Entropy of surroundings

There exists a state function S, called the entropy of a system, related to the heat Dq absorbedfrom the surroundings during an infinitesimal change by the relations... 

Hq and Sq = enthalpy and entropy of the same stream at equiUbrium with the surroundings and Tq = temperature of the surroundings (sink). 

Ion-Dipole Forces. Ion-dipole forces bring about solubihty resulting from the interaction of the dye ion with polar water molecules. The ions, in both dye and fiber, are therefore surrounded by bound water molecules that behave differently from the rest of the water molecules. If when the dye and fiber come together some of these bound water molecules are released, there is an increase in the entropy of the system. This lowers the free energy and chemical potential and thus acts as a driving force to dye absorption. 

The second law reqmres that the entropy of an isolated system either increase or, in the limit, where the system has reached an equilibrium state, remain constant. For a closed (but not isolated) system it requires that any entropy decrease in either the system or its surroundings be more than compensated by an entropy increase in the other part or that in the Emit, where the process is reversible, the total entropy of the system plus its surroundings be constant. 

To understand the flow in turbomachines, an understanding of the basic relationships of pressure, temperature, and type of flow must be acquired. Ideal flow in turbomachines exists when there is no transfer of heat between the gas and its surroundings, and the entropy of the gas remains unchanged. This type of flow is characterized as a rever.sible adiabatic flow. To describe this flow, the total and static conditions of pressure, temperature, and the concept of an ideal gas must be understood. 

The entropy of the system plus surroundings is unchanged by reversible processes the entropy of the system plus surroundings increases for irreversible processes. 

Thus, in adiabatic processes the entropy of a system must always increase or remain constant. In words, the second law of thermodynamics states that the entropy of a system that undergoes an adiabatic process can never decrease. Notice that for the system plus the surroundings, that is, the universe, all processes are adiabatic since there are no surroundings, hence in the universe the entropy can never decrease. Thus, the first law deals with the conservation of energy in any type of process, while the sec-... 

In some cases, an alternative explanation is possible. It may be assumed that any very complex organic counterion can also interact with the CP matrix with the formation of weak non-ionic bonds, e.g., dipole-dipole bonds or other types of weak interactions. If the energy of these weak additional interactions is on the level of the energy of the thermal motion, a set of microstates appears for counterions and the surrounding CP matrix, which leads to an increase in the entropy of the system. The changes in Gibbs free energy of this interaction may be evaluated in a semiquantitative way [15]. 

If any cyclic process is performed with a given material system, the entropy of all the surrounding bodies which have in any way been involved in the process, either as emitters or absorbers of heat, either remains unchanged, if the cycle is reversible, or else increases, if the cycle is performed irreversibly. 

Adiabatic processes are examples of (d). If a mole of ideal gas is allowed to expand adiabatically into an evacuated bulb to twice its initial volume, the entropy of the gas increases by 5.76 J K-1 mol-1. No entropy change occurs in the surroundings, since there is no exchange of heat. Hence, 5.76 J K-1 mol-1 is the net increase in entropy in the universe. 

Entropy is an important concept in chemistry because we can use it to predict the natural direction of a reaction. However, not only does the entropy of the reaction system change as reactants form products, but so too does the entropy of the surroundings as the heat produced or absorbed by the reaction enters or leaves them. Both the entropy change of the system and that of the surroundings affect the direction of a reaction, because both contribute to the entropy of the universe. We explore the contribution of the system in this section and the contribution of the surroundings in the next section. 

An example of the role of the surroundings in determining the spontaneous direction of a process is the freezing of water. We can see from Table 7.2 that, at 0°C, the molar entropy of liquid water is 22.0 J-K 1-mo -1 higher than that of ice... 

This important formula, which can be derived more formally from the laws of thermodynamics, applies when any change takes place at constant pressure and temperature. Notice that, for a given enthalpy change of the system (that is, a given output of heat), the entropy of the surroundings increases more if their temperature is low than if it is high (Fig. 7.16). The explanation is the sneeze in the street analogy mentioned in Section 7.2. Because AH is independent of path, Eq. 10 is applicable whether the process occurs reversibly or irreversibly. 

EXAMPLE 7.10 Sample exercise Calculating the change in entropy of the surroundings... 

SOLUTION We can expect the entropy of the surroundings to increase when water freezes because the heat released stirs up the thermal motion of the atoms in the surroundings (Fig. 7.17). To calculate the change in entropy of the surroundings when water freezes we write AHfrceze = —AHflls = —6.0 kj-mol 1 and T = ( — 10. + 273) K = 263 K. Then... 

Note that the increase in entropy of the surroundings is indeed greater than the decrease in entropy of the system, —22.0 J-K, the value at 0°C (the value at —10.°C is similar), so the overall change is positive and the freezing of water is spontaneous at -10.°C. 

FIGURE 7.17 (a) In an exothermic process, heat escapes into the surroundings and increases their entropy, (b) In an endothermic process, the entropy of the surroundings decreases. The red arrows represent the transfer of heat between system and surroundings, and the green arrows indicate the entropy change of the surroundings. 

As we have already emphasized, to use the entropy to judge the direction of spontaneous change, we must consider the change in the entropy of the system plus the entropy change in the surroundings ... 

In an exothermic reaction, such as the synthesis of ammonia or a combustion reaction, the heat released by the reaction increases the disorder of the surroundings. In some cases, the entropy of the system may decrease, as when a gaseous reactant is converted into a solid or liquid. However, provided that AH is large and negative, the release of energy as heat into the surroundings increases their entropy so much that it dominates the overall change in entropy and the reaction is spontaneous (Fig. 7.18). 

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