Why is molecular shape important

The electron-pair geometries will be the same as the molecular structures when there are no lone electron pairs around the central atom, but they will be different when there are lone pairs present on the central atom.

VSEPR structures like the one in Figure 4 are often drawn using the wedge and dash notation, in which solid lines represent bonds in the plane of the page, solid wedges represent bonds coming up out of the plane, and dashed lines represent bonds going down into the plane.

For example, the methane molecule, CH 4 , which is the major component of natural gas, has four bonding pairs of electrons around the central carbon atom; the electron-pair geometry is tetrahedral, as is the molecular structure Figure 4. On the other hand, the ammonia molecule, NH 3 , also has four electron pairs associated with the nitrogen atom, and thus has a tetrahedral electron-pair geometry.

One of these regions, however, is a lone pair, which is not included in the molecular structure, and this lone pair influences the shape of the molecule Figure 5. Figure 5. As seen in Figure 5, small distortions from the ideal angles in Figure 6 can result from differences in repulsion between various regions of electron density.

VSEPR theory predicts these distortions by establishing an order of repulsions and an order of the amount of space occupied by different kinds of electron pairs. The order of electron-pair repulsions from greatest to least repulsion is:. This order of repulsions determines the amount of space occupied by different regions of electrons. A lone pair of electrons occupies a larger region of space than the electrons in a triple bond; in turn, electrons in a triple bond occupy more space than those in a double bond, and so on.

The order of sizes from largest to smallest is:. Consider formaldehyde, H 2 CO, which is used as a preservative for biological and anatomical specimens Figure 1. This molecule has regions of high electron density that consist of two single bonds and one double bond. In the ammonia molecule, the three hydrogen atoms attached to the central nitrogen are not arranged in a flat, trigonal planar molecular structure, but rather in a three-dimensional trigonal pyramid Figure 5 with the nitrogen atom at the apex and the three hydrogen atoms forming the base.

The ideal bond angles in a trigonal pyramid are based on the tetrahedral electron pair geometry. Again, there are slight deviations from the ideal because lone pairs occupy larger regions of space than do bonding electrons. Figure 6 illustrates the ideal molecular structures, which are predicted based on the electron-pair geometries for various combinations of lone pairs and bonding pairs. Figure 6. The molecular structures are identical to the electron-pair geometries when there are no lone pairs present first column.

For a particular number of electron pairs row , the molecular structures for one or more lone pairs are determined based on modifications of the corresponding electron-pair geometry. According to VSEPR theory, the terminal atom locations Xs in Figure 6 are equivalent within the linear, trigonal planar, and tetrahedral electron-pair geometries the first three rows of the table.

It does not matter which X is replaced with a lone pair, because the molecules can be rotated to convert positions. The following procedure uses VSEPR theory to determine the electron pair geometries and the molecular structures:. The following examples illustrate the use of VSEPR theory to predict the molecular structure of molecules or ions that have no lone pairs of electrons.

In this case, the molecular structure is identical to the electron pair geometry. Predict the electron-pair geometry and molecular structure for phosgene, COCl 2 , a chemical warfare agent used during World War I. This shows us three regions of high electron density around the carbon atom—each single bond counts as one region, and the double bond counts as another region, and there are no lone pairs on the carbon atom. The electron-pair geometry and molecular structure are identical because all the groups are bonding, and COCl 2 molecules are trigonal planar.

Our ability to draw structural formulas for molecules is remarkable. To see how this is done Click Here. Formula Analysis.

Although structural formulas are essential to the unique description of organic compounds, it is interesting and instructive to evaluate the information that may be obtained from a molecular formula alone. Three useful rules may be listed: The number of hydrogen atoms that can be bonded to a given number of carbon atoms is limited by the valence of carbon.

The origin of this formula is evident by considering a hydrocarbon made up of a chain of carbon atoms. Here the middle carbons will each have two hydrogens and the two end carbons have three hydrogens each. Thus, when even-valenced atoms such as carbon and oxygen are bonded together in any number and in any manner, the number of remaining unoccupied bonding sites must be even. If these sites are occupied by univalent atoms such as H, F, Cl, etc.

If the four carbon atoms form a ring, two hydrogens must be lost. Similarly, the introduction of a double bond entails the loss of two hydrogens, and a triple bond the loss of four hydrogens.

By rule 2 m must be an even number, so if m The presence of one or more nitrogen atoms or halogen substituents requires a modified analysis. The above formula may be extended to such compounds by a few simple principles: The presence of oxygen does not alter the relationship.

All halogens present in the molecular formula must be replaced by hydrogen. Each nitrogen in the formula must be replaced by a CH moiety. However, the structures of some compounds and ions cannot be represented by a single formula.

For clarity the two ambiguous bonds to oxygen are given different colors in these formulas. If only one formula for sulfur dioxide was correct and accurate, then the double bond to oxygen would be shorter and stronger than the single bond. This averaging of electron distribution over two or more hypothetical contributing structures canonical forms to produce a hybrid electronic structure is called resonance.

Likewise, the structure of nitric acid is best described as a resonance hybrid of two structures, the double headed arrow being the unique symbol for resonance. The above examples represent one extreme in the application of resonance. Here, two structurally and energetically equivalent electronic structures for a stable compound can be written, but no single structure provides an accurate or even an adequate representation of the true molecule.

In cases such as these, the electron delocalization described by resonance enhances the stability of the molecules, and compounds or ions composed of such molecules often show exceptional stability. The electronic structures of most covalent compounds do not suffer the inadequacy noted above. Nevertheless, the principles of resonance are very useful in rationalizing the chemical behavior of many such compounds. For example, the carbonyl group of formaldehyde the carbon-oxygen double bond reacts readily to give addition products.

The course of these reactions can be explained by a small contribution of a dipolar resonance contributor, as shown in equation 3. Here, the first contributor on the left is clearly the best representation of this molecular unit, since there is no charge separation and both the carbon and oxygen atoms have achieved valence shell neon-like configurations by covalent electron sharing. If the double bond is broken heterolytically, formal charge pairs result, as shown in the other two structures.

The preferred charge distribution will have the positive charge on the less electronegative atom carbon and the negative charge on the more electronegative atom oxygen.

Therefore the middle formula represents a more reasonable and stable structure than the one on the right. The application of resonance to this case requires a weighted averaging of these canonical structures.

The double bonded structure is regarded as the major contributor, the middle structure a minor contributor and the right hand structure a non-contributor. Since the middle, charge-separated contributor has an electron deficient carbon atom, this explains the tendency of electron donors nucleophiles to bond at this site.

The basic principles of the resonance method may now be summarized. These are the canonical forms to be considered, and all must have the same number of paired and unpaired electrons. The following factors are important in evaluating the contribution each of these canonical structures makes to the actual molecule.

The stability of a resonance hybrid is always greater than the stability of any canonical contributor. Consequently, if one canonical form has a much greater stability than all others, the hybrid will closely resemble it electronically and energetically. This is the case for the carbonyl group eq. On the other hand, if two or more canonical forms have identical low energy structures, the resonance hybrid will have exceptional stabilization and unique properties.

This is the case for sulfur dioxide eq. To illustrate these principles we shall consider carbon monoxide eq. In each case the most stable canonical form is on the left.

For carbon monoxide, the additional bonding is more important than charge separation. Furthermore, the double bonded structure has an electron deficient carbon atom valence shell sextet. A similar destabilizing factor is present in the two azide canonical forms on the top row of the bracket three bonds vs.

The bottom row pair of structures have four bonds, but are destabilized by the high charge density on a single nitrogen atom. All the examples on this page demonstrate an important restriction that must be remembered when using resonance. No atoms change their positions within the common structural framework.

Only electrons are moved. A more detailed model of covalent bonding requires a consideration of valence shell atomic orbitals. The spatial distribution of electrons occupying each of these orbitals is shown in the diagram below. Very nice displays of orbitals may be found at the following sites: J. Gutow, Univ. Wisconsin Oshkosh R. You can use Lewis structures to predict molecular geometry, but it's best to verify these predictions experimentally. Several analytical methods can be used to image molecules and learn about their vibrational and rotational absorbance.

Examples include x-ray crystallography, neutron diffraction, infrared IR spectroscopy, Raman spectroscopy, electron diffraction, and microwave spectroscopy. The best determination of a structure is made at low temperature because increasing the temperature gives the molecules more energy, which can lead to conformation changes. The molecular geometry of a substance may be different depending on whether the sample is a solid, liquid, gas, or part of a solution.

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