Chap 5 - Pressure All gases exert pressure Think of gas...

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Pressure All gases exert pressure Think of gas molecules as tiny balls whizzing around a container Naturally, as they bounce against the container’s wall, they have a tendency to “push” against it. This is the basic concept of pressure The more molecules you put into a given volume, the more pressure there is
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Measuring pressure The barometer: a pool of liquid with a column rising from it, used for finding pressure of ambient air The manometer: used for finding differences in pressure The surrounding air pressure forces the liquid up the tube Traditionally, mercury is used because it is a dense liquid. Unfortunately, mercury is poisonous and people are clumsy, so other alternatives are used.
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Units of pressure Since we have always used mercury in the passed, the units of measuring pressure is expressed in the number of mm’s of mercury that is raised by the air pressure Atmospheric pressure is defined to be 760mm of Hg, or 760 torrs. Pascals are also used, which is 1 newton divided by one meter squared (pressure = force/area) 1 atm = 760 mmHg = 760 torr = 101,325 Pa = 29.92 in Hg = 14.7 psi
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Boyle’s Law PV= k What that means: given a container with a variable volume and with a constant number of gas molecules, the pressure times the volume will always be some constant volume Implications: If you double the volume, you will half the volume, or mathematically, P = k/V Why: We said gas exerts pressure by “pushing” if you increase the volume, you spread the molecules out and the force they exert per area decreases proportionally.
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Charles’ Law V = bT What this means: volume goes up
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Chap 5 - Pressure All gases exert pressure Think of gas...

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