I don't understand the difference between exothermic and endothermic :(

[BH 90210]

Can someone explain this -

Endothermic - bonds broken during the reaction are of higher energy than the bonds formed

On the reaction diagram, aren't the reactants at a lower energy?

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[BH 90210]

bump

 

[mcloaf]

In an endothermic reaction the products are higher energy than the reactants. For the reaction to proceed you need to input thermal energy, which will be taken from the surrounding environment. Thus an endothermic reaction in a beaker would feel cold to the touch as the thermal energy that was in the room temp beaker is being converted into chemical bond energy in the higher energy product(s).

An exothermic reaction is the opposite. The products are lower energy than the reactants, so the excess chemical bond energy released by the reaction is given off as thermal energy. An exothermic reaction heats up its surrounding environment as it releases this thermal energy and will feel warm to the touch.

 

[Endure]

Can someone explain this -

Endothermic - bonds broken during the reaction are of higher energy than the bonds formed

On the reaction diagram, aren't the reactants at a lower energy?

Your book's correct, though I hate the phrasing.

What they're referring to is the strength of the bond being formed, not necessarily the thermodynamic state. The reactants in an endothermic reaction are thermodynamically favored in that they are stronger and more stable, whereas the products are (generally) less stable and in a state of greater potential energy. Consider ATP. Many textbooks mistakenly suggest that the "high-energy" phosphate groups make up a strong bond (phosphoanhydride bond), when in reality what they're referring to is a thermodynamically unstable state.


If that doesn't help, think about it intuitively. The reactants are more stable (happy), so more energy (the activation energy) is needed to drive the reaction forward.

 

[MD Odyssey]

Your book's correct, though I hate the phrasing.

What they're referring to is the strength of the bond being formed, not necessarily the thermodynamic state. The reactants in an endothermic reaction are thermodynamically favored in that they are stronger and more stable, whereas the products are (generally) less stable and in a state of greater potential energy. Consider ATP. Many textbooks mistakenly suggest that the "high-energy" phosphate groups make up a strong bond (phosphoanhydride bond), when in reality what they're referring to is a thermodynamically unstable state.


If that doesn't help, think about it intuitively. The reactants are more stable (happy), so more energy (the activation energy) is needed to drive the reaction forward.

This is wrong. The responder is confusing free energy change with enthalpy change.

Exothermic reactions have a negative enthalpy change and heat is transferred to the surroundings. Endothermic reactions have a positive enthalpy change and heat is transferred from the surroundings to the reactants. That's all there is to it.

The broader discussion, which you should definitely make yourself familiar with is the relationship between the enthalpy change, which for the MCAT you can almost always think of as heat, and the free energy change. The free energy change is what determines whether or not a reaction will be spontaneous or non-spontaneous for a given set of conditions. In particular, you should understand how the entropy change for a given reaction can influence the free energy. This is a key idea and something that will absolutely be on both science sections of the MCAT.

 

[Endure]

This is wrong. The responder is confusing free energy change with enthalpy change.

Just for my own clarification, how does my description fail to explain the definition OP was given? I admit that I spoke generally in that I didn't mention free energy change and its dependance on entropy, but it is absolutely true that endothermic reactions yield high-energy products – which is essentially what was being asked.

 

[MD Odyssey]

Just for my own clarification, how does my description fail to explain the definition OP was given? I admit that I spoke generally in that I didn't mention free energy change and its dependance on entropy, but it is absolutely true that endothermic reactions yield high-energy products – which is essentially what was being asked.

Speaking about 'high energy' vs 'low energy' products is really unwise, as you implied in your original post. It makes more sense to discuss the relative stability of the reactants and products, which shows up in the relative concentrations of both at equilibrium. Any discussion of stability is a discussion of the free energy for the process, not whether or not it is exothermic or endothermic. There are plenty of endothermic processes which are spontaneous and yield more stable products than their reactants. An example of this is the reaction of solid ammonium nitrate to form gaseous ammonia and water vapor.

NH4NO3 + Δ→ N2O + 2H2O​

This reaction is endothermic, but forms very energetically stable products because of the huge gain in entropy for the process. This is pretty straightforward stuff and you seem to be getting confused by the difference between heat transferred, which is sort of what enthalpy is, and energetics. It is very improper to discuss enthalpy leading to stability, which is a common misunderstanding. I guarantee that the MCAT will test your understanding of these two topics and will lay traps like the one I showed up above for you to make. You should really nail this topic down before your exam.

 

[Endure]

Speaking about 'high energy' vs 'low energy' products is really unwise, as you implied in your original post. It makes more sense to discuss the relative stability of the reactants and products, which shows up in the relative concentrations of both at equilibrium. Any discussion of stability is a discussion of the free energy for the process, not whether or not it is exothermic or endothermic. There are plenty of endothermic processes which are spontaneous and yield more stable products than their reactants. An example of this is the reaction of solid ammonium nitrate to form gaseous ammonia and water vapor.

NH4NO3 + Δ→ N2O + 2H2O​

This reaction is endothermic, but forms very energetically stable products because of the huge gain in entropy for the process. This is pretty straightforward stuff and you seem to be getting confused by the difference between heat transferred, which is sort of what enthalpy is, and energetics. It is very improper to discuss enthalpy leading to stability, which is a common misunderstanding. I guarantee that the MCAT will test your understanding of these two topics and will lay traps like the one I showed up above for you to make. You should really nail this topic down before your exam.

Sorry for not getting back to you! Insane exam schedule this past week.

What I was referring to, and feel free to correct me if I'm wrong, was the bond dissociation energy of the products in an endothermic reaction, which is what I believe OP's book was describing. What you're insisting I don't understand is a reaction's free energy, which I would argue is beyond the scope of the question. Yes, of course, plenty of endothermic reactions are rendered energetically favorable because of an increase in the thermal energy dispersed (entropy). This is precisely why reactions that are endothermic can be simultaneously exergonic. However, generally speaking, endothermic reactions result in bonds of weaker dissociation energy that are often described in a state of "high-energy" because of their instability.

In short, the answer to the question is this: endothermic reactions generally result in weaker bonds, which explains the large activation energy, their position on an energy diagram, and the resultant bond dissociation energies.

EDIT: Anything that I've spoken about thus far can be cross-checked with TBR Chem Part II (p. 146) and EK.