, there are many more possible geometries. What geometries are actually observed? What determines which geometry will be observed in a particular molecule? We seek a model which allows us to understand the observed geometries of molecules and thus to predict these geometries.
Once we have developed an understanding of the relationship between molecular structure and chemical bonding, we can attempt an understanding of the relationship of he structure and bonding in a polyatomic molecule to the physical and chemical properties we observe for those molecules.
Observation 1: Geometries of molecules
The geometry of a molecule includes a description of the arrangements of the atoms in the molecule. At a simple level, the molecular structure tells us which atoms are bonded to which. At a more detailed level, the geometry includes the lengths of all of these bonds, that is, the distances between the atoms which are bonded together, and the angles between pairs of bonds. For example, we find that in water, H2O, the two hydrogens are bonded to the oxygen and each O-H bond length is 95.72pm (where 1pm=10-12m). Furthermore, H2O is a bent molecule, with the H-O-H angle equal to 104.5°. (The measurement of these geometric properties is difficult, involving the measurement of the frequencies at which the molecule rotates in the gas phase. In molecules in crystalline form, the geometry of the molecule is revealed by irradiating the crystal with x-rays and analyzing the patterns formed as the x-rays diffract off of the crystal.)
Not all triatomic molecules are bent, however. As a common example, CO2 is a linear molecule. Larger polyatomics can have a variety of shapes, as illustrated in Figure 1. Ammonia, NH3, is a pyramid-shaped molecule, with the hydrogens in an equilateral triangle, the nitrogen above the plane of this triangle, and a H-N-H angle equal to 107°. The geometry of CH4 is that of a tetrahedron, with all H-C-H angles equal to 109.5°. (See also Figure 2(a).) Ethane, C2H6, has a geometry related to that of methane. The two carbons are bonded together, and each is bonded to three hydrogens. Each H-C-H angle is 109.5° and each H-C-C angle is 109.5°. By contrast, in ethene, C2H4, each H-C-H bond angle is 116.6° and each H-C-C bond angle is 121.7°. All six atoms of ethene lie in the same plane. Thus, ethene and ethane have very different geometries, despite the similarities in their molecular formulae.
Figure 1 Molecular Structures 
  
We begin our analysis of these geometries by noting that, in the molecules listed above which do not contain double or triple bonds (H2O, NH3, CH4and C2H6), the bond angles are very similar, each equal to or very close to the tetrahedral angle 109.5°. To account for the observed angle, we begin with our valence shell electron pair sharing model, and we note that, in the Lewis structures of these molecules, the central atom in each bond angle of these molecules contains four pairs of valence shell electrons. For methane and ethane, these four electron pairs are all shared with adjacent bonded atoms, whereas in
ammonia or water, one or two (respectively) of the electron pairs are not shared with any other atom. These unshared electron pairs are called lone pairs . Notice that, in the two molecules with no lone pairs, all bond angles are exactly equal to the tetrahedral angle, whereas the bond angles are only close in the molecules with lone pairs
One way to understand this result is based on the mutual repulsion of the negative charges on the valence shell electrons. Although the two electrons in each bonding pair must remain relatively close together in order to form the bond, different pairs of electrons should arrange themselves in such a way that the distances between the pairs are as large as possible. Focusing for the moment on methane, the four pairs of electrons must be equivalent to one another, since the four C-H bonds are equivalent, so we can assume that the electron pairs are all the same distance from the central carbon atom. How can we position four electron pairs at a fixed distance from the central atom but as far apart from one another as possible? A little reflection reveals that this question is equivalent to asking how to place four points on the surface of a sphere spread out from each other as far apart as possible. A bit of experimentation reveals that these four points must sit at the corners of a tetrahedron, an equilateral triangular pyramid, as may be seen in Figure 2(b). If the carbon atom is at the center of this tetrahedron and the four electron pairs at placed at the corners, then the hydrogen atoms also form a tetrahedron about the carbon. This is, as illustrated in Figure 2(a), the correct geometry of a methane molecule. The angle formed by any two corners of a tetrahedron and the central atom is 109.5°, exactly in agreement with the observed angle in methane. This model also works well in predicting the bond angles in ethane.
Figure 2 Tetrahedral Structure of Methane 
(a) The dotted lines illustrate that the hydrogens form a tetrahedron about the carbon atom. 
(b) The same tetrahedron is formed by placing four points on a sphere as far apart from one another as possible. 
We conclude that molecular geometry is determined by minimizing the mutual repulsion of the valence shell electron pairs. As such, this model of molecular geometry is often referred to as the valence shell electron pair repulsion (VSEPR) theory . For reasons that will become clear, extension of this model implies that a better name is the Electron Domain (ED) Theory .
This model also accounts, at least approximately, for the bond angles of H2O and NH3. These molecules are clearly not tetrahedral, like CH4, since neither contains the requisite five atoms to form the tetrahedron. However, each molecule does contain a central atom surrounded by four pairs of valence shell electrons. We expect from our Electron Domain model that those four pairs should be arrayed in a tetrahedron, without regard to whether they are bonding or lone-pair electrons. Then attaching the hydrogens (two for oxygen, three for nitrogen) produces a prediction of bond angles of 109.5°, very close indeed to the observed angles of 104.5° in H2O and 107° in NH3.
Note, however, that we do not describe the geometries of H2O and NH3 as "tetrahedral," since the atoms of the molecules do not form tetrahedrons, even if the valence shell electron pairs do. (It is worth noting that these angles are not exactly equal to 109.5°, as in methane. These deviations will be discussed later.)
We have developed the Electron Domain model to this point only for geometries of molecules with four pairs of valence shell electrons. However, there are a great variety of molecules in which atoms from Period 3 and beyond can have more than an octet of valence electrons. We consider two such molecules illustrated in Figure 3.
Figure 3 More Molecular Structures 
First, PCl5 is a stable gaseous compound in which the five chlorine atoms are each bonded to the phosphorous atom. Experiments reveal that the geometry of PCl5 is that of a trigonal bipyramid : three of the chlorine atoms form an equilateral triangle with the P atom in the center, and the other two chlorine atoms are on top of and below the P atom. Thus there must be 10 valence shell electrons around the phosphorous atom. Hence, phosphorous exhibits what is called an expanded valence in PCl5. Applying our Electron Domain model, we expect the five valence shell electron pairs to spread out optimally to minimize their repulsions. The required geometry can again be found by trying to place five points on the surface of a sphere with maximum distances amongst these points. A little experimentation reveals that this can be achieved by placing the five points to form a trigonal bipyramid. Hence, Electron Domain theory accounts for the geometry of PCl5.
Second, SF6 is a fairly unreactive gaseous compound in which all six fluorine atoms are bonded to the central sulfur atom. Again, it is clear that the octet rule is violated by the sulfur atom, which must therefore have an expanded valence. The observed geometry of SF6, as shown in Figure 3, is highly symmetric: all bond lengths are identical and all bond angles are 90°. The F atoms form an octahedron about the central S atom: four of the F atoms form a square with the S atom at the center, and the other two F atoms are above and below the S atom. To apply our Electron Domain model to understand this geometry, we must place six points, representing the six electron pairs about the central S atom, on the surface of a sphere with maximum distances between the points. The requisite geometry is found, in fact, to be that of an octahedron, in agreement with the observed geometry.
As an example of a molecule with an atom with less than an octet of valence shell electrons, we consider boron trichloride, BCl3. The geometry of BCl3 is also given in Figure 3: it is trigonal planar , with all four atoms lying in the same plane, and all Cl-B-Cl bond angles equal to 120°. The three Cl atoms form an equilateral triangle. The Boron atom has only three pairs of valence shell electrons in BCl3. In applying Electron Domain theory to understand this geometry, we must place three points on the surface of a sphere with maximum distance between points. We find that the three points form an equilateral triangle in a plane with the center of the sphere, so Electron Domain is again in accord with the observed geometry.
We conclude from these predictions and observations that the Electron Domain model is a reasonably accurate way to understand molecular geometries, even in molecules which violate the octet rule.
Observation 2: Molecules with Double or Triple Bonds
In each of the molecules considered up to this point, the electron pairs are either in single bonds or in lone pairs. In current form, the Electron Domain model does not account for the observed geometry of C2H4, in which each H-C-H bond angle is 116.6° and each H-C-C bond angle is 121.7° and all six atoms lie in the same plane. Each carbon atom in this molecule is surrounded by four pairs of electrons, all of which are involved in bonding, i.e. there are no lone pairs. However, the arrangement of these electron pairs, and thus the bonded atoms, about each carbon is not even approximately tetrahedral. Rather, the H-C-H and H-C-C bond angles are much closer to 120°, the angle which would be expected if three electron pairs were separated in the optimal arrangement, as just discussed for BCl3.
This observed geometry can be understood by re-examining the Lewis structure. Recall that, although there are four electron pairs about each carbon atom, two of these pairs form a double bond between the carbon atoms. It is tempting to assume that these four electron pairs are forced apart to form a tetrahedron as in previous molecules. However, if this were this case, the two pairs involved in the double bond would be separated by an angle of 109.5° which would make it impossible for both pairs to be localized between the carbon atoms. To preserve the double bond, we must assume that the two electron pairs in the double bond remain in the same vicinity. Given this assumption, separating the three independent groups of electron pairs about a carbon atom produces an expectation that all three pairs should lie in the same plane as the carbon atom, separated by 120° angles. This agrees very closely with the observed bond angles. We conclude that the our model can be extended to understanding the geometries of molecules with double (or triple) bonds by treating the multiple bond as two electron pairs confined to a single domain. It is for this reason that we refer to the model as Electron Domain theory.
Applied in this form, Electron Domain theory can help us understand the linear geometry of CO2. Again, there are four electron pairs in the valence shell of the carbon atom, but these are grouped into only two domains of two electron pairs each, corresponding to the two C=O double bonds. Minimizing the repulsion between these two domains forces the oxygen atoms to directly opposite sides of the carbon, producing a linear molecule. Similar reasoning using Electron Domain theory as applied to triple bonds correctly predicts that acetylene, HCCH, is a linear molecule. If the electron pairs in the triple bond are treated as a single domain, then each carbon atom has only two domains each. Forcing these domains to opposite sides from one another accurately predicts 180° H-C-C bond angles.
Observation 3: Distortions from Expected Geometries
It is interesting to note that some molecular geometries (CH4, CO2, HCCH) are exactly predicted by the Electron Domain model, whereas in other molecules, the model predictions are only approximately correct. For examples, the observed angles in ammonia and water each differ slightly from the tetrahedral angle. Here again, there are four pairs of valence shell electrons about the central atoms. As such, it is reasonable to conclude that the bond angles are determined by the mutual repulsion of these electron pairs, and are thus expected to be 109.5°, which is close but not exact.
One clue as to a possible reason for the discrepancy is that the bond angles in ammonia and water are both less than 109.5°. Another is that both ammonia and water molecules have lone pair electrons, whereas there are no lone pairs in a methane molecule, for which the Electron Domain prediction is exact. Moreover, the bond angle in water, with two lone pairs, is less than the bond angles in ammonia, with a single lone pair. We can straightforwardly conclude from these observations that the lone pairs of electrons must produce a greater repulsive effect than do the bonded pairs. Thus, in ammonia, the three bonded pairs of electrons are forced together slightly compared to those in methane, due to the greater repulsive effect of the lone pair. Likewise, in water, the two bonded pairs of electrons are even further forced together by the two lone pairs of electrons.
This model accounts for the comparative bond angles observed experimentally in these molecules. The valence shell electron pairs repel one another, establishing the geometry in which the energy of their interaction is minimized. Lone pair electrons apparently generate a greater repulsion, thus slightly reducing the angles between the bonded pairs of electrons. Although this model accounts for the observed geometries, why should lone pair electrons generate a greater repulsive effect? We must guess at a qualitative answer to this question, since we have no description at this point for where the valence shell electron pairs actually are or what it means to share an electron pair. We can assume, however, that a pair of electrons shared by two atoms must be located somewhere between the two nuclei, otherwise our concept of "sharing" is quite meaningless. Therefore, the powerful tendency of the two electrons in the pair to repel one another must be significantly offset by the localization of these electrons between the two nuclei which share them. By contrast, a lone pair of electrons need not be so localized, since there is no second nucleus to draw them into the same vicinity. Thus more free to move about the central atom, these lone pair electrons must have a more significant repulsive effect on the other pairs of electrons.
These ideas can be extended by more closely examining the geometry of ethene, C2H4 . Recall that each H-C-H bond angle is 116.6° and each H-C-C bond angle is 121.7°, whereas the Electron Domain theory prediction is for bond angles exactly equal to 120°. We can understand why the H-C-H bond angle is slightly less than 120° by assuming that the two pairs of electrons in the C=C double bond produce a greater repulsive effect than do either of the single pairs of electrons in the C-H single bonds. The result of this greater repulsion is a slight "pinching" of the H-C-H bond angle to less than 120°.
The concept that lone pair electrons produce a greater repulsive effect than do bonded pairs can be used to understand other interesting molecular geometries. Sulfur tetrafluoride, SF4, is a particularly interesting example, shown in Figure 4.
Figure 4 Molecular Structure of SF4 
Note that two of the fluorines form close to a straight line with the central sulfur atom, but the other two are approximately perpendicular to the first two and at an angle of 101.5° to each other. Viewed sideways, this structure looks something like a seesaw.
To account for this structure, we first prepare a Lewis structure. We find that each fluorine atom is singly bonded to the sulfur atom, and that there is a lone pair of electrons on the sulfur. Thus, with five electron pairs around the central atom, we expect the electrons to arrange themselves in a trigonal bipyramid, similar to the arrangement in PCl5 in Figure 3. In this case, however, the fluorine atoms and the lone pair could be arranged in two different ways with two different resultant molecular structures. The lone pair can either go on the axis of the trigonal bipyramid (i.e. “above” the sulfur) or on the equator of the bipyramid (i.e. “beside” the sulfur).
The actual molecular structure in Figure 4 shows clearly that the lone pair goes on the equatorial position. This can be understood if we assume that the lone pair produces a greater repulsive effect than do the bonded pairs. With this assumption, we can deduce that the lone pair should be placed in the trigonal bipyramidal arrangement as far as possible from the bonded pairs. The equatorial position does a better job of this, since only two bonding pairs of electrons are at approximately 90° angles from the lone pair in this position. By contrast, a lone pair in the axial position is approximately 90° away from three bonding pairs. Therefore, our Electron Domain model assumptions are consistent with the observed geometry of SF4. Note that these assumptions also correctly predict the observed distortions away from the 180° and 120° angles which would be predicted by a trigonal bipyramidal arrangement of the five electron pairs.  —Preceding unsigned comment added by 83.54.18.125 (talk) 17:41, 1 February 2010 (UTC)
thread_title: Electron Domain

message: This is redirect from "Molecular structure". Many standard texts, e.g. Physical Chemistry, 9th edition, by Peter Atkins and Julio dePaula, Freeman, New York, 2009, interpret "molecular structure" as "molecular electronic structure" which includes strengths and electronic structure of bonds as well as geometry. Several methods for determining structure of paramount importance are ignored. Selection of topics is unsatisfactory. I am trying to work on another article that refers to this and cannot spend more time on this now. Michael P. Barnett (talk) 17:39, 30 January 2011 (UTC)
thread_title: Reason for warning

message: Could somebody please explain how come the ΔE is given as eg 500 cm-1? I have heard about many energy units from Joule to erg to electron volts, but cm-1? For all I know there is a convention about that in some branch of physics or chemistry, but the article should be readable for a wider audience.
Calculating backwards, I find that the unit cm-1 as used in the article should be equivalent to about 1.99×10-23J, which is considerably less than an electron Volt. What is it, anyone?Cacadril (talk) 13:54, 23 October 2011 (UTC)
replies: text: I don't know how widespread is this practice, but it should rely on E=hcn where n=1/lambda is the wavenumber, and hc=6.6e-34 * 3e8 gives the factor you mentioned. Materialscientist (talk) 00:01, 24 October 2011 (UTC)
thread_title:  Energy measured in cm<sup>-1</sup>? 

message: In today's edits, the results of VSEPR theory seem to have been replaced by the results of orbital hybridization theory. The edit summary of Officer781 (08:53 2 Dec 2012 UT) reads change angle prediction to use orbital hybridisation theory instead [of VSEPR] as that one is a less "crude" way of predicting. In fact most current general chemistry books mention both theories as both have their advantages. VSEPR is easier to understand at an introductory level since it can be presented without mentioning quantum mechanics, but is indeed necessarily a crude model. Orbital hybridisation can be presented as a crude quantum-mechanical model, but can also be systematically refined by reference to MO calculations. In an encyclopedia article, I think we should follow the textbooks and discuss both, starting with a phrase such as Two simple theories are widely used to describe the arrangement of atoms bonded to a central atom. Then I think VSEPR should be discussed first as it is simpler (no quantum mechanics), followed by orbital hybridisation.
One problem is that complete tables of all geometries using both theories would make the article too long. My solution would be to include here only short illustrative tables of a few geometries (perhaps 5) for each method, with links to the articles on VSEPR and Orbital hybridization for the reader who wants more complete tables. Dirac66 (talk) 23:44, 2 December 2012 (UTC)
replies: text: Hmm. I've reverted my edits with a short mention of orbital hybridisation for transition metals and a more comprehensive description. Personally, I would reserve VSEPR for usage as a "crude" but simple model, and reserve orbital hybridisation to reflect MO and quantum chemical calculations (ie more advanced).--Officer781 (talk) 02:35, 7 December 2012 (UTC)
thread_title:  VSEPR or Orbital hybridization? I suggest both plus short tables. 

message: The Image column in the VSEPR table is misleading because several entries (SO2, NH3, H2O, SF4, ClF3, XeF2) now show the electron arrangement on the central atom including both bonds and lone pairs. As shown at VSEPR theory#AXE method (second table) which has two columns, the electron arrangement does include lone pairs, but the geometry does not. Since this article is on Molecular geometry and the table has only one Image column, it would be preferable to show the arrangement of the bonds only, corresponding to the results of experimental structural determinations by  spectroscopic, crystallographic or other methods. The columns Lone pairs and Electron domains can still make clear that the bond arrangement depends on the unobserved lone pairs, with a link to the VSEPR article for the details. Dirac66 (talk) 14:26, 9 July 2013 (UTC)
replies: text: Fixed today. Dirac66 (talk) 20:45, 18 November 2013 (UTC)
thread_title:  VSEPR table should not show lone pairs 

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thread_title:  External links modified