Because carbon plays such a significant role in organic chemistry, we will be using it as an example here. Carbon's 2s and all three of its 2p orbitals hybridize to form four sp 3 orbitals. These orbitals then bond with four hydrogen atoms through sp 3 -s orbital overlap, creating methane. The resulting shape is tetrahedral, since that minimizes electron repulsion. Lone Pairs: Remember to take into account lone pairs of electrons. These lone pairs cannot double bond so they are placed in their own hybrid orbital.
This is why H 2 O is tetrahedral. We can also build sp 3 d and sp 3 d 2 hybrid orbitals if we go beyond s and p subshells. The frontal lobes align themselves in the trigonal planar structure, pointing to the corners of a triangle in order to minimize electron repulsion and to improve overlap.
The remaining p orbital remains unchanged and is perpendicular to the plane of the three sp 2 orbitals. Hybridization of an s orbital with two p orbitals p x and p y results in three sp 2 hybrid orbitals that are oriented at o angle to each other Figure 3. Sp 2 hybridization results in trigonal geometry. In aluminum trihydride, one 2s orbital and two 2p orbitals hybridize to form three sp 2 orbitals that align themselves in the trigonal planar structure.
The three Al sp 2 orbitals bond with with 1s orbitals from the three hydrogens through sp 2 -s orbital overlap. Similar hybridization occurs in each carbon of ethene. For each carbon, one 2s orbital and two 2p orbitals hybridize to form three sp 2 orbitals. These hybridized orbitals align themselves in the trigonal planar structure. For each carbon, two of these sp orbitals bond with two 1s hydrogen orbitals through s-sp orbital overlap. The remaining sp 2 orbitals on each carbon are bonded with each other, forming a bond between each carbon through sp 2 -sp 2 orbital overlap.
This leaves us with the two p orbitals on each carbon that have a single carbon in them. These orbitals form a? Because a double bond was created, the overall structure of the ethene compound is linear. However, the structure of each molecule in ethene, the two carbons, is still trigonal planar.
This formation minimizes electron repulsion. Because only one p orbital was used, we are left with two unaltered 2p orbitals that the atom can use.
These p orbitals are at right angles to one another and to the line formed by the two sp orbitals. Figure 1: Notice how the energy of the electrons lowers when hybridized.
These p orbitals come into play in compounds such as ethyne where they form two addition? This only happens when two atoms, such as two carbons, both have two p orbitals that each contain an electron. An sp hybrid orbital results when an s orbital is combined with p orbital Figure 2.
So what happens in those guys? Well one of those bonds within a multiple bond is called a sigma bond and again don't forget sigma bonds are hybridized so one of those bonds is going to be hybridized. The rest of those bonds are called pi bonds, those pi bonds are just P orbitals overlapping each other, they're only P orbitals they're a little higher in energy, they actually are different in energy.
So we're going actually keep then separated, so we have 1 sigma and 1 pi. So let's look at carbon dioxide here, we have a double bond alright of these is going to be a sigma bond and we're going to denote that with a sigma and one of the bonds is going to be a pi bond we'll denote it with pi. These are only P orbitals, these are hybridized orbitals we're just talking about the carbon right now. Okay so carbon we already know looks like this we're going to save 2 of the P orbitals and I don't care which 2 I save it doesn't really matter, they're all the same in energy I don't care.
I'm just going to save this just for practical purposes, these are going to be the ones to use in pi bonding, so I'm going to save those so they hybridize 1S and the other P.
Let's look at the oxygen, oxygen also has a sigma bond, a pi bond but notice it has lone pair. So the sigma bond and the lone pair are going to be hybridized but not the pi bond we're going to leave this lone.
So we need 3 hybrid orbital, 1 from S and 2 from P, so it's going to be SP2. It doesn't matter that these guys have different types of orbitals, we just want to make sure that the orbitals within the atom itself are the same. So let's look at triple bond, triple bond one is sigma don't forget.
So we're going to say this is sigma and these 2 are pi. Okay, so these 2 are just P we're going to ignore them and this nitrogen has, needs 2, 1 for the lone pair and 1 for the sigma hybridized orbitals.
Okay look at Ozone, O3 all 3 of these are a little bit different so this guy has 1, 2, 3 lone pair and 1 sigma bond so it needs 4, 1 from S, 3 from P so it's going to be SP3. Oxygen is going to need, it has 2 sigma bonds 1 from here, 1 from here and lone pair so it needs 3 so it's going to be SP2 1 from S, 2 from P. This guy over here has 1 sigma bonds and 2 lone pairs again 3 hybridized orbitals, so it's going to be SP2, 1 from S, 2 from P. So hopefully that made it a little bit easier for you to figure out hybridized orbitals.
All Chemistry videos Unit Chemical Bonds. Previous Unit The Periodic Table. Next Unit Chemical Reactions. Se the explanation portion of this answer. Hybridization is a theory which starts from the consideration that when atoms combine to form a molecule, the orbitals combine to create an entirely new and different molecular orbital that does not resemble the original atomic orbitals. This new hybrid molecular orbital "belongs" to the molecule as a whole, but its geometry is determined by the types of atomic orbitals in the central atom that were involved in the bonding of the molecule.
Thus is the case of carbon, the atom uses its 2 s and all three 2 p orbitals to create four identical bonding orbitals. How and why does hybridization occur? Dec 18,
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