Molecular Orbital MO Theory Simplified for Sigma and Pi Bonds

Summary

The video transcript by Leah4Sci.com explains Molecular Orbital Theory in a simplified manner, focusing on Sigma and Pi Bonds. It emphasizes the difference between atomic and molecular orbitals, and how electrons in bonding molecular orbitals contribute to stability, while those in antibonding orbitals indicate instability. The analogy of relationships is used to illustrate the concepts, making it more relatable. The video also touches on the energy diagrams and the significance of Highest Occupied Molecular Orbital (HOMO) and Lowest Unoccupied Molecular Orbitals (LUMO) in understanding molecular behavior.

Takeaways

• 🌟 Molecular Orbital Theory is a method to understand the electronic structure of molecules, particularly focusing on sigma and pi bonds.
• 📚 In Organic Chemistry, the focus is on simple concepts rather than complicated math and physics behind Molecular Orbital Theory.
• 🔴 Atomic Orbitals describe the regions where electrons are likely to be found around an atom, with different types like sp3, sp2, sp, and p orbitals housing different electrons.
• 🔵 Molecular Orbitals represent the distribution of electrons across the entire molecule, showing how atoms are bound together in a chemical bond.
• ⚡️ The energy levels of Atomic and Molecular Orbitals differ and depend on the specific situation in the molecule.
• 🤝 In H2 gas, the bonding of two hydrogen atoms results in the formation of a low-energy bonding molecular orbital and a high-energy antibonding molecular orbital.
• 🌀 Constructive and destructive interferences lead to the formation of bonding (sigma) and antibonding (sigma*) molecular orbitals, respectively.
• 💑 Electrons in bonding molecular orbitals are stable and unreactive, akin to a happy and stable relationship, while those in antibonding orbitals are high-energy and less stable, like a tumultuous fight.
• 📌 Pi bonds are formed by the overlap of non-hybridized p-orbitals above and below the plane of the molecule, with electrons able to move around the bond with a node in the middle.
• 🔄 Resonance structures in molecules with pi bonds can involve the shifting of electrons between bonding and antibonding molecular orbitals, leading to different stable and unstable contributing structures.
• 🔓 The concepts of Highest Occupied Molecular Orbital (HOMO) and Lowest Unoccupied Molecular Orbitals (LUMO) are introduced as key factors in understanding complex molecular systems.

Q & A

What is Molecular Orbital Theory?

Molecular Orbital Theory is a method used to understand the electronic structure of molecules, particularly in relation to bonding. It involves the interaction and combination of atomic orbitals to form molecular orbitals, which can be bonding or antibonding, and are associated with sigma and pi bonds.

Why is Molecular Orbital Theory important in Organic Chemistry?

Molecular Orbital Theory is important in Organic Chemistry because it provides a simplified way to understand and predict the behavior of molecules, especially in terms of their bonding, stability, and reactivity, without getting into the complex mathematical and physics aspects.

What are the differences between Atomic Orbitals and Molecular Orbitals?

Atomic Orbitals refer to the regions around an atom where electrons are most likely to be found, while Molecular Orbitals represent the distribution of electrons across the entire molecule, showing how the atoms are bound together. Atomic Orbitals are specific to individual atoms, whereas Molecular Orbitals combine the characteristics of the atomic orbitals from the bonded atoms.

How does the formation of a bond between carbon and hydrogen in CH4 illustrate the concept of Molecular Orbitals?

In CH4, the carbon atom is sp3 hybridized and each hydrogen atom has an s-orbital. The bond between carbon and hydrogen is formed by the overlap of the sp3 hybrid orbital of carbon and the s-orbital of hydrogen. This bond formation results in the creation of a molecular orbital that represents the shared electrons between the two atoms, indicating their bonded state.

What are the two possible outcomes when two hydrogen atoms combine to form a molecular orbital?

When two hydrogen atoms combine to form a molecular orbital, there are two possible outcomes: constructive interference leading to a low-energy bonding molecular orbital (sigma bond), and destructive interference resulting in a high-energy antibonding molecular orbital (sigma star).

How can the relationship between people be used as an analogy to explain the concept of Antibonding Molecular Orbitals?

Antibonding Molecular Orbitals can be likened to a couple in a fight. While still in a relationship, they have opposing views and high energy, which is less stable than their single days. This high energy state represents the antibonding molecular orbital where the 'relationship' is strained and the 'atoms' are not holding on to each other as tightly.

What is the significance of the node in the antibonding molecular orbital?

The node in the antibonding molecular orbital represents a region where there is no probability of finding the bonding electrons. It symbolizes a temporary 'rift' or separation in the 'relationship' of the two electrons, indicating a higher energy and less stable state.

How does the concept of pi bonds differ from sigma bonds?

Sigma bonds result from a direct overlap of atomic orbitals, like the overlap between the sp3 hybrid orbital of carbon and the s-orbital of hydrogen in CH4. Pi bonds, on the other hand, are formed by the lateral overlap of unhybridized p-orbitals, which are above and below the plane of the molecule, as seen in molecules like ethylene.

What happens to the energy levels of electrons when they form a pi bonding molecular orbital?

When electrons form a pi bonding molecular orbital, they occupy a lower energy state that is more stable and represents the 'happy' state of the electrons in a bonding relationship. If excited, they can temporarily move to the pi star antibonding molecular orbital, which is a higher energy and less stable state.

What are the Highest Occupied Molecular Orbital (HOMO) and Lowest Unoccupied Molecular Orbital (LUMO), and why are they significant?

The Highest Occupied Molecular Orbital (HOMO) is the molecular orbital with the highest energy that is fully occupied by electrons, while the Lowest Unoccupied Molecular Orbital (LUMO) is the next higher energy orbital that is unoccupied. They are significant because they play crucial roles in determining the reactivity and stability of molecules, as well as their participation in chemical reactions.

How does the concept of resonance help in understanding the behavior of molecules with pi bonds?

Resonance allows molecules with pi bonds to distribute their electrons over multiple possible structures, which can lower the overall energy and increase stability. It represents the 'back and forth' movement of electrons between the bonding and antibonding molecular orbitals, with the electrons preferring to reside in the lower energy, more stable bonding orbital.

Outlines

00:00

📚 Introduction to Molecular Orbital Theory

This paragraph introduces the concept of Molecular Orbital Theory, focusing on Sigma and Pi Bonds. It emphasizes the need to understand the simple takeaways in Organic Chemistry rather than getting lost in complex Math and Physics. The discussion begins with Atomic Orbitals, explaining their role in determining the location of electrons around an atom and how they contribute to the formation of chemical bonds. The example of a CH4 molecule illustrates the bond formation between carbon and hydrogen atoms. The key difference between atomic and molecular orbitals is highlighted, with the atomic orbitals associated with individual atoms and molecular orbitals representing the electrons in the entire molecule. The energy levels of these orbitals vary depending on the situation. The concept is further explained using the H2 gas molecule, where the bonding and antibonding molecular orbitals are introduced, along with the idea of constructive and destructive interference. The paragraph concludes by emphasizing the importance of understanding the basic principles without delving into unnecessary mathematical details.

05:03

💑 The Dynamics of Bonding and Antibonding Orbitals

This paragraph delves into the dynamics of bonding and antibonding molecular orbitals, using the analogy of a relationship between two people to explain the concepts. It describes how two electrons, initially in atomic orbitals, form a bond and occupy the low-energy bonding molecular orbital, resulting in a stable and happy state. However, just like in relationships, conflicts can arise, leading to the electrons temporarily occupying the high-energy antibonding molecular orbital, which represents a state of higher energy and less stability. The paragraph explains that the electrons prefer to return to the comfortable bonding molecular orbital once the source of excitement or conflict is removed. The concept of the antibonding node is introduced, illustrating the temporary separation in the bond despite the atoms still being technically connected. The summary encourages the viewers to focus on the simple understanding of these concepts rather than getting overwhelmed by the mathematical and physical complexities.

10:07

🌟 Understanding Pi Bonds and Molecular Orbitals

This paragraph focuses on the Molecular Orbital Theory related to pi bonds, which are the second bond in a double bond. The explanation begins with the structure of the Ethyne or Ethylene molecule, highlighting the difference between sigma and pi bonds. The sigma bond is formed by the overlap of hybrid orbitals, while the pi bond involves the overlap of non-hybridized p-orbitals above and below the plane of the molecule. The energy diagram for pi bonds is similar to that of hydrogen, with a pi bonding molecular orbital and a pi star antibonding molecular orbital. The paragraph uses the ethylene molecule as an example to explain how the pi bond can be thought of as a resonance structure, with one possible less stable contributing structure where the electrons are not shared equally between the two carbon atoms, leading to one carbon having a negative charge and the other a positive charge. The explanation concludes by introducing the Highest Occupied Molecular Orbital (HOMO) and Lowest Unoccupied Molecular Orbitals (LUMO), which will be discussed in more detail in the next video, accessible through the provided link.

Keywords

💡Molecular Orbital Theory

Molecular Orbital Theory is a method used to understand the electronic structure of molecules, particularly in explaining the formation of sigma and pi bonds. It involves the mathematical combination of atomic orbitals to form molecular orbitals, which represent the probability of finding electrons in different regions around a molecule. In the video, the theory is simplified to help viewers grasp the concept without delving into complex math and physics, focusing instead on the practical implications for understanding organic chemistry.

💡Sigma Bonds

Sigma bonds are the first bonds formed between atoms in a covalent structure, resulting from the head-on overlap of atomic orbitals, such as the sp3 hybrid orbital of carbon and the s-orbital of hydrogen in methane (CH4). These bonds are characterized by their lower energy and stability, and they form the backbone of molecular structures. In the video, the formation of a sigma bond is likened to a happy and stable relationship between two people, emphasizing the low energy and stability associated with such bonds.

💡Pi Bonds

Pi bonds are a type of covalent bond that forms in addition to sigma bonds in double or triple bonded molecules. They occur when the unhybridized p-orbitals of atoms overlap sideways, above and below the plane of the molecule. Pi bonds are higher in energy and less stable than sigma bonds, and they contribute to the molecule's reactivity and physical properties. The video simplifies the concept by explaining pi bonds in the context of ethylene (CH2=CH2), where one of the two bonds between the carbon atoms is a pi bond.

💡Atomic Orbitals

Atomic orbitals are regions around an atom's nucleus where electrons are most likely to be found. They are defined by their shape, size, and energy, and are designated by quantum numbers. The video discusses how atomic orbitals like sp3, sp2, and p orbitals house different electrons in molecules, and how these orbitals hybridize to form molecular orbitals upon bonding.

💡Hybridization

Hybridization is a concept in chemistry that describes the process by which atomic orbitals combine to form new equivalent orbitals, known as hybrid orbitals, which are suitable for bonding with other atoms. This process helps to explain the bonding in molecules and the geometry around the central atom. In the video, hybridization is used to explain how carbon atoms form different types of bonds with other atoms, such as in methane (CH4) where carbon is sp3 hybridized.

💡Constructive Interference

Constructive interference is a phenomenon in wave mechanics where two waves meet and their amplitudes add together to form a wave of greater intensity. In the context of molecular orbital theory, constructive interference occurs when atomic orbitals overlap in a way that the electron probability densities add together, leading to the formation of a lower energy bonding molecular orbital. This results in a stable and strong bond between atoms.

💡Destructive Interference

Destructive interference is the opposite of constructive interference and occurs when two waves meet and their amplitudes subtract from each other, leading to a reduction or cancellation of the wave's intensity. In molecular orbital theory, destructive interference happens when atomic orbitals overlap in such a way that the electron probability densities cancel each other out, resulting in a higher energy antibonding molecular orbital. This leads to a weaker, less stable bond that is more likely to break.

💡Antibonding Molecular Orbital

An antibonding molecular orbital is a higher energy orbital that results from destructive interference between atomic orbitals. Electrons in an antibonding orbital are less stable and more likely to be involved in chemical reactions because the bond is weaker and the atoms are less tightly held together. In the context of the video, antibonding molecular orbitals are represented with an asterisk ("*") and are associated with higher energy and less stability.

💡Linear Combination of Atomic Orbitals (LCAO)

The Linear Combination of Atomic Orbitals (LCAO) is a mathematical method used in quantum chemistry to describe molecular orbitals as linear combinations of atomic orbitals. This approach helps in understanding the formation of molecular orbitals and their energies by combining atomic orbitals in a linear manner. However, the video emphasizes that the focus should be on the practical understanding of molecular orbitals rather than the complex mathematical aspects of LCAO.

💡Highest Occupied Molecular Orbital (HOMO)

The Highest Occupied Molecular Orbital (HOMO) is the molecular orbital that is filled with electrons and has the highest energy level among the occupied orbitals in a molecule. It plays a crucial role in determining the molecule's reactivity, as electrons from the HOMO can be involved in chemical reactions. The concept of HOMO is used to discuss the stability and reactivity of molecules, and it is closely related to the concept of the Lowest Unoccupied Molecular Orbital (LUMO).

💡Resonance

Resonance is a concept in chemistry that describes the distribution of electrons in a molecule over two or more possible structures, each of which can contribute to the actual structure of the molecule. Resonance structures are used to represent the delocalization of electrons, particularly in molecules with multiple bonds or charged species. The video uses the concept of resonance to explain the movement of electrons between bonding and antibonding molecular orbitals, and how this affects the molecule's stability.

Highlights

Molecular Orbital Theory provides a simple take away for understanding sigma and pi bonds in organic chemistry.

Atomic Orbitals define where an electron is located around an atom, with different types like sp3, sp2, sp, and p orbitals.

In CH4, the carbon atom is sp3 hybridized, and the bond with hydrogen is an overlap between the sp3 hybrid orbital of carbon and the s-orbital of hydrogen.

Molecular Orbitals represent the electrons on the entire molecule, as opposed to Atomic Orbitals which refer to a single atom.

The energy difference between Atomic and Molecular Orbitals depends on the specific situation and can result in bonding or antibonding molecular orbitals.

In H2 gas, two hydrogen atoms with lone electrons form a bond resulting in a new molecular orbital, combining their atomic orbitals.

The Linear Combination of Atomic Orbitals (LCAO) is used to mathematically combine atomic orbitals to form molecular orbitals.

Electrons in a bonding molecular orbital are in a stable, low-energy state, whereas those in an antibonding molecular orbital are in a high-energy, unstable state.

The relationship between electrons in bonding and antibonding molecular orbitals can be likened to people in a relationship, with low energy representing happiness and stability.

In a double bond, such as in ethylene (CH2=CH2), there is one sigma bond and one pi bond, with the pi bond involving non-hybridized p-orbitals.

The pi bond is formed by the overlap of p-orbitals above and below the plane of the molecule, with electrons able to move around the bond with a node in the middle.

The energy diagram for pi bonds is similar to that of sigma bonds, with a pi bonding molecular orbital and a pi antibonding molecular orbital.

Resonance structures in molecules can involve the shifting of electrons between bonding and antibonding molecular orbitals, leading to different contributing structures.

The Highest Occupied Molecular Orbital (HOMO) and Lowest Unoccupied Molecular Orbitals (LUMO) are concepts that will be discussed in the next video.

Understanding Molecular Orbital Theory can help students grasp complex concepts in organic chemistry and molecular bonding.

The video provides a non-mathematical, intuitive understanding of molecular orbitals, making it accessible for students without a strong background in physics or math.

Transcripts

Leah here from Leah4sci.com and in this video  we're going to look at Molecular Orbital Theory  

for Sigma and Pi Bonds. If you look  up what is Molecular Orbital Theory,  

you'll find some complicated Math and Physics  explanation of Quantum wave function that overlap  

and all really nice except that in Organic  Chemistry, we're looking to understand the  

simple take away rather than the complicated over  your head don't need to know this information. And  

that's what I wanna look at today. So let's back  up and talk about Atomic Orbitals. Atomic Orbitals  

as you remember tells us where an electron is  located or specifically in a type of orbital  

that an electron is located around an atom. And  if you go back to my Orgo Basics video series,  

you'll see how we have the sp3, sp2, sp and even  p orbitals that house different electrons. If  

you need a refresher, visit the link below or  go to my website, leah4sci.com/OrgoBasics.  

For example, in a molecule CH4, we have a  carbon atom bound to 4 hydrogen atoms. The  

carbon is sp3 hybridized and each of the hydrogen  atoms have an s-orbital only so no hybridization.  

The bond between carbon and hydrogen is an  overlap between the sp3 hybrid orbital of carbon  

and the s-orbital of hydrogen. But when that bond  forms we're no longer looking at the individual  

atomic orbital because now, these two orbitals  have fused together to make a molecular orbital  

to show that the atoms are bound together. The  difference between atomic and molecular orbitals  

is that the atomic orbitals refers to just the  atom where the molecular orbital now shows us  

the electrons on the entire molecule. The energy  between an atomic and molecular orbital is very  

different and depends on the specific situation.  So let's take a look at a simple molecule, H2 gas.  

This is made up of 2 hydrogen atoms that each  have a lone electron. If we plot this on an  

energy diagram, then each of the hydrogen atoms  has its energy somewhere in the middle. They're  

not happy but they're not unhappy, this  is how they are and they're okay with it.  

This one electron refers to the atomic orbital for  each hydrogen sitting in the 1S orbital. When they  

come together to form a bond, we no longer have  the atomic orbitals because now the electrons  

are sitting in a new molecular orbital. Another  way to visualize this is the two hydrogen atoms,  

we have the one atomic orbital, the second atomic  orbital. Now they're overlapping and you get one  

giant molecular orbital. When the two atoms  come together to form that molecular orbital,  

what's happening is they get combined  mathematically using the LCAO, the Linear  

Combination of Atomic Orbitals. No, no, no! This  is too much mathematical mambo jumbo that we don't  

care about. What we do care about, the takeaway  is that we get two different options. The first  

is constructive interference or should we  say low energy bonding molecular orbital.  

And the second is destructive interference  or should I say High Energy Antibonding  

Molecular Orbital. You'll often see an asterisk  at the Antibonding Molecular Orbital. So where  

does this go on the energy diagram? We have  the low energy bonding molecular orbital  

and we have the high energy antibonding molecular  orbital where the bonding is a sigma bond and the  

antibonding is a sigma star. Remember that we  only had two electrons and that means we have  

the options of putting them into the low energy  bonding molecular orbital where they're happy  

or bump them up into the antibond. But they're  not going to occupy both molecular orbitals at  

the same time. So how does this make sense?  I like to think of the electrons as people  

in or out of the relationship. We start with the  atomic orbital where we have a single electron.  

Here we have a single person  and another single person,  

they don't know each other, they're  just flipping life and happy enough.  

And then one day they meet and fall in love.  Together, they are so happy and so stable even  

more stable than their single days giving them the  lowest possible energy because remember, happy,  

stable, unreactive. And if energy represents anger  and temper, there's not a lot going on there.  

But as with most couples, every now and then they  get into a fight. They're still in a relationship,  

they're still together, but right now, they're  so mad at each other, they have opposing views,  

differing opinions, there's yelling and screaming  and very very high energy, that together,  

in a fight, there's much higher energy meaning  much less stable than their single days.  

As a reminder, the two electrons are sitting in  a bonding molecular orbital, if they get excited  

by something upsetting, these electrons  can temporarily come up to the antibonding  

but if that source of extra energy goes away,  they prefer to sit here in the comfortable,  

happy, stable, bonding molecular orbital. This  temporary rift between them, this fight, is called  

the antibonding node, the thing that separates  them even though they're technically together.  

Turning our stick figure back into hydrogen,  we have almost a single bubble where the  

two electrons exist together in the bonding  molecular orbital and this right here for the  

antibonding molecular orbital where there's  a clear node between them showing that the  

atoms while still together are definitely  not holding on to each other as well.  

This is a topic I use to struggle with and my TA  told me, oh don't worry, this is just mathematical  

physics beyond what you need to understand, which  made sense that I didn't need to understand it but  

I was so confused. So even though I'm telling  you about the same thing, don't worry about the  

math and physics behind all this, I still want  you to have a simple enough takeaway of what  

the heck is going on here. So if you're with me  so far, make sure to give this video a thumbs up  

and let me know what's your biggest  takeaway in the comments below  

and then let move on to pi bonds. Molecular Orbital for pi bonds start  

out very similar to sigma but can get  complicated, so let's take a step back.  

A pi bond or a double bond doesn't refer to two  bonds, it's actually the second in a double bond.  

For example if we look at the molecule Ethyne or  Ethylene which is CH2 CH2, the simple drawing for  

this is a carbon double bound to another carbon  atom and each carbon has two additional carbon  

atom sigma bound to the side. The double  bond between them has two lines that don't  

differentiate so it makes it look like it's two of  the same when in actuality we have one sigma bond  

and one pi bond. Remember that a sigma bond like  we saw at the hydrogen atoms is a simple overlap,  

and don't forget the carbon to hydrogen bonds  are also all sigma. The bond would be shown  

with the Sp2 hybrid from carbon overlapping with  the S non-hybrid from hydrogen to form a sigma  

bond. But since our interest is the pi bond, let's  simplify this as follows. We have carbon single  

bound to carbon and almost imagine that this  is drawn just tilted off the plane on the page  

so that for each carbon we have one hydrogen  coming forward out of the page and one hydrogen  

going back into the page. That lets us put our  eye right here and looks straight on to see  

about the pi bond. The pi bond is made with a non  hybridized P-orbital that individually sets just  

above and below the plane of the molecule, in our  case the plane of the page. When those p-orbitals  

come together and overlap, the pi bonds formed.  Now, let's bring them a bit closer together.  

What this shows us is the electrons are free  to go up and down, around and around between  

the two atoms with a node in the middle so  they're sitting either at the top half or  

bottom half of that pi bond. The energy diagram  for this will look exactly the same as the energy  

diagram for hydrogen. We start out with some  in-between neutral energy for the 2p orbital.  

Each of which has a single electron. The  2p orbital comes from the hybridization  

of carbon which as we remember is 1s², 2s²,  2p². When the 2s and 2p electrons hybridize  

we get three 2 sp², so let me tell you  what the numbers mean. We have three, 1,2,3  

sp2 hybrids in row 2, let's put all the numbers,  so we can take away that too. And then on the side  

we also have that single p but because it's coming  from the second row, it's coming from a 2p, that's  

what this 2p stands for. When the electrons come  together, we get a pi bonding molecular orbital,  

but if they get excited, they jump up into  the pi star antibonding molecular orbital  

which we can imagine something like  this. As you can see, it looks very  

much like our initial ethylene molecule  but with a very obvious node between them  

because right now, the carbons in this  relationship are not getting along very well.  

If you're following up to this point, but you're  trying to think how to make sense of this in the  

grand scheme of molecules and reactions, here is  how I like to think of it. We have our ethylene  

molecule, CH2 CH2 which if we try it out, we  have a pi bond between the two carbon atoms.  

I like to think of the bonding molecular orbital  as a pi bond. Remember on resonance structures  

one type of unstable resonance is to take the two  electrons and collapse it onto one of the atoms.  

Show the double headed arrow and this  new less stable contributing structure.  

We still have the sigma bond between them, but now  we have one lone pair on the carbon on the right,  

that means it's got extra electrons and a negative  charge. The carbon on the left, loss the carbon,  

got nothing in return is now deficient and  therefore gets a positive formal charge.  

Remember, this is not the mathematical  and physics explanation of bonding and  

antibonding but instead, this is how I like  to think of it, in bonding, they're together,  

they're happy. In antibonding, they're  separated, there's the burden of charge,  

one is negative, one is positive, they don't like  that burden, that's what makes them so unstable.  

As students, you take away that extra energy,  the electrons were able to come back, able to  

resonate back to the comfortable structure, back  to the happy low energy bonding molecular orbital.  

And to be honest, once I started thinking of it  this way. that's when the topic finally clicked.  

But this is just one pi bond with two electrons.  What about a more complex system with 4,6,  

or more electrons and lots of resonance involved?  Remembering that not all of the molecular orbitals  

will be occupied brings us to the concept of  Highest Occupied Molecular Orbital or HOMO  

and Lowest Unoccupied Molecular Orbitals or  LOMO. And this is exactly what we'll discuss  

in the next video which you can find  on my website by clicking the link  

below or going to leah4sci.com/MOtheory  . The link again, leah4sci.com/MOtheory.