Ch.10 - Molecular Shapes & Valence Bond TheoryWorksheetSee all chapters
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Ch.1 - Intro to General Chemistry
Ch.2 - Atoms & Elements
Ch.3 - Chemical Reactions
BONUS: Lab Techniques and Procedures
BONUS: Mathematical Operations and Functions
Ch.4 - Chemical Quantities & Aqueous Reactions
Ch.5 - Gases
Ch.6 - Thermochemistry
Ch.7 - Quantum Mechanics
Ch.8 - Periodic Properties of the Elements
Ch.9 - Bonding & Molecular Structure
Ch.10 - Molecular Shapes & Valence Bond Theory
Ch.11 - Liquids, Solids & Intermolecular Forces
Ch.12 - Solutions
Ch.13 - Chemical Kinetics
Ch.14 - Chemical Equilibrium
Ch.15 - Acid and Base Equilibrium
Ch.16 - Aqueous Equilibrium
Ch. 17 - Chemical Thermodynamics
Ch.18 - Electrochemistry
Ch.19 - Nuclear Chemistry
Ch.20 - Organic Chemistry
Ch.22 - Chemistry of the Nonmetals
Ch.23 - Transition Metals and Coordination Compounds
Sections
Electron Geometry
Bond Angles
Hybridization
Molecular Orbital Theory
MO Theory: Heteronuclear Diatomic Molecules
Additional Practice
Delocalized Electrons
Additional Guides
Molecular Geometry (IGNORE)
VSEPR Theory (IGNORE)

How would an s orbital combine with a p orbital to form a covalent bond? The answer is through hybridization where a hybrid oribital is formed. 

Atomic Orbitals vs. Hybrid Orbitals

Concept #1: Understanding Overlapping Atomic Orbitals 

Transcript

What we should realize is that covalent bonds are formed when atomic orbitals basically just overlap with their electrons. Now for electrons to overlap, they have to be in the same type of orbital. So an s orbitals has to overlap with an s orbital. A p orbital has to overlap with a p orbital.
So if we take a look at this first example, here we have two Hydrogens. They want to combine to give us H2. Now, their electron configurations are 1s1. So they each have 1 electron in their first shell. Now, they can overlap to form a bond, because here, what’s going to be forming the bond are their s orbitals. Each s orbital needs one more electron in order to get to 1s2, so that each Hydrogen can be just like Helium. Now, again since they are both s, they can overlap. So that’s what they do. And the electrons are in here. And in this way they get to share the electron of the other Hydrogen, and that way they have two electrons around them, and they are satisfied.
Now if we take a look at two Chlorines, they can also overlap their p orbitals, because the electron configurations will be Argon, or actually [Ne]3s23p5, [Ne]3s23p5. Remember, chlorines want to gain one more electron to become just like Argon. Where can that one electron go? It’s going to go to the p orbitals, because remember, p can hold up to 6 electrons. So there’s an electron in here, and there’s an electron in here. And they are going to overlap with one another, so that they can share those electrons.
So here we’re going to say is where they form their bond. We’re going to say this is a direct overlapping of their atomic orbitals, we say that this is called a sigma bond. A sigma bond is a direct overlapping of atomic orbitals. They share electrons with each other in order to complete their octet, to be just like closest noble gas. 

A sigma bond is formed by the direct end to end overlapping of atomic orbitals. 

Concept #2: Understanding Hybrid Orbitals 

Transcript

Now, we’re going to say that the opposite of a sigma bond is a pi bond. We’re going to say a pi bond is a sideways overlapping of p orbitals. And we’re going to say if you have elements single bonded to each other, you have 1 sigma. Every bond you make is always a sigma bond.
C-C
Let’s say compound has two bonds together. It’s a double bond. Well, it’s going to have 1 sigma, and 1 pi.
C=C
And let’s say compound has three bonds, triple bonds. Again, it’s going to have 1 sigma, and 2 pi bonds. Every extra bond you add, that’s what pi bond be added.
C=C
So just remember, here we’re talking about direct overlapping of atomic orbitals, so we’re talking about the formation of a bond. A sigma bond. And then we’re going to say if we have a double bond or a triple bond, those will be pi bonds that we’re adding on top of the sigma bond. They are kind of like bread that goes around the meat of the bun. The meat portion will be the sigma bond, bread part adds extra strength to that sigma bond.
Now, what would happen though if we’re trying to form a covalent bond with different atomic orbitals?
For example, we have BeCl2. Now BeCl2 looks like this.
Cl-Be-Cl
Looks like Beryllium is in group 2A, so there are two bound electrons, Chlorine is in group 7A, so it has seven. Halogens are not in the center, only make one bond. But the problem here is, electron configuration of Beryllium is Helium, 2s2, and Chlorine is [Ne]3s23p5. And here’s the thing. An s and a p cannot mix together. There are different atomic orbitals. They have to both be the same letter for them to successfully mix. This is where hybridization comes into play. Now, hybridization is a way of us mixing our s, our p, and sometimes our d orbitals together, to make hybrid orbital. That hybrid orbital will have a little bit of s character, p character, and d character. That will allow the central element to form bonds with basically any other element.
So what’s going to happen here is, for the Beryllium to form a bond with the Chlorine, it needs to somehow gain a p orbital. So what’s going to wind up happening here is, in this 2s electron orbital, we’re going to have a promotion of one of its electrons. So this electron here is actually going to get promoted, and jump up to the 2p orbital. And as a result, we have now one in the 2s, one in the 2p.
And we’re going to say here, now the way it looks visually is this. And I’m going to move myself away from shock guide so we have more room to write stuff. So visually this is the way it looks. We have our electrons in the 2s, remember the shape of s is a sphere. We want to promote it to the p, and p looks like a dumbbell. Like that. What’s going to happen like I said is this electron is going to jump over here, and what that really means is that this p and this s shape are going to mix together, and they are kind of going to spit out a hybrid baby. This hybrid orbital is going to have a little s character, a little bit of p character. And as you can notice, it kind of looks like a dumbbell but one side is a lot bigger than the other side. It’s a hybrid of these two guys here. These are the parents, and this is their offspring. And because we promoted one electron to the two p region, we’re going to say that this becomes an sp hybrid orbital. Because it has one that’s in the s, and one that went to the p. That’s why it’s called sp. But remember. p has how many electron orbitals? p has three electron orbitals. We only had enough electron to place only one of them, so these two over here are not being utilized. We’re going to say those two over here, those two electron orbitals are called unhybridized orbitals.
This is the way bonds form, and if the atomic orbitals are different letters, they cannot mix together. So to get around this, we’re going to mix up those orbitals together, forming hybrid orbitals. Since this now has s and p characteristics, it can form a bond, with Chlorine. Because Chlorine wants something with p characteristics involved. Since it is hybrid, it does have p characteristics involved so that’s why we were able to do this. Now hybridization is a little conceptual, really conceptual, but just remember the basics of this. When it comes to answering questions, which we’re going to do next, it’s really easy and simple to follow, as long as you look at it in a very simple way. So just get down to the fundamentals of this. In order for different atomic orbitals, different letters to connect together, we have to mix and hybridize the atomic orbitals of the central element. And in that way we can form bonds. 

For an s orbital to bond with a p orbital the central element must first undergo hybridization. 

Determining Hybridization

Concept #3: Simplifying Hybridization 

To determine the hybridization of a central element we look at the number of groups around it. 

Example #1: For each of the given covalent compounds draw out the Lewis Structure and answer the questions.  

CH2Cl2                   Hybridization,                                                          

                               Unhybridized Orbitals,                                             

                               Bonding orbitals (C – H).


XeCl5+                   Hybridization,                                                          

                               Unhybridized Orbitals,                                             

                               Bonding orbitals (Xe – Cl).

Practice: For the given covalent compound draw out the Lewis Structure and answer the questions. 

IF2:  Hybridization,

       Unhybridized Orbitals, 

       Bonding orbitals (I-F).

Practice: For the given covalent compound draw out the Lewis Structure and answer the questions. 

CH3+:  Hybridization: 

         Unhybridized Orbitals:

         Bonding orbitals (C - H):