Melting of ice

Variations in cooling load can be provided from the latent heat of melting of ice or a frozen eutectic. Ice can be formed by allowing it... 

The melting of ice absorbs heat from the surroundings, which might be the water in a glass of iced tea. The temperature of the surroundings drops, perhaps from 25°C to 3°C, as they give up heat to the system. 

On the other hand, this simple rule fails for many familiar phase changes. An example of a spontaneous reaction that is not exothermic is the melting of ice. This takes place spontaneously at 1 atm above 0°C, even though it is endothermic... 

As we saw in Section 5.1, a single substance can exist in different phases, or physical forms. The phases of a substance include its solid, liquid, and gaseous forms and its different solid forms, such as the diamond and graphite phases of carbon. In one case—helium—two liquid phases are known to exist. The conversion of a substance from one phase into another, such as the melting of ice, the vaporization of water, and the conversion of graphite into diamond, is called a phase transition (recall Section 6.11). 

Spontaneous processes result in the dispersal of matter and energy, hi many cases, however, the spontaneous direction of a process may not be obvious. Can we use energy changes to predict spontaneity To answer that question, consider two everyday events, the melting of ice at room temperature and the formation of ice in a freezer. 

The melting of ice is the reverse of the freezing of water. Energy becomes more constrained as it is transferred from the air in the room to the melting ice. At the same time, the molecules in the ice cube become less constrained, because they are free to move about in the liquid phase. Melting disperses matter but constrains energy. 

Curve A After cooling at 30 °C/min down to -100 °C, the DSC plots have been recorded during rewarming with 5 °C/min. Tg approx. -85 °C, respectively -88 °C. At approx. -48 °C, respectively -44 °C, ice crystallization starts clearly, followed by the beginning of the melting of ice. (During freezing only a part of the water has been crystallized.)... 

The variation of the phase transition temperature with pressure can be calculated from the knowledge of the volume and enthalpy change of the transition. Most often both the entropy and volume changes are positive and the transition temperature increases with pressure. In other cases, notably melting of ice, the density of the liquid phase is larger than of the solid, and the transition temperature decreases... 

The aqueous fluids formed by melting of ices in asteroids reacted with minerals to produce a host of secondary phases. Laboratory studies provide information on the identities of these phases. They include hydrated minerals such as serpentines and clays, as well as a variety of carbonates, sulfates, oxides, sulfides, halides, and oxy-hydroxides, some of which are pictured in Figure 12.15. The alteration minerals in carbonaceous chondrites have been discussed extensively in the literature (Zolensky and McSween, 1988 Buseck and Hua, 1993 Brearley, 2004) and were most recently reviewed by Brearley (2006). In the case of Cl chondrites, the alteration is pervasive and almost no unaltered minerals remain. CM chondrites contain mixtures of heavily altered and partially altered materials. In CR2 and CV3oxb chondrites, matrix minerals have been moderately altered and chondrules show some effects of aqueous alteration. For other chondrite groups such as CO and LL3.0-3.1, the alteration is subtle and secondary minerals are uncommon. In some CV chondrites, a later thermal metamorphic overprint has dehydrated serpentine to form olivine. 

Asteroids that formed beyond the snowline represent rock and ice accreted inside the orbit of Jupiter. The most distant asteroids may still contain ices, but many asteroids have been heated. Melting of ice produced aqueous fluids, which reacted with chondritic minerals at low temperatures to form secondary minerals (phyllosilicates, carbonates, sulfates, oxides). The alteration minerals can be discerned in asteroid spectra and characterized by analyses of chondrites derived from these bodies. 

Various thermal transitions can occur in rapidly cooled solutions, including glass transition, devitrification (ice formation on warming a rapidly-frozen solution) and melting of ice. The relationship between temperature, weight fraction of solids, solubility and glass transition of lactose is shown in Figure 7.16. 

Around 1800, experimental challenges to caloric theory were being presented by Count Rumford (cannon boring) and Humphrey Davy (melting of ice by friction). It became apparent that heat could be produced from a body in unlimited quantity by friction, further stretching its credibility as a substance. By about 1840, caloric theory was overturned by the modem kinetic molecular theory of heat (Sidebar 2.7), which identified heat with the energy of random molecular motions. 

Endothermic processes are common in our everyday experience. Examples include the melting of ice cubes in a glass of water, or the evaporation of sweat from our skin. In these cases, the greater entropy of the products favors the utilization of energy to allow these processes to occur (Case III). A dramatic example is a chemical cooling pack that contains a salt (usually ammonium nitrate,... 

To avoid the use of the ambiguous term "heat" in connection with "heat content," it is customary to use the term enthalpy. At a given temperature and pressure, every substance possesses a characteristic amount of enthalpy (H), and the heat changes associated with chemical and physical changes at constant pressure are called changes in enthalpy (AH) AHT is the enthalpy of transition. Two common enthalpies of transition are AHf = 1435 cal/mole for the enthalpy of fusion (melting) of ice at 0°C, and AH, = 9713 cal/mole for the enthalpy of vaporization of water at 100°C. 

Here q is an infinitesimal quantity of heat absorbed from the surroundings by the system and T is measured in kelvins (K). For a reversible phase transition such as the melting of ice at constant pressure and temperature, the change in entropy of the H20 is just AH/T. 

Entropy is measured in units of joules per kelvin (or °C) or calories per K, the latter sometimes being abbreviated as e. u. (entropy units). Since the melting of ice at 0°C is a reversible process, the second law asserts that the entropy of the surroundings decreases by the same amount that the entropy of the water increases. The value of T AS is numerically equal to the heat of fusion, 6.008 kj mol1 in the case of water at 0°C. Thus, the entropy increase in the ice as it melts at 0°C is 6.008 x 103 J/273.16 K = 22.0 J K1. 

We have seen that while many spontaneous processes, e.g., combustion of organic compounds, are accompanied by liberation of heat (negative AH), others are accompanied by absorption of heat from the surroundings (positive AH). An example of the latter is the melting of ice at a temperature just above 0°C, during which there is a large increase in the entropy of the water. As we have seen, at 0°C at equilibrium T AS is just equal to -AH (Eq. 6-7). 

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