Question about bond energy and strength.

[RainDog102]

If shorter bonds (sp2-sp2) require more energy to break, then why are they lower in energy than (sp3-sp3 bonds)? I know that the shorter the bond, the more stable it is and thus the lower energy it is but what confuses me is that if something is low energy (stable), doesn't that mean it requires less energy to break it than something with high energy?

---

Members don't see this ad.

 

[RainDog102]

Or am I getting the terms Bond Energy and Bond Dissociation energy confused? So bonds that are low in energy have higher bond dissociation energy bc they are more stable and harder to break?

 

[aldol16]

BDE and bond energy are basically equivalent. What you're confusing here is which bond you're talking about. sp2-sp2 bonds are actually two bonds - one sigma and one pi. sp3-sp3 bonds only have one type of bond - sigma. So one sigma + one pi is harder to break than one sigma for the simple fact that there are two bonds to break instead of one. But once you break the first pi bond, you basically have the same situation as the sp3-sp3 case.

 

[NextStepTutor_1]

Hi @billy201x -

Excellent question! I think it may help to step back and clarify the terminology a little bit. "High-energy" and "low-energy" can be confusing terms sometimes because we have to think about "high" and "low" relative to something. In general, thermodynamics tells us that energy likes to be distributed throughout the environment, so something that "contains" a high amount of energy is going to be relatively unstable. You can use an analogy with gravitational potential energy here -- you certainly can lift a heavy object above your head, giving it high gravitational PE, but that's not the most stable situation. Therefore, reduced energy (relative to the surroundings) = more stable. Based on your initial post, it seems like you have that intuition down, but it never hurts to review it.

The details of the underlying physics go beyond the MCAT, but covalent bonds form in the first place because it is energetically favorable for them to do so. Thus, a molecule of methane (CH4) is lower-energy and more stable than its components in isolation. Therefore, if you want to break those bonds, you will have to invest energy. The more stable the bond, the more energy you will have to invest to break it. This is usually referred to as the bond dissociation energy, and if you just see the term "bond energy", that is often what they're referring to. So lower-energy bonds have less energy than their non-bonded components, and are more stable, so more energy has to be put in to break the bonds.

Another way of thinking about this is that "low-energy" bonds refers to the energy that is tied up, so to speak, in the bond. To break these bonds, you have to invest a lot of energy from the outside, which is the bond dissociation energy. The analogy with gravitational potential energy might help here as well. Imagine that we have five books on a shelf, and one falls to the floor. The book that fell is now "low energy" compared to the books on the shelf, and we'd have to invest energy to get it back up there. The further it falls, the "lower energy" it becomes, and the more energy we have to use to get it back up. For molecules, when bonds are formed, that's like the book that fell to the floor, and breaking the bonds is like putting the book back on the shelf.

Hope this is helpful, & best of luck studying!

 

[aldol16]

Excellent question! I think it may help to step back and clarify the terminology a little bit. "High-energy" and "low-energy" can be confusing terms sometimes because we have to think about "high" and "low" relative to something. In general, thermodynamics tells us that energy likes to be distributed throughout the environment, so something that "contains" a high amount of energy is going to be relatively unstable. You can use an analogy with gravitational potential energy here -- you certainly can lift a heavy object above your head, giving it high gravitational PE, but that's not the most stable situation. Therefore, reduced energy (relative to the surroundings) = more stable. Based on your initial post, it seems like you have that intuition down, but it never hurts to review it.

The details of the underlying physics go beyond the MCAT, but covalent bonds form in the first place because it is energetically favorable for them to do so. Thus, a molecule of methane (CH4) is lower-energy and more stable than its components in isolation. Therefore, if you want to break those bonds, you will have to invest energy. The more stable the bond, the more energy you will have to invest to break it. This is usually referred to as the bond dissociation energy, and if you just see the term "bond energy", that is often what they're referring to. So lower-energy bonds have less energy than their non-bonded components, and are more stable, so more energy has to be put in to break the bonds.

Another way of thinking about this is that "low-energy" bonds refers to the energy that is tied up, so to speak, in the bond. To break these bonds, you have to invest a lot of energy from the outside, which is the bond dissociation energy. The analogy with gravitational potential energy might help here as well. Imagine that we have five books on a shelf, and one falls to the floor. The book that fell is now "low energy" compared to the books on the shelf, and we'd have to invest energy to get it back up there. The further it falls, the "lower energy" it becomes, and the more energy we have to use to get it back up. For molecules, when bonds are formed, that's like the book that fell to the floor, and breaking the bonds is like putting the book back on the shelf.

Chemists do not use these definitions for bond energies versus bond dissociation energies. These terms were invented by people who were not chemists and their continued use has fed into the pre-med confusion about the concept of bond energy. Bond energy is actually quite similar to BDE. Here's the IUPAC definition: IUPAC Gold Book - bond energy (mean bond energy). In other words, BDE refers to a specific bond at a specific location in a molecule whereas bond energy refers to the average of all such bonds in a molecule. Very similar concepts and for pre-med purposes, equivalent as in this case.

 

[aldol16]

Thus, a molecule of methane (CH4) is lower-energy and more stable than its components in isolation. Therefore, if you want to break those bonds, you will have to invest energy. The more stable the bond, the more energy you will have to invest to break it.

Another way of thinking about this is that "low-energy" bonds refers to the energy that is tied up, so to speak, in the bond. To break these bonds, you have to invest a lot of energy from the outside, which is the bond dissociation energy.

It's also more accurate to say that energy is released when methane is formed from its components in solution. Chemistry operates on relative energies, not absolute energies, so the energy of methane's components in isolation is meaningless and if it's used as one end of a comparison, students can get easily confused. So it's clearer to say that in the process: methane components ---> methane, energy is released. In order to go in the reverse direction, i.e. bond breaking, energy must therefore be put into the system again. The more stable methane is relative to its components, the more energy is released and therefore the more energy must be put in to break those bonds again. In other words, it's not the energy tied up in methane's bonds that make it stable per se but rather the energy that must be put in to the system to break those bonds.

 

[NextStepTutor_1]

Hi @aldol16 -

Yes, I understand what you're saying, and thanks for the input. Unfortunately, I think this point of confusion is deeply embedded in the language that is often used to talk about energy in general and bond energy in particular -- when someone hears the words "high-energy bond" or "low-energy bond" it is very easy to think that the bond "has" a certain amount of energy, even though that's misleading at best. However, in my experience, it can be helpful to address such misconceptions by initially working within the framework of that language (which is why I try to be careful by using hedges like "so to speak" when using metaphors like energy being "tied up" in a bond) and using it as a jumping off point to thinking about relative energy & energy flow -- including by exploring analogies with more intuitively familiar forms of energy like gravitational PE.

But yes, the takeaway point is of course that stability should be thought of in terms of the energy that needs to be added to the system to break bonds.

 

[RainDog102]

Hey @NextStepTutor_1 and @aldol16 thank you for your wonderful explanations!