In the region beyond the critical point no pressure or temperature change can convert a supercritical fluid to a gas or liquid. The fusion curve of carbon dioxide has a positive slope, while for water, the slope is negative. This is an atypical feature of water. Increasing the pressure favors a liquid-to- solid transition in carbon dioxide but a solid-to-liquid transition in water.
The higher pressure favors the denser solid form of carbon dioxide. In the case of water, the denser liquid form is favored. A phase diagram combines plots of pressure versus temperature for the liquid-gas, solid-liquid, and solid-gas phase-transition equilibria of a substance.
These diagrams indicate the physical states that exist under specific conditions of pressure and temperature and also provide the pressure dependence of the phase-transition temperatures melting points, sublimation points, boiling points. Regions or areas labeled solid, liquid, and gas represent single phases, while lines or curves represent two phases coexisting in equilibrium or phase change points.
The triple point indicates conditions of pressure and temperature at which all three phases coexist. In contrast, a critical point indicates the temperature and pressure above which a single phase—whose physical properties are intermediate between the gaseous and liquid states—exists.
A phase diagram identifies the physical state of a substance under specified conditions of pressure and temperature. To illustrate the utility of these plots, consider the phase diagram of water, shown below. Curve BC is the liquid-vapor curve separating the liquid and gaseous regions of the phase diagram and provides the boiling point for water at any pressure.
The physical properties of water under these conditions are intermediate between those of its liquid and gaseous phases. This unique state of matter is called a supercritical fluid. The solid-vapor curve labeled AB, indicates the temperatures and pressures at which ice and water vapor are in equilibrium. These temperature-pressure data pairs correspond to the sublimation, or deposition, points for water.
Note that this curve exhibits a slight negative slope, indicating that the melting point for water decreases slightly as pressure increases. Water is an unusual substance in this regard, as most substances exhibit an increase in melting point with increasing pressure.
The point of intersection of all three curves—labeled B—is the triple point of water, where all three phases coexist in equilibrium. At pressures lower than the triple point, water cannot exist as a liquid, regardless of the temperature.
The solid-liquid curve exhibits a positive slope, indicating that the melting point for CO 2 increases with pressure as it does for most substances. Notice that the triple point is well above 1 atm, indicating that carbon dioxide cannot exist as a liquid under ambient pressure conditions.
Instead, cooling gaseous carbon dioxide at 1 atm results in its deposition into the solid-state. Likewise, solid carbon dioxide does not melt at 1 atm pressure but instead sublimes to yield gaseous CO 2.
Finally, the critical point for carbon dioxide is observed at a relatively modest temperature and pressure in comparison to water. This text is adapted from Openstax, Chemistry 2e, Section To learn more about our GDPR policies click here.
If you want more info regarding data storage, please contact gdpr jove. Your access has now expired. Provide feedback to your librarian. If you have any questions, please do not hesitate to reach out to our customer success team. We can also focus on the lines that divide the diagram into states, which represent the combinations of temperature and pressure at which two states are in equilibrium.
The points along the line connecting points A and B in the phase diagram in the figure above represent all combinations of temperature and pressure at which the solid is in equilibrium with the gas. At these temperatures and pressures, the rate at which the solid sublimes to form a gas is equal to the rate at which the gas condenses to form a solid.
The solid line between points B and C is identical to the plot of temperature dependence of the vapor pressure of the liquid. It contains all of the combinations of temperature and pressure at which the liquid boils. At every point along this line, the liquid boils to form a gas and the gas condenses to form a liquid at the same rate.
The solid line between points B and D contains the combinations of temperature and pressure at which the solid and liquid are in equilibrium. At every point along this line, the solid melts at the same rate at which the liquid freezes. The BD line is almost vertical because the melting point of a solid is not very sensitive to changes in pressure. For most compounds, this line has a small positive slope, as shown in the figure above.
Obviously, pressure will vary wildly if you go to smaller planets or larger planets, or have thicker atmospheres, or if we're just doing different types of applications dealing with gases and liquids and solids. So what I've drawn here is a phase diagram. Let me write that down. And there are many forms of phase diagrams. This is the most common form that you might see in your chemistry class or on some standardized test, but what it captures is the different states of matter and when they transition according to temperature and pressure.
This is the phase diagram for water. So just to understand what's going on here, is that on this axis, I have pressure. On the x-axis, I have temperature, and at any given point, this diagram will tell you whether you're dealing with a solid, so solid will be here, a liquid will be here, or a gas. For example, if I told you that I was at 0 degrees, let's say 0 degrees is right there, if I'm at 0 degrees Celsius and 1 atmosphere, where am I? So 0 degrees, 1 atmosphere, I'm right at that point right there.
So I'm at a boundary point between solids and liquids at 1 atmosphere of pressure, right? This is when we're at 1 atmosphere of pressure. So this coincides with our traditional notion of when ice freezes or when it melts at 0 degrees. If we made the pressure higher, what happens? Well, then ice starts melting at a lower temperature, right?
So this is pressure going up, so pressure going up, let's say-- I don't know what this is. This is maybe 10 atmospheres, ten times Earth's atmospheric pressure at sea level, then all of a sudden, the temperature at which solid turns into liquid-- this transition is solid to liquid --the temperature at which that happens will go down.
Likewise, if we lower the pressure, if we go to Denver and it's a mile high, pressure is lower because we have less of the atmosphere above us, then all of a sudden, the freezing point increases, so the freezing point will be something above 1 degree. This isn't drawn completely to scale, but the idea is your ice would actually freeze a little bit faster and would freeze at a higher temperature in Denver than it would at the bottom of the Dead Sea or in Death Valley at some below sea level point on the planet.
Now, this transition is the transition between anything and gas. And we're very familiar, this is 1 atmosphere. And remember, this is water we're dealing with. This is the diagram for water, so at 1 atmosphere, this is kind of the stuff that we're used to seeing. Let me draw a line here. So at 1 atmosphere, 0 degrees is where solid, or ice, turns into liquid water. And then we go up here, so we keep going at a higher, higher, higher temperature, and then here, this would be, since we're at 1 atmosphere, this is degrees Celsius right there.
And that's the point at 1 atmosphere of pressure where liquid turns into gas, or water vaporizes, or the liquid boils.
All of those are acceptable ways to think about that. But what happens when we go to low pressure? Once again, let's take our little trip to Denver. So that's Denver right there. It's not that drastic. I'm just doing that for education purposes. Or even better, let's say Mount Everest. Mount Everest, very low pressure there.
Then our freezing point, we already said that goes up when you lower the pressure, and your boiling point goes down, so it's much easier to boil something on the top of Mount Everest than it is to boil it at the bottom or at the lowest point in Death Valley or the Dead Sea.
The intuition behind that is if I have a liquid, a bunch of molecules in liquid form, and they're touching each other, but they have enough kinetic energy to move past each other, so they're flowing past each other, they're kind of rubbing up against each other, one of the reasons why they don't just evaporate, why this guy doesn't just jump up there, is that there's air above him. There's air pressure. And air pressure, we've learned about this when we did PV nRT. That's a bunch of gas molecules, and the pressure they're creating is essentially caused by their temperature and their kinetic energy.
And they sit there, and they bounce, and they essentially keep these heavier molecules from going up. They keep them from essentially separating from each other and turning into a gas. So the more pressure you have, the harder it is for these guys to escape. On the other hand, if we're in a vacuum, if we're doing this on the surface of the moon and there's none of these guys there, then just a little slight bump. Even though this guy's still a little bit attracted to over here, they're still attracted to each other.
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