СЕМЕСТРОВАЯ РАБОТА
(10000 ЗНАКОВ)
Выполнила:
студент гр. ВАУ-227
Гранкина Н.В.
Проверила:
ст. преподаватель
Галицына Т.А.
ВОЛЖСКИЙ 2008
THERMODYNAMICS AND THERMOCHEMISTRY
Thermodynamics and Energy.
The energy of a body may be defined broadly as its capacity for doing work. This energy may take various forms, such as kinetic energy of a body in motion, potential energy due to position, heat energy as measured by the temperature, electrical energy, chemical energy, etc. Chemical and physical processes are almost invariably accompanied by energy changes, and results of considerable importance have been obtained studying the laws underlying these changes. It is this study of energy transformation which constitutes the subject matter of thermodynamics. Although thermodynamics may appear to be somewhat theoretical in nature, the two laws have led to results of fundamental importance to chemistry, as well as to physics.
Conservation of Energy: The First Law of Thermodynamics.
Many attempts have been made from time to time to realize "perpetual motion", that is, the continuous production of mechanical work without supplying an equivalent amount of energy from another source. The failure of all such efforts has led to the universal acceptance of the principle of conservation of energy. This principle has been stated in many forms, but essentially they amount to the fact that although energy can be converted from one form to another, it cannot be created or destroyed or, alternatively, whenever a quantity of one kind of energy is produced, an exactly equivalent amount of other kinds must disappear. It is evident that perpetual motion, in the generally accepted sense of the term, would be contrary to this principle, for it would involve the creation of energy. Further, the exact equivalence of mechanical or electrical work and heat, as found by Joule and others, is a necessary consequence of the same principle.
The law of conservation of energy is purely the result of experience, no exception to it having as yet been found. The assumption that it is of universal applicability is the basis of the first law of thermodynamics. This law can be stated in any of the ways given above for the principle of the conservation of energy, or else it may be put in the following form. The total energy of a system and its surroundings must remain constant, although it may be changed from one form, to another.
Heat Changes in Chemical Reactions.
The subject of thermochemistry deals with the heat changes accompanying chemical reactions. As will be seen shortly the laws of thermochemistry are based-largely on the principle of the conservation of energy or the first law of thermodynamics. Different substances have different amounts of internal (chemical) energy, and so the total energy of the products of a reaction is generally different from that of the reactants; hence, the chemical change will be accompanied by the liberation or absorption of energy, which may appear in the form of heat. If heat is liberated in the reaction the process is said to be exothermic, but if heat is absorbed it is described as endothermic. The majority of, although not all, chemical reactions which go to virtual completion at ordinary temperatures are exothermic in character, since they are accompanied by an evolution of heat. If a chemical reaction is associated with a volume change, as is particularly the case for many processes involving the combination of gases, the magnitude of the heat change will depend on whether the reaction is carried out at constant pressure or at constant volume. Since many reactions normally occur at constant (atmospheric) pressure it is the usual practice to record heat changes by quoting the value of qp, the heat absorbed at constant pressure; this may, of course, be identified with ΔH, theincrease of heat content under the same conditions. This quantity is often referred to as the heat of reaction; it represents the difference in the heat contents of the reaction products and of the reactants, at constant pressure and at definite temperature, with every substance in a definite physical state. From the value of qp (or ΔH) the value of gv (or ΔE) can be readily determined if the volume change ΔV at the constant pressure P is known as will be seen below.
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The heat change accompanying a reaction, for example, that between solid carbon (graphite) and gaseous oxygen to yield carbon dioxide gas, is represented in the form of a thermochemical equation, as follows:
C(s) + 02(g) = C02 (g) ΔH = -94.00 kcal.
This means that when 12.01 grams of solid carbon (graphite) and 32 grams of gaseous oxygen react completely to yield 44.01 grams of gaseous carbon dioxide, at constant pressure, there is a decrease in heat content, since ΔH is negative, of 94 kilocalories (kcal.), i. e., 94,000 calories. It is the general practice in modern thermochemical work to express results in kilocalories because the statement of heat changes in calories implies an accuracy greater than is usually attainable experimentally. It should be noted, incidentally, that the ΔH (or ΔE) values always refer to completed reactions, appropriate allowance having been made, if necessary, if the process does not normally go to completion.
The symbols g, /, and s, placed in parentheses after the formula indicate whether the substance taking part in the reaction is gas, liquid or solid. Reactions taking place in aqueous solution are indicated by the symbol aq;thus,
HCl (aq) + NaOH (aq) = NaCl (aq) + H2O)
ΔH = 13.70 kcal.
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Strictly speaking the use of aq implies that the reaction is occurring in such dilute solution that the addition of further water causes no detectable heat change.
A negative value of ΔH, as in the two instances quoted above, means that the reaction is accompanied by a decrease in heat content; that is to say, the heat content of the products is less than that of the reactants at a specified temperature, in other words, the reaction at the given temperature is associated with an evolution of heat. It follows, therefore, that when ΔH is negative the reaction is exothermic; similarly, if ΔH is positive the process is endothermic. The same conclusions can be reached directly from the fact that qp, which is equal to ΔH, the heat absorbed in the reaction; hence, when ΔH is negative heat is actually evolved.
Spontaneous Processes.
The second law of thermodynamics has led to results which are of considerable importance to chemistry, physics and engineering, but to the chemist its greatest value probably lies in the fact that it provides a means of foretelling whether a particular reaction can occur, and if so to what extent. However, thermodynamics can only indicate if the reaction is possible or not; other considerations which lie outside thermodynamics are necessary to determine whether the process will take place slowly or rapidly. Even with this limitation in mind, it must be admitted that information concerning the fundamental possibility of a reaction, apart from its speed, would be of great interest to the chemist. At one time it was believed that chemical changes always occurred spontaneously in the direction of heat evolution that is, in the direction leading to a decrease in the heat content. This conclusion is, however, manifestly incorrect, as is evident from the fact that many reactions which take place spontaneously are known to involve an absorption of heat.
The question being considered resolves itself into the problem of understanding the conditions under which spontaneous processes in general take place. It is convenient in this connection to examine some physical processes that are of spontaneous occurrence; the conclusions drawn are found to be applicable to all changes that tend to take place without external influence. Consider, for example, a bar of metal that is hot at one end and cold at the other; heat will be conducted spontaneously along the bar from the hot end to the cold end until the temperature is uniform. It is important to note, however, that this process is not found to reverse itself spontaneously; it has not been observed that a metal bar of uniform temperature spontaneously becomes hotter at one end and colder at the other.