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# spontaneous processes watch

1. It says in my book 'a process is spontaneous if a chemical system becomes more stable and its overall energy decreases. The overall energy decrease results from contributions from both enthalpy and entropy'. So is △G the overall energy in the system plus surroundings? Isn't enthalpy the only measure of energy? Is entropy a measure of energy, and how and why does it contribute to the total energy?
2. The △G value determines if a reaction is feasible or not. if the △G value is >0 or =0 then the reaction is not feasible.
3. (Original post by Nitrogen)
The △G value determines if a reaction is feasible or not. if the △G value is >0 or =0 then the reaction is not feasible.
Yeah but why does it say in my textbook that a process is spontaneous (i.e. negative △G value) if a chemical system becomes more stable and its overall energy decreases? I don't understand how △G is a measure of the total energy, and why entropy contributes to this measure of total energy
4. There are several thermodynamic 'potentials', all with units of energy.

The Gibbs potential (G) is commonly used because it is the one which is minimised when a system at constant temperature and pressure reaches equilibrium.

I guess you have seen this, but G is defined as:

G = H - TS

Where H is the enthalpy, T is the temperature and S the entropy.

You can see that minimising enthalpy and maximising entropy (at a constant temperature) will both mimimise the Gibbs free energy.

For spontaneous processes (at constant temperature and pressure), the Gibbs free energy does decrease, which can be due to a reduction in energy, an increase in entropy, or both.
5. (Original post by celina10)
Yeah but why does it say in my textbook that a process is spontaneous (i.e. negative △G value) if a chemical system becomes more stable and its overall energy decreases? I don't understand how △G is a measure of the total energy, and why entropy contributes to this measure of total energy
A negative △G value means that the reaction is spontaneous. Something that has a high energy means that there is more disorder. For example a solid has less energy/less disorder than a gas. As you can observe solids are a lot more stable than gasses because its particles are tightly packed. On the other hand, the gas particles moves in different directions; its more disordered.
6. (Original post by celina10)
Yeah but why does it say in my textbook that a process is spontaneous (i.e. negative △G value) if a chemical system becomes more stable and its overall energy decreases? I don't understand how △G is a measure of the total energy, and why entropy contributes to this measure of total energy
ΔG is a measure of the free energy of the UNIVERSE, not the system under study. You can think of free energy as that portion of the energy of the universe that is available to do work.

The universal entropy is always increasing and this decreases the available free energy.

Gibbs free energy equation relates the free energy of the universe to the systems enthalpy and entropy change:

ΔG = ΔH - TΔS

When the value of ΔG is negative this means that the universal free energy has decreased and hence the universal entropy has increased. The process is consequently favourable (thermodynamically spontaneous)
7. (Original post by charco)
ΔG is a measure of the free energy of the UNIVERSE, not the system under study. You can think of free energy as that portion of the energy of the universe that is available to do work.

The universal entropy is always increasing and this decreases the available free energy.

Gibbs free energy equation relates the free energy of the universe to the systems enthalpy and entropy change:

ΔG = ΔH - TΔS

When the value of ΔG is negative this means that the universal free energy has decreased and hence the universal entropy has increased. The process is consequently favourable (thermodynamically spontaneous)
But when entropy increases isn't there a dispersal energy, so why can't this energy be used to do work? Also what is entropy? I don't understand what they mean by 'degree of randomness'
8. (Original post by Kerch)
There are several thermodynamic 'potentials', all with units of energy.

The Gibbs potential (G) is commonly used because it is the one which is minimised when a system at constant temperature and pressure reaches equilibrium.

I guess you have seen this, but G is defined as:

G = H - TS

Where H is the enthalpy, T is the temperature and S the entropy.

You can see that minimising enthalpy and maximising entropy (at a constant temperature) will both mimimise the Gibbs free energy.

For spontaneous processes (at constant temperature and pressure), the Gibbs free energy does decrease, which can be due to a reduction in energy, an increase in entropy, or both.
That formula confuses me, to me this would make more sense:
G = TS - H
so that an increase in enthalpy and entropy causes an increase in the free energy available for the universe to do work. Why is not like this? I'm so confused
9. (Original post by celina10)
That formula confuses me, to me this would make more sense:
G = TS - H
so that an increase in enthalpy and entropy causes an increase in the free energy available for the universe to do work. Why is not like this? I'm so confused
It is impossible to increase the free energy of the universe.

Gibbs free energy is actually derived as being equal to: -TΔSuniverse

So when the universal entropy increases (a spontaneous process) then the Free energy decreases...

The derivation goes like this:

We agree that the universe is made up of the system under study and its surroundings.

Therefore the total entropy in the universe must be equal to the entropy of the system + the surroundings, and any change in the universal entropy must be due to a change in entropy of system, surroundings or both. i.e.

ΔSuniverse = ΔSsystem + ΔSsurroundings

The surroundings can only be affected by heat exchange with the system. For an exothermic reaction the surroundings receives energy that increases its disorder (entropy) by a factor of q/T where q (in terms of the system) must be ΔH/T

Hence universal entropy can be written:

ΔSuniverse = ΔSsystem + ΔH/Tsurroundings

But when the surroundings increases in energy so the system decreases in energy, therefore in terms of the system:

ΔSuniverse = ΔSsystem - ΔH/Tsystem

Multiply through by -T to remove fractions:

-TΔSuniverse = -TΔSsystem + ΔHsystem

and rearrange:

-TΔSuniverse = ΔHsystem -TΔSsystem

Gibbs lumped together the -TΔSuniverse term and called it ΔG, Gibbs free energy.

So when Gibbs free energy is negative the entropy change of the universe must be positive. This can be measured by looking at the enthalpy and entropy change of the system under study:

ΔG = ΔHsystem -TΔSsystem
10. (Original post by charco)
It is impossible to increase the free energy of the universe.

Gibbs free energy is actually derived as being equal to: -TΔSuniverse

So when the universal entropy increases (a spontaneous process) then the Free energy decreases...
Why can't the free energy of the universe increase? I mean if the Gibbs free energy for the universe doesn't increase or decrease then why do we get positive and negative numbers using the formula i.e. shouldn't it be constant? And free energy for what (bold bit)?
11. (Original post by celina10)
Why can't the free energy of the universe increase? I mean if the Gibbs free energy for the universe doesn't increase or decrease then why do we get positive and negative numbers using the formula i.e. shouldn't it be constant? And free energy for what (bold bit)?
Because entropy is nature's arrow. Overall, everything tends towards increasing disorder as this is overwhelmingly the most probable state for particles and associated heat energy ...

Free energy is just a term given to the function -TΔS(universe)

It has been interpreted as being the energy that could be used to do work. All of this came about with the study of engines. We know that energy tends to flow from high temperature to ambient. It cannot go in reverse. So a region of high energy can be used to do work, such as boiling water to drive a steam engine. But the same energy spread out evenly throughout the universe cannot do any work. All it succeeds in doing is increasing the average heat of the universe by a miniscule fraction.

Energy does not have to exist as heat. It could also be chemical potential energy. In that state it could be transformed into heat energy and do work, before dissipating and being useless.

So the idea of 'Free energy' is energy that is still available to do work. Energy that has been released and dissipated has increased the entropy of the universe and is no longer 'useful'.

Hence an exothermic reaction releases heat energy, this increases universal entropy and decreases the free energy of the universe.
12. (Original post by charco)
Because entropy is nature's arrow. Overall, everything tends towards increasing disorder as this is overwhelmingly the most probable ...
Oh so it rarely ever occurs
13. (Original post by celina10)
Oh so it rarely ever occurs
It never occurs.

It would be like the same man winning the euromillions lottery every week for billions of years (assuming he lives that long). It is conceivable, but statistically so remote a possibility as to be an effective impossibility...
14. (Original post by charco)
It never occurs.

It would be like the same man winning the euromillions lottery every week for billions of years (assuming he lives that long). It is conceivable, but statistically so remote a possibility as to be an effective impossibility...
I think I understand, can you check if it's right?

ΔG = ΔH - TΔS
So basically ΔH is a measure of how much energy is available in the chemical system and TΔS is a measure of how much energy has now been dispersed and can no longer be used to do work. So the reason why you subtract TΔS from ΔH is because you want to find out how much energy is left to do work

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