The Valence Shell Electron Pair Resulsion theory, or VSEPR theory, is used to determine the three-dimensional shape of bonds and lone pairs around an interior atom. It's based on the idea that electron groups repel other electron groups. An electron group could be a lone pair, single bond, double bond, triple bond, or an unpaired electron; each of these counts as one electron group. The three-dimensional shape that results positions these electron groups as far away from each other as geometrically possible. The simplest shape we can have results from just two electron groups. The farthest we can put two electron groups away from each other is 180° apart, or at the opposite sides of a line through the central atom. This geometry is called linear. Both BeCl2 and CO2 have a linear molecular geometry. Notice that in carbon dioxide each double bond accounts as just one electron group, so that there are two electron groups on our central carbon atom. If we add a third electron group on our central atom, as in boron trifluoride, the farthest we can put the three electron groups away from each other is 120° apart, or at the corners of a triangle in a plane. This geometry is called trigonal planar. In the formaldehyde molecule, we also have three electron groups, but two of them are single bonds and one is a double bond. The double bond takes up more room around the atom than the single bonds do because the double bond has more electrons. So our bond angles vary slightly from 120° as the two single bonds are pushed together. If we add a fourth electron group on our central atom, the farthest we can put the four electron groups away from each other is 109.5° apart, or at the corners of the tetrahedron; this geometry is called tetrahedral. Methane, or CH4, has a tetrahedral molecular geometry. If we add a fifth electron group on our central atom, the farthest we can put the five electron groups away from each other is to the corners of a trigonal bipyramid; this geometry is called trigonal bipyramidal. We classify the five electron groups into two categories: axial and equatorial. There are two axial groups which are 180° apart from one another, much like the linear geometry. There are three equatorial groups, which are 120° apart from one another, going to the corners of a triangle, much like trigonal planar. Any given equatorial group is 90° away from an axial group. Phosphorus pentachloride has a trigonal bipyramidal molecular geometry. If we add a sixth electron group to our central atom, the farthest we can put the six electron groups away from each other is to the corners of an octahedron. This geometry is called octahedral. The bond angle between any two groups in octahedral geometry is 90°. Sulphur hexafluoride has an octahedral molecular geometry. The presence of a lone pair of electrons changes the name of the molecular geometry. For example, if we have three bonding groups and one lone pair of electrons on the central atom, as we do in ammonia, our molecular geometry becomes trigonal pyramidal. Just like a multiple bond, a lone pair of electrons can take up more room around the central atom and change the other bond angles slightly. In ammonia, instead of bond angles of 109.5°, the lone pair of electrons repels the other electron groups so that they are 107° apart. If we have two lone pairs of electrons and two bonding groups, the molecular geometry is called bent. Water (or H2O) has a bent molecular geometry. Here again, our lone pairs of electrons repel the single bonds more so that the bond angle is 104.5° instead of 109.5°. So if we have four electron groups, the molecular geometry is different depending on the number of electron groups that are lone pairs. If we have four bonding groups and one lone pair on the central atom, our molecular geometry is called seesaw. This is the case for sulfur tetrafluoride. Notice that the lone pair is in the equatorial position instead of the axial position. This gives the lone pair of electrons the most room possible because its bond angles with the other equatorial groups are about 120°. If we have three bonding groups and two lone pairs of electrons on the central atom, our molecular geometry is called T-shaped. This is the case for bromine trifluoride. Notice again that the two lone pairs are put in the equatorial positions to give them the most room. If we have two bonding groups and three lone pairs of electrons of the central atom, our molecular geometry becomes linear. Xenon difluoride has a linear molecular geometry. All three loan pairs of electrons are again in the equatorial positions. If we have five bonding groups and one lone pair of electrons on the central atom, our molecular geometry is called square pyramidal. This is the case for bromine pentafluoride. And finally, if we have four bonding groups and two lone pairs of electrons on the central atom, our molecular geometry is called square planar. Xenon tetrafluoride has a square planar molecular geometry. Let's practice assigning molecular geometries, starting with xenon trioxide. First, we'll draw the Lewis structure for xenon trioxide. We have three triple bonds and one lone pair on xenon in our best Lewis structure, which minimizes formal charges. From our Lewis structure, we see that xenon has three bonding groups and one lone pair of electrons. When determining the molecular geometry around a particular atom, we only count the lone pairs on the atom in question. When we have three bonding groups and one lone pair, our molecular geometry is called trigonal pyramidal. Next, let's determine the molecular geometry for nitrate. When we draw the nitrate ion, we have three equivalent resonance structures, so we know the nitrate ion exists as an average of these three structures. All three of the structures have three bonding groups and zero lone pairs, so the molecular geometry of nitrate is trigonal planar.