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In looking at the lewis structure for CO2, you see that C has two regions of electron density and, therefore, has the hybridized orbital of sp, making one sigma sp bond with each oxygen. You know that there are two more unhybridized p orbitals because there are three total p orbitals for any energy level and one has already been hybridized. In general, you can tell how many remaining unhybridized orbitals there are based on the total number of orbitals in that subshell minus the number of orbitals in the same subshell that has already been hybridized. Oxygen, on the other hand, has three regions of electron density and has a hybridized orbital of sp2, leaving one unhybridized p orbital. This leaves the carbon two unhybridized p orbitals to make a pi p bond between each oxygen and their one remaining unhybridized p orbital.
In the a table in the textbook, it says that a molecule with an electron arrangement that is linear, the hybridization of the central atom is sp. If it is trigonal planar, it is sp^2. If the electron arrangement is tetrahedral, it is sp^3. If the electron arrangement is trigonal bipyramidal, it is sp^3d. And if it is octahedral, it is sp^3d^2. The overarching rule is that N atomic orbitals always produce N hybrid orbitals.
Carbon has 4 valence electrons that can form bonds. In CO2, there are 2 regions of electron density in which the sp hybrid orbitals participate. There are still two more electrons, therefore there are two non hybridized orbitals.
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