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I've been thinking about refrigeration technology and am a bit confused about two common answers. Specifically, the part where the expansion valve releases the pressurized fluid and stuff gets real cold.

One is that refrigeration works by lower pressure = lower temperature. This makes sense to me because if there is lower pressure, I can imagine it as the opposite of higher pressure = higher temperature. Less pressure means particles are freer to move apart and thus eventually boil and lose energy as they travel further away from each other.

The other is the enthalpy of vaporization, which I understand as meaning that some amount of energy is required for a phase change. This also makes sense to me: when the refrigerant enters the low pressure side of the valve, the particles are more free to spread apart, begin to boil, and thus suck up surrounding energy to break apart from each other. Although, it seems a bit more magical than the lower pressure = lower temperature explanation (it seems odd that already hot liquid particles "suck up" more heat).

Could someone please help me understand this better? Thanks!

NOTE: As you can see, I think of this in very layman terms and am currently reading books like Feynman's lectures. My background is engineering, not physics, so I tend to understand things best in a much more physical-visualization, implementation-details kind of way.

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  • $\begingroup$ The second reason is correct: free expansion of a nonideal gas absorbs energy. As does evaporation, BTW, which is why we cool down as perspiration evaporates. $\endgroup$ Mar 20 '14 at 11:46
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Enthalpy (H) is the internal energy (U) of a material PLUS the product of pressure (P) and volume (V).

$H = U + PV$ by definition

When something boils, the gas phase takes up more volume than the liquid phase. So unless the boiling is in a vacuum, work is being done by expanding against a pressure, such as atmospheric pressure. This represents a change in the PV term of enthalpy. At a given temperature and pressure (say boiling water at atmospheric pressure), PV is greater for the gas phase than the liquid phase.

Additionally, intermolecular forces hold molecules of liquids together. For example, molecules of a particular compound (say fluoromethane) may have a permanent net dipole moment, and the positive end of one molecule is attracted to the negative end of the other. There are other types of intermolecular forces such as hydrogen bonding and London forces. It takes energy to pull the molecules apart against such forces. This represents a change in the internal energy (U) term of enthalpy. At a given temperature and pressure (say boiling water at atmospheric pressure), U is greater for the gas phase than the liquid phase.

In summary, a desired region is cooled because some of its energy is transferred to the refrigerant to increase its enthalpy. Part of the energy goes to expanding against a pressure (the PV term) and part goes to the increase in internal energy (the U term).

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  • $\begingroup$ thank you for the explanation, but in the case of a refrigerant, doesnt it boil because the lack of pressure on the other side of the expansion valve? if something boils because we put it in a lower pressure environment, it will pull in heat? im having a tough time visualizing how a boiling liquid (particles moving a lot) takes in heat. EDIT: i understand it takes energy to break the bond to turn something into a gas, but what drives (or motivates) already moving particles to pull in more energy to break themselves apart? $\endgroup$
    – tau
    Mar 20 '14 at 19:31
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    $\begingroup$ yes, you have to consider both pressure and temperature to decide if the substance will boil. I'd have to explain entropy and Gibbs free energy for a full explanation. I was giving constant temperature and constant pressure examples for simplicity. In any case, it takes energy to pull apart the molecules. This is reflected in the temperature of the evaporating liquid dropping or heat from the surrounding being absorbed, or a combinaiton thereof. $\endgroup$
    – DavePhD
    Mar 20 '14 at 19:47
  • $\begingroup$ so what i imagine in my head is a drop of freon coming out of the expansion valve. it is now in a lower pressure environment, so its boiling point lowers. this is due to the fact that in a lower pressure environment, the freon particles are able to move more freely. the particles then use the ambient heat (thus cooling things) to break their bonds and move freely as a gas. why dont the particles just remain bonded but under less pressure? if the vacuum (ie, lack of pressure) pulls them apart, then why does it need heat? thank you very, very much!! $\endgroup$
    – tau
    Mar 20 '14 at 20:07
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    $\begingroup$ The parameter for determining whether the gas or the liquid phase is more stable is Gibbs Free Energy (G). G = H - TS, where S is entropy. Entropy is a measure of disorder. The phase that has the lower G is the more favored (stable) phase at a particular T & P. At equilibrium G is the same for both phases. When the freon boils in response to the pressure drop this represents the gas phase reaching a lower G than the liquid phase under the new T & P conditions. $\endgroup$
    – DavePhD
    Mar 20 '14 at 20:30

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