Properties of gases and the gas laws

The few examples illustrated above demonstrate how gases and the gas laws affect our lives on a daily basis. We are surrounded by gas and our lives are dependent on a steady supply of gas. Humans can survive a few weeks without food, a few days without water, but only a few minutes without a steady supply of air. In this chapter, we have focused on some of the unique properties of gases and the gas laws. We continue our exploration of gases when we examine phase changes in Chapter 10 and solubility in Chapter 11. 

Summaries of the properties of gases, particularly the variation of pressure with volume and temperature, are known as the gas laws. The first reliable measurements of the properties of gases were made by the Anglo-Irish scientist Robert Boyle in 1662 when he examined the effect of pressure on volume. A century and a half later, a new pastime, hot-air ballooning, motivated two French scientists, Jacques Charles and Joseph-Louis Gay-Lussac, to formulate additional gas laws. Charles and... 

In this chapter we have considered several properties of gases and have seen how the relationships among these properties can be expressed by various laws written in the form of mathemafical equafions. The most useful of these is the ideal gas equation, which relates all the important gas properties. However, under certain conditions gases do not obey the ideal gas equation. 

Chapter 11, Gases, describes the properties of a gas and calculates changes in gases using the gas laws and the ideal gas law. The amounts of gases required or produced in chemical reactions are calculated. 

Pressure, temperature, and volume are properties of gases that are completely interrelated. Boyle s law and Charles law may be combined into one equation that is referred to as the ideal gas law. This equation is always true for ideal gases and is true for real gases under certain conditions. 

In this chapter, you learned about the properties of gases. You learned that you can use the combined gas law, the ideal gas law, or the individual gas laws to calculate certain gas quantities, such as temperature and pressure. You also learned that these equations could also be useful in reaction stoichiometry problems involving gases. You learned the postulates of the Kinetic-Molecular... 

While studying gases in this chapter you will consider four main physical properties—volume, pressure, temperature, and amount—and their interrelationships. These relationships, commonly called gas laws, show up quite often on the AP exam, so you will spend quite a bit of time working problems in this chapter. But before we start looking at the gas laws, let s look at the Kinetic Molecular Theory of Gases, the extremely useful model that scientists use to represent the gaseous state. 

Unlike solids and liquids, different gases show remarkably similar physical behavior regardless of their chemical makeup. Helium and fluorine, for example, are vastly different in their chemical properties yet are almost identical in much of their physical behavior. Numerous observations made in the late 1600s showed that the physical properties of any gas can be defined by four variables pressure (P), temperature (T), volume (V), and amount, or number of moles in). The specific relationships among these four variables are called the gas laws, and a gas whose behavior follows the laws exactly is called an ideal gas. 

Every chemistry student is familiar with the ideal gas equation PV = nRT. It turns out that this equation is a logical consequence of some basic assumptions about the nature of gases. These simple assumptions are the basis of the kinetic theory of gases, which shows that the collisions of individual molecules against the walls of a container creates pressure. This theory has been spectacularly successful in predicting the macroscopic properties of gases, yet it really uses little more than Newton s laws and the statistical properties discussed in the preceding chapters. 

The ideal gas law has been used in many examples in earlier chapters, and some of the important physical properties of gases (the one-dimensional velocity distribution, average speed, and diffusion) were presented in Chapter 4. This chapter puts all of these results into a more comprehensive framework. For example, in Section 7.3 we work out how the diffusion constant scales with pressure and temperature, and we explore corrections to the ideal gas law. 

From the earliest days of quantitative inquiry, scientists have sought to uncover the mathematical relationships that describe natural phenomena, including the properties of gases. Because there are four fundamental properties of a gas, namely, P, T, V, and n, discovering the relationship between any two requires that the other two properties be kept constant. Some of the earliest quantitative studies of gases were reported in the mid-1600s by British chemist Robert Boyle, who found that for a fixed amount of a gas at a specific temperature (i.e., constant n and T), the volume was inversely proportional to the applied pressure. This V-P relationship, known as Boyle s law, is represented as... 

It is desirable to compare the predictions of the theory presented here with experimental results obtained on some systems in which an independent computation of the gas-surface potential function can be carried out. A calculation of the potential functions for the adsorption of rare gases on solid rare gases involves the least number of unknown parameters. The rare gases crystallize into face-centered cubic solids with known lattice constants. Furthermore, the parameters appearing in the Lennard-Jones potential functions for the gas-gas and the gas-solid atom interaction can be estimated to a good degree of accuracy from experiments on the gas properties as well as from the empirical combining laws for potential parameters. Furthermore, some experimental results have already been reported for these adsorption systems (18, 20). 

Several gas laws have been introduced in this chapter, but no explanation as to why those laws apply to all gases has been proposed. This section introduces the kinetic molecular theory of gases, which explains the gas laws and when extended, also explains some properties of liquids and solids. Five postulates explain why gases behave as they do ... 

The early chapters in this book deal with chemical reactions. Stoichiometry is covered in Chapters 3 and 4, with special emphasis on reactions in aqueous solutions. The properties of gases are treated in Chapter 5, followed by coverage of gas phase equilibria in Chapter 6. Acid-base equilibria are covered in Chapter 7, and Chapter 8 deals with additional aqueous equilibria. Thermodynamics is covered in two chapters Chapter 9 deals with thermochemistry and the first law of thermodynamics Chapter 10 treats the topics associated with the second law of thermodynamics. The discussion of electrochemistry follows in Chapter 11. Atomic theory and quantum mechanics are covered in Chapter 12, followed by two chapters on chemical bonding and modern spectroscopy (Chapters 13 and 14). Chemical kinetics is discussed in Chapter 15, followed by coverage of solids and liquids in Chapter 16, and the physical properties of solutions in Chapter 17. A systematic treatment of the descriptive chemistry of the representative elements is given in Chapters 18 and 19, and of the transition metals in Chapter 20. Chapter 21 covers topics in nuclear chemistry and Chapter 22 provides an introduction to organic chemistry and to the most important biomolecules. 

The simphcity of the relationship between the thermodynamic scale and the gas thermometer scale is due principally to the simple properties of rarefied gases, and also to the fortunate choice of mercury as thermometric substance by Celsius and Reaumur before the discovery of the gas laws. The coefficient of expansion of mercury happens to be almost exactly proportional to the coefficient of expansion of rarefied gases. All our thermodynamical relationships would have been very much more comphcated had water or alcohol, for example, or the resistance of a metal, been used for the definition of the practical scale of temperature. Their strict validity, however, would not have been affected. 

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