free energy
Do all exothermic and endothermic reactions result in a greater stability or is it just spontaneous processes? Also how does it become more stable?
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Exothermic reaction increase stability, because energy is released into the surroundings. Endothermic reactions also increase stability, however this is because entropy and not energy. This is for spontaneous reactions.
So I don't think ALL exo and endo reaction do, however I am not 100% certain tbh
This probably didn't help what so ever
So I don't think ALL exo and endo reaction do, however I am not 100% certain tbh
This probably didn't help what so ever
Original post by DudeBoy
Exothermic reaction increase stability, because energy is released into the surroundings. Endothermic reactions also increase stability, however this is because entropy and not energy. This is for spontaneous reactions.
So I don't think ALL exo and endo reaction do, however I am not 100% certain tbh
This probably didn't help what so ever
So I don't think ALL exo and endo reaction do, however I am not 100% certain tbh
This probably didn't help what so ever
So this is just for spontaneous reactions? But I thought endothermic reactions would cause more 'disorder' so wouldn't it be less stable?
Original post by tazmaniac97
So this is just for spontaneous reactions? But I thought endothermic reactions would cause more 'disorder' so wouldn't it be less stable?
Increased disorder means MORE stability.
The universe tends towards disorder (entropy) as this is the most probable state.
This disorder is a function of both the number of particles and the available energy that can be spread on the particles.
Exothermic reactions spread more energy around increasing disorder, while endothermic reactions actually increase the total number of particles on which the energy can be spread.
Gibbs Free energy relates the enthalpy of the system and the entropy of the system to the overall entropy of the universe. If Gibbs free energy change is negative then the overall entropy of the universe is increasing.
Original post by charco
Increased disorder means MORE stability.
The universe tends towards disorder (entropy) as this is the most probable state.
This disorder is a function of both the number of particles and the available energy that can be spread on the particles.
Exothermic reactions spread more energy around increasing disorder, while endothermic reactions actually increase the total number of particles on which the energy can be spread.
Gibbs Free energy relates the enthalpy of the system and the entropy of the system to the overall entropy of the universe. If Gibbs free energy change is negative then the overall entropy of the universe is increasing.
The universe tends towards disorder (entropy) as this is the most probable state.
This disorder is a function of both the number of particles and the available energy that can be spread on the particles.
Exothermic reactions spread more energy around increasing disorder, while endothermic reactions actually increase the total number of particles on which the energy can be spread.
Gibbs Free energy relates the enthalpy of the system and the entropy of the system to the overall entropy of the universe. If Gibbs free energy change is negative then the overall entropy of the universe is increasing.
Exothermic reactions cause more disorder? But I thought less moles on the right hand side meant there was more order?
Also, why do exothermic reactions give out energy as heat? I tried doing some research and apparently it's because energy is given out because the electrons are at a lower energy level. But I don't know what they mean
Original post by tazmaniac97
Exothermic reactions cause more disorder? But I thought less moles on the right hand side meant there was more order?
Also, why do exothermic reactions give out energy as heat? I tried doing some research and apparently it's because energy is given out because the electrons are at a lower energy level. But I don't know what they mean
Also, why do exothermic reactions give out energy as heat? I tried doing some research and apparently it's because energy is given out because the electrons are at a lower energy level. But I don't know what they mean
You must differentiate between the system under study and the surroundings.
Remember that system + surroundings = universe
Gibbs free energy treats things from the point of view of the system under study, and doing so relates the system to the universe.
In terms of the entropy of the SYSTEM the moles of gas give you a big clue as to how it changes. This is where more moles of gas on the RHS side means an increase in entropy.
Exothermic reactions send energy to the SURROUNDINGS, this does not change the entropy of the SYSTEM, but it increases the entropy of the SURROUNDINGS. This is why exothermic reactions are usually favourable.
If questions ask you about entropy of a reaction you can forget the surroundings and focus only on the reaction (the system). This is where you look for moles of gas change...
The most important thing is:
The Gibbs free energy (G) of the system decreases in a spontaneous process.
Endothermic and exothermic usually refer to the enthalpy change of the reaction, but this is not the only factor in G. G is also dependant upon entropy, which becomes more important at higher temperatures. Of course the 2nd law applies, so the entropy of the universe must increase, but that does not mean a closed system cannot decrease in entropy, only isolated systems cannot.
The Gibbs free energy (G) of the system decreases in a spontaneous process.
Endothermic and exothermic usually refer to the enthalpy change of the reaction, but this is not the only factor in G. G is also dependant upon entropy, which becomes more important at higher temperatures. Of course the 2nd law applies, so the entropy of the universe must increase, but that does not mean a closed system cannot decrease in entropy, only isolated systems cannot.
Original post by charco
You must differentiate between the system under study and the surroundings.
Remember that system + surroundings = universe
Gibbs free energy treats things from the point of view of the system under study, and doing so relates the system to the universe.
In terms of the entropy of the SYSTEM the moles of gas give you a big clue as to how it changes. This is where more moles of gas on the RHS side means an increase in entropy.
Exothermic reactions send energy to the SURROUNDINGS, this does not change the entropy of the SYSTEM, but it increases the entropy of the SURROUNDINGS. This is why exothermic reactions are usually favourable.
If questions ask you about entropy of a reaction you can forget the surroundings and focus only on the reaction (the system). This is where you look for moles of gas change...
Remember that system + surroundings = universe
Gibbs free energy treats things from the point of view of the system under study, and doing so relates the system to the universe.
In terms of the entropy of the SYSTEM the moles of gas give you a big clue as to how it changes. This is where more moles of gas on the RHS side means an increase in entropy.
Exothermic reactions send energy to the SURROUNDINGS, this does not change the entropy of the SYSTEM, but it increases the entropy of the SURROUNDINGS. This is why exothermic reactions are usually favourable.
If questions ask you about entropy of a reaction you can forget the surroundings and focus only on the reaction (the system). This is where you look for moles of gas change...
Oh I get it, what about exothermic reactions? Why do they give out heat energy and what have electrons got to do with it?
Original post by tazmaniac97
Oh I get it, what about exothermic reactions? Why do they give out heat energy and what have electrons got to do with it?
The energy is given out in exothermic reactions as the bonds formed by the products are stronger than the bonds broken in the reactants.
Electrons are needed to form the bonds - they 'glue' the positive nuclear centres to one another by electrostatic attraction ...
Original post by charco
You must differentiate between the system under study and the surroundings.
Remember that system + surroundings = universe
Gibbs free energy treats things from the point of view of the system under study, and doing so relates the system to the universe.
In terms of the entropy of the SYSTEM the moles of gas give you a big clue as to how it changes. This is where more moles of gas on the RHS side means an increase in entropy.
Exothermic reactions send energy to the SURROUNDINGS, this does not change the entropy of the SYSTEM, but it increases the entropy of the SURROUNDINGS. This is why exothermic reactions are usually favourable.
If questions ask you about entropy of a reaction you can forget the surroundings and focus only on the reaction (the system). This is where you look for moles of gas change...
Remember that system + surroundings = universe
Gibbs free energy treats things from the point of view of the system under study, and doing so relates the system to the universe.
In terms of the entropy of the SYSTEM the moles of gas give you a big clue as to how it changes. This is where more moles of gas on the RHS side means an increase in entropy.
Exothermic reactions send energy to the SURROUNDINGS, this does not change the entropy of the SYSTEM, but it increases the entropy of the SURROUNDINGS. This is why exothermic reactions are usually favourable.
If questions ask you about entropy of a reaction you can forget the surroundings and focus only on the reaction (the system). This is where you look for moles of gas change...
Strictly, this doesn't seem to be entirely correct (at least to me).
An exothermic reaction is one that lowers the internal energy U of the reagents/system, but this is very difficult to measure. Enthalpy H is therefore used as a measure as it is easier to measure.
The important factor to weather a reaction is spontaneous is the free energy (gibbs G under constant pressure, or helmholtz A under constant volume)
If the free energy is lowered then the reaction is spontaneous.
At absolute zero, the gibbs free energy is equal to the enthalpy change as:
DeltaG=DeltaH-TDeltaS
So therefore at T>0 we must consider the entropy. An exothermic reaction may not be spontaneous at room T because dG is positive due to a disfavorable dS
(I'm sure you know this, it's for the OP)
Original post by JMaydom
Strictly, this doesn't seem to be entirely correct (at least to me).
An exothermic reaction is one that lowers the internal energy U of the reagents/system, but this is very difficult to measure. Enthalpy H is therefore used as a measure as it is easier to measure.
The important factor to weather a reaction is spontaneous is the free energy (gibbs G under constant pressure, or helmholtz A under constant volume)
If the free energy is lowered then the reaction is spontaneous.
At absolute zero, the gibbs free energy is equal to the enthalpy change as:
DeltaG=DeltaH-TDeltaS
So therefore at T>0 we must consider the entropy. An exothermic reaction may not be spontaneous at room T because dG is positive due to a disfavorable dS
(I'm sure you know this, it's for the OP)
An exothermic reaction is one that lowers the internal energy U of the reagents/system, but this is very difficult to measure. Enthalpy H is therefore used as a measure as it is easier to measure.
The important factor to weather a reaction is spontaneous is the free energy (gibbs G under constant pressure, or helmholtz A under constant volume)
If the free energy is lowered then the reaction is spontaneous.
At absolute zero, the gibbs free energy is equal to the enthalpy change as:
DeltaG=DeltaH-TDeltaS
So therefore at T>0 we must consider the entropy. An exothermic reaction may not be spontaneous at room T because dG is positive due to a disfavorable dS
(I'm sure you know this, it's for the OP)
I don't see where you are saying anything different to my post?
Original post by charco
The energy is given out in exothermic reactions as the bonds formed by the products are stronger than the bonds broken in the reactants.
Electrons are needed to form the bonds - they 'glue' the positive nuclear centres to one another by electrostatic attraction ...
Electrons are needed to form the bonds - they 'glue' the positive nuclear centres to one another by electrostatic attraction ...
But how do the electrons go to a lower energy level?
Original post by JMaydom
Strictly, this doesn't seem to be entirely correct (at least to me).
An exothermic reaction is one that lowers the internal energy U of the reagents/system, but this is very difficult to measure. Enthalpy H is therefore used as a measure as it is easier to measure.
The important factor to weather a reaction is spontaneous is the free energy (gibbs G under constant pressure, or helmholtz A under constant volume)
If the free energy is lowered then the reaction is spontaneous.
At absolute zero, the gibbs free energy is equal to the enthalpy change as:
DeltaG=DeltaH-TDeltaS
So therefore at T>0 we must consider the entropy. An exothermic reaction may not be spontaneous at room T because dG is positive due to a disfavorable dS
(I'm sure you know this, it's for the OP)
An exothermic reaction is one that lowers the internal energy U of the reagents/system, but this is very difficult to measure. Enthalpy H is therefore used as a measure as it is easier to measure.
The important factor to weather a reaction is spontaneous is the free energy (gibbs G under constant pressure, or helmholtz A under constant volume)
If the free energy is lowered then the reaction is spontaneous.
At absolute zero, the gibbs free energy is equal to the enthalpy change as:
DeltaG=DeltaH-TDeltaS
So therefore at T>0 we must consider the entropy. An exothermic reaction may not be spontaneous at room T because dG is positive due to a disfavorable dS
(I'm sure you know this, it's for the OP)
How are reactions exothermic in the first place? I mean how do they give out heat? I read that the electrons are a lower energy level, but how does this happen?
Original post by tazmaniac97
How are reactions exothermic in the first place? I mean how do they give out heat? I read that the electrons are a lower energy level, but how does this happen?
The internal enegy of a substance is a function of its inherent nature and its relative position in the universe.
The second one can never be known, but the first one is demonstrated by a lowering of energy when bonds are formed.
You could think of gravity as an analogy. An object in outer space has no apparent potential energy, but if it comes into the sphere of influence of a massive body then it will be attracted to it and accelerate to the surface. In doing so it loses the potential energy caused by its position and transforms it first into kinetic energy and then into heat (deformation)
Atoms can stick together because they are attracted to one another.
If they have the chance they do so, they lose their inherent chemical energy (or a part of it) in the process. This is transformed into kinetic energy and we say that the particles have got 'hotter'.
We register this as heat energy being produced.
Original post by charco
The internal enegy of a substance is a function of its inherent nature and its relative position in the universe.
The second one can never be known, but the first one is demonstrated by a lowering of energy when bonds are formed.
You could think of gravity as an analogy. An object in outer space has no apparent potential energy, but if it comes into the sphere of influence of a massive body then it will be attracted to it and accelerate to the surface. In doing so it loses the potential energy caused by its position and transforms it first into kinetic energy and then into heat (deformation)
Atoms can stick together because they are attracted to one another.
If they have the chance they do so, they lose their inherent chemical energy (or a part of it) in the process. This is transformed into kinetic energy and we say that the particles have got 'hotter'.
We register this as heat energy being produced.
The second one can never be known, but the first one is demonstrated by a lowering of energy when bonds are formed.
You could think of gravity as an analogy. An object in outer space has no apparent potential energy, but if it comes into the sphere of influence of a massive body then it will be attracted to it and accelerate to the surface. In doing so it loses the potential energy caused by its position and transforms it first into kinetic energy and then into heat (deformation)
Atoms can stick together because they are attracted to one another.
If they have the chance they do so, they lose their inherent chemical energy (or a part of it) in the process. This is transformed into kinetic energy and we say that the particles have got 'hotter'.
We register this as heat energy being produced.
So do you mean that the electrons are closer to the nucleus so have a lower energy level? If so how do they get closer to the nucleus?
(edited 11 years ago)
Original post by tazmaniac97
So do you mean that the electrons are closer to the nucleus so have a lower energy level? If so how do they get closer to the nucleus?
The electrons from one atom get closer to the nucleus of ANOTHER atom...
Bonding in hydrogen
Original post by charco
Oh so electrons aren't just in specific energy levels in the atoms they belong to , but all surrounding atoms too. That makes sense!
Original post by tazmaniac97
Oh so electrons aren't just in specific energy levels in the atoms they belong to , but all surrounding atoms too. That makes sense!
When bonded the electrons are in orbitals called a molecular orbital. This is of lower energy than the original atomic orbitals...
Original post by tazmaniac97
Oh so electrons aren't just in specific energy levels in the atoms they belong to , but all surrounding atoms too. That makes sense!
Although the model of electrons orbiting the nucleus is 'nice' it just isn't true. Electrons do not have a defined exact position, as defined by quantum mechanics. QM does however describe molecular orbitals in which the electrons are confined. These describe the energy of the electron and also a probability density of finding the electron within a region of space relative to the nucleus.
Now, a lot of that is totally unnecessary to your understanding, the important part is that the electrons have to occupy defined quantised energy levels in the atom. They can make transitions between these levels, usually promoted by energy from a photon.
In molecules, there are other energy levels to consider (translational motion also applies to atoms). It is the result of quantum mechanics that virtually everything is quantised, i.e. there is not a continuous range of values available for a variable. This includes the magnitude of the translational motion (velocity), vibrational motion and rotational motion for the molecule.
These energy levels are more closely spaced than the electronic energy levels and so the molecule can be promoted to a higher vib/rot/trans level by the thermal energy. This stores energy but can equally be released by relaxing a particle from a higher energy level to a lower one.
Original post by charco
When bonded the electrons are in orbitals called a molecular orbital. This is of lower energy than the original atomic orbitals...
Sorry, think you are right. There is very little different. I think it's a similar thread where somebody wrote a bad response and i got mixed up.... Soz
Original post by charco
When bonded the electrons are in orbitals called a molecular orbital. This is of lower energy than the original atomic orbitals...
So was what I said wrong? Sorry I don't know what you mean by molecular and atomic orbitals, how are they different and what exactly are they? (never come across them before- actually I've heard of atomic orbitals but not molecular)
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