Hybridization is a theory that takes a "holistic" view of the bonds in a molecule. 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. It only has 2 unpaired electrons even though we know that carbon can form 4 bonds.
Re: reason for hybridization Post by harperlacroix1a » Sat Nov 24, am It exists because hybridized orbitals are lower in energy which results in a more stable compound. Re: reason for hybridization Post by Harshita Talkad 4L » Sat Nov 24, pm Hybridization occurs because hybridized orbitals have lower energy than when they are unhybridized, so the compounds that they form are more stable.
Re: reason for hybridization Post by Manas Jinka » Sun Nov 25, am Hybridized orbitals have less energy than the unhybridized orbitals. The compound is more stable when hybridization occurs. The unhybridized orbitals morph to become new hybridized orbitals which have been experimentally observed. So we're going to have 4 new orbitals and we're going to call them the 1S and the 3 of them are P so we're going to call it SP3, 1 from S 3 from P.
And we're going to spread these out just like the rule tells us to and we're going to say okay we have 4 electrons which gives us 4 equal places for chlorine to come in and actually bond with that carbon. So chlorine is coming here, here, here and here and make 4 of those bonds like you see in the picture. Alright so what actually would get hybridized? What do we create to actually mix it so that these equal orbitals are necessary? So we know all single bonds are going to be hybridized because a single bond there's not one that's more energetic than the other.
So all single bonds are going to be hybridized. Because they're hybridized bonds we're going to now call single bonds sigma bonds, this is just the way they overlap, the way that orbitals overlap we're going to call them, denote them sigma bonds. And we also want to say that low in pairs are also going to be hybridized because they're not higher or lower in energy than those bonds either. So let's look at ammonia as an example, ammonia if you look at nitrogen within ammonia it has these 2 lone pair of electrons.
So ammonia before had the same thing, ammonia has 5 valence electrons so 2, 3, 4, 5, this should be the same I'm sorry they're kind of uneven, they should actually be the same in energy and we have the 5 electrons. If we're going to hybridize all of them we need to have 1, 2, 3 of these are the same along with this fourth one so we need to have all 4 of these is the same, so we're goingto have again 4 equal in energy we're going to call it SP3, 1 from S, 3 from P 1, 2, 3, 4, 5 here's our lone pair and here's the hydrogens that are going to come in and bond with them all equal in energies so we have this new hybrid orbitals.
Okay about when we have multiple bonds? So in different cases we may have multiple bonds, double bonds and triple bonds. 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.
Double and triple bonds still count as being only bonded to one atom. Use this method to go over the above problems again and make sure you understand it. It's a lot easier to figure out the hybridization this way. Introduction Carbon is a perfect example showing the value of hybrid orbitals. Carbon's ground state configuration is: According to Valence Bond Theory , carbon should form two covalent bonds, resulting in a CH 2 , because it has two unpaired electrons in its electronic configuration.
That would give us the following configuration: Now that carbon has four unpaired electrons it can have four equal energy bonds. Energy changes occurring in hybridization.
Energy changes occurring in hybridization 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.
Example: sp 2 Hybridization in Aluminum Trihydride 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. Example: sp 2 Hybridization in Ethene Similar hybridization occurs in each carbon of ethene. Energy changes occurring in hybridization Figure 1: Notice how the energy of the electrons lowers when hybridized.
Example: sp Hybridization in Magnesium Hydride In magnesium hydride, the 3s orbital and one of the 3p orbitals from magnesium hybridize to form two sp orbitals. Hybridization Example: sp Hybridization in Ethyne The hybridization in ethyne is similar to the hybridization in magnesium hydride.
References John Olmsted, Gregory M. Carey Advanced Organic Chemistry Springer Wade, Jr. Problems Using the Lewis Structures , try to figure out the hybridization sp, sp 2 , sp 3 of the indicated atom and indicate the atom's shape. The carbon. The oxygen. The carbon on the right. Answers 1. Hybridization Lone Pairs: Remember to take into account lone pairs of electrons.
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