“The Battle of”
Medical Gas, NFPA 99 and Plumbing Codes
Introduction
If you study chemistry and matters, you will easily bump into these chapters in the book that deals with intra and intermolecular forces with gases, the gas laws (Ideal, Charles, and Boyle, Avagadro), the density, the Dalton’s law of partial pressure, grahams’ law of diffusion and effusion.
Without taking the reader to a four month Journey in Chemistry, the basics needed to understand NFPA 99 and Plumbing Codes.
Fundamentals of Gas Laws
On earth the matter is subdivided into three categories, solids, with fixed volume and weight occupying known boundaries, the liquid, where the boundaries must be contain in an open container, and gas, where the boundaries are defined within a closed container.
The Gas Laws
The nature of gas can be determined with a simple gas law. We need to assume that gases are ideal and not “real”.
Ideal Gases
Based in ideal gas assumption, the following properties.
- All random motions of gas particles are in constant motion,
- The conservation of energy on impact of the gas particle holds true. All particles are elastic.
- All gas molecules do not hat attractive or repulsive interactions based on magnetic field.
- The average kinetic energy of gas is proportional to its temperature, where all molecules have same temperature.
The property of gases are simple: Temperature, T (Kelvin), volume, V, and pressure, P. Based on quantity (n=number of moles). A reference STP stands for Standard temperature and pressure at 273 degrees K (or zero degree Celsius) and one atmosphere.
In an Ideal gas law, the equation can be written as:
PV=nRT
P is the absolute gauge pressure, T is in Kelvin, n is the number of molecules, R is a constant, and V is the volume.
Pressure:
Measuring the Pressure of a Gas
Gas pressure is a gauge of the number and force of collisions between gas particles and the walls of the container that holds them. The SI unit for pressure is the pascal Pa), but other pressure terms include atmospheres atms), millimeters of mercury (mmHg), and torr. The following is a list of all of the standard pressure in every unit for pressure. Memorize these for the exam so you can convert units where necessary:
760 mmHg
760 torr
1.00 atm
101,325 Pa
101.325 kPa
The piece of lab equipment specifically designed to measure the pressure of gases is known as the barometer. A barometer uses the height of a column of mercury to measure gas pressure in millimeters of mercury or torr (1 mmHg = 1 torr). The mercury is pushed up the tube from the dish until the pressure at the bottom of the tube (due to the mass of the mercury) is balanced by the atmospheric pressure.
When using a barometer, you calculate gas pressure with the following equation:
Gas pressure = atmospheric pressure – h (height of the mercury)
The open-tube manometer is another device that can be used to measure pressure. The open-tube manometer is used to measure the pressure of a gas in a container.
The pressure of the gas is given by h (the difference in mercury levels) in units of torr or mmHg. Atmospheric pressure pushes on the mercury from one direction, and the gas in the container pushes from the other direction. In a manometer, since the gas in the bulb is pushing more than the atmospheric pressure, you add the atmospheric pressure to the height difference:
gas pressure = atmospheric pressure + h
There is one other possibility for a manometer question that could appear on the SAT II Chemistry test: they could ask you about a closed-tube manometer. Closed-tube manometers look similar to regular manometers except that the end that’s open to the atmospheric pressure in a regular manometer is sealed and contains a vacuum. In these systems, the difference in mercury levels (in mmHg) is equal to the pressure in torr.
Boyle’s Law
Boyle’s law is simply when in an experiment the temperature and number of molecules are constant, where the ideal gas law equation is simplified to:
PV = Constant. Boyles’ Law
As the pressure increases, the volume reduces and as the volume increases, the pressure decreases. This inverse proportionality of pressure and volume is known as Boyles Law. In an enclosed balloon, as the volume increases, the pressure decreases.
Boyles law between two states can be written as:
P1V1= P2V2 Boyles’ Law
Where the 1 and 2 identifies the properties of gas in state 1 and in state 2.
Charles’s Law
Charles’s law is similar to Boyles law where one property remains constant, the pressure and number of molecules. Ideal gas law equation reduces to:
V/T= Constant Charles’ Law
In Charles’ law the relation between the temperature and volume is directly proportional. That is if temperature increases, the volume will increase. The ideal gas law between two gas state becomes:
V1/T1= V2/T2 Charles’ Law
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