echecs16.info Gratis INTRODUCTION TO THERMODYNAMICS PDF

# INTRODUCTION TO THERMODYNAMICS PDF

thermodynamic equilibrium is considered. Keywords: thermodynamics, expansion of gas, heat machines, energy, phase transition. 1. Introduction. COURSE INTRODUCTION. Course Learning Objectives: To be able to use the First Law of Thermodynamics to estimate the potential for thermo- mechanical. Introduction to Thermodynamics. 1. Introduction. History of Thermodynamics. 2. The First Law. Microscopic view. Joule. 3. The Second Law. Microscopic View. Author: JACQUILINE COPLIN Language: English, Spanish, German Country: Switzerland Genre: Technology Pages: 122 Published (Last): 02.03.2015 ISBN: 247-5-57515-724-4 ePub File Size: 19.74 MB PDF File Size: 8.16 MB Distribution: Free* [*Register to download] Downloads: 45261 Uploaded by: TRACY Thermodynamics. Training Centre / Centre de formation. Introduction to Thermodynamics. Training Objectives. The participant will be introduced to: basic. INTRODUCTION. What is thermodynamics? Thermodynamics is the science which has evolved from the original investiga- tions in the 19th century into the. APPLICATIONS TO NON-HOMOGENEOUS SYSTEMS. PART III. THE SECOND FUNDAMENTAL PRINCIPLE OF. THERMODYNAMICS. I. INTRODUCTION.

However, it would be too cumbersome to avoid referring to these scales until then, though formally that could be done, and everyone is aware of the Celsius centigrade and Fahrenheit scales. As we shall see, however, there is an absolute zero of temperature, a zero that cannot be passed and where negative 7 The zeroth law: The concept of temperature 3. Three common temperature scales showing the relations between them. The vertical dotted line on the left shows the lowest achievable temperature; the two dotted lines on the right show the normal freezing and boiling points of water The Laws of Thermodynamics temperatures have no meaning except in a certain formal sense, not one that depends on the technology of the time see Chapter 5. It is therefore natural to measure temperatures by setting 0 at this lowest attainable zero and to refer to such absolute temperatures as the thermodynamic temperature.

From a molecular viewpoint, the raising of a weight corresponds to all its atoms moving in the same direction. Thus, when a block of iron is raised, all the iron atoms move upwards uniformly. When the block is lowered—and does work on the system, like compressing a spring or a gas, and increases its internal energy—all its atoms move downwards uniformly. Work is the transfer of energy that makes use of the uniform motion of atoms in the surroundings Figure 8.

Now, the molecular nature of heat. Once inside, the energy is stored as the kinetic energy the energy due to motion and the potential energy the energy due to position of the constituent atoms, and that energy can be withdrawn either as heat or as work. The distinction between work and heat is made in the surroundings: the system has no memory of the mode of transfer nor is it concerned about how its store of energy will be used.

This blindness to the mode of transfer needs a little further explanation. Thus, if a gas in an adiabatic container is compressed by a falling weight, the incoming piston acts like a bat in a microscopic game of table-tennis. When a molecule strikes the piston, it is accelerated. In more pictorial terms, a block of iron at high temperature consists of atoms that are oscillating vigorously around their average positions.

At low temperatures, the atoms continue to oscillate, but with less vigour. If a hot block of iron is put in contact with a cooler block, the vigorously oscillating atoms at the edge of the hot block jostle the less vigorously oscillating atoms at the edge of the cool block into more vigorous motion, and they pass on their energy by jostling their neighbours. There is no net motion of either block, but energy is transferred from the hotter to the cooler block by this random jostling where the two blocks are in contact.

That is, heat is the transfer of energy that makes use of the random motion of atoms in the surroundings Figure 8. The Laws of Thermodynamics sample of gas is heated, the random jostling of the atoms in the surroundings stimulates the gas molecules into more vigorous motion, and the acceleration of the molecules at the thermally conducting walls is quickly distributed over the entire sample. The result within the system is the same.

## Introduction to Chemical Engineering Thermodynamics

We can now return to the faintly enigmatic remark made earlier that an electric heater is better regarded as an electric worker. The electrons of that current collide with the atoms of the wire and cause them to wobble around their mean positions. That is, the energy—and the temperature—of the coil of wire is raised by doing work on it.

So, although we do work on the heater itself, that work is translated into heating the system: worker has become heater. Fire preceded the harnessing of fuels to achieve work.

Work is energy tamed, and requires greater sophistication to contrive. Introducing reversibility The originators of thermodynamics were subtle people, and quickly realized that they had to be careful when specifying how a process is carried out.

The expansion is reversible in the colloquial sense but not in the thermodynamic sense. The transfer of energy as heat is not reversible in the thermodynamic sense in this instance.

However, now consider the case in which the external pressure matches the pressure of the gas in the system exactly. As we saw in Chapter 1, we say that the system and its surroundings are in mechanical equilibrium.

The expansion is reversible in the thermodynamic sense. Likewise, consider a system at the same temperature as the surroundings. In everyday language, a reversible process is one that can be reversed. Thus, the rolling of a wheel can be reversed, so in principle a journey can be traversed in reverse. The compression of a gas can be reversed by pulling out the piston that effected the compression.

The Laws of Thermodynamics thermal equilibrium. The transfer of energy as heat is reversible in the thermodynamic sense in this instance.

The greatest amount of work is done if the expansion of a gas is reversible at every stage. The pressure of the gas falls a little because it now occupies a greater volume.

This process of effectively matching the external pressure to the falling pressure of the gas continues until the piston has moved out a desired amount and, through its coupling to a weight, has done a certain amount of work.

## Introduction: Why Thermodynamics?

That is, by ensuring that at every stage the expansion is reversible in the thermodynamic sense, the system does maximum work. This conclusion is general: reversible changes achieve maximum work. We shall draw on this generalization in the following chapters. Introducing enthalpy Thermodynamicists are also subtle in their discussion of the quantity of heat that can be extracted from a system, such as when a fuel burns. We can appreciate the problem as follows. As the fuel burns it produces carbon dioxide and water vapour, which occupy more space than the original fuel and oxygen, and as a result the piston is driven out to 28 Thermodynamicists have developed a clever way of taking into account the energy used to do work when any change, and particularly the combustion of a fuel, occurs, without having to calculate the work explicitly in each case.

To do so, they switch attention from the internal energy of a system, its total energy content, to a closely related quantity, the enthalpy symbol H.

It turns out that the energy released as heat by a system free to expand or contract as a process occurs, as distinct from the total energy released in the same process, is exactly equal to the change in enthalpy of the system. This expansion requires work. That is, when a fuel burns in a container that is free to expand, some of the energy released in the combustion is used to do work. If the combustion takes place in a container with rigid walls, the combustion releases the same amount of energy, but none of it is used to do work because no expansion can occur.

In other words, more energy is available as heat in the latter case than in the former. To calculate the heat that can be produced in the former case, we have to account for the energy that is used to make room for the carbon dioxide and water vapour and subtract that from the total change in energy. This is true even if there is no physical piston—if the fuel burns in a dish—because, although we cannot see it so readily, the gaseous products must still make room for themselves.

The Laws of Thermodynamics actually by mathematics—the leakage of energy from a system as work is automatically taken into account by focusing on the change in enthalpy.

In other words, the enthalpy is the basis of a kind of accounting trick, which keeps track invisibly of the work that is done by the system, and reveals the amount of energy that is released only as heat, provided the system is free to expand in an atmosphere that exerts a constant pressure on the system. It follows that if we are interested in the heat that can be generated by the combustion of a fuel in an open container, such as a furnace, then we use tables of enthalpies to calculate the change in enthalpy that accompanies the combustion.

Then we identify that change with the heat generated by the system. As an actual example, the change of enthalpy that accompanies the combustion of a litre of gasoline is about 33 megajoules 1 megajoule, written 1 MJ, is 1 million joules. Therefore we know without any further calculation that burning a litre of gasoline in an open container will provide 33 MJ of heat. A deeper analysis of the process shows that in the same combustion, the system has to do about kJ where 1 kilojoule, written 1 kJ, is one thousand joules of work to make room for the gases that are generated, but that energy is not available to us as heat.

We could extract that extra kJ, which is enough to heat about half a litre of water from room temperature to its boiling point, if we prevent the gases from expanding so that all the energy released in the combustion is liberated as heat.

One way to achieve that, and to obtain all the energy as heat, would be to arrange for the combustion to take place in a closed container with rigid walls, in which case it would be unable to expand and hence would be unable to lose any energy as work. In practice, it is technologically much simpler to use furnaces that are open to the atmosphere, 30 and in practice the difference between the two cases is too small to be worth the effort.

However, in formal thermodynamics, which is a precise and logical subject, it is essential to do all the energy accounting accurately and systematically. In formal thermodynamics the differences between changes in internal energy and enthalpy must always be borne in mind. This energy is commonly supplied in the form of heat—that is, by making use of a temperature difference between the liquid and its surroundings.

The scalding effect of steam is an illustration. The enthalpy of vaporization of 1 g of water is close to 2 kJ. The condensation of 1 g of steam therefore releases 2 kJ of heat, which may be enough to destroy the proteins of our skin where it comes in contact. Gram-for-gram, the enthalpy of fusion is much less than the enthalpy of vaporization, and we do not get scalded by touching water that is freezing to ice. The Laws of Thermodynamics As the temperature of a system is raised and the Boltzmann distribution acquires a longer tail, populations migrate from states of lower energy to states of higher energy.

Consequently, the average energy rises, for its value takes into account the energies of the available states and the numbers of molecules that occupy each one. In other words, as the temperature is raised, so the internal energy rises. The slope of a graph of the value of the internal energy plotted against temperature is called the heat capacity of the system symbol C.

The supply of 1 J of energy as heat to 1 g of water results in an increase in temperature of about 0. Substances with a high heat capacity water is an example require a larger amount of heat to bring about a given rise in temperature than those with a small heat capacity air is an example. For instance, if the heating takes place under conditions of constant pressure with the sample free to expand, then some of the energy supplied as heat goes into expanding the sample and therefore to doing work.

Less energy remains in the sample, so its temperature rises less than when it is constrained to have a constant volume, and therefore we report that its heat capacity is higher. Heat capacities vary with temperature.

In most cases, the spread of populations increases with increasing temperature, so the heat capacity typically increases with rising temperature, as is observed.

## Introduction to thermodynamics of irreversible processes

However, the relationship is a little more complex than that because it turns out that the role of the spread of populations decreases as the temperature rises, so although that spread increases, the heat capacity does not increase as fast.

Heat capacity is a dissipation term: it is a measure of the ability of a substance to absorb the energy supplied to it as heat. At higher temperatures, the populations are spread over a range of states and hence the heat capacity is non-zero, as is observed.

The Laws of Thermodynamics to the heat capacity, and the heat capacity settles into a constant value. This is the case for the contribution of all the basic modes of motion: translation motion through space , rotation, and vibration of molecules, all of which settle into a constant value.

Broadly speaking, the energy levels lie close together when the atoms are heavy. As we saw in Chapter 1, we say that the system and its surroundings are in mechanical equilibrium. The expansion is reversible in the thermodynamic sense. Likewise, consider a system at the same temperature as the surroundings.

In everyday language, a reversible process is one that can be reversed. Thus, the rolling of a wheel can be reversed, so in principle a journey can be traversed in reverse. The compression of a gas can be reversed by pulling out the piston that effected the compression.

The Laws of Thermodynamics thermal equilibrium. The transfer of energy as heat is reversible in the thermodynamic sense in this instance.

The greatest amount of work is done if the expansion of a gas is reversible at every stage. The pressure of the gas falls a little because it now occupies a greater volume. This process of effectively matching the external pressure to the falling pressure of the gas continues until the piston has moved out a desired amount and, through its coupling to a weight, has done a certain amount of work. That is, by ensuring that at every stage the expansion is reversible in the thermodynamic sense, the system does maximum work.

This conclusion is general: reversible changes achieve maximum work. We shall draw on this generalization in the following chapters. Introducing enthalpy Thermodynamicists are also subtle in their discussion of the quantity of heat that can be extracted from a system, such as when a fuel burns.

We can appreciate the problem as follows. As the fuel burns it produces carbon dioxide and water vapour, which occupy more space than the original fuel and oxygen, and as a result the piston is driven out to 28 Thermodynamicists have developed a clever way of taking into account the energy used to do work when any change, and particularly the combustion of a fuel, occurs, without having to calculate the work explicitly in each case.

To do so, they switch attention from the internal energy of a system, its total energy content, to a closely related quantity, the enthalpy symbol H. It turns out that the energy released as heat by a system free to expand or contract as a process occurs, as distinct from the total energy released in the same process, is exactly equal to the change in enthalpy of the system. This expansion requires work.

That is, when a fuel burns in a container that is free to expand, some of the energy released in the combustion is used to do work. If the combustion takes place in a container with rigid walls, the combustion releases the same amount of energy, but none of it is used to do work because no expansion can occur.

In other words, more energy is available as heat in the latter case than in the former. To calculate the heat that can be produced in the former case, we have to account for the energy that is used to make room for the carbon dioxide and water vapour and subtract that from the total change in energy.

This is true even if there is no physical piston—if the fuel burns in a dish—because, although we cannot see it so readily, the gaseous products must still make room for themselves. The Laws of Thermodynamics actually by mathematics—the leakage of energy from a system as work is automatically taken into account by focusing on the change in enthalpy. In other words, the enthalpy is the basis of a kind of accounting trick, which keeps track invisibly of the work that is done by the system, and reveals the amount of energy that is released only as heat, provided the system is free to expand in an atmosphere that exerts a constant pressure on the system.

It follows that if we are interested in the heat that can be generated by the combustion of a fuel in an open container, such as a furnace, then we use tables of enthalpies to calculate the change in enthalpy that accompanies the combustion. Then we identify that change with the heat generated by the system. As an actual example, the change of enthalpy that accompanies the combustion of a litre of gasoline is about 33 megajoules 1 megajoule, written 1 MJ, is 1 million joules.

Therefore we know without any further calculation that burning a litre of gasoline in an open container will provide 33 MJ of heat. A deeper analysis of the process shows that in the same combustion, the system has to do about kJ where 1 kilojoule, written 1 kJ, is one thousand joules of work to make room for the gases that are generated, but that energy is not available to us as heat.

We could extract that extra kJ, which is enough to heat about half a litre of water from room temperature to its boiling point, if we prevent the gases from expanding so that all the energy released in the combustion is liberated as heat.

One way to achieve that, and to obtain all the energy as heat, would be to arrange for the combustion to take place in a closed container with rigid walls, in which case it would be unable to expand and hence would be unable to lose any energy as work.

In practice, it is technologically much simpler to use furnaces that are open to the atmosphere, 30 and in practice the difference between the two cases is too small to be worth the effort. However, in formal thermodynamics, which is a precise and logical subject, it is essential to do all the energy accounting accurately and systematically.

In formal thermodynamics the differences between changes in internal energy and enthalpy must always be borne in mind. This energy is commonly supplied in the form of heat—that is, by making use of a temperature difference between the liquid and its surroundings. The scalding effect of steam is an illustration.

The enthalpy of vaporization of 1 g of water is close to 2 kJ. The condensation of 1 g of steam therefore releases 2 kJ of heat, which may be enough to destroy the proteins of our skin where it comes in contact. Gram-for-gram, the enthalpy of fusion is much less than the enthalpy of vaporization, and we do not get scalded by touching water that is freezing to ice. The Laws of Thermodynamics As the temperature of a system is raised and the Boltzmann distribution acquires a longer tail, populations migrate from states of lower energy to states of higher energy.

Consequently, the average energy rises, for its value takes into account the energies of the available states and the numbers of molecules that occupy each one. In other words, as the temperature is raised, so the internal energy rises. The slope of a graph of the value of the internal energy plotted against temperature is called the heat capacity of the system symbol C.

The supply of 1 J of energy as heat to 1 g of water results in an increase in temperature of about 0. Substances with a high heat capacity water is an example require a larger amount of heat to bring about a given rise in temperature than those with a small heat capacity air is an example.

For instance, if the heating takes place under conditions of constant pressure with the sample free to expand, then some of the energy supplied as heat goes into expanding the sample and therefore to doing work. Less energy remains in the sample, so its temperature rises less than when it is constrained to have a constant volume, and therefore we report that its heat capacity is higher.

Heat capacities vary with temperature. In most cases, the spread of populations increases with increasing temperature, so the heat capacity typically increases with rising temperature, as is observed.

However, the relationship is a little more complex than that because it turns out that the role of the spread of populations decreases as the temperature rises, so although that spread increases, the heat capacity does not increase as fast.

Heat capacity is a dissipation term: it is a measure of the ability of a substance to absorb the energy supplied to it as heat. At higher temperatures, the populations are spread over a range of states and hence the heat capacity is non-zero, as is observed. The Laws of Thermodynamics to the heat capacity, and the heat capacity settles into a constant value. This is the case for the contribution of all the basic modes of motion: translation motion through space , rotation, and vibration of molecules, all of which settle into a constant value.

Broadly speaking, the energy levels lie close together when the atoms are heavy. Moreover, translational energy levels are so close together as to form a near continuum, the rotational levels of molecules in gases are further apart, and vibrational energy levels—those associated with the oscillations of atoms within molecules—are widely separated.

Thus, as a gaseous sample is heated, the molecules are readily excited into higher translational states in English: they move faster and, in all practical cases, they quickly spread over many rotational states in English: they rotate faster. In each case the average energy of the molecules, and hence the internal energy of the system, increases as the temperature is raised. The molecules of solids are free neither to move through space nor to rotate. However, they can oscillate around their average positions, and take up energy that way.

These collective wobblings of the entire solid have much lower frequencies than the oscillations of atoms within molecles and so they can be excited much more readily. As energy is supplied to a solid sample, these oscillations are excited, the populations of the higher energy states increase as the Boltzmann distribution reaches to higher levels, and we report that the temperature of the solid has risen. Similar remarks apply to liquids, in which molecular motion is less constrained than in solids.

Water has a very high heat capacity, which means that to raise its temperature takes a lot of energy. Conversely, hot water stores a lot of energy, which is why it is such a good medium for central heating systems as well as being 34 cheap , and why the oceans are slow to heat and slow to cool, with important implications for our climate.

Conservation laws—laws that state that a certain property does not change—have a very deep origin, which is one reason why scientists, and thermodynamicists in particular, get so excited when nothing happens. Thus, conservation laws are based on various aspects of the shape of the universe we inhabit. In the particular case of the conservation of energy, the symmetry is that of the shape of time. It is much harder to give a molecular interpretation of enthalpy because it is a property contrived to do the bookkeeping of expansion work and is not as fundamental a property as internal energy.

For the purposes of this account, it is best to think of the enthalpy as a measure of the total energy, but to bear in mind that that is not exactly true. In short, as the temperature of a system is raised its molecules occupy higher and higher energy levels and as a result their mean energy, the internal energy, and the enthalpy all increase. Precise fundamental molecular interpretations can be given only of the fundamental properties of a system, its temperature, its internal energy, and—as we shall see in the next chapter—the entropy. Time is a uniformly structured coordinate.

If time were to bunch up and spread out, energy would not be conserved. I hope that you will see in the course of this chapter why I take that view, and perhaps go so far as to agree with me. Indeed, the novelist and former chemist C. Snow is famous for having asserted in his The Two Cultures that not knowing the second law of thermodynamics is equivalent to never having read a work by Shakespeare. I actually have serious doubts about whether Snow understood the law himself, but I concur with his sentiments.

The second law is of central importance in the whole of science, and hence in our rational understanding of the universe, because it provides a foundation for understanding why any change occurs.

Thus, not only is it a basis for understanding why engines run and chemical reactions occur, but it is also a foundation for understanding those most exquisite consequences of chemical reactions, the acts of literary, artistic, and musical creativity that enhance our culture.

Likewise, the second law implies the existence of another thermodynamic property, the entropy symbol S. We shall elaborate this interpretation and show its consequences in the rest of the chapter. Hence it was numbered the zeroth law. The importance of the law as a foundation to the earlier laws is that it allows the definition of temperature in a non-circular way without reference to entropy, its conjugate variable.

Such a temperature definition is said to be 'empirical'. The first law of thermodynamics may be stated in several ways:. Combining these principles leads to one traditional statement of the first law of thermodynamics: Or more briefly, a perpetual motion machine of the first kind is impossible. The second law of thermodynamics indicates the irreversibility of natural processes, and, in many cases, the tendency of natural processes to lead towards spatial homogeneity of matter and energy, and especially of temperature.

It can be formulated in a variety of interesting and important ways. It implies the existence of a quantity called the entropy of a thermodynamic system.

In terms of this quantity it implies that. When two initially isolated systems in separate but nearby regions of space, each in thermodynamic equilibrium with itself but not necessarily with each other, are then allowed to interact, they will eventually reach a mutual thermodynamic equilibrium.

The sum of the entropies of the initially isolated systems is less than or equal to the total entropy of the final combination. Equality occurs just when the two original systems have all their respective intensive variables temperature, pressure equal; then the final system also has the same values. The second law is applicable to a wide variety of processes, reversible and irreversible.

All natural processes are irreversible. Reversible processes are a useful and convenient theoretical fiction, but do not occur in nature.

## Introduction to Thermodynamics of Mechanical Fatigue - CRC Press Book

A prime example of irreversibility is in the transfer of heat by conduction or radiation. It was known long before the discovery of the notion of entropy that when two bodies initially of different temperatures come into thermal connection, then heat always flows from the hotter body to the colder one. The second law tells also about kinds of irreversibility other than heat transfer, for example those of friction and viscosity, and those of chemical reactions. The notion of entropy is needed to provide that wider scope of the law.

Entropy may also be viewed as a physical measure of the lack of physical information about the microscopic details of the motion and configuration of a system, when only the macroscopic states are known. This lack of information is often described as disorder on a microscopic or molecular scale. The law asserts that for two given macroscopically specified states of a system, there is a quantity called the difference of information entropy between them.

This information entropy difference defines how much additional microscopic physical information is needed to specify one of the macroscopically specified states, given the macroscopic specification of the other — often a conveniently chosen reference state which may be presupposed to exist rather than explicitly stated.

A final condition of a natural process always contains microscopically specifiable effects which are not fully and exactly predictable from the macroscopic specification of the initial condition of the process. This is why entropy increases in natural processes — the increase tells how much extra microscopic information is needed to distinguish the final macroscopically specified state from the initial macroscopically specified state.

The third law of thermodynamics is sometimes stated as follows:. At zero temperature the system must be in a state with the minimum thermal energy. This statement holds true if the perfect crystal has only one state with minimum energy.

Entropy is related to the number of possible microstates according to:. A more general form of the third law that applies to a system such as a glass that may have more than one minimum microscopically distinct energy state, or may have a microscopically distinct state that is "frozen in" though not a strictly minimum energy state and not strictly speaking a state of thermodynamic equilibrium, at absolute zero temperature:.

The constant value not necessarily zero is called the residual entropy of the system.