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# gas laws in anesthesia

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### gas laws in anesthesia

1. 1. Dr Bikash Subedi Moderator: Prof. Dr Baburaja Shrestha The Gas Laws & Clinical Applications
2. 2. Elements that exist as gases at 250C and 1 atmosphere
3. 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. 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. 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. 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. 7. Gas Laws
8. 8. Boyle’s Law • Robert boyle,1662 • At constant temperature, V  1/P • PV = K (constant) • P1V1 = P2V2 P V
9. 9. Calculation of Amount of gas in a cylinder 1300 Litres @ 1 atm pressure 10x130= V x 1 bar 10 litre O2 cylinder @ 130 bar
10. 10. 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
11. 11. • Gases expand when heated, become less dense , thus hot air rises >> convection • LMA inflatable cuff expands in an autoclave
12. 12. 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
13. 13. • Hydrogen thermometer or constant volume gas thermometer • Internal combustion engines
14. 14. 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
15. 15. 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
16. 16. 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
17. 17. 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
18. 18. • 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 !
19. 19. • 2 g of Hydrogen or 32 g of Oxygen or 44 g of Carbon dioxide occupy 22.4 litres at STP
20. 20. 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.
21. 21. 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
22. 22. 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 .
23. 23. 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 + ...
24. 24. AIR
25. 25. O2 – 104 CO2 – 40 H2O – 47 N2 - 569 Pulmonary capillary vein arteryO2- 40 CO2- 46 O2- 100 CO2- 40 Alveolus at 760 mm Hg
26. 26. At everest • Atm pressure almost one third at sea level • Thus, alveolar O2 pressure about 35 mm Hg • Supplemental Oxygen
27. 27. Henry’s Law • William Henry in 1803 • At constant temperature Solubility of gas  Partial Pressure of gas
28. 28. Solubility of gas : • Depends on type of gas and liquid • Decreases with increase in temperature • Caisson’s disease/ decompression sickness
29. 29. 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)
30. 30. 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
31. 31. 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
32. 32. 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
33. 33. 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
34. 34. Pseudocritical temperature • Deals with gas mixture • Temperature at which gas mixture may separate out into constituents • Entonox 50% O2 50% N2O