1. Dr Bikash Subedi
Moderator: Prof. Dr Baburaja Shrestha
The Gas Laws
& Clinical
Applications

2. Elements that exist as gases at 250C and 1 atmosphere

3. Physical Characteristics of Gases
• Assume the volume and shape of their
containers.
• most compressible state of matter
• mix evenly and completely when confined to
the same container
• much lower densities than liquids and solids.

4. Ideal gas
• Theoretical
• Negligible intermolecular forces
• collisions between atoms or molecules are
perfectly elastic
• Obeys universal gas law PV= nRT at all temp &
pressures
• Where P = Pressure V = Volume n = Numbers of moles R = Universal gas constant = 8.3145
J/mol K T = Temperature

5. Real gas
• Real gases H2, N2 , O2
• exhibit properties that cannot be explained
entirely using the ideal gas law
• Behave like ideal gas at STP
• Air at atmospheric pressure is a nearly ideal
• STP = standard temperature and pressure

6. Standard Temperature & Pressure
(STP)
• Variable
• IUPAC has, since 1982
• standard reference conditions as being 0 °C
and 100 kPa (1 bar), in contrast to its old
standard of 0 °C and 101.325 kPa (1 atm)

7. Gas Laws

8. Boyle’s Law
• Robert boyle,1662
• At constant temperature, V  1/P
• PV = K (constant)
• P1V1 = P2V2
P
V

10. Calculation of Amount of gas in a
cylinder
1300 Litres @ 1 atm pressure
10x130= V x 1 bar
10 litre O2 cylinder @ 130 bar

11. Charles law
• Jacques Charles, 1678
• At constant pressure volume of a given mass
of gas varies directly with temperature , that is
V  T ( in kelvin)
or V/T = Constant (k2)
V
T

12. • Gases expand when
heated, become
less dense , thus hot
air rises >> convection
• LMA inflatable cuff
expands in an autoclave

13. 3rd gas law /Gay-Lussac’s law
• At constant volume the absolute pressure of a
given mass of gas varies directly with the
absolute temperature
P  T
or P/T = Constant
P
T

14. • Hydrogen thermometer or constant volume
gas thermometer
• Internal combustion engines

15. Combined Gas Law
• Boyle’s + Charle’s + Gay Lussac’s law
• P1V1 / T1 = P2V2 / T2
• useful for converting gas volumes collected
under one set of conditions to a new volume
for a different set of conditions

16. Spirometry
• BTPS 37 0 C and H2O pressure 47 mm of Hg
• Standard room temp & H2O pressure
250C
• spirometer records volume under room air, not
body conditions
• Thus, conversion factor or 1.07
• BTPS – body temp & pressure

17. Avogadro’s Hypothesis
• For a given mass of an ideal gas
volume  amount (moles) of the gas
if temperature and pressure are constant
Amedeo Avogadro, 1811

18. 1 MOLE OF A SUBSTANCE
• Quantity of a substance containing the same
number of particles as there are atoms in
0.012kg of carbon12
• There are 6.022 x 1023 atoms in 12 g of carbon
12. This is called Avogadro’s Number

19. • equal volumes of gases at the same
temperature and pressure contain equal
numbers of molecules
• One mole of any gas at STP occupies
22.4litres !

20. • 2 g of Hydrogen or 32 g of Oxygen or 44
g of Carbon dioxide occupy 22.4 litres at
STP

21. Calculating the volume of nitrous
oxide in a cylinder :
• A nitrous oxide cylinder contains 3.4 kg of
nitrous oxide .
• The molecular weight of nitrous oxide is 44
• One mole is 44 g
• At STP , 44 g occupies 22.4 Litres . Therefore
3,400 g occupies 22.4 x 3,400/44 = 1730
litres.

22. Ideal Gas Equation
Charles’ law: V a T (at constant n and P)
Avogadro’s law: V a n (at constant P and T)
Boyle’s law: V a (at constant n and T)
1
P
V a
nT
P
V = constant x = R
nT
P
nT
P
R is the gas constant
PV = nRT
Practical application ; use of pressure gauges to assess the contents of a
cylinder

23. Dalton’s Law of Partial Pressures
V and T are
constant
P1 P2 Ptotal = P1 + P2
John Dalton , 1801
In a mixture of gases , pressure exerted by each gas is the same as that which it
would exert if it alone occupied the container .

24. Dalton’s Law
• The total pressure of a mixture of gases equals
the sum of the partial pressures of the
individual gases.
Ptotal = P1 + P2 + ...

25. AIR

26. O2 – 104
CO2 – 40
H2O – 47
N2 - 569
Pulmonary capillary
vein arteryO2- 40
CO2- 46
O2- 100
CO2- 40
Alveolus at 760 mm Hg

27. At everest
• Atm pressure almost one third at sea level
• Thus, alveolar O2 pressure about 35 mm Hg
• Supplemental Oxygen

28. Henry’s Law
• William Henry in 1803
• At constant temperature
Solubility of gas  Partial Pressure of gas

30. Solubility of gas :
• Depends on type of gas and liquid
• Decreases with increase in temperature
• Caisson’s disease/ decompression sickness

31. Adiabatic changes of state in a gas
• If the state of a gas is altered without a change in
heat energy , it is said to undergo adiabatic change
• Heat energy neither received nor given to surrounding
• If a gas is rapidly compressed ; its temperature rises
(the Joule – Kelvin principle).
• Conversely , If a compressed gas expands rapidly,
cooling occurs (cryoprobe)

32. Application
• Compression of air rapidly in compressor >>
↑ temp >> need of coolant
• Cylinder connected to an anesthetic
machine rapidly turned on >> ↑↑
temperature in gauges & pipelines >> fire
or explosion

33. Cryoprobe
• Rapidly expanding gas through a capillary tube
causes cooling
• N2O, He, Argon, N2
• Cooling causes degeneration, necrosis
• Wart, mole removal. Nerve degeneration for
pain

34. Critical temperature
• Temperature above which a gas cannot be
liquefied
• No matter how much pressure!
• For N2O 36.5 0C, - 119 0C for O2
• for CO2 = 31.1oC

35. Critical Pressure
• Minimum pressure that causes liquefaction
of a gas at its critical temperature
(for CO2 pc = 73 atmospheres)
• So CO2 liquefies ↓ 73 atm at 31.1 0C

36. Pseudocritical temperature
• Deals with gas mixture
• Temperature at which gas mixture may
separate out into constituents
• Entonox 50% O2 50% N2O