What is Critical Point?
In chemistry, a critical point is known to be the point at which the minimum temperature (TCR , the critical temperature) and pressure (PCR , the critical pressure) that are needed to create a supercritical fluid are achieved. The point itself is defined as being located at the end of the phase equilibrium curve that occurs between liquids and gasses. Once the critical point is achieved or surpassed, a gas will not undergo condensation to become a liquid. Instead, it will phase into a supercritical fluid – a substance that has properties of both a liquid and a gas.
What is a Supercritical Fluid?
Supercritical fluids are unique in that they are not quite a liquid but are also not quite a gas. Despite this fact, they undeniably possess properties of both of these types of matter. A supercritical fluid has a low density that is often compared to that of a gas and can pass through certain materials. Additionally, supercritical fluids are able to dissolve certain materials similar to the manner in which water and other liquids would do so.
This combination of abilities makes supercritical fluids ideal for a number of different studies and real life applications. One of the most notable uses of supercritical fluids has been to extract certain materials and compounds from substances that would otherwise prove difficult to alter. This practice is employed frequently, such as the way caffeine is separated from coffee beans to create decaffeinated coffee.
Understanding Critical Point
Understanding the importance of critical point is largely dependent on having a comprehensive understanding of phase change and the way in which substances are depicted throughout their many different states in a pressure-temperature phase diagram. These tools help to express how molecules and compounds transition from state to state, and how different temperatures and pressures influence their change in state.
There are many different changes in state that can occur between elements. These changes in state are known as phase changes and help to determine what types of interactions molecules will have with one another. There are many different types of phase changes – all of which can be demonstrated in pressure-temperature phase diagrams.
The process in which a solid phases into a liquid is known as melting. This phase change is often associated with an increase in temperature, because the internal energy of the solid must increase before the transition can occur. Although this process is more commonly accomplished with the application of heat, an increase in pressure can sometimes have the same effect.
The process in which a liquid phases into a solid is known as freezing. This phase change tends to occur when there is a decrease in temperature. The lowered temperature causes the thermal energy in the liquid to slowly be removed, causing the substance to transition into a solid state. A similar effect can also be accomplished simply by increasing pressure. This process is known as solidification. Although the transformation from liquid to solid is the same, this is technically a different process. It is important that it is recognized, however, because the term is often used interchangeably with freezing due to their similarities.
The process in which liquid phases into a gas is known as vaporization. This phase change occurs when there is an increase in temperature at certain pressures. This temperature can be below the boiling point of a substance (resulting in evaporation) or at or above the boiling point of a substance.
The process in which a gas transitions to a liquid state is known as condensation. This phase change occurs when the gas particles are exposed to cool or cold conditions in a decrease in temperature. This usually happens when the gas comes into contact with a liquid or a solid surface.
The process in which a solid phases into a gas without first entering into a liquid state is known as sublimation. This state can only be achieved at temperatures and pressures below the triple point on a pressure-temperature phase diagram – specifically it must occur below the lowest pressure that is necessary for a liquid to exist. An example of this occurrence can be seen with dry ice, which will phase directly from its solid state to a gaseous state at room temperature.
The process in which a gas phases directly into a solid without first entering into a liquid state is known as deposition. Because this process is the reverse of sublimation, it is also sometimes called desublimation. In order for this phase change to occur, all the thermal energy must be removed from a gas. Although this may seem like a rare occurrence, this is actually the process that is responsible for the formation of snow in clouds.
Pressure-Temperature Phase Diagrams
Pressure-temperature phase diagrams are an excellent way to make quick observations about substances and the way they respond in different environments. These charts help to show the various states that elements phase in and out of and how they respond to certain temperatures and pressures.
Below, we’ll examine the structure of a pressure temperature diagram and the different phases it helps to distinguish.
The y axis of the diagram (the vertical axis) represents pressure and the x axis of the diagram (the horizontal axis) represents temperature. These two variables are responsible for adding or taking away energy from molecules and compounds, enabling them to manifest themselves into different states of matter.
The state of a substance can easily be identified by simply looking at the different lines on the graph. Solids are always identified in the space that is closest to the y axis, above the melting point line and the boiling point curve.
Liquids occupy the middle space of the graph past the triple point (the blue dot). Liquids will always occupy the middle space, because while solids and gasses are able to exist in a wide variety of temperatures, liquids are only able to exist in spaces past the triple point. Liquids are separated from solids by the melting point line. They are also separated from gases by the boiling point curve. Once the critical point has been reached, however, liquids cease to exist.
Gasses occupy the space that is closest to the x axis of the graph. They are located underneath the line that separates the solid state from the gaseous state, as well as the boiling point curve (separating liquids from gasses). It is also important to note that like liquids, the gaseous state ceases to exist after the critical point has been achieved.
The critical point is the point at which the critical pressure (PCR) and the critical temperature (TCR) converge on the graph together. As such, it can also be seen as the point at which the phase equilibrium curve between liquids and gasses ends. Once these two values intersect and form the critical point, there is an important shift. When liquids and gasses are exposed to critical pressure and critical temperature, they transition into what is known as a supercritical fluid, as seen below.
Supercritical fluids are unique substances that have qualities of both liquids and gasses but are classified as a different state of matter. Like gasses, they have low densities and can pass through certain materials. Like liquids, they are able to dissolve certain particles. Because supercritical fluids have this unique property, they are often used to extract certain properties from compounds that we would otherwise find difficult to alter.
Uses of Critical Point and Supercritical Fluids
Because supercritical fluids are able to dissolve certain compounds thanks to their partially liquid nature, they are often compared to water. This is essentially because they act as solvents that can easily be separated from the compounds they dissolve later on. The way in which supercritical fluids behave can easily be understood by examining the process of extracting salt from certain objects by using water.
To extract a salt from an object – say a piece of iron coated in salt – one simply needs to drop the iron into a container filled with water. The salt will easily dissolve into the water due to the nature of its bonds. The iron will remain, however, because its bonds are too strong to be broken by the water molecule. The iron can then be removed, leaving only the water and salt solution. Next, the water is easily separated from the salt by heating it to its boiling point and then waiting for the water molecules to evaporate. Salt has a higher boiling point than water, which means that all the salt molecules that were bonded to the water molecules will be left behind after the water molecules reach their boiling point and transition to a gaseous state.
Similarly, when a supercritical fluid is used, it can dissolve certain elements in compounds, which then mix in with the supercritical fluid. However, because the supercritical fluid state is dependent on both critical temperature and critical pressure, dropping the critical pressure will cause the supercritical fluid to transition back into its gaseous state. Because gasses are not able to dissolve compounds in the same manner that liquids can, the extracted molecules that were previously bonded to the supercritical fluid molecules of the compound will be left behind – just like salt molecules are left behind when water is heated into its gaseous state. When this process is done in a closed environment, the extracted molecules will drop to the bottom of the chamber or container.
How Critical Point and Supercritical Fluids are Used to Extract Caffeine from Coffee
There are many different real life applications of the aforementioned extraction process, but one of the most common uses of this process is the extraction of caffeine from coffee. It happens to be that caffeine is soluble in carbon dioxide (CO2). Because carbon dioxide cannot dissolve caffeine in its gaseous state, it becomes necessary to transition the substance into its supercritical fluid state in order to extract the caffeine particles.
To accomplish this, coffee beans are put into a chamber with dry ice (the solid form of CO2). The chamber is then pressurized and the temperature is increased until a critical point is reached and the dry ice transitions into a supercritical fluid. When this phase change occurs, the supercritical fluid begins to dissolve and absorb the caffeine particles in the coffee beans.
Once the caffeine particles have been dissolved, the pressure in the chamber is lowered while the temperature is maintained. This allows the CO2 to phase back into a gas state. However, because gas cannot dissolve particles in the same way that liquids and supercritical fluids can, the caffeine particles fall to the bottom of the chamber where they can be collected and used for other products.