Gases: The Empirical Gas Laws: Pressure

Quantitative measurements on gases were first made by the English chemist Robert Boyle (1627 - 1691). Boyle used two instruments to measure pressure: the manometer, which measures differences in pressure, and the barometer, which measures the total pressure of the atmosphere.


The operation of a manometer, which is simply a bent piece of tubing, preferably glass with one end closed. When the fluid level in both arms is the same, the pressure of the sample of gas inside the closed end must equal the pressure of the external atmosphere since the downward force on the two columns of liquid is then equal. When the liquid levels are unequal, the pressures must differ. The difference in pressure can be measured in units of length of the vertical column of liquid. The mmHg, or its modern version the torr, originated in this use of the manometer. Mercury is particularly convenient for use in manometers (and barometers) because at room temperature it has low vapor pressure, does not wet glass, and has a high density. Other liquids such as linseed oil or water have also been used in manometers.
 
The barometer was invented by Torricelli, one of Galileo's students. It is a device for measuring the total pressure of the atmosphere. A Torricellian barometer can easily be constructed by taking a glass tube about a meter long, sealing one end, filling the tube completely with mercury, placing your thumb firmly over the open end, and carefully inverting the tube into an open dish filled with mercury. The mercury will fall to a height independent of the diameter of the tube and a vacuum will be created above it. The height of the mercury column will be the height which the atmospheric pressure can support. The standard atmospheric pressure, one atmosphere (atm), is 760 mmHg but the actual atmospheric pressure varies depending upon altitude and local weather conditions. For this reason barometers can be used to help predict the weather. A falling barometer indicates the arrival of a low pressure air system, which often means stormy weather. A rising barometer indicates the arrival of a high pressure air system, and that often means clear weather.
 
While mercury is again the most convenient liquid for use in barometers it is by no means the only liquid which can be used. Preparation of a water barometer, and many of the early barometers did use water, requires use of a vacuum pump (or arms 13 meters long).
 

Units of Pressure

Units of pressure were originally all based on the length of the column of liquid, usually mercury, supported in a manometer or barometer. By far the most common of these units was the mmHg, however, the modern SI unit of pressure is derived from the fundamental units of the SI.  Pressure is force per unit area, and force is the product of mass times acceleration, so the SI unit of pressure is the kg m s-2/m2 or newton/m2, which is called the pascal (Pa).

All of the older units of pressure have now been redefined in terms of the pascal. One standard atmosphere, the pressure of the atmosphere at sea level, is by definition exactly 101,325 Pa. The torr, named in honor of Torricelli, is defined as 1/760 of a standard atmosphere or as 101,325/760 Pa. The mmHg can be considered identical to the torr. The term bar is used for 100000 Pa, which is slightly belowone standard atmosphere.
 
Robert Boyle and his Law
Boyle used the manometer and barometer to study the pressures and volumes of different samples of different gases. The results of his studies can be summarized in a simple statement which has come to be known as Boyle's law:
 
At any constant temperature, the product of the pressure and the volume of any size sample of any gas is a constant.
 
For a particular sample of any gas, Boyle's law can be shown graphically as is done in the Figure below. It is more common to express it mathematically as P1V1 = P2V2.
 
The pressure and the volume vary inversely; as the pressure increases the volume of the sample of gas must decrease.
 
Example:  A sample of gas occupies a volume of 47.3 cm3 at 20oC when the pressure is 30 cm of mercury.  If the pressure is increased to 75 cm of mercury, the sample will occupy a volume of 47.3 cm3 (30 cmHg/75 cmHg) = 18.9 cm3.
 
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