Molecular geometry, the bond lengths and angles, are determined experimentally. Lewis structures can give us an approximate measure of molecular bonding. There is a simple procedure that allows us to predict overall geometry is the VSEPR, Valence Shell Electron Pair Repulsion. The concept is that valence shell electron pairs are involved in bonding, and that these electron pairs will keep as far away from each other, due to electron-electron repulsion.
When two groups try to get as far away from each other as possible, a linear shape is formed.
The M represents the central atom. The pairs of dots (:) represent two pairs of valence electrons forming bonds to the central atom. The angle between the lone pairs and the central atom is 180 o.
When three electrons pairs get as far apart from each other, a trigonal planar structure is formed, as shown below.
The bond angle in this structure is 120o.
Four electron pairs form a tetrahedron when they are separated as far as possible from each other. The bond angle in a tetrahedron is 109.5o.
Five electron pairs form a trigonal bipyramid when they are separated in space.
This shape consists of a triangle with two electron pairs directly above and below it. The bond angles of the middle, triangular portion are 120o. The groups above and below the central triangle are at a 90o angle to the triangle.
Six pairs of electrons separated in space will form an octahedron, with 90o between all of the valence electrons.
This is just geometry. Now, we will apply it to Chemistry.
Linear Molecular Geometries
What is the Lewis structure of BeCl2 ? Be has 2 valence electrons, and each Cl contributes 7 valence electrons, for a total of 16 electrons. The structure is shown below.
This is a molecule with fewer than 8 valence electrons. The central atom, Be, has two pairs of bonding electrons. It will have a linear geometry, because there are two groups (the two Cl atoms) separated in space.
What is the Lewis structure of carbon dioxide, CO2? There are 16 valence electrons to work with.
The central carbon only has a share in 4 valence electrons, so we can move a lone pair from each oxygen, to form two double bonds between C and O.
Here again, there are two groups bonded to the central atom, so it is again a linear molecule. Single and double bonds are treated the same in VSEPR.
Trigonal Planar Molecular Geometries
We have already looked at the structure of boron trifluoride, BF3. It is another example of a molecule with fewer than 8 valence electrons.
There are three groups bonded to the central atom, B. This molecule will be triangular planar (flat) with bond angles of 120o.
What is the Lewis structure of sulfur dioxide? Sulfur and oxygen are both Group 6. So, there are 18 valence electrons.
The sulfur doesn't have a share in 8 electrons, so one lone pair from O will be used to form a double bond. It gives a resonance structure.
The central atom, S, has three groups bonded to it, two oxygens and a lone pair. The molecular geometry of SO2 will be trigonal planar. It would be drawn as shown below:
The lone pair of electrons actually occupy a relatively large volume, since the are only held by one atom. In fact, they will compress the bond angle between the oxygens and sulfur to less than 120o.
The molecular shape of SO2 will not be trigonal planar (molecular geometry). In determining the molecular shape, we only consider the position of the atoms, not the lone pairs. So, the molecular shape of SO2 is called bent and would be represented as:
The bond angle between the O and S will still be less than 120o because of the presence of the one pair on S.
Tetrahedral Molecular Geometries
The Lewis structure of methane, CH4, follows the rules for molecules with 8 valence electrons. It is shown below.
There are four groups (H) bonded to the central atom (C). They will be arranged in a tetrahedron around the central atom, with bond angles of 109.5o.
The Lewis structure of ammonia, NH3, is shown below. There are again, 8 valence electrons to work with.
There are four groups around the central N atom, 3 H and a lone pair. So, the molecular geometry will be tetrahedral, with bond angles close to 109.5o. However, the lone pair electrons will again distort this angle, because they occupy a larger volume than the bonding pairs of electrons.
The geometric shape is tetrahedral, but the molecular shape is said to be a trigonal pyramid. In defining molecular shapes, we only look at the position of the atoms, not the lone pairs. So, ammonia would look like:
The Lewis structure of water, H2O, is shown below.
Again, the central O has four groups around it, two H and two lone pairs. So, the molecular geometry is tetrahedral, with 109.5o bond angles. However, the lone pairs distort this ideal angle to about 104o. The molecular shape is said to be bent and is shown below.
Trigonal bipyramidal Molecular Geometries
In these molecules, the central atom is surrounded by 5 groups, a combination of atoms and lone pairs. A compound with 5 surrounding atoms, and no lone pairs, is phosphorous pentachloride, PCl5. It has the Lewis structure shown below:
Shown below is the molecular geometry of PCl5. There are three Cl atoms in a triangle around the middle of the molecule. They are separated by 120o bond angles. These are called the equatorial positions. There are also two Cl atoms directly above and below the central P atom. They are at 180o to each other and at 90o to the equatorial Cl. These are called the axial positions. This geometry, with atoms at all 5 positions around the central atom is trigonal bipyramidal.
Another compound with 5 groups around the central atom is sulfur tetrafluoride, SF4. The Lewis dot structure is shown below. This is another compound that is an exception to the octet rule, since it has more than 8 valence electrons. This is possible because sulfur is in period 3, and has d orbitals to accommodate additional electrons. This molecule has 6 + (4x7) = 34 e-. The extra pair of electrons, that is not used in bonding the F to the S and is not needed to complete the octets on F is added as a lone pair on the central S atom.
In this molecule, with trigonal bipyramidal geometry, one position is occupied by a lone pair. Because lone pairs occupy more space than bonded electron pairs, it will go to the position with the most room in this molecule. The equatorial positions are more "roomy" so lone pairs will always go to the equatorial position in trigonal pyramidal compounds. Shown below is the molecular structure of SF4:
Notice that the bonds angles are slightly distorted due to the lone pair in the equatorial position. The axial fluorides are not quite at 180o to each other, and the remaining equatorial fluorides are at slightly less than 120o to each other. Viewed on its side, some think that this looks like a seesaw, so that is the molecular shape.
Chlorine trifluoride is another compound with 5 groups around the central atom. There are 28 valence electrons in the Lewis structure, shown below:
Chlorine has d orbitals and can accommodate 5 pairs of electrons. The two lone pairs will both occupy equatorial positions, to minimize electron-electron repulsion. The molecular shape is shown below. There are F in both axial positions (180o from each other) and a F in one of the equatorial positions (at 90o to the other two F). The shape is described as a T.
Finally, we can look at the iodine dichloride ion, ICl2-. There are 7 + (2x7) + 1 = 22 valence e-. The Lewis structure is shown below.
The three lone pairs are in the equatorial positions, leaving the Cl and I in a linear molecular shape.
Octahedral Molecular Geometries
Finally, we compounds with 6 groups around the central atoms. A compound with 6 atoms surrounding the central atoms will have an octahedral geometric shape and an octahedral molecular shape. An example of this is sulfur hexafluoride, SF6. The Lewis structure is shown below:
All of the bond lengths and angles (90o) are identical in this structure. The molecular shape, octahedral, is shown below:
When one atom is replaced by a lone pair in octahedral geometries, it doesn't matter which position it goes to, since they are all equivalent. An example of this is bromine pentafluoride, BrF5.
Its molecular shape is shown below. It is called square pyramidal.
Finally, when two atoms are replaced by lone pairs in octahedral geometry, a new molecular shape is formed. An example of this
xenon tetrafluoride, one of the few Noble gas compounds that have been made. The Lewis structure is shown below.
In this case, the two lone pairs take positions as far away from each other as possible, which leaves a square planar molecular shape, as shown below.
We will be learning which atomic orbitals are used to arrive at these molecular shapes in the next lecture.