Molecular Shape Important
The structure of a molecule is not complete until you know the molecular shape. Molecules which have seemingly an identical molecular shape often have very different chemical properties.
The featured image above shows a flow chart you can use to deduce the electron geometry and molecular geometry of simple molecules.
The shape of a molecule influences if a molecule has a dipole. If a molecule has a dipole it interacts with other molecules with dipoles. It affects the melting point and boiling point of liquids.
Surface Area and Boiling Point
For instance, compare two five-carbon hydrocarbons with exactly the same formula weight, Figure 1. Linear n-pentane in a straight line has a higher boiling point then the more sphere like neopentane.
The straight chain pentane has a larger surface to area ratio to interact with other pentane molecules. Neopentane, on the other hand has a much smaller surface area to interact with neighboring molecules.
Shape and Stacking
Molecular shape controls how molecules pack together.
Linear stearic acid allows efficient packing compared to oleic acid, which has a bend in the molecule. As a result, oleic acid melts at a much lower temperature than stearic acid, Figure 2.
Molecular recognition comes about from molecular shape. An example is the ability of psychoactive drugs to effect the nervous system. If a molecule shares a molecular shape similar to normal substances found in the body, it activates a response.
Figure 3 shows an amino acid accepted by the surface of a cell, whereas CH2ClF gets rejected on the basis of molecular shape.
This article covers molecular shape up to the point where the central atom has up to four atoms attached. This forms the most useful group of atoms for readers interested in learning the basics of molecular shape and what will be the most useful for those reviewing molecular shapes who intend to use this information to help with organic chemistry.
Simple examples are used where the molecular shape and geometry is described for a molecule which has a central atom. The molecular geometry results from the immediate environment around the central atom.
The basis of molecular shape resides in the negative charge of electrons. All electrons repel each other, Figure 4.
Whether the electrons take the form of shared electrons between bonds or lone pairs of nonbonding electrons.
Atoms and Electrons Around Single Point
When several atoms are held around a central point, the electrons spread as far apart as possible to minimize the force created by electron-electron repulsion.
For instance, if a central atom has three atoms attached to it without any lone pairs of electrons, the atoms arrange themselves so they are all as far apart as possible, Figure 5.
The way electron-electron repulsion leads to molecular shape is called Valence Shell Electron Pair Repulsion, VESPR.
Deduce the molecular shape through a three step process:
The process to find the molecular shape is in Figure 6.
The Lewis dot structure shows the connections between atoms. This is a two-dimensional structure.
The three-dimensional structure considers electron-electron repulsions. The electron geometry shows how the bonds and lone pairs of electrons map into a three-dimensional space.
In the final step, consider the position of the atoms. The molecular geometry describes the molecular shape in terms of where the atoms map out with respect to each other.
The details of how to arrive at a correct Lewis structure are not included, but required at a starting point to figure out a correct molecular shape. It is especially important to include the “dots” which represent lone pairs of electrons.
The term electron groups includes the electrons within a bond between two atoms and the electrons which reside in a lone pair of electrons. For the purpose of finding molecular shapes, the important factor is electrons repel each other. This means bonding electrons and nonbonding electrons must both be considered. Collectively they are both electron groups.
Electron geometry comes from electron-electron repulsion. If all bonds between atoms with shared electrons and lone pairs of electrons are treated as being equal, the molecular shape which minimize electron-electron repulsion is adopted.
The further apart the electrons the better.
Common examples of linear molecules are shown in Figure 7.
Diatomic molecules connect with chemical bonds. Any combination of single, double, or triple bonds results in a linear molecular shape.
In the case of more advanced molecules like CO2 and HCN, the molecules are linear because the central atom, carbon, has no lone pairs of electrons to influence the electron geometry.
Trigonal planar electron geometry comes from a central atom which has three atoms attached or two atoms and one lone pair of electrons, Figure 8.
BF3 shows a simple molecule with three single bonds and no lone pairs. Electron-electron repulsion force the atoms to form the points of a triangle around the central boron.
COCl2 is like the previous example, except the central carbon has a double bond between carbon and oxygen. Notice double bonds behave the same way as single bonds for finding the electron geometry.
O3 and SO2 are different. Instead of the central atom having three atoms attached, there are two atoms and a lone pair of electrons. Treat the lone pairs of electrons the same way you consider an atoms surrounded by electrons.
To assign a molecule’s electron geometry, the electrons in a chemical bond and the electrons in lone pairs are equivalent. This makes the central atoms in O3 and SO2 have three substituents.
Central atoms with three electron groups adopt a trigonal planar geometry.
Molecules with a central atom and four electron groups take a tetrahedral electron geometry, Figure 9.
This is easy to see when carbon is the central atom. Carbon has no lone pairs of electron and four atoms attached, CH4 and CH2ClF.
The optimal molecular spreads the four atoms and electrons to the four points of a tetrahedron.
Both H2O and NH3 have lone pairs of electrons which act as electron groups. The electron groups each occupy places at the points of a tetrahedron.
The final way you describe the molecular shape of a molecule is with the molecular geometry. Once you find the electron geometry, the molecular geometry corresponds to how the atoms instead of electrons. Keep the same position you found from the electron geometry, but now examine where the atoms reside without electrons.
Examples of molecules which have a linear molecular geometry are shown in Figure 10.
Any diatomic molecules will always remain linear. Both CO2 and HCN both have carbon as a central atom. Carbon does not have any lone pairs of electrons. This means any molecule with a linear electron geometry also has a linear molecular geometry.
Bent molecular geometry always comes from an electron geometry which differs from the molecular geometry.
All molecules that have a bent molecular geometry have lone pairs of electrons present in the electron geometry Figure 11.
O3 and SO2 have three electron groups around the central atom. One of the electron groups is a lone pair of electrons. To consider the molecular geometry, the lone pair of electrons is removed.
In that way, the electron geometry which is trigonal planar, becomes a molecular geometry which is bent.
When a central atom like boron or carbon has three electron groups, the molecule adopts a trigonal planar electron geometry. Because neither carbon or boron have lone pairs, the electron geometry and the molecular geometry remain the same. This results in COCl2 and BF3 have a molecular geometry which is trigonal planar, Figure 12.
The trigonal pyramidal molecular geometry comes from a central atom with four electron groups, Figure 13.
Three of the electron groups are bonds connected to atoms. The fourth electron groups comes from a lone pair of electrons. This makes the electron geometry of NH3 tetrahedral.
If carbon has four atoms attached, it takes a tetrahedral electron geometry. A tetrahedral molecular geometry comes from a tetrahedral electron geometry.
Any Group 14 (IV A) element will take a tetrahedral molecular geometry, like CH4 and CH2ClF, Figure 14.
Only the molecular geometry of atoms which obey the octet rule, (except boron), have been treated here. These are the most common molecules you encounter in a general chemistry course and what you need the most in order to move onto organic chemistry.
Elements from Period 3 and heavier elements can have more than four electron groups attached. This means there are more electron geometries and molecular geometries. These will be treated elsewhere.