If you prefer to watch it, see Video HERE, or catch the entire Organic Chemistry Basics Series
Leah here from leah4sci.com and in this video we’ll talk about hybridization, bond angles on geometry as it’ll show up in your Organic Chemistry course. You can find my entire Orgo Basic Series by visiting my website https://leah4sci.com/OrgoBasics.
To truly appreciate and understand Hybridization, you must know your Orbitals and your Electron Configuration. If you’re not confident with these topics, go back to my Intro to Orgo video series which you can find on my website at https://leah4sci.com/OrganicChemistry.
The idea of Hybridization is taking orbitals that want to bond, orbitals that are not capable of bonding and combining them into some hybrid that is able to create a bond with another atom. Let’s look at carbon to understand this in more details.
The electron configuration for carbon is is 1S^2, 2S^2, 2p^2. 1S^2 is the kernel or the internal shell which we don’t look at. It’s only the valence shell that participates in bonding and in this case we have 2 electrons in the 2 S Orbital and two electrons in the two P orbital.
If we draw this out in terms of where the electrons are located and their energy, we have 1s Orbital and 3p Orbitals. The s Orbital is full with 2 electrons, the p Orbital has just 2 electrons between the three sub orbitals. In order to create a covalent bond, another atom will have to take one of its electrons and combine it with a free or an available electron.
If we look at carbon the way it is, we can have an electron bind with this p electron, another electron bind here giving me a total of two bonds. But we know that carbon as the core atom in Organic Chemistry is capable of forming 4 bonds but how is that possible?
In order to form the 4 bonds we have to somehow combine the Orbitals and Electrons from S and P and create a hybrid as follows: Bonding requires energy and that’s lipid of energy will take the low energy s-orbital and raise it up to become equivalent with the p-orbital. The electrons are now considered to be degenerate meaning of the same energy. And if we have degenerate electrons, they will prefer to be spread out over the individual orbitals instead of being doubled up if they’re still an empty orbital available.
As I explained in my Orbital video, the idea is, would you share a room with your sister if you have an empty room in the house? No! I personally would want to have my own room if I didn’t have to share with a sibling.
What we have now is some mix of 4 sub-orbitals containing a total of 4 equivalent electrons. To create this combination we have to combine 1s and 2p orbitals and the designation for this would be s times p times p. In Math, when you have something times itself it becomes squared times itself again becomes cube and so the hybrid orbital is sp cube or simply sp3. An example of this would be the molecule methane. We have carbon single bound to 4 hydrogen atoms.
When carbon is an sp3 hybridization, we have four equivalent sp3 hybrid orbitals and each hydrogen atom is capable into binding to one of them. If you’ll look at Methane and had to figure out the hybridization, here’s the trick; Count your bonds or group, in this case we have 1, 2, 3, 4 – draw four lines and then start counting from s to p knowing that there’s only 1 s, 3p anything beyond that becomes p. We fill it in we have S and then we followed a p, p and p and that gives me sp3.
In order for 4 groups to be equidistant from each other, we get a bond angle or feda of 109.5 degrees. This is a number that you do have to memorize. SP3 hybridization is typically 109.5. Sometimes the number will be slightly greater or slightly less. If there a difference in polarity between the atoms that bind that central atom, that’s the number to memorize.
You can have atoms that are not carbon that will also undergo hybridization. But a little harder to see. For example if we look at Nitrogen NH3, Nitrogen is bound to 3 hydrogen atoms, it also has 2 valence electrons. If we look at the electron configuration for Nitrogen, we have 1s^2, 2s^2, 2p3. 1s^2’s a kernel, we get rid of it, 2s^2 and 2p^3 are the valence electrons.
We have the 2s with 2 electrons and 2p with 3 electrons. As it is, it looks like the Nitrogen is cable of binding to the hydrogens without any change or hybridization. But when you have lone pairs of electrons in a molecule, always think of it as the electron is bound to the other electron and therefore it has to be at the same or very similar energy level to the regular bonds.
So in this case, we’ll still elevate the s to the level of p and have the combine hybrid orbital of sp3. The way you want to envision this is as follows: Remember for carbon we draw our shape like this showing 4 equal units with the tetrahedral geometry. We’ll bind that 3 hydrogens because we know that’s there. For the 4th one, instead of putting the lone pair, just imagine as it the electron is bound to the other electron. It’s not really, they’re in there together but this is how you can understand that even though we have only 3 bonds it’s still a tetrahedral electronic configuration.
The reason I stress electronic is there are two ways to look at this shape. For the electronic geometry, we’re looking at where the electrons are located. We’re not looking at the atoms that they’re bound to. So for example Nitrogen bound to hydrogen, we’re only looking at the bond which is an electron on electron. Just like with the lone pair, we have the electron and electron.
The electronic geometry which comes from the sp3 hybridization has 4 different groups giving me a tetrahedron as the noun or tetrahedral as the verb, same thing. But if we look at just the molecule, meaning only the visible atom we have Nitrogen bound in a weird pyramid to the three hydrogen atoms. Ignoring the effect of the electrons, the Molecular Geometry is actually Trigonal Pyramidal. The trigonal comes from the 3 atoms which is a triangle shape but the pyramid tells us that it’s bent slightly downward from the Nitrogen and forms a pyramid rather than being flat. This is why you should know about the electronic and molecular.
Even though we have the triangle shape from the hydrogens, that lone pair of electrons being more negative compared to the atoms that are sharing is going to repel the bonds downward forcing them into a pyramid rather than allowing them to be flat as we see later with sp2 hybridization.
Let’s look at one more example of an SP3 atom. This time, we’ll look at water where oxygen is bound to 2 hydrogen atoms but also has 2 lone pairs. We’ll draw the Oxygen in a tetrahedral shape given that it’s sp3 and has 4 equal group surrounding it. Two of the groups will get a hydrogen atom where we have the bonds and two of the groups will get an electron bound to an electron as our way of understanding how it’s tetrahedral. Notice that even though we have the two atoms, we still have that overall sp3 hybridization.
Since this is sp3, the electronic geometry is still tetrahedral but the molecular geometry which only looks at the molecule visible and not the electrons surrounding it is going to be simply bend because the oxygen has two hydrogen atoms and they’re bent slightly towards each other. The bond angles are still approximately 109.5 but because the lone pairs are considered to be slightly more negative they’re going to push on the hydrogen atoms forcing them closer together. If this was an exam and you didn’t memorize the exact number, you can write slightly less than 109.5 for the angle between the hydrogen atoms and slightly greater than 109.5 for the angle between the two lone pairs. I think the angle is somewhere near 105 but unless you’re told to memorize it, don’t worry about it.
But be sure to join me in the next video where I take you through sp2 and sp hybridization and then show you a fun trick for how to quickly recognize the hybridization on wacky molecules including drugs that professors love to put on exams.
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