From the Lewis structure of CO2 and VSEPR, we can determine that this is a linear molecule.
Let's look at some of the physical properties of CO2.
Property
CO2
boiling point
195 (sublimes)
Hf (298 K)
-393.5 kJ/mol
bond energy
806 kJ/mol
C-O bond distance
1.16
dipole moment
0 D
The C-O bonds in carbon dioxide are polar and yet the dipole moment is zero because the 2 bond dipoles cancel each other.
One thing that we can understand by looking at the structure of CO2, is that the carbon center of the molecule must be electrophilic. An electrophile (electron-lover) is a center that is electron poor and will be attracted to centers that are electron-rich.
Even though the total electron count around the carbon is 8, this overestimates the electron density. This carbon is bonded only to highly electronegative oxygen atoms. The bonding electrons will all be more closely associated with oxygen than with carbon.
Bonding in Carbon Dioxide
From the Lewis structure we can see that the carbon in CO2 must make 2 sigma bonds and it has no lone pairs. This atom will be 2sp hybridized with remaining 2px and 2py atomic orbitals.
Each oxygen makes 1 sigma bond and also needs 2 orbitals for lone pairs of electrons. These must each be 2sp2 hybridized with a remaining 2p orbital. One of the oxygens will have a 2px orbital to combine with the carbon 2px orbital. The other oxygen will have a 2py orbital that can combine with the other p orbital on carbon.
A 2sp2 orbital on O1 combines with a 2sp orbital on C to make a sigma bonding and a sigma antibonding molecular orbital. The other 2sp orbital on C combines with a 2sp2 orbital on O2 to make another set of sigma bonding and sigma antibonding molecular orbitals.
The remaining 2sp2 from the oxygen atoms become non-bonding molecular orbitals.
The O1 2px combines with the C 2px to make a pi bonding and pi antibonding molecular orbital. The O2 2py combines with the C 2py to make another set of pi bonding and pi antibonding molecular orbitals.
The 16 valence electrons fill through the 2 pi bonding orbitals so there is a full double bond between carbon and each oxygen.
Oxidation States
As you saw above, the total electron count around the atoms in carbon dioxide seriously overestimates the electron density around the carbon atom. It doesn't help us predict the reactivity of this atom. The oxidation state formalism can give us a better idea about the electron density around an atom and its tendency to add electrons and become reduced.
To find the oxidation state of atoms in CO2,
Draw the Lewis structure.
Break the bonds giving all of the bonding electrons to the more electronegative of the 2 atoms. (When the atoms are the same, give each atom 1/2 of the bonding electrons.)
Count the electrons around each atom and compare the number of electrons to the number of valence electrons, just as you do for formal charge.
Use Roman numerals instead of numbers to designate the oxidation state.
From the oxidation states, we see that the carbon center is very electron poor and in its highest possible oxidation state. It should be susceptible to reduction.
Reductive Coupling
Sodium metal has a single electron in its valence shell. It has a strong tendency to lose that electron and become oxidized. Carbon in CO2 is in it's highest oxidation state. It should have a tendency to gain an electron and become oxidized.
The unpaired electrons on two of the reduced carbon centers can combine to form a covalent bond in the product, sodium oxalate.
Hydroxide Addition
The carbon of CO2 is electrophilic (electron-poor). The oxygen in hydroxide ion, HO-, has excess electron density on oxygen. An electron-rich center that can form a bond with an electron-poor carbon atom is called a nucleophile (positive charge-lover). The purple arrows in the reaction scheme indicate the flow of electrons in the reaction.
Professor Patricia Shapley, University of Illinois, 2012
Hf (298 K)
