Phase-changes

A phase is a homogeneous part of a system that is separated from the rest of the system by a well-defined boundary. When an ice cube floats in a glass of water, for example, the liquid water is one phase and the solid water (the ice cube) is another. Although the chemical properties of water are the same in both phases, the physical properties of a solid are different from those of a liquid. 

TABLE 12.5 Compositiini. .tnd Propertifs of Threit ryiir s uf Gliss 

Pure quartz glass 100% Si02 Low thermal expansion, transparent to a wide 

Pyrex glass 60%-80% SiOi, Low thermal expansion transparent to visible 

Soda-lime glass 75% Si02, Easily attacked by chemicals and sensitive to 

Lecture 5 Heat Capacity, Phase Change, and Colucative Properties 83 

You need to know the names of the types of phase changes melting-freezing vaporizatioii-condense ion sublimation 

The slope of the heating curve, where not zero, is proportional to the inverse of the Each phase of a substance has its ovun 

Evaporation occurs when the partial pressure above a liquid is less than the liquid s vapor pressure, but the atmospheric pressure is greater than the vapor pressure. Under these conditions, the liquid evaporates rather than boils. 

I Explain how the addition and removai of energy can cause a phase change. 

Real-World Reading Link Have you ever wondered where the matter in a soiid air freshener goes The day it is opened and put in a room, it is a solid, fragrant mass. Day-by-day, the solid gets smaller and smaller. Finally, almost nothing is left and it is time to put a new one out. You never observe a puddle of liquid like you would see if it had melted. 

Most substances can exist in three states depending on the temperature and pressure. A few substances, such as water, exist in all three states under ordinary conditions. States of a substance are referred to as phases when they coexist as physically distinct parts of a mixture. Ice water is a heterogeneous mixture with two phases, solid ice and liquid water. 

Determine what phase changes occur between solids and liquids. 

Describe how the curve would look for the same liquid at 30°C. 

Consider the following sequence of phase changes solid liquid vapor. Solids must absorb heat in order to melt, and Uquids must absorb heat in order to vaporize. Both are endothermic processes. When a soUd melts, there isn t very much change in volnme (maybe 10 percent or so), so the work done when the solid expands is usually ignored. (Of course, in the case of water, the solid contracts as it melts—more about that in Chapter 24.) On the other hand, there is a big change in volume when a liquid vaporizes—the increase in volume is on the order of a factor of 1,000. 

A liquid must do work of expansion against the surrounding atmosphere as the liquid vaporizes and the vapor expands. Consequently, a liquid needs a lot of energy in order to vaporize. A liquid must absorb much more heat to convert to the vapor phase than is required for a solid to convert to the liquid phase. Ice, for example, requires 333 joules of heat per gram to melt at 0°C, in contrast to liquid water, which requires 2,257 joules per 

The reverse sequence of phase changes—vapor liquid solid—are all exothermic. In each case, heat must be given off. Gases must give offbeat in order to condense into the liquid phase, and liquids must give offbeat in order to solidify or crystallize. This brings up an important point to remember The reverse of any endothermic process is an exothermic process, and vice versa. 

Heat of fusion is the amount of heat required to melt a solid. 

When liquid water is heated, some molecules escape from the liquid and enter the gas phase. If a substance is usually a liquid at room temperature (as water is), the gas phase is called a vapor. Vaporization is the process by which a liquid changes into a gas or vapor. When vaporization occurs only at the surface of a hquid, the process is called evaporation. 

Vapor pressure is the pressure exerted by a vapor over a liquid. As temperature increases, water molecules gain kinetic energy and vapor pressure increases. When the vapor pressure of a liquid equals atmospheric pressure, the liquid has reached its boiling point, which is 100°C for water at sea level. At this point, molecules throughout the hquid have the energy to enter the gas or vapor phase. 

Solving Problems A Chemistry Handbook Chemistry Matter and Change 133 

The process by which a solid changes directly into a gas without first becoming a liquid is called sublimation. Solid air fresheners and dry ice are examples of solids that sublime. At very low temperatures, ice will sublime in a short amount of time. This property of ice is used to preserve freeze-dried foods. 

Liquid water left uncovered in a glass eventually evaporates. An ice cube left in a warm room quickly melts. Solid CO2 (sold as a product called dry ice) sublimes at room temperature that is, it changes directly from solid to gas. In general, each state of matter—solid, 

The enormous differences in temperature one can create depend (disregarding energy dissipation) on heat output, rate of reactioo, and on the physical properties of the ingredients and the products. Of these factors, the phases of the products—solid, liquid, or gaseous —are undoubtedly the most important. We need only think of a certain amount of thermite compared with the same amount of an ammonium nitrate propellant, both of identical heat output per unit weight, to appreciate the extremes of behavior due to phase.changes. 

In this example, the division seems to be fairly clear cut, since the products from the propellant are unequivocally gaseous even in the standard state or only slightly above it, including water vapor. But in a flare we have to ascertain if or to what extent an oxide or chloride is volatilized, and even in a highly refractory system such as thermite when used in technical quantities, some of the aluminum oxide whose bulk is concentrated in a liquid slag may appear as Rauch (smoke),  

Dislocation of heat is also caused by the small amounts of gas 

The liquid state in the reaction zone causes, as a rule, little concern from the viewpoint of heat loss, since the latent heat of fusion is small. It can, however, cause a form of convective heat loss that may have unfortunate consequences. In a flare that burns with the flame upward, such as a railroad fusee, a molten slag is sometimes normally permitted to drop off. By faulty formulation, an excessive liquefaction may take place so that not only the burned-out slag but the molten material in the burning zone itself may slide off and prematurely end the burning of the flare. Similar conditions may occur with flame-down burning candles under violent motion. 

Melting and boiling points of substances at high temperatures ate quoted among others from Brewer, and from a book by Kuba-schewski and Evans, the latter providing many useful tables on heat of formation and heat of transformation, heat capacities, and information on calorimetry. 

Melting or fusion is the change of a solid to a liquid. Heat is absorbed as intermolecular forces are partially overcome, allowing the solid to melt to a liquid. 

Example Heat + H2Ofe) — HzO(1) (melting ice to cool a glass of tea) 

Although we are not aware of it happening, ice and snow sublime slowly on winter days. Water molecules escape directly from the surface of ice and enter the atmosphere as a gas. Dry ice gets its name from the fact that as a solid at -78°C, it can disappear (sublime) at normal atmospheric pressure without becoming a liquid. Freeze-dried foods are made by freezing them and removing the water by sublimation at low pressures. 

Freezing is the change of a liquid to a solid. Heat energy is released as a liquid freezes. 

Note that the diagram has three general areas corresponding to the three states of matter solid, liquid, and gas. The line from A to C represents the change in vapor pressure of the solid with temperature for the sublimation (going directly from a solid to a gas without first becoming a liquid) equilibrium. The A to 

D line represents the variation in the melting point with pressure. The A to B line represents the variation of the vapor pressure of a liquid with pressure. This B point shown on this phase diagram is the critical point of the substance, the point beyond which the gas and liquid phases are indistinguishable from each other. At or beyond this critical point, no matter how much pressure is applied, it is not possible to condense the gas into a liquid. Point A is the triple point of the substance, the combination of temperature and pressure at which all three states of matter can exist. 

In looking at phase diagrams, be careful when moving from point to point to pay attention to any phase changes which might occur. 

Intermolecular forces can affect phase changes. Strong intermolecular forces require more kinetic energy to convert a liquid into a gas. Stronger intermolecular forces, make it easier to condense a gas into a liquid. 

Facts do not cease to exist because they are ignored. 

Changing from the solid state to the liquid state is called fusion or melting, and from liquid to solid is called freezing. Because both liquids and solids are less energetic than liquids and gases, the amount of heat required to melt a solid is normally less than the amount of heat required to evaporate a liquid. This state transition is similar to the liquid-to-gas transition in that it occurs at a constant temperature for a pure substance. 

Liquid solutions usually freeze at lower temperatures than the pure liquids, with the freezing point depression (for aqueous solutions) being directly related to the amount of solute dissolved. Certain proportions of two metals often freeze at a temperature much lower than for other proportions. The proportions of these two metals that freeze at the minimum temperature is called the eutectic mixture. One useful eutectic mixture is tin/lead in a 60/40 mixture. Pure lead melts at 327°C pure tin melts at 232°C. The 60/40 mixture, called solder, melts at 190°C, lower than for either metal alone. For this reason, 60/40 solder is often used for constructing electric circuits. 

Some solid molecules can attain enough energy to change directly from solid to gas. Such a transformation is called sublimation. Because of sublimation, there is a vapor pressure of the material that is present over a solid. The water vapor pressure of ice is normally considered to be dependent only on temperature. The heat of sublimation for ice to steam is slightly greater than the sum of the heat of fusion for ice to water plus the heat of vaporization from water to steam. 

State change is not the only possibility with solids. There are also phase changes that occur under different conditions of temperature and pressure. Phase changes may be crystal to glass, or one crystalline structure to another crystalline structure. When these occur, there are definite changes of the physical properties of the material, and heat may either be evolved or absorbed. 

If the inside surface of each tube were coated with wax, would the general shape of the water meniscus change Would the general shape of the mercury meniscus change  

Because adhesive cohesive, H2O molecules touching glass adhere to the wall more than to each other, forming concave surface 

Melting is called (somewhat confusingly)/ sion. The increased freedom of motion of the peirticles requires energy, measured by the heat of fusion or enthalpy affusion, AHfus. The heat of fusion of ice, for excunple, is 6.01 kj/mol  

How is energy evolved in deposition related to those for condensation and freezing  

TABLE 11.5 Composition and Properties of Three Types of Glass 

Low thermal expansion transparent to visible and infrared, but not to UV, radiation. Used mainly in laboratory and household cooking glassware. 

What happens at the molecular level during evaporation In the beginning, the traftic is only one way Molecules are moving from the liquid to the empty space. Soon the molecules in the space above the liquid establish a vapor phase. As the concentration of molecules in the vapor phase increases, some molecules condense, that is, they return to the liquid phase. Condensation, the change from the gas phase to the liquid phase, occurs because a molecule strikes the liquid surface and becomes trapped by intermolecular forces in the hquid. 

Apparatus for measuring the vapor pressure of a liquid (a) before the evaporation begins and (b) at equilibrium, when no further change is evident. In (b) the number of molecules leaving the liquid is equal to the number of molecules returning to the liquid. The difference in the mercury levels (h) gives the equilibrium vapor pressure of the liquid at the specified temperature. 

It is important to note that the equilibrium vapor pressure is the maximum vapor pressure a liquid exerts at a given temperature and that it is constant at 

Comparison of the rates of evaporation and condensation at constant temperature. 

The increase in vapor pressure with temperature for three liquids. The normal boiling points of the liquids (at 1 atm) are shown on the horizontal 

Unless otherwise noted, all art on this page is Cengage Learning 2014. 

In most cases, changes in phase (solid liquid, liquid gas, solid gas) occur under experimental conditions of constant pressure, so that the heat involved, q, is also equal to AH. For example, for the melting of ice at its normal melting point of 0°C 

HjO can exist at 0°C as either a solid or a liquid. Because there is no AT, equation 2.9 does not apply. Instead, the amount of heat involved is proportional to the amount of material. The proportionality constant is called the heat of fusion, AfusH, so that we have a simpler equation  

The word fusion is a synonym for melting. If amount m is given in units of grams, AfusH has units of J/g. If the amount is given in units of moles, equation 2.50 is more properly written as 

FIGURE 2.12 Qualitative behavior of the thermal energy and the heat capacity for water from 0 K to over 373 K. Both graphs start at ay value of zero. 

The metal iridium (Ir) crystallizes with a face-centered cubic unit cell. Given that the length of the edge of a unit cell is 383 pm, determine the density of iridium in g/cw . 

Strategy A face-centered metalhc crystal contains four atoms per unit cell [8 X (comers) and 6 X i(faces)]. Use the number of atoms per cell and the atomic mass to determine the mass of a unit cell. Calculate volume using the edge length given in the problem statement. Density is then mass divided by volume (d = m/V). Be sure to make all necessary unit conversions. 

Setup The mass of an Ir atom is 192.2 amu. The conversion factor from amu to grams is 

Practice Problem A Aluminum metal crystallizes in a face-centered cubic unit cell. If the length of the cell edge is 404 pm, what is the density of aluminum in g/cm  

Gf Practice Problem B Copper crystallizes in a face-centered cubic lattice. If the density of the metal is 8.96 g/cm, what is the length of the unit cell edge in picometers  

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As the conditions of pressure and temperature vary, the phases in which hydrocarbons exist, and the composition of the phases may change. It is necessary to understand the initial condition of fluids to be able to calculate surface volumes represented by subsurface hydrocarbons. It is also necessary to be able to predict phase changes as the temperature and pressure vary both in the reservoir and as the fluids pass through the surface facilities, so that the appropriate subsurface and surface development plans can be made. 

The application of load in materials produces internal modifications such as crack growth, local plastic deformation, corrosion and phase changes, which are accompanied by the emission of acoustic waves in materials. These waves therefore contain information on the internal behaviour of the material and can be analysed to obtain this information. The waves are detected by the use of suitable sensors, that converts the surface movements of the material into electric signal. These signals are processed, analysed and recorded by an appropriate instrumentation. 

So in order to improve selective characteristics of eddy current testing one should minimize phase change under interference factors influence. Analysis of the above characteristics has indicated that in case of interacting under-surface defects, there is an optimal frequency providing the best sensitivity to defect in amplitude. 

A much better way would be to use phase contrast, rather than attenuation contrast, since the phase change, due to changes in index of refraction, can be up to 1000 times larger than the change in amplitude. However, phase contrast techniques require the disposal of monochromatic X-ray sources, such as synchrotrons, combined with special optics, such as double crystal monochromatics and interferometers [2]. Recently [3] it has been shown that one can also obtain phase contrast by using a polychromatic X-ray source provided the source size and detector resolution are small enough to maintain sufficient spatial coherence. 

After amplification both signals change their initial phases due to the delay r of the amplifier unblank (r = 0.1 - 0.5 ms), phase shift in it and wave propagation in passive vibrator s elements. All the mentioned phase changes are proportional to the frequency. The most contribution of them has unblank delay z. Thus frequency variations changes the initial phases) f/, and j(/c) of both signals and their difference A - Vi ... 

Above 81.5 K the C(2x 1) structure becomes the more stable. Two important points are, first, that a change from one surface structure to another can occur without any bulk phase change being required and, second, that the energy difference between dtemative surface structures may not be very large, and the free energy difference can be quite temperature-dependent. 

On the other hand, as applied to the submonolayer region, the same comment can be made as for the localized model. That is, the two-dimensional non-ideal-gas equation of state is a perfectly acceptable concept, but one that, in practice, is remarkably difficult to distinguish from the localized adsorption picture. If there can be even a small amount of surface heterogeneity the distinction becomes virtually impossible (see Section XVll-14). Even the cases of phase change are susceptible to explanation on either basis. 

Restructuring of a surface may occur as a phase change with a transition temperature as with the Si(OOl) surface [23]. It may occur on chemisorption, as in the case of oxygen atoms on a stepped Cu surface [24]. The reverse effect may occur The surface layer for a Pt(lOO) face is not that of a terminal (100) plane but is reconstructed to hexagonal symmetry. On CO adsorption, the reconstruction is lifted, as shown in Fig. XVI-8. 

As LEED studies have shown, the stmcture of a chemisorbed phase can change with 6. In terms of transition state theory, we can write A = (I/tq) and a common observation is that while E may change with a phase change, AS will tend to change also, and similarly. The result, again known as a compensation effect, is that the product remains relatively constant... 

Berry R S 1999 Phases and phase changes of small systems Theory of Atomio and Moleoular Clusters ed J Jelllnek (Berlin Springer)... 

If, in going from 0 K to T, a substance undergoes phase changes (fusion, vaporization, etc) at and Tg with molar enthalpies of transition AHy, one can write... 

The preparation of the reflecting silver layers for MBI deserves special attention, since it affects the optical properties of the mirrors. Another important issue is the optical phase change [ ] at the mica/silver interface, which is responsible for a wavelength-dependent shift of all FECOs. The phase change is a fimction of silver layer thickness, T, especially for T < 40 mn [54]. The roughness of the silver layers can also have an effect on the resolution of the distance measurement [59, 60]. 

Farreii B, Baiiey A i and Chapman D 1995 Experimentai phase changes at the mica-siiver interface iiiustrate the experimentai accuracy of the centrai fiim thickness in a symmetricai three-iayer interferometer App/. Opt. 34 2914-20... 

Calorimetry is the basic experimental method employed in thennochemistry and thennal physics which enables the measurement of the difference in the energy U or enthalpy //of a system as a result of some process being done on the system. The instrument that is used to measure this energy or enthalpy difference (At/ or AH) is called a calorimeter. In the first section the relationships between the thennodynamic fiinctions and calorunetry are established. The second section gives a general classification of calorimeters in tenns of the principle of operation. The third section describes selected calorimeters used to measure thennodynamic properties such as heat capacity, enthalpies of phase change, reaction, solution and adsorption. 

Unlike melting and the solid-solid phase transitions discussed in the next section, these phase changes are not reversible processes they occur because the crystal stmcture of the nanocrystal is metastable. For example, titania made in the nanophase always adopts the anatase stmcture. At higher temperatures the material spontaneously transfonns to the mtile bulk stable phase [211, 212 and 213]. The role of grain size in these metastable-stable transitions is not well established the issue is complicated by the fact that the transition is accompanied by grain growth which clouds the inteiyDretation of size-dependent data [214, 215 and 216]. In situ TEM studies, however, indicate that the surface chemistry of the nanocrystals play a cmcial role in the transition temperatures [217, 218]. 

The ability to control pressure in the laboratory environment is a powerful tool for investigating phase changes in materials. At high pressure, many solids will transfonn to denser crystal stmctures. The study of nanocrystals under high pressure, then, allows one to investigate the size dependence of the solid-solid phase transition pressures. Results from studies of both CdSe [219, 220, 221 and 222] and silicon nanocrystals [223] indicate that solid-solid phase transition pressures are elevated in smaller nanocrystals. 

In three-dimensional (3D) applications the overall phase change over a cycle may therefore be expressed as a surface integral, analogous to Eq. (43), namely. 

Single-surface calculation with phase change. 

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