To explain the bonding in simple molecules using atomic orbital hybridization theory.
To correctly draw molecular orbital diagrams of sp3, sp2 and sp hybridized atomic orbitals.
To explain how orbital hybridization theory is consistent with the observed properties of carbon-carbon bonds.
When we use atomic orbitals to explain bonding, a few observations can be made with elements and their compounds. Take carbon for example:
Carbon's ground state electron configuration is [1s2] 2s2 2p2. The two 2s electrons are paired up and two p electrons are unpaired. This would suggest that two electrons are in a different energy state than the other two. Two examples of carbon compounds:
Methylene (:CH2) is a type of compound called a carbene. Carbenes have two electrons on carbon that are not shared; the other two are making covalent bonds. However, methylene is very reactive and unstable – most simple carbenes are.
In methane (CH4), carbon is bonded to four hydrogens while the bond angles and strengths are all equal. Methane is a stable compound that is easily stored at standard conditions.
This suggests these electrons in methane are actually in an equivalent state. Another possible way of thinking about and explaining this bonding is the s and p atomic orbitals containing the four electrons are hybridized – mixed together. We call this idea atomic orbital hybridization. Like a hybrid animal or hybrid car, hybrid orbitals are mixes of different atomic orbitals that show characteristics of both.
In methane for example, orbital hybridization explains that carbon's four valence electrons do not bond as in their ground state (2s2 2p2). The atomic orbitals mix together to create four equivalent orbitals that are equal in energy.
This is done by a 2s electron being excited to a slightly higher 2p orbital that is currently empty.
This results in four singly occupied orbitals; three 2p orbitals equal in energy and one 2s orbital slightly lower in energy. Now carbon has the ability to form four bonds.
These four orbitals mix together equally to create four hybrid orbitals of equal energy.
In methane for example, the lone 2s orbital and three 2p orbitals create four sp3hybrid orbitals. The hybrids are proportional in character to the AOs that made them, so they are 25% 's character' and 75% 'p character'.
These hybrid orbitals then combine with orbitals in other atoms (like the 1s AOs in hydrogen, for example). These combine head on so are considered sigma bonds.
This would explain the observation of equal bond angles and bond strength in methane. These hybrid orbitals can be drawn as a mix of s and p orbitals, where they do still look similar to the p orbitals they have 75% of the character of, with the s character shown as one lobe of the p orbitals is enlarged, the other shrunk. See below for a diagram:
Hybridization can be used to explain the bonding involving double bonds too. Ethene (C2H4), for example, has its two carbon atoms bonded to three other atoms; the other carbon atom and two hydrogens.
How does atomic orbital hybridization explain its bonding?
Both carbon atoms have two bonds to hydrogen atoms and one to the other carbon atom. So after s-p mixing, only three (not four like in CH4) atomic orbitals need to be hybridized. This takes the 2s orbital and two p orbitals to make three in-plane sigma bonds, leaving one 2p orbital alone. Because these hybrid orbitals are formed from one s orbital and two p orbitals, the orbitals in ethene are called sp2hybrid orbitals. They are one-third s character and two-thirds p character – you can draw them more-or-less like the sp3 hybrids you did for methane!
The last unhybridized 2p orbital will form the pi bond between the two carbon atoms. However, this 2p orbital is higher in energy than the hybridized sp2 orbitals, so the MO diagram looks a little 'top heavy'. See the diagram below.
Hybridization can be used to show the bonding for a carbon triple bond, too. We say carbon in ethyne (the simplest alkyne) is sp hybiridized. How many orbitals do you think are hybridized here?
Ethyne has only two atoms bonded to each carbon atom – it only needs to hybridize for a C-C and C-H sigma bond, taking the 2s orbital and one 2p orbital, hence it the label sp and 50:50 s and p character. The two perpendicular unhybridized p orbitals form the two pi bonds between the two carbon atoms to make the triple bond overall. See the diagram below.
The type of hybridization a carbon atom has with its orbitals is easily found by the number of atoms the carbon center is bonded to:
Number of atoms bonded (to carbon)
Shape around carbon
Along with molecular orbital theory, orbital hybridization helps explain the observed difference in properties of compounds. Here are some examples:
The difference in bond enthalpy – effectively how strong a bond is – shows:
A C=C double bond is not twice as strong as a C-C single bond.
The carbon-carbon triple bond is not three times as strong as a C-C single bond.
This is because the double and triple 'parts' of these bonds are the pi bonds, where orbitals overlap side on and are thus higher in energy than the head-on interactions of orbitals in a sigma bond, which is the single bond 'part' of the bond.
Mixing atomic orbitals - hybridization.
Atomic orbital hybridization
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