Changes in Physical State - phase transitions - Laura Qiu
|Solid Liquid Gas Plasma|
|Solid||Solid-Solid Transformation Melting Sublimation -|
|Liquid||Freezing N/A Boiling/Evaporation -|
|Gas||Deposition Condensation N/A Ionization|
|Plasma||- - Recombination/Deionization N/A|
Evaporation is the process by which molecules in a liquid state (e.g. water) spontaneously become gaseous (e.g. water vapor), without being heated to boiling point. It is the opposite of condensation. Generally, evaporation can be seen by the gradual disappearance of a liquid, when exposed to a significant volume of gas.
The reason a liquid evaporates is that its molecules are all in motion in nearly random directions and speeds, and the energy of that movement can be compared to the heat needed to boil that liquid.
Condensation is the change in matter of a substance to a dense phase, such as a gas (or vapor) to a liquid. Condensation commonly occurs when a vapor is cooled to a liquid, but can also occur if a vapor is compressed (i.e., pressure on it increased) into a liquid, or undergoes a combination of cooling and compression.
Water vapor from air which naturally condensed on cold surfaces into liquid water is called dew. Water vapor will only condense onto another surface when that surface is cooler than the temperature of the water vapor, or when the water vapor equilibrium in air, i.e. saturation humidity, has been exceeded.
Melting is a process that results in the phase change of a substance from a solid to a liquid. The internal energy of a solid substance is increased (typically by the application of heat) to a specific temperature (called the melting point) at which it changes to the liquid phase
The melting point of a substance is equal to freezing point.
** List of elements by melting point ** http://en.wikipedia.org/wiki/List_of_elements_by_melting_point (the picture is too big to put here)
In physics and chemistry, freezing is the process whereby a liquid turns to a solid. The freezing point is the temperature at which this happens. Melting, the process of turning a solid to a liquid, is the opposite of freezing. All known liquids undergo freezing when the temperature is lowered with the sole exception of helium, which remains fluid at abusolute zero and can only be solidified under pressure.
For most substances, the melting and freezing points are the same temperature, however, certain substances possess differing solid-liquid transition temperatures. For example, agar melts at 85 °C (185 °F) and solidifies from 31 °C to 40 °C (89.6 °F to 104 °F); this process is known as thermal hysteresis.
Sublimation of an element or compound is a transition from the solid to gas phase with no intermediate liquid stage. Sublimation is a phase transition that occurs at temperatures and pressures below the triple point.
The opposite of sublimation is deposition. The formation of frost is an example of meteorological deposition.
In chemistry, deposition is the settling of particles (atoms or molecules) or sediment from a solution, suspension mixture or vapor onto a pre-existing surface. Deposition generally results in growth of new phase and is of fundamental importance in a large number of scientific disciplines and practical applications, the most obvious ones being in material science, geology, meteorology and chemical engineering.
A phase change may be written as a chemical reaction. The transition from liquid water to steam, for example, may be written as
H2(l) = H2(g)
The equilibrium constant for this reaction (the vaporization reaction) is
where Pw is the partial pressure of the water in the gas phase when the reaction is at equilibrium. This pressure is often called the vapor pressure. The vapor pressure is literally the partial pressure of the compound in the gas.
This equilibrium may be established at any temperature. Because vaporization reactions are endothermic, an increase in temperature will shift the equilibrium to the right. Thus at low temperatures the vapor pressure of the liquid is very low and at high temperatures the vapor pressure is quite large.
* Flash about Vapor Pressure * http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/vaporv3.swf
Boiling vs. Evaporation:
Ordinary evaporation is a surface phenomenon - since the vapor pressure is low and since the pressure inside the liquid is equal to atmospheric pressure plus the liquis pressure, bubbles of water vapor cannot form. But at the boiling point, the saturated vapor pressure is equal to atmospheric pressure, bubbles form, and the vaporization becomes a volume phenomena.
Normal Boiling Temperature ( Boiling Point):
Boiling Point is defined as the temperature at which the saturated vapor pressure of a liquid is equal to the surrounding atmospheric pressure.
*For water, the vapor pressure reaches the standard sea level atmospheric pressure of 760mmHg at 100°C.
Since the vapor pressure increases with temperature, it follows that for pressure greater than 760 mmHg (e.g., in a pressure cooker), the boiling point is above 100°C and for pressure less than 760 mmHg (e.g., at altitudes above sea level), the boiling point will be lower than 100°C. As long as a vessel of water is boiling at 760 mmHg, it will remain at 100°C until the phase change is complete. Rapidly boiling water is not at a higher temperature than slowly boiling water. The stability of the boiling point makes it a convenient calibration temperature for temperature scales.
Boiling Point Variation:
The standard boiling point for water at 100°C is for standard atmospheric pressure, 760 mmHg. It is the experience of high altitude hikers that it takes longer to cook food at altitude because the boiling point of water is lower. On the other hand, food cooks more quickly in a pressure cooker because the boiling point is elevated. Raising or lowering the pressure by about 28 mmHg will change the boiling point by 1°C.
Boiling Point Variation Near 100 C
Values were taken from the saturated vapor pressure table for water near 100 degrees Celsius. An empirical fit to these data values was made, and the formula obtained is shown on the diagram. It could be considered to be reasonably valid only for a few degrees above and below 100 °C since the curve is very non-linear.
Vapor Pressure Curves
temperature increases, the amount of vapor generated by a liquid in a closed container increases. This occurs because as the liquid gains kinetic energy, the molecules can overcome the intermolecular forces of attraction that are prevalent in the liquid phase.
Example: Vapour-pressure curve of carbon dioxide
Concept: A diagram showing the various phases of a system is called a phase diagram.
Phase diagrams for a pure compound such as phase diagrams for water and carbon dioxide are phase diagrams for a single component system. In these diagrams, pressure (P) and temperature (T) are usually the coordinates. The phase diagrams usually show the (P, T) conditions for stable phases.
Phase Diagram of Water
Figure: Phase diagram of water. Blue area represents the gas phase, yellow area is the liquid phase, cyan area is the solid. the green region represents the critical region in the phase diagram where it is not possible to distinguish between gas and liquid. Words in italics represent the process (double arrows). letters in Ariel mark special points on the graph and normal test labels the regions and the axes. This plot is not to scale! It is meant to be an illustration only. There are several features of such a phase diagram that need discussion.
Phase Diagram of Carbon Dioxide
Beyond a specific temperature and pressure (the critical point) carbon dioxide becomes a supercritical fluid, a state that is neither a gas nor a liquid, but has properties of both.
Phase diagrams are useful for material engineering and material applications. With their aid, scientists and engineers understand the behavior of a system which may contain more than one component (compounds). Multicomponent phases diagrams show the conditions for the formation of solutions and new compounds. Thus, phase equilibria is still a field of research, and there is a Journal of Phase Equilibria for the publication of these research results.
** Flash of Phase Diagram * * http://www.chm.davidson.edu/ChemistryApplets/PhaseChanges/PhaseDiagram.html
Skills: Relating Phase Behavior to Molecular Structure
The effects of the molecular structure of the interface and continuous phase on the solubilization capacity of water in water/oil microemulsions have been studied from both theoretical considerations and experimental observations. From the consideration of the thermodynamic stability of microemulsion systems, we have shown that the growth of droplets during the solubilization process is limited at least either by the spontaneous curuature of the interface or by the attractive interaction among the microemulsion droplets. The influence of the chemical structure of the components on the solubilization capacity was therefore analyzed on the basis of consideration of their effects on the curvature and attraction just mentioned. Experimentally, the solubilization of water in water/oil microemulsions was studied by changing the following variables: molecular volume of oil, chain length, polar head and concentrations of cosurfadants, and salinity. For the systems where the solubilization capacity is limited by the radius of the spontaneous curvature of interface, it was found that the solubilization can be improved by any change in the above variables leading to the decrease of the curvature. For the system where the solubilization capacity is limited by the attractive interaction between microemulsion droplets, any change resulting in the decrease of attractive interaction increases the solubilization capacity. We have further shown that any change in the abovementioned variables can have opposite effects on the curvature and attraction. Therefore, maximum solubilization is observed as a result of the compromise between these two opposite effects.