What’s the term or word for the temperature where frozen (let’s say water) stays frozen and liquid stays liquid? Slightly below and the liquid freezes; slightly above, the frozen (water) thaws.
Is that the “critical point”?
The critical point is the point on a vapor-pressure curve at which the liquid and vapor phases of a substance are in equilibrium. That is, as a specific temperature and pressure, both liquid and gas phases of the substance will be present.
Aside from “below freezing” or “above freezing,” I can’t recall a special term used for temperatures slightly above or below that of the phase change for solids and liquids. Are you perhaps thinking of the “triple point,” where all three phases of a substance are in equilibrium?
About the only thing I can think of right now is steam/freezing points. Water at 0 C is at it’s freezing point, at -1 C it’s just below it’s freezing point, 1 C is just above it’s freezing point. Not sure if there’s a specific word to describe one substance in isolation just getting close to crossing these points. But you might use water and air being in thermal equilibrium if they are both at say -5 C. If the air is 5 C and ice -5 C, then they’re not in thermal equilibrium and the ice will start to melt (well… duh:D)
This is not the critical point:
“The critical point is the point on a vapor-pressure curve at which the liquid and vapor phases of a substance are in equilibrium. That is, as a specific temperature and pressure, both liquid and gas phases of the substance will be present.”
For example, start with a bottle completely filled with liquid. Then suck half the liquid out, without letting anything else in. Now the bottom half of the bottle contains liquid and the top half contains only the vapor of that substance (which would be at a low pressure, but not absolutely zero pressure). This is not at the critical point.
The critical point is the maximum temperature and pressure at which gas and liquid are separate phases. Above this temperature and pressure you have a “supercritical fluid”. If temperature is above the critical temperature and pressure is below the critical temperature, you have a gas, but squeezing it will not make it condense. If pressure is above the critical pressure but temperature is below the critical temperature, you have a liquid, but heating it will not make it evaporate.
Supercritical fluids are somewhat weird and unfamiliar to most of us, but you may have noticed something close to its critical point. The propane inside a propane torch (you know, the blue Bernz-O-Matic job) is not all that far from being supercritical. If you shake the thing, you feel stuff sloshing around inside, but it doesn’t slosh very hard - feels more like soap suds than a liquid. Imagine that as you heat this container up, the liquid would get less dense and the vapor would get more dense until they would sort of run together and you could no longer sense two separate phases in there. The sloshing has dwindled away to nothing. That’s what the critical point is about.
I think it’s just called the melting point. (Or, if you approach it from the other direction, the freezing point.) And water doesn’t thaw at just above the melting point, it thaws at the melting point.
Say you’ve got a block of ice. You apply energy to it, for instance by putting it in teh sun. The ice is at whatever temperature it was at in your freezer, let’s say -2 C. The heat applied to the ice raises its temperature from -2 to -1 to 0. Then, it stays at 0 for a while. But the same amount of energy is being applied. Where does it go? It’s used to transform the ice into liquid water. For some time, liquid and solid water are present at the same time, while the additional heat works on transforming the remaining ice. Eventually, all the ice is melted, but all the water is still at 0 degrees. After that, the sun continues to apply heat and the water therefore rises to 1, 2, 3 degrees C. If you keep applying enough heat, it keeps rising until a similar process occurs at 100 degrees.
Of course it’s actually more complex than this, but I haven’t taken a science class in 10 years, so I’ve forgotten a lot.
–Cliffy
I’ve heard it refered to as “phase interface”. At 0, water freezes, but ice also melts. I think.
Water freezes at 0 °C assuming two things:
- It’s (for the most part) pure, and
- The air pressure is 1 atm. (There is a slight dependence on air pressure.)
General Chemistry, 5th ed.; Darrell D. Ebbing, Wayne State University
Page 444: "The temperature above which the liquid state of a substance no longer exists regardless of the pressure is called the critical temperature. The vapor pressure at the critical temperature is called the critical pressure. It is the minimum pressure that must be applied to a gas at the critical temperature to liquefy it.
On a phase diagram, {the preceding experiment} corresponds to following the vapor-pressure curve where the liquid and vapor are in equilibrium. Note that this curve {in Figure 11.12A} ends at a point at which the temperature and pressure have their critical values. This is the critical point."
Chemistry, 3rd ed.; Steven S. Zumdahl, University of Illinois.
Page 481: "As the temperature increases, something happens {that we did not see in the first three experiments}: the liquid gradually changes into a vapor but goes through an intermediate “fluid” region, which is neither true liquid nor vapor.
This unusual behavior occurs because the conditions are beyond the critical point… The critical temperature can be defined as the temperature above which the vapor cannot be liquefied no matter what pressure is applied. The critical pressure is the pressure required to produce liquifications at the critical temperature. Together, the critical temperature and the critical pressure define the critical point.
Note that the liquid/vapor line on the phase diagram ends at the critical point. Beyond this point the transition from one state to another involves the intermediate “fluid” region just described."
xcheopis summmation:
“The critical point is the point on a vapor-pressure curve at which the liquid and vapor phases of a substance are in equilibrium.”
Which is, indeed, correct.
Actually, there is a difference between the “freezing point (FP) temperature” and “melting point (MP) temperature” of a substance (including water). This is because, for most elements and compounds, there is a characteristic curve (temperature vs. time) for the MP and FP. In other words, when a pure substance is melting, the temperature does not remain constant. Likewise, the temperature does not remain constant during freezing. And these two curves look different. Research has shown that FP curves tend to be an order-of-magnitude flatter than MP curves, thus FP’s are typically used to define “fixed points” on the International Temperature Scale (ITS-90) for temperatures above 29 °C (Gallium is the only exception). Below 29 °C triple points and vapor pressure points are used.
Sorry, but you’ve misunderstood the texts. The liquid and vapor phases are in equilibrium along the entire liquid/vapor curve in the phase diagram; that’s what the curve MEANS. Napier is correct.
Interesting, Crafter_Man… Can you give us any more details? Just how much does the temperature change while ice is melting, or water freezing? Are we talking a tenth of a Kelvin, or a ten millionth of a Kelvin? And which point on the curve is defined as “the” freezing point? Any cites?
Nope. I understood the text just fine. Doc Nickel asked if the critical point was the point at which the solid/liquid phase of a substance was no longer in equilibrium with the liquid/solid phase. The answer is “No, the critical point refers to the liquid/gas phase and at the critical point there is still an equilibrium.”
Sorry if that wasn’t clear before.
You’re correct to speculate that the drift is very slight… how does 50 to 100 millionths of a degree (°C) over 10 hours sound? But it’s there, and it’s measurable.
NIST is the authority. Go here for some sample curves:
http://www.cstl.nist.gov/div836/836.05/thermometryc/miscelaneous/faq.htm
I’ve visited their Precision Thermometry Lab where they calibrate Standard Platinum Resistance Thermometers (SPRTs). (I used to operate a temperature calibration lab at a DOE facility, so I was in touch with these guys frequently.) It’s quite a site.
“No, the critical point refers to the liquid/gas phase and at the critical point there is still an equilibrium.”
If you plot density as a function of temperature and pressure, there will be a discontinuity curve between gas and liquid phases, and this discontinuity will stop at the critical point. But every point on the graph represents an equilibrium condition. And in systems where there is not an equilibrium, you can still have gas, liquid, supercritical, and other phases. The issue of equilibrium is separate from the issue of criticality.
Read the texts carefully - the reference to equilibrium is not the central point. One could have bolded any word in the text, but that would not have changed the meaning. Your first quoted paragraph explains it clearly enough - the second paragraph tells you how to find the point on the phase diagram - and neither states that equilibrium provides the test of criticality.