Why is cooling considered to slow reactions?

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Hello everyone :smile:

I'm just wondering why is cooling considered to slow reactions. I mean f you want condensation heating would slow the reaction. I like to know exactly how condensation occur at molecular level. I know when you heat gas it expands. So when you cool a gas does it contract, and where does the energy for change phase from gas to liquid come from. Does cooling provide energy too like heat. Why is that when cooling in this case there is not much kinetic energy but a reaction (condensation) takes place. Your help would much appreciated thanks!!
 

Answers and Replies

  • #2
Mapes
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Hello everyone :smile:

I'm just wondering why is cooling considered to slow reactions.
A lower temperature means that a substance's atoms and molecules have a lower energy distribution. This generally makes them less likely to participate in reactions, all else being equal (i.e., if the reaction is spontaneous at either temperature).

When a phase change is involved, though, the reaction may not be spontaneous at the high temperature. In this case, cooling enables the reaction rather than suppressing it, and continued cooling can increase the reaction rate up to a point.

I like to know exactly how condensation occur at molecular level.
Below a certain temperature (e.g., 100°C for water), the liquid phase is more stable. But unlike gases, liquids and solids need to start by nucleating, or forming many-atom clusters. It's just a matter of chance for atoms/molecules to associate into a cluster. Once this happens, the cluster (a drop of liquid water) can grow by the attachment of additional atoms/molecules.

I know when you heat gas it expands. So when you cool a gas does it contract, and where does the energy for change phase from gas to liquid come from.
Yes, it contracts. No energy is needed to condense; in fact, energy needs to be taken away. Gas has a higher enthalpy than liquid (the difference, quantitatively, is the well-known heat of vaporization).

So why would gas ever be the stable phase if it requires more energy? The answer is that at constant temperature and pressure, the most stable phase is the one with the lowest Gibbs free energy (which is different from internal energy and enthalpy). The Gibbs free energy [itex]G=H-TS[/itex] rewards higher-entropy substances at higher temperatures, and a gas has a higher entropy than a liquid.

Does cooling provide energy too like heat.
No; it removes energy, but the lower temperature can cause a particular reaction (like condensation or freezing) to be spontaneous, as described above.
 
  • #3
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The idea that heat increases reactions really applies to chemical bonding, not condensation!:

When two atoms get near each other, the electron clouds surrounding those atoms act like "shields" trying to push the atoms apart, because those clouds have like charges (both negative), and two negative charges repel.

But if the atoms are pushed hard enough together, then other effects become more prominent, like the interaction (or lack of interaction) between two electrons of opposite spin, which let's the attraction of the nucleus of one atom with an electron of the other atom become the prominent force. So if the atoms are pushed hard enough together, then, depending on the type of atoms, they might get locked together by sharing electrons, for example.

When everything is hot, more atoms are moving fast enough to get close enough to get locked together. Likewise, fast moving molecules are more likely to smack into an existing molecule and break it apart.
 
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Thanks a lot for your very informative reply Mapes :smile: Your obviously very advanced in chemistry and I haven't learned many of the concepts you described above. So I get bits and pieces of what you said. I'll ask few questions to simplify this

1. So in different temperatures stability of water is different? Am I right?
2. If Gas has the highest enthalpy that means it is less stable so the reason it doesn't release this energy and become liquid in high temperature is becaus of the Gibbs energy (Concept I have not learned). Am I right?
3. So in cooling the gas molecules must get closer together and form liquid. Wouldn't it take time for these gas molecules to get closer together. I mean when you heat somethings molecules collide fast but when you cool it collides slowly. So how does condensation does not take time. I mean how is making these clusters happen quickly without high heat to speed up the molecules.

Thanks once again. Thanks for Fleem as well but I that wasn't what I was asking.
 
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  • #5
Mapes
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I understand spontaneous reaction is something that does not need an external force to happen. So why does it matter if it is spontaneous or not in a phase change.
My point here is that a phase change reaction such as [itex]H_2O(g)\rightarrow H_2O(l)[/itex] is spontaneous at less than 100°C only. So increasing the temperature from 98°C to 104°C, for example, doesn't speed up the reaction, it prevents it! So phase changes can be an exception to the general principle that heat transfer speeds up reactions.

Is this simply because when bonds are formed energy is released . Is that the reason cooling does not need energy because it makes bonds.
Yes, energy is released when a substance condenses or freezes. Now, it takes energy to operate a refrigerator to cool an object, but obviously this is a different issue.
 
  • #6
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I just edited my previous reply Mapes it has more questions. So please look into that as well. Thank you :smile:
 
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My point here is that a phase change reaction such as [itex]H_2O(g)\rightarrow H_2O(l)[/itex] is spontaneous at less than 100°C only. So increasing the temperature from 98°C to 104°C, for example, doesn't speed up the reaction, it prevents it! So phase changes can be an exception to the general principle that heat transfer speeds up reactions.



Yes, energy is released when a substance condenses or freezes. Now, it takes energy to operate a refrigerator to cool an object, but obviously this is a different issue.
lol. Yeah thanks for clearing it up :smile: . Now I understand much better. However I would like to clear up these two questions. I just edited my post at the same time as your reply so it got bit confusing. Here are my questions.

2. If Gas has the highest enthalpy that means it is less stable so the reason it doesn't release this energy and become liquid in high temperature is becaus of the Gibbs energy (Concept I have not learned). Am I right?
3. So in cooling the gas molecules must get closer together and form liquid. Wouldn't it take time for these gas molecules to get closer together. I mean when you heat somethings molecules collide fast but when you cool it collides slowly. So how does condensation does not take time. I mean how is making these clusters happen quickly without high heat to speed up the molecules.

Thanks once again for your time so far :smile:
 
  • #8
Mapes
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2. If Gas has the highest enthalpy that means it is less stable so the reason it doesn't release this energy and become liquid in high temperature is becaus of the Gibbs energy (Concept I have not learned). Am I right?
Right; the Gibbs free energy is the final word on stability for systems at constant pressure and temperature. Lower is more stable.

3. So in cooling the gas molecules must get closer together and form liquid. Wouldn't it take time for these gas molecules to get closer together. I mean when you heat somethings molecules collide fast but when you cool it collides slowly. So how does condensation does not take time. I mean how is making these clusters happen quickly without high heat to speed up the molecules.
It does take a finite time, but have you ever calculated how fast gas atoms/molecules move? The time is short by human standards.

But now what about in the liquid state, when the atoms are moving much slower? Your question gets right to the heart of some very interesting physical effects. For example, it's possible to cool a material below the freezing point (where it would normally form a well-organized crystal) and have it stay in an unorganized amorphous state because the atoms are too slow to attach to the crystalline structure in any reasonable time. It's a kinetic limitation, as you describe. It's easier to accomplish with some materials than others; amorphous metals are difficult to make, but all window glass is amorphous (the crystalline form is quartz).

Thanks once again for your time so far :smile:
My pleasure; I think your intuition is good. If you're really interested in these topics, you might want to explore materials science, which includes the study of how macroscale material properties arise from atomic and molecular interactions. It's full of concepts of energy, entropy, phase changes and crystal structure.
 
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That cleared everything up :smile: Thanks a lot once again Mapes :smile: Yes I might think about exploring these fields in the future. Thanks again :smile:
 
  • #10
So in sense, the moisture holding capacity of air depends on temperature because when the water molcules are ordered in a more random state there is more space between each molecule resulting in the molecules to be pushed up against one another (on the opposite side) forming a bond resulting in liquid water. Correct?
 
  • #11
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My point here is that a phase change reaction such as [itex]H_2O(g)\rightarrow H_2O(l)[/itex] is spontaneous at less than 100°C only. So increasing the temperature from 98°C to 104°C, for example, doesn't speed up the reaction, it prevents it!
You seem to be confusing equilibrium thermodynamics with kinetics.

It's entirely possible (and relatively common) to speed up a reaction while at the same time reducing the amount of products formed--you just have to speed up the reverse reaction by a greater amount.
 
  • #12
Mapes
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Point taken. As I wrote in the very next line, my intention was just to provide a counter-example to the idea that a temperature increase always accelerates processes. I can appreciate that condensation is accelerated as we pass 100°C even as it is dominated by evaporation.
 

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