The document presents an overview of gas laws and their clinical implications in anaesthesia, detailing definitions, properties of gases, and key gas laws including Boyle’s, Charles’, and Gay-Lussac’s laws. Each law is explained with its clinical relevance, specifically regarding the behavior of gases under various conditions and their applications in medical settings, such as determining gas volumes and pressures. Additionally, the document discusses concepts like the ideal gas law, Dalton’s law of partial pressures, and the Joule-Thomson effect, highlighting their importance in anaesthetic practices.

1. Gas Laws,
Clinical Implications and
Importance in Anaesthesia
Presenter: Dr. Suresh Pradhan

2. Outline
• introduction
• definitions
• gas laws and clinical implications

3. Introduction
• origins of the Gas laws came out of experimental
work conducted during the seventeenth and
eighteenth centuries by several people
• these experiments ultimately gave us the three gas
laws
• the theoretical construct that arises from the gas
laws is the concept of an ideal gas

4. • this is a gas that does obey the laws completely
under all circumstances
• as the laws break down at the extremes of
temperature and pressures, there is no gas that
obeys all the laws perfectly
• in our practical day-to-day use at room
temperature they can be assumed to do so and this
simplifies our consideration of them

5. Definitions
 Gas
• a substance that is in its gaseous phase, but is
above its critical temperature
• Critical temperature is the temperature above
which a gas cannot be liquefied no matter how
high the pressure applied is
• a vapour in contrast is a substance in the
gaseous phase but is below its critical
temperature

6. Elements that exist as gases at 25oC and 1 atmosphere

7. • 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

8.  Ideal gas
• theoretical
• negligible intermolecular forces
• collisions between atoms or molecules are
perfectly elastic
• obeys universal gas law PV= nRT at all temp &
pressures

9.  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

10.  Kinetic theory of Gases
• this model assumes that the molecules have
very small sizes relative to the distance between
them and that the molecules are in constant,
random motion related to their kinetic energy
• is given by e=½ mv2, where m is the mass of the
molecules and v is its velocity
• the molecules frequently collide with each other
and with the walls of the container holding the
gas molecules

12.  like all molecules, gas molecules have physical
properties of mass, velocity, momentum, and
energy
 at the macroscopic level these properties are
related to properties of density, pressure, and
temperature
 the temperature of a gas is related to the mean
kinetic energy of the gas
 the higher the temperature, the greater the
molecular motion

13. Pressure
• is defined as "the force per unit area acting at
right angles to the surface under consideration”
• Pressure = Force/Area
• unit of pressure is the Pascal
• pressure is the consequence of molecular
bombardment of the surface by the gas
• Kinetic energy is transferred to the surface and a
force is produced that creates the pressure

14. • if the volume falls, the pressure goes up because
the area for collisions fall and so more kinetic
energy transfer per unit area, and so an increase
in pressure
• other units of pressure
• atmospheres
• mm Hg
• cm H20
• PSI (pounds per square inch)
• dynes/cm2

15. Different units of Pressure

16.  Temperature
• is a physical quantity expressing the subjective
perceptions of hot and cold
• is a measure of the average kinetic energy in a
system
• denotes the degree of hotness or coldness of
the system

17. • is measured with a thermometer, historically
calibrated in various temperature scales
and units of measurement
• the most commonly used scales are the Celsius
scale, denoted in °C, the Fahrenheit scale (°F),
and the Kelvin scale
• the kelvin (K) is the unit of temperature in
the International System of Units (SI), in which
temperature is one of the seven
fundamental base units

18. • the divisions of the Kelvin and Celsius scale are
the same but the start points differ
• 0oC is 273K, so body temperature is 310K on this
scale
• the lowest theoretical temperature is absolute
zero, at which the thermal motion of all
fundamental particles in matter reaches a
minimum

19. • although classically described as motionless,
particles still possess a finite zero-point
energy in the quantum mechanical description
• absolute zero is denoted as 0 K on the Kelvin
scale, −273.15 °C on the Celsius scale, and
−459.67 °F on the Fahrenheit scale

21.  Standard Temperature & Pressure (STP)
• IUPAC has changed the definition in 1982
• until 1982, STP was defined as a temperature of
273.15 K (0 °C, 32 °F) and an absolute
pressure of exactly 1 atm (1.01325 × 105 Pa)
• Since 1982, STP is defined as a temperature of
273.15 K (0 °C, 32 °F) and an absolute pressure
of exactly 105 Pa (100 kPa, 1 bar)

22.  Volume
• space occupied by a substance measured in
three dimensions by cubic cm or cubic mm
• common units used to express volume include
liters, cubic meters, gallons, milliliters

23. GAS LAWS

24. Boyle’s Law
• Boyle–Mariotte law or Mariotte's law
• Robert Boyle, 1662
• at a constant temperature, the volume of a given
mass of gas is inversely proportional to the
absolute pressure
• at constant temperature, V ᾳ 1/P
• PV = K (constant)
• P1V1 = P2V2

27. Clinical Implication
 Calculation of Amount of gas in a cylinder
• oxygen cylinder of volume 10 L, Pressure = 138
bars
• So how much oxygen is stored?
• P1V1=P2V2
• 138X10=1XV2
• i.e. V2=1380L
• so, if we use oxygen@3l/m, the cylinder will last
for about 460 mins

28. Gas Laws and
Anaesthetic Implications
…..contd……

29. Charle’s Law
• also known as the law of volumes
• describes how gases tend to expand when heated
• Jacques Charles, 1787
• at constant pressure, volume of a given mass of gas
varies directly with temperature, that is
V ᾳ T ( in kelvin)
or V/T = Constant (k2)
or V1/T1=V2/T2
• gases expand when heated, become less dense,
thus hot air rises >> convection

32. Clinical Implication
• respiratory gas measurements of tidal volume &
vital capacity etc. are done at ambient temperature
while these exchanges actually take place in the
body at 37OC
• one way of heat loss from the body is that air next
to the body surface gets warmer and moves up and
thus our patient loses heat this way (esp. important
in pediatric anaesthesia)

33. Gay Lussac’s Law
• also known as third gas law, Amontons' law or the
pressure law
• Joseph Louis Gay-Lussac, 1809
• at constant volume the absolute pressure of a
given mass of gas varies directly with the absolute
temperature,
i.e. P ᾳ T
or, P/T = constant
or, P1/T1=P2/T2

34. Clinical Implication
• medical gases are stored in cylinders having a
constant volume and high pressures (138 Barr in a
full oxygen / air cylinder); if these are stored at high
temperatures, pressures will rise causing explosions
• molybdenum steel can withstand pressures till 210
bars. Weakening of metal in damaged cylinders are
at a greater risk of explosion due to rise in
temperature

35. Combined Gas Law
• Boyle’s + Charle’s + Gay Lussac’s law
• PV/T=k
• 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
• Clinical application:
• in spirometry, measurement of volumes is done at
ambient condition of T and P; so correction should be
done by a factor of 1.07

36. Avogadro’s Hypothesis/Law
• an experimental gas law relating volume of a gas to
the amount of substance of gas present
• equal volumes of all gases, at same temperature
and pressure, have the same number of molecules
• can also be defined as one mole of a gas contains
6.023x1023 (avogadro’s number) molecules and
occupies 22.4L at STP

37. • for a given mass of an ideal gas, the volume and
amount (moles) of the gas are directly proportional
if the temperature and pressure are constant
• which can be written as:
Vᾳ n
or, V/n= k where, V=volume of gas
or, V1/n1=V2/n2 n is the amount of substance of the gas
(measured in moles)
k is a constant

38. Clinical Implication

39. Ideal Gas Law
• by combined gas law, PV/T=k
• or P1V1 / T1 = P2V2 / T2
• combining with Avogadros Law, combined gas law
can be restated as,
PV/T=nRT where
P is pressure
V is volume
n is the number of moles
R is the universal gas constant
T is temperature (K)

40. • the equation are exact only for an ideal gas, which
neglects various intermolecular effects
• however, the ideal gas law is a good approximation
for most gases under moderate pressure and
temperature

41. This law has the following important consequences:
i. if temperature and pressure are kept constant,
then the volume of the gas is directly
proportional to the number of molecules of gas
ii. if the temperature and volume remain constant,
then the pressure of the gas changes is directly
proportional to the number of molecules of gas
present

42. iii. if the number of gas molecules and the
temperature remain constant, then the
pressure is inversely proportional to the volume
iv. if the temperature changes and the number of
gas molecules are kept constant, then either
pressure or volume (or both) will change in
direct proportion to the temperature

43. Clinical Implication
• this equation may be used in anaesthesia when
calculating the contents of an oxygen cylinder
- constant room temp
- fixed internal volume
- R is a constant
• Only variables are P and n so that
P ∝ n
• therefore, pressure gauge acts as a content gauge
for gases – measure of amount of O2 left in a
cylinder

44. BUT,
• we cannot use a nitrous oxide cylinder pressure
gauge in the same way as these cylinders contain
both vapour & liquid and so the gas laws do not
apply

45. Dalton’s Law of Partial Pressures
• 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

47. • the total pressure of a mixture of gases equals
the sum of the partial pressures of the individual
gases

49. Adiabatic compression or
expansion of gases
• if the state of a gas is altered without a change in
heat energy , it is said to undergo adiabatic change
• adiabatic, when applied to expansion or
compression of a gas, means that energy is not
added or removed when the changes occur.
⌐ Compression of gas – temperature rises
⌐ Expansion of gas – temperature falls
 joule thompson effect states that when a gas is
allowed to escape through a narrow opening,
there is a sudden temperature drop

50. Clinical Implication
• compression of gases will require added cooling
• in cyroprobe, expansion of gas in the probe – low
temperature in probe tip
• 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

51. 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

52. • manufacture of oxygen:
• when air is cooled by external cooling and is
made to suddenly expand, it loses further
temperature as energy is spent in order to hold
the molecules together (Joule Thomson’s Effect)
• when this is repeated many times the
temperature reduces to less than -1830C and
through fractional distillation, liquid oxygen
collected in the lower part is separated from
nitrogen with a boiling point of -197oC which
collects at the top of the container

53. Henry’s Law
• William Henry in 1803
• Henry’s law states that for a gas-liquid interface the
amount of the gas that dissolves in the liquid is
proportional to its partial pressure
• so Henry’s law helps to predict how much gas will
be dissolved in the liquid
• at constant temperature,
Solubility of gas ᾳ Partial Pressure of gas

55. Graham’s Law
• states that the rate of diffusion of gases is inversely
proportional to the square root of its molecular
weight
• so the larger the molecule, the slower it diffuses

56. Clinical Implication
• explains the second gas effect when using nitrous
and a volatile anaesthetic in oxygen
• for example, halothane is more massive than
nitrous oxide, Graham’s law will indicate that the
nitrous will diffuse quicker and so raise the
concentration of the halothane in the alveolus.

57. Fick’s Law of diffusion
• states that the rate of diffusion of a gas across a
membrane is proportional to the membrane area
(A) and the concentration gradient (C1-C2) across
the membrane and inversely proportional to the
thickness (D)

58. Clinical Implication
• anaesthetic vapour diffusing into breathing
circuits and later acting as vaporizers at the time
of discontinuation of anaesthetic agents
• N2O diffusion into cuff of ETT
• diffusion of N2O into air filled cavities

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

60. 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

61. Pseudocritical temperature
• when two gases, one of high and another of low
critical temperature are mixed in a container, the
critical temperature of the gas with a high critical
temperature will decrease to a lower level (pseudo
critical temperature) and the mixture will remain
as a gas above this pseudo critical temperature
• this effect is called as Poynting effect

62. • is the temperature of a gas mixture at which the
gas mixture may separate out into constituents
gases
• Entonox
• N2O 50% / O2 50% = - 5.5°C for cylinders (most
likely at 117 bar)
• N2O 50% / O2 50% = - 30° C for piped gas

63. Thank you!!!

64. • venturi effect
• coanada effect