Hybrid Orbitals explained - Valence Bond Theory | Crash Chemistry Academy

Welcome to a video about hybrid orbitals often called Valence Bond Theory developed in the 1930s by the great chemist Linus Pauling as a model of bonding to understand the there is only a small group of atoms in the second period that the model really works for. But among those are carbon nitrogen and oxygen which make up the vast majority of molecules that exist on earth. So the model applies to a limited number of elements, but it applies by far to the majority of molecules. Let's take a look at the word hybrid. It is a blending of two varieties If you get a horse and donkey together, you get a mule, a blending or hybrid of a horse and a donkey. Let's get to hybrid orbitals using carbon as our model. As a single atom not bonded to anything carbon has two electrons in 1s two electrons in 2s and two electrons in 2p. This electron configuration is the energy arrangement of carbon's electrons. However carbon rarely exists in nature as an individual atom except momentarily while undergoing chemical reactions. Carbon exists with its valence electrons bonded to other atoms. When carbon bonds to four other atoms carbon's four bonds are experimentally seen to be equivalent, and so when the carbon atom finds itself in a bonding situation, its bonding electrons themselves exist at equivalent energies which requires that they hybridize to an energy that is intermediate between the 2s energy and the 2p energy. Or you can think of it as a blending or hybridization of the two energies, the s and p energies. And since the energies of these electrons have now changed, the shape of the orbitals they occupy are different as well, Which we will see momentarily, and those are called hybrid orbitals. They are named 2sp3 hybrid orbitals. The naming often confuses students so before we go any further Let's take a look at where the 2sp3 name comes from. The 2 comes from the second principle energy level that the valence orbitals are in The s comes from the 2s orbital contributing to the hybridization and the p comes from the 2p orbitals contributing to the hybridization and the ^3 comes from the number of 2p orbitals used in hybridization. once hybridized, the 2s and 2p orbitals no longer exist, and so we have Before we look at the shape of hybrid orbitals it would be helpful to briefly review the atomic orbitals The 1s orbital is a sphere the 2s orbital is a larger sphere surrounding 1s, and here we will get rid of 1s since we are only concerned with the valence electrons each 2p orbital is a 2 lobed shape converging at the nucleus So there are the three 2p orbitals. However, when hybridization occurs the s and p orbitals cease to exist and the 2sp3 orbitals have an entirely different shape We can see that The carbon atom only hybridizes when it is in a bonding situation Here four hydrogen atoms bond to carbon by overlapping their orbitals with carbon's hybrid orbitals So what would be the reason for this? if we go back and see that both carbon and Hydrogen have unpaired electrons The overlap allows electrons to pair and thus go to a lower potential energy. The Illustration here contains the valence electrons of both carbon and hydrogen and since everyone likes to visualize atoms as spheres we can do the same. Here is our carbon atom and here are the hydrogen's with the overlapping spheres indicating the overlapping orbitals that constitute the bond The bonds are more readily discernible in a ball and stick model which also makes the bond angle more visible The bonds in hybridization also have their own nomenclature The overlapping orbitals are called sigma bonds which represents the single bond occupied by a single pair of electrons What about double bonds? How does the hybridization model explain double bonds? We will use ethene, C2H4, to see what happens in a double bond The single bonds we know are sigma bonds, and the double bond also has a sigma bond But the second bond of a double bond is a pi bond. Let's see how hybridization and orbital overlap can explain a double bond. The two carbon atoms in ethene are equivalent So let's look at one of the carbon atoms first. Leaving a p orbital unhybridized for the pi bond. The hybrid orbital is called 2sp^2, the superscript ^2 denoting that only two 2p orbitals have contributed to the hybridization. The 2sp2 hybrid orbitals exist in a plane perpendicular to the unhybridized 2p orbital. Let's remove the 2p orbital to more easily see that the 2p hybrid orbitals are spread out at a 120 degree angle which means that they exist in a plane the plane is perpendicular to the unhybridized 2p orbital So this is what both carbon atoms do when bonding occurs in ethene: each carbon atom is sp2 hybridized. The sigma bond occurs with 2sp2 orbital overlap. What about the pi bond, the second bond of the double bond? Previously we said that it comes from the unhybridized p orbitals, which we see here from both carbon atoms The top and bottom lobes of the 2p orbitals overlap above and below the axis of the sigma bond, forming a single pi bond The space in which the now paired electrons move around The ball and stick model shows this double bond with two dashes To summarize: The ethene molecule also bonds to four hydrogen atoms by overlapping with both carbons' other 2sp2 hybrid orbitals creating for more sigma bonds In the ball-and-stick model we can readily see that each carbon has a trigonal planar geometry and thus the whole molecule exists in a plane with the single pi bond above and below that plane Now let's look at how hybridization can be a model for the triple bond using ethyne, C2H2 The carbon-hydrogen bonds are sigma bonds and the triple bond is one sigma bond and two pi bonds Let's see how hybridization can accommodate this Since Pi bonds come from p orbitals and we need two pi bonds then two 2p orbitals have to remain unhybridized and so the remaining single 2s orbital and a single 2p orbital will hybridize to two 2sp hybrid orbitals and there is also the violet 2px and the blue 2py orbitals Here each green lobe is a single orbital and so they each have an electron while both violet lobes constitute the single 2px orbital with a single electron, and both blue lobes constitute the single 2 py orbital with a single electron The other carbon and C2H2 also has that same triple bond and so it has the same hybridization Let's see what happens during bonding sp orbitals from both carbons overlap, forming a sigma bond The upper lobes of the blue 2p orbitals overlap as do the lower two lobes creating the first of the two pi bonds. Can you guess where the second of the two pi bonds comes from? Yes that's right... It is the overlapping of the violet 2p lobes Let's get rid of the sigma bond for a moment to take a look at something interesting Each pi bond lies on a separate plane and the two planes are perpendicular to each other and So the two pi bonds are perpendicular to each other Finally two hydrogen's will overlap with the remaining sp hybrid orbitals Creating the C2H2 molecule. The overlap of the space-Filling model reflects the overlapping orbitals which is also represented by the ball-And-stick model. In the remainder of the video we will look at hybridization of nitrogen and oxygen using NH3, ammonia, as our model for nitrogen hybridization and H2O, water, for our oxygen model. In NH3, nitrogen has three sigma bonds and a lone pair So how does hybridization account for this? The hybridization is 2sp3, and nitrogen has 5 valence electrons, So one of the four 2sp3 hybrid orbitals has a pair of electrons The 3 sigma bonds come from the sp3 orbitals with a single electron, so they can pair up, and so the remaining electron pair is a lone pair, an unbonded pair of electrons. As with sp3 hybridization in carbon, nitrogen's hybrid orbitals spread out in a tetrahedral shape. And lastly water. here, oxygen has two sigma bonds and two lone pairs. In water oxygen is also 2sp3 hybridized, but with six valence electrons two of the sp3 orbitals have paired electrons you can probably guess that the SP3 orbitals with a single electron will overlap with hydrogen and the remaining two pairs are unbonded: they are lone pairs. Again, oxygen's hybrid orbitals spread out in a tetrahedral shape. That's it for hybridization the product of a mad scientist. Seeya!

Channel: Crash Chemistry Academy

Views: 495,920

Rating: 4.9573693 out of 5

Keywords: hybrid, hybrid orbitals, valence, valence electrons, bonding, bond, bonds, valence bond theory, Linus Pauling, bondind orbitals, organic chemistry, chemistry, find, fun, easy, beginners, calculate, math, table, s orbitals, p orbitals, methane, ethene, ethylene, acetyline, ethyne, ammonia, water, tetrahedral, trigonal planar, metallic, crash, how to, funny, academy, song, problems, crash chemistry

Id: vHXViZTxLXo

Length: 11min 58sec (718 seconds)

Published: Thu May 04 2017

Reddit Comments

[u/myball_]

Wow it’s like god is watching and posted this for me, through you

Thank you

[u/daal-chawal]

Same here lol bio major but literally only understood most of chem after graduating when I started studying for the mcat

[u/aspiringdoc74]

V needed rn, thank you!

[u/UniqueInterview]

I didn’t even know I had this gap of knowledge before watching this lol thank you!

[u/minutestomidnight]

THey don't teach this stuff in Bio. It's chem, baby.