Bond Length and Bond Energy

Summary

In this chemistry essentials video, Mr. Andersen explores bond length and bond energy, starting with Kathleen Lonsdale's pioneering work on hexamethylbenzene. He explains how electron-proton attraction and atomic repulsion determine bond strength and length, with bond energy being the energy required to break bonds. As bond numbers increase from single to double to triple, bond strength and energy rise, while bond length decreases. The video uses examples like boron, rhenium, ethane, and acetylene to illustrate these concepts, aiming to help viewers understand the fundamental energies in chemical bonding.

Takeaways

• 🔬 Kathleen Lonsdale pioneered the measurement of bond lengths using x-ray crystallography, contributing significantly to the understanding of molecular structures like hexamethylbenzene.
• 📏 Bond length is the distance between two atoms within a molecule, and it can be determined through techniques such as x-ray diffraction.
• 🔗 The attractive force between atoms is due to the electrons of one atom being attracted to the protons of another, while too close proximity leads to repulsion.
• 📊 An energy-distance graph illustrates the relationship between the energy required to separate atoms and the physical distance between them, helping to define bond energy and length.
• ⚡ Bond energy is defined as the energy needed to break the bond between two atoms, with bond formation releasing an equivalent amount of energy.
• 🔋 The strength of a bond is influenced by the charges of the atoms involved, with larger charges leading to greater bond energy.
• 🔄 As the number of electrons in a bond increases (from single to double to triple bond), the bond strength increases, pulling atoms closer and decreasing the bond length.
• 📍 The optimal bond length, akin to Goldilocks' 'just right' point, balances the attractive and repulsive forces between atoms for maximum stability.
• 🚀 Bond energy is influenced by the electronegativity of the atoms in a molecule, with higher electronegativity contributing to stronger bonds.
• 🌌 Atomic radius can be used to estimate bond length, but actual measurements may vary due to the specific interactions within a molecule.
• 🔑 Increasing the bond number, as seen in the transition from ethane to acetylene, results in shorter bond lengths and higher bond energies due to greater electron sharing.

Q & A

Who was the first scientist to measure bond lengths using x-ray crystallography?

Kathleen Lonsdale was the first scientist to measure bond lengths using x-ray crystallography.

What did Kathleen Lonsdale's work with x-rays reveal about the structure of hexamethylbenzene?

Kathleen Lonsdale's work revealed that hexamethylbenzene had a flat hexagonal structure and allowed her to measure the aromatic bonds between the carbons.

What is the source of the attraction between two atoms forming a bond?

The attraction between two atoms forming a bond comes from the electrons of one atom being attracted to the protons of another atom.

What is the definition of bond energy?

Bond energy is the energy required to break the bond between two atoms.

How is bond energy related to the formation of a bond?

When a bond is formed between two atoms, energy is released, which is the negative of the bond energy required to break the bond.

What factor contributes to the strength of a chemical bond?

The strength of a chemical bond is influenced by the charges of the atoms involved and the number of electrons in the bond, which increases from single to double to triple bonds.

Why does increasing the bond number decrease the bond length?

Increasing the bond number, such as from a single to a triple bond, increases the bond strength, which pulls the atoms closer together, thus decreasing the bond length.

What is the significance of the 'Goldilocks' point in the context of bond formation?

The 'Goldilocks' point refers to the optimal distance between two atoms where the attractive and repulsive forces balance, resulting in the highest energy and most stable bond.

How does the atomic radius of an element relate to the expected bond length between two atoms of that element?

The expected bond length between two atoms of the same element is typically twice the atomic radius of the element, assuming no other atoms are involved.

What can be inferred about the bond strength if the measured bond length in a molecule is less than the sum of the atomic radii of the two atoms?

If the measured bond length is less than the sum of the atomic radii, it indicates that the atoms are closer together than expected, suggesting a strong bond due to greater charge overlap.

How does the presence of a triple bond between two carbon atoms in acetylene compare to a single bond in ethane in terms of bond strength and bond length?

In acetylene, the presence of a triple bond results in a stronger bond and a shorter bond length compared to the single bond in ethane, due to the increased number of shared electrons.

Outlines

00:00

🔬 Bond Length and Energy Fundamentals

This paragraph introduces the concept of bond length and energy, highlighting the pioneering work of Kathleen Lonsdale, an x-ray crystallographer who first measured bond lengths using x-rays. Lonsdale's work on hexamethylbenzene led to the understanding of resonance in aromatic compounds. The paragraph explains the forces of attraction and repulsion between atoms and how these can be visualized and measured using an energy-distance graph. It defines bond energy as the energy required to separate bonded atoms and notes that bond strength is influenced by the charges of the atoms and the number of electrons shared, leading to a decrease in bond length as bond strength increases.

05:03

📏 The Relationship Between Bond Number and Length

This paragraph delves into the relationship between the number of bonds between atoms and the resulting bond strength and length. It explains that as the number of bonds increases, from single to double to triple, the bond length decreases and the bond energy increases due to greater electron sharing and charge. The paragraph uses the example of ethane and acetylene to illustrate how a triple bond results in a shorter bond length and higher bond energy compared to a single bond. It also discusses how atomic radii can be used to predict bond lengths and the implications of observed bond lengths on the strength of the bond, using the examples of boron and rhenium atoms.

Keywords

💡Bond Length

Bond length refers to the distance between the nuclei of two atoms that are bonded together. In the video, it is discussed in the context of how bond length can be measured using x-ray crystallography, as demonstrated by Kathleen Lonsdale with hexamethylbenzene. The script explains that as bond strength increases, bond length decreases, which is exemplified by comparing single, double, and triple bonds.

💡Bond Energy

Bond energy is the amount of energy required to break a bond between two atoms. The video describes it as being measured on an energy-distance graph, where the energy holding atoms together is plotted against the physical separation between them. The script also mentions that the formation of a bond releases energy, which is the negative of the bond energy.

💡X-ray Crystallography

X-ray crystallography is a technique used to determine the atomic and molecular structure of a crystal, which was pivotal in measuring bond lengths. Kathleen Lonsdale's work with x-ray crystallography is highlighted in the script as a significant contribution to the understanding of molecular structures, such as the structure of hexamethylbenzene.

💡Resonance

Resonance is a concept in chemistry that describes the phenomenon where a molecule can be represented by two or more Lewis structures that are in equilibrium. The script mentions that Lonsdale's measurements of bond lengths in benzene led to the development of the idea of resonance, which explains the delocalized electrons in the benzene ring.

💡Electronegativity

Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. The video script explains that electronegativity contributes to bond energy, with the atoms involved in a molecule having an impact on the strength of the bond formed between them.

💡Atomic Radius

Atomic radius is defined as the distance from the nucleus to the outermost shell of an atom. The script uses the example of boron and rhenium to illustrate how atomic radius can be used to predict bond length, with the actual bond length in a molecule sometimes differing from the expected distance based on atomic radii.

💡Goldilocks Area

The term 'Goldilocks Area' in the script metaphorically refers to the optimal distance between two atoms where the attractive forces are balanced with repulsive forces, resulting in the highest energy and the most stable bond. This concept is used to describe the equilibrium point in an energy-distance graph.

💡Attractive Force

Attractive force is the force that pulls atoms together, as described in the script when discussing how electrons of one atom are attracted to protons of another. The script explains that as atoms get closer, the attractive force increases until a certain point, after which repulsion sets in.

💡Repulsion

Repulsion is the force that pushes atoms apart when they are too close to each other. In the script, it is mentioned in the context of atoms getting too close and the resulting force that counteracts the attractive force, leading to the concept of the optimal bond distance.

💡Triple Bond

A triple bond is a type of chemical bond where three pairs of valence electrons are shared between two atoms. The script contrasts ethane, which has a single bond between carbon atoms, with acetylene, which has a triple bond, to illustrate how increasing the number of bonds decreases bond length and increases bond energy.

💡Nitrogen

Nitrogen is an element that is used in the script to demonstrate the relationship between bond type and bond length. The script explains that as you move from a single to a double to a triple bond in nitrogen, the bond length decreases and the bond energy increases.

Highlights

Kathleen Lonsdale pioneered the measurement of bond lengths using x-ray crystallography, contributing to the understanding of molecular structures.

Lonsdale's work on hexamethylbenzene led to the concept of resonance in aromatic compounds.

Bond length and bond energy are fundamental concepts in chemistry, describing the distance and energy involved in atomic interactions.

The bond energy is defined as the energy required to break atoms apart and is measured using an energy-distance graph.

Bond strength is influenced by the charges of the atoms and increases with the number of shared electrons.

As bond strength increases, atoms are pulled closer together, resulting in a decrease in bond length.

The concept of a 'Goldilocks' zone in chemistry refers to the optimal distance for maximum bond energy.

Electronegativity of atoms contributes to bond energy, with more electronegative atoms forming stronger bonds.

Atomic radius is a key factor in predicting bond length, with larger atoms typically having longer bond lengths.

Actual bond lengths can differ from expected values based on atomic radii, indicating bond strength variations.

Increasing the number of bonds between atoms, such as from single to triple, decreases bond length and increases bond energy.

The comparison between ethane and acetylene illustrates the effect of bond type on bond strength and length.

Nitrogen molecules with varying bond types (single, double, triple) demonstrate a consistent decrease in bond length with increasing bond number.

The video aims to educate viewers on describing the energies involved in both breaking and forming chemical bonds.

Lonsdale's achievements in chemistry and her role in breaking barriers as a female scientist are highlighted.

The video concludes with a hope that viewers have gained a better understanding of bond length and energy.

Transcripts

Hi. It's Mr. Andersen and this is chemistry essentials video 52. It's on bond length and

bond energy. Some of the first bond lengths ever measured were measured by scientist Kathleen

Lonsdale who was an x-ray crystallographer. She would put a solid crystal in front of

x-rays. Measure the diffraction. It told her a lot about the structure of that molecule

and she could even measure the distance between atoms. And that allowed her to unlock the

structure of hexamethylbenzene, that famous benzene ring. She was able to determine that

it was a hexagon. That it was flat. And she even measured the aromatic bonds. In other

words the single, double, single, double, single double bonds between the carbons. And

by measuring the length of that, that eventually led to this whole idea of resonance. So not

only was she an amazing chemist, but she broke down a lot of walls. She was one of the first

two female fellows elected to the royal society. First tenured professor. She was an amazing

scientist. But she really got us down this road of bond length and bond energy. And so

if we have two atoms, they're both going to be attracted to one another. And where is

that attraction come from? It's the electrons of one atom being attracted to the protons

of another. And so that's going to pull atoms together. But if they get too close to one

another, then there's going to be repulsion. And we can measure these forces just using

an energy distance graph where we've got the energy that's pulling those or pushing those

apart and then how far they are physically separated. And using a graph like that we're

able to determine the bond energy, the energy holding them together. And also measure the

bond length. How far those atoms are going to be apart. And so the bond energy, if we

were to define it, is simply the energy required to break those atoms apart. Now likewise when

we form a bond between the two we're going to release energy. And that's just simply

the negative bond energy. Now the strength of that bond, or the bond strength is going

to be built upon the charges of the atoms themselves. The bigger the charges are the

greater that bond energy is going to be. And as we increase the number of electrons in

a bond, so as we move from a single to a double to a triple bond, we're actually increasing

the charges. And therefore increasing that bond strength. And one interesting thing that

you should know is that as we increase that strength we're going to pull those atoms together

and we're going to actually decrease the bond length. And so if we look at these two representative

atoms right here, as we bring them close together there's going to be attraction. In other words

as we shorten the distance between the atoms, as they get closer and closer together, there's

going to be a greater attractive force. Now as they start to get really close together,

there's actually going to be repulsion between those two atoms. It's going to push them apart.

And so we can find what's called that Goldilock's. That point where they're not too close, not

too far away and we're going to have the highest amount of energy right here. And they're just

going to kind of vibrate at that space. So if we were to look at two atoms, like two

atoms of hydrogen we would find that that energy to distance graph would look something

like this. There's going to be attractive force between these two hydrogen atoms out

here, but as we bring them closer and closer together, that energy of attraction is going

to get greater. As we move them even closer then we're going to move into this whole idea

of repulsion. And so this is going to be that Goldilock's area right here. And so what we

can measure then is the bond energy. The energy holding those atoms together. And likewise,

since we've got distance here on the x-axis, we can measure the bond length. How far those

atoms are apart. And so if we define what is bond energy it's just two atoms connected

together. And so bond energy is the energy absorbed when we break those atoms apart.

Now likewise there's going to be energy that's released as we bond those together. And that's

going to be the negative bond energy. It's going to be exactly the same. The amount of

energy to break it apart is the same amount of energy that we get when those two atoms

are going to be attracted together. Now what contributes to that bond energy? It's going

to be the electronegativity of all the atoms involved in that molecule. And so let's say

we were to look at boron. Boron is going to have an atomic radius of 83.0 picometers.

What does that mean? The distance from the center of that nucleus to the outside is going

to be 83.0 picometers. And so if we were to connect a boron to a boron, we would anticipate

that that bond length, the distance between the two atoms is going to be twice that. It's

going to be 166 picometers. But if we were to measure the actual bond length in a molecule,

let's throw one up here, and it ends up being 175 picometers. In other words they're farther

apart then we would expect. That means that we're going to have a very weak bond between

these two atoms. And so basically as we increase that bond length, that distance between the

atoms, we're decreasing the bond strength. Now if we were to use another example, let's

say we're looking a rhenium which has a radius of 137.5. If we had two rhenium atoms we would

expect that that bond length is going to be 275 picometers. If we actually measure it

in a molecule like this and it's less than that, that means that they're really overlapping.

There's greater charges holding it together and so we would call that a strong bond. Again,

the higher your bond energy is the closer those atoms are going to be and the shorter

that bond length is going to be. Now what happens as we increase the bond number? Let's

say we're looking at two molecules that look essentially the same. We've got ethane which

is carbon attached to carbon and then hydrogen around the outside and then acetylene. And

so this is the same thing. Carbon attached to carbon with hydrogen around the outside.

But in this case we've got a triple bond right here. How is that going to affect the bond

strength and the bond length? And so if we look at these two carbons, we've got one that's

a single bond and one that's a triple bond. If we were to measure their bond length we

would find that as we increase the number of bonds we're actually decreasing that bond

length. They're getting closer and closer together. Well what does that tell us about

the bond energy? The bond energy is increasing over time. So again, the more electrons that

we're sharing, the more charges there are, the greater that bond energy and therefore

the shorter the bond length is. If we were to measure the same things with nitrogen,

single, double, triple bond, we'd find that we're decreasing the bond length again. So

what's going to happen to our bond energy over here? It's going to increase over time.

And so did you learn to describe the energies involved in both breaking and forming chemical

bonds? I hope so. And I hope that was helpful.