Chapter VSEPR - Molecular Geometry - Chemistry LibreTexts
Chemical bonding - Molecular shapes and VSEPR theory: There is a sharp distinction The methane molecule, CH4, can be used to illustrate the procedure for. According to the VSEPR model, the H - C - H bond angle in methane should be The difference between the predicted and observed bond angles can be. Framework model - 1CRN Methane has 4 regions of electron density around the central carbon atom (4 bonds, no lone The resulting shape is a regular tetrahedron with H-C-H angles of °. Click the structures to load the molecules.
Notice that, in the Lewis structure of these molecules, the central atom s bonds with only two other atoms and has no unshared electrons. Only two electron clouds emerge from that central atom. The VSEPR theory says, then, that the geometry around an atom that has only two bonds and no unshared electrons is a straight line. In these molecules, each central atom has three electron clouds emanating from it.
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In sulfur dioxide, the sulfur atom is bonded to two oxygen atoms and has one unshared pair of electrons. In formaldehyde and ethylene, each carbon atom has two single bonds to hydrogen, a double bond to another atom, and no unshared pair. The sulfur atom in sulfur dioxide and the carbon atom in ethylene and formaldehyde is surrounded by three clouds of high electron density.
The structure will be trigonal planar. The central atom will be in the center of the triangle, and the ends of the electron clouds at the corners of the triangle.
If you experiment with a marshmallow as the central atom and three toothpicks as electron clouds, you can prove to yourself that the toothpicks are farthest apart when using a trigonal planar structure.
Although the electron clouds of these molecules give a trigonal planar shape around each carbon atom, one describes the geometry of a molecule only on the basis of the relationships between its atoms. A formaldehyde molecule is trigonal planar because it has an atom at the end of each electron cloud.
The ethylene molecule has trigonal planar geometry around each of its carbon atoms. The whole molecule is planar, and its shape resembles two triangles joined point to point. In sulfur dioxide, there are three electron clouds around the sulfur.
Only two of these connect two atoms. The molecule is bent. A central atom surrounded by three clouds of high electron density will have trigonal planar geometry if it is bonded to three atoms. Its geometry will be called bent if it is bonded to two atoms and also has an unshared pair of electrons.
Structures with Four Regions of High Electron Density around the Central Atom The following Lewis structures show three molecules whose central atom is surrounded by four clouds of high electron density: These molecules are alike in that each central atom is surrounded by four pairs of electrons, but they differ in the number of unshared electron pairs on the central atom.
Remember that, although we have drawn them in a plane, the molecules are three-dimensional and atoms may be in front of or behind the plane of the paper. The central atom, carbon, contributes four valence electrons, and each hydrogen atom has one valence electron, so the full Lewis electron structure is 2.
There are four electron groups around the central atom. As shown in Figure 6. All electron groups are bonding pairs, so the structure is designated as AX4. With four bonding pairs, the molecular geometry of methane is tetrahedral Figure 6.
Chapter 6.3: VSEPR - Molecular Geometry
Methane has a three-dimensional, tetrahedral structure. The hybridization of the C atom orbitals is sp3. In ammonia, the central atom, nitrogen, has five valence electrons and each hydrogen donates one valence electron, producing the Lewis electron structure 2.
There are four electron groups around nitrogen, three bonding pairs and one lone pair. Repulsions are minimized by directing each hydrogen atom and the lone pair to the corners of a tetrahedron. With three bonding pairs and one lone pair, the structure is designated as AX3E.
This designation has a total of four electron pairs, three X and one E. We expect the LP—BP interactions to cause the bonding pair angles to deviate significantly from the angles of a perfect tetrahedron. There are three nuclei and one lone pair, so the molecular geometry is trigonal pyramidal. In essence, this is a tetrahedron with a vertex missing Figure 6. However, the H—N—H bond angles are less than the ideal angle of The hybridization of the N atom orbitals is sp3. Oxygen has six valence electrons and each hydrogen has one valence electron, producing the Lewis electron structure 2.
There are four groups around the central oxygen atom, two bonding pairs and two lone pairs. Repulsions are minimized by directing the bonding pairs and the lone pairs to the corners of a tetrahedron Figure 6. With two bonding pairs and two lone pairs, the structure is designated as AX2E2 with a total of four electron pairs. With two hydrogen atoms and two lone pairs of electrons, the structure has significant lone pair interactions.
There are two nuclei about the central atom, so the molecular shape is bent, or V shaped, with an H—O—H angle that is even less than the H—N—H angles in NH3, as we would expect because of the presence of two lone pairs of electrons on the central atom rather than one.
This molecular shape is essentially a tetrahedron with two missing vertices. The hybridization of the O atom orbitals is sp3. Five Electron Groups In previous examples it did not matter where we placed the electron groups because all positions were equivalent.
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In some cases, however, the positions are not equivalent. We encounter this situation for the first time with five electron groups. Phosphorus has five valence electrons and each chlorine has seven valence electrons, so the Lewis electron structure of PCl5 involves an expanded octet.
In the following section we will describe how this expanded octet is formed by the combination of a d orbital with an s and three p orbitals.
There are five bonding groups around phosphorus, the central atom. The structure that minimizes repulsions is a trigonal bipyramid, which consists of two trigonal pyramids that share a base Figure 6. All electron groups are bonding pairs, so the structure is designated as AX5.
There are no lone pair interactions. The molecular geometry of PCl5 is trigonal bipyramidal, as shown in Figure 6. The molecule has three atoms in a plane in equatorial positions and two atoms above and below the plane in axial positions. The axial and equatorial positions are not chemically equivalent, as we will see in our next example.
There are five groups around sulfur, four bonding pairs and one lone pair. With five electron groups, the lowest energy arrangement is a trigonal bipyramid, as shown in Figure 6. However, because the axial and equatorial positions are not chemically equivalent, where do we place the lone pair?
We also expect a deviation from ideal geometry because a lone pair of electrons occupies more space than a bonding pair. With four nuclei and one lone pair of electrons, the molecular structure is based on a trigonal bipyramid with a missing equatorial vertex; it is described as a seesaw.
The bromine atom has seven valence electrons, and each fluorine has seven valence electrons, so the Lewis electron structure is Once again, we have a compound that is an exception to the octet rule. There are five groups around the central atom, three bonding pairs and two lone pairs and the hybridization of orbitals on the Br atom will be sp3d.
We again direct the groups toward the vertices of a trigonal bipyramid. With three bonding pairs and two lone pairs, the structural designation is AX3E2 with a total of five electron pairs. Bonding electrons, however, must be simultaneously close to two nuclei, and only a small region of space between the nuclei satisfies this restriction. Because they occupy more space, the force of repulsion between pairs of nonbonding electrons is relatively large.
The force of repulsion between a pair of nonbonding electrons and a pair of bonding electrons is somewhat smaller, and the repulsion between pairs of bonding electrons is even smaller. The figure below can help us understand why nonbonding electrons are placed in equatorial positions in a trigonal bipyramid.
If the nonbonding electrons in SF4 are placed in an axial position, they will be relatively close 90o to three pairs of bonding electrons. But if the nonbonding electrons are placed in an equatorial position, they will be 90o away from only two pairs of bonding electrons.
As a result, the repulsion between nonbonding and bonding electrons is minimized if the nonbonding electrons are placed in an equatorial position in SF4. When the nonbonding pair of electrons on the sulfur atom in SF4 is placed in an equatorial position, the molecule can be best described as having a see-saw or teeter-totter shape. Repulsion between valence electrons on the chlorine atom in ClF3 can be minimized by placing both pairs of nonbonding electrons in equatorial positions in a trigonal bipyramid.
When this is done, we get a geometry that can be described as T-shaped. The Lewis structure of the triiodide I3- ion suggests a trigonal bipyramidal distribution of valence electrons on the central atom.