gas laws in anesthesia
Post on 13-Jul-2015
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Dr Bikash Subedi
Moderator: Prof. Dr Baburaja Shrestha
The Gas Laws & Clinical
Applications
Elements that exist as gases at 250C and 1 atmosphere
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.
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
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
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)
Gas Laws
Boyle’s Law
• Robert boyle,1662
• At constant temperature, V 1/P
• PV = K (constant)
• P1V1 = P2V2
P
V
Calculation of Amount of gas in a cylinder
1300 Litres @ 1 atm pressure
10x130= V x 1 bar
10 litre O2 cylinder @ 130 bar
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
• Gases expand when heated, become less dense , thus hot air rises >> convection
• LMA inflatable cuff expands in an autoclave
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
• Hydrogen thermometer or constant volume gas thermometer
• Internal combustion engines
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
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
Avogadro’s Hypothesis
• For a given mass of an ideal gasvolume amount (moles) of the gasif temperature and pressure are constant
Amedeo Avogadro, 1811
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
• 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 !
• 2 g of Hydrogen or 32 g of Oxygen or 44 g of Carbon dioxide occupy 22.4 litres at STP
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.
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 anT
P
V = constant x = RnT
P
nT
PR is the gas constant
PV = nRT
Practical application ; use of pressure gauges to assess the contents of acylinder
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 .
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 + ...
AIR
O2 – 104CO2 – 40H2O – 47N2 - 569
Pulmonary capillary
vein arteryO2- 40CO2- 46
O2- 100CO2- 40
Alveolus at 760 mm Hg
At everest
• Atm pressure almost one third at sea level
• Thus, alveolar O2 pressure about 35 mm Hg
• Supplemental Oxygen
Henry’s Law
• William Henry in 1803
• At constant temperatureSolubility of gas Partial Pressure of gas
Solubility of gas :
• Depends on type of gas and liquid
• Decreases with increase in temperature
• Caisson’s disease/ decompression sickness
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)
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
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
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
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
Pseudocritical temperature
• Deals with gas mixture
• Temperature at which gas mixture may separate out into constituents
• Entonox 50% O2 50% N2O
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