pneumatic introduction, gas laws

1. PNEUMATIC AND ELECTRO PNEUMATIC SYSTEMS

2. PNEUMATIC AND ELECTRO PNEUMATIC SYSTEMS
Properties of air – Perfect Gas Laws – Compressor – Filters, Regulator, Lubricator, Muffler, Air
control Valves, Quick Exhaust Valves, Pneumatic actuators, Design of Pneumatic circuit –
Cascade method – Electro Pneumatic System – Elements – Ladder diagram – Problems,
Introduction to fluidics and pneumatic logic circuits.

3. Properties of Air, GAS LAWS

4. 1) Gases are highly compressible
• An external force compresses the gas sample and decreases its volume, removing the external force
allows the gas volume to increase.
2) Gases are thermally expandable
•When a gas sample is heated, its volume increases, and when it is cooled its volume decreases.
3) Gases have high viscosity
• Gases flow much easier than liquids or solids.
4) Most Gases have low densities
• Gas densities are on the order of grams per liter whereas liquids and solids are grams per cubic cm, 1000
times greater.
5) Gases are infinitely miscible
• Gases mix in any proportion such as in air, a mixture of many gases.
Properties of Gases

5. Properties of Gases
Pressure
Volume
Amount (moles)
Temperature

6. Pressure
• Units of Pressure
• 1 pascal (Pa) = 1 N/m2
• 1 atm = 760 mmHg = 760 torr
• 1 atm = 101,325 Pa
Pressure = Force
Area

7. Volume
• Volume is the three-dimensional space inside the container
holding the gas. The SI unit for volume is the cubic meter, m3. A
more common and convenient unit is the liter, l.
• Think of a 2-liter bottle of soda to get an idea of how big
a liter is.
• (OK, how big two of them are…)

8. Amount (moles)
• Amount of substance is tricky.As we’ve already learned, the SI unit for amount of
substance is the mole, mol. Since we can’t count molecules, we can convert
measured mass (in kg) to the number of moles, n, using the molecular or formula
weight of the gas.
By definition, one mole of a substance contains approximately 6.022 x 1023 particles
of the substance.

9. Temperature
Temperature is the measurement with which you’re probably most familiar
(and the most complex to describe completely). For these lessons, we will be
using temperature measurements in Kelvin, K.

10. Boyle’s Law
This law is named for Charles Boyle, who studied the
relationship between pressure, p, and volume, V, in the
mid-1600s.
He determined that for the same amount of a gas at
constant temperature,
p * V = constant
This defines an inverse relationship:
when one goes up, the other comes
down.
Doubling the pressure reduces the volume by
half.
Conversely, when the volume doubles, the
pressure decreases by half.

11. Application of Boyle’s Law
Boyle’s Law can be used to predict the interaction of pressure and volume.
p1 * V1 = p2* V2
p1= initial pressure
V1 = initial volume
p2= final pressure
V2 = final volume

12. Charles’ Law
This law is named for Jacques Charles, who
studied the relationship volume, V, and
temperature, T,around the turn of the 19th century.
He determined that for the same amount of a gas at
constant pressure,
V / T = constant
This defines a direct relationship: an
increase in one results in an increase
in the other.
As the temperature increases, the volume increases.
Conversely, when the temperature decreases, volume
decreases.

13. Partial Pressure
Partial Pressure
Pressure each gas in a mixture would exert if it were the only
gas in the container
Dalton's Law of Partial Pressures
The total pressure exerted by a gas mixture is the sum of the
partial pressures of the gases in that mixture.
PT = P1 + P2 + P3 + .....

14. Application of Charles’ Law
Charles’ Law can be used to predict the interaction of temperature
and volume.
V1 / T1= V2/ T2
V1 = initial volume
T1= initial temperature V2 = final
volume
T2= final temperature

15. Charles’ Law: Summary
Volume / Temperature = Constant V1/ T1= V2/ T2
With constant pressure and amount of gas, you can use
these relationships to predict changes in temperature and
volume.

16. Avogadro’s Law
V  number of moles (n)
V = constant x n
V1/n1 = V2/n2
Constant temperature
Constant pressure

17. Ideal Gas Equation
Boyle’s law: V  1 (at constant n and T)
P
Charles’ law: V  T (at constant n and P)
Avogadro’s law: V  n (at constant P and T)
V 
nT
P
nT
P
V = constant x
nT
= R
P
R is the gas constant
PV = nRT

18. Dalton’s Law of Partial Pressures
V and T
are
constant
P1 P2 Ptotal = P1 + P2

19. The Nature of Gases
Three basic assumptions of the kinetic theory as it applies to gases:
1. Gas is composed of particles- usually molecules or atoms
– Small, hardspheres
– Insignificant volume; relatively farapart from each other
– No attraction or repulsionbetween particles
2. Particles in a gas move rapidly in constant random motion
– Move in straight paths, changingdirection only when colliding with one another or other objects
– Average speed of O2 in air at 20 oC is an amazing 1660 km/h! (1.6km=1mile)
3. Collisions are perfectly elastic- meaning kinetic energy is transferred without loss from one particle to
another- the total kinetic energy remains constant.