The Gas Laws and Their Experimental Foundations

The physical behavior of a sample of gas can be described completely by four variables pressure (P), volume (V), temperature (T), and amount (number of moles, n). The variables are interdependent any one of them can be determined by measuring the other three. We know now that this quantitatively predictable behavior is a direct outcome of the structure of gases on the molecular level. Yet, it was discovered, for the most part, before Dalton s atomic theory was published  

Three key relationships exist among the four gas variables—Boyle s, Charles s, and Avogadro s laws. Each of these gas laws expresses the effect of one variable on another, with the remaining two variables held constant. Because gas volume is so easy to measure, the laws are expressed as the effect on gas volume of a change in pressure, temperature, or amount of gas. 

These three laws are special cases of an all-encompassing relationship among gas variables called the ideal gas law. This unifying observation quantitatively describes the state of a so-called ideal gas, one that exhibits simple linear relationships among volume, pressure, temperature, and amount. Although no ideal gas actually exists, most simple gases, such as N2, O2, H2, and the noble gases, show nearly ideal behavior at ordinary temperatures and pressures. We discuss the ideal gas law after the three special cases. 

The generalization of Boyle s observations is known as Boyle s law at constant temperature, the volume occupied by a fixed amount of gas is inversely proportional to the applied (external) pressure, or 

The constant is the same for the great majority of gases. Thus, tripling the external pressure reduces the volume to one-third its initial value halving the external pressure doubles the volume and so forth. 

Gases exert pressure (force/area) on all surfaces with which they make contact. A barometer measures atmospheric pressure in terms of the height of the mercury column that the atmosphere can support (760 mmHg at sea level and 0°C). Chemists measure pressure in units of atmospheres (atm), torrs (equivalent to mmHg), and pascals (Pa, the SI unit). 

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Measuring Atmospheric Pressure Units of Pressure The Gas Laws and Their Experimental Foundations... 

Clearly, it is not possible to have reliable, detailed experimental data at hand on all the useful pure compounds and mixtures that exist in the world. Consequently, in the absence of experimental information, we estimate (predict) properties based on modifications of well-established principles, such as the ideal gas law, or based on empirical correlations. Thus the foundation of the estimation methods ranges from quite theoretical to completely empirical, and their reliability ranges from excellent to terrible. 

It was during the first half of the 17th century that scientists began to study chemical reactions experimentally. Jan Baptista van Helmont laid the foundations of the law of conservation of mass. Van Helmont showed that in a number of reactions an aerial fluid was liberated which he defined as a gas. A new class of substances with their own physical properties was shown to exist. A kinetic-molecular theory of gases began to develop. Notable in this field were the experiments of Robert Boyle whose studies, later known as Boyle s law, provided an equation describing the inverse relation between pressure and volume of gas (see the ideal gas law in Chapter 3). 

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