Gas laws and anaesthetic implications

Gas laws and anaesthetic implications

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 25 o C and 1 atmosphere

7. a ssume 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 t heoretical n egligible 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 H 2 , N 2 , O 2 exhibit properties that cannot be explained entirely using the ideal gas law behave like ideal gas at STP a ir 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 t he molecules frequently collide with each other and with the walls of the container holding the gas molecules

11. 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

12. 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

13. 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

14. Different units of Pressure

15. 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

16. 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

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

18. a lthough classically described as motionless, particles still possess a finite zero-point energy in the quantum mechanical description a bsolute 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

19. Standard Temperature & Pressure (STP) IUPAC has changed the definition in 1982 u ntil 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 × 10 5   Pa) Since 1982, STP is defined as a temperature  of 273.15  K (0 °C, 32 °F) and an absolute pressure of exactly 10 5   Pa (100   kPa , 1 bar )

20. Volume s pace occupied by a substance measured in three dimensions by cubic cm or cubic mm c ommon units used to express volume include liters, cubic meters, gallons, milliliters

21. GAS LAWS

22. 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

23. Clinical Implication Calculation of Amount of gas in a cylinder o xygen 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

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

25. 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 g ases expand when heated , become less dense, thus hot air rises >> convection

26. 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 37 O C o ne 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 )

27. Gay Lussac’s Law also known as third gas law, Amontons ' law or the pressure law  Joseph Louis Gay-Lussac, 1809 a t 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

28. 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

29. 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

30. 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

31. f or 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: V ᾳ 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

32. Clinical Implication c alculating the volume of nitrous oxide in a cylinder a nitrous oxide cylinder contains 3.4 kg of nitrous oxide t he molecular weight of nitrous oxide is 44 o ne mole is 44 g a t STP , 44 g occupies 22.4 Litres therefore 3,400 g occupies 22.4 x 3,400/44 = 1730 litres

33. Clinical Implication

34. 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 )  

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

36. This law has the following important consequences: if temperature and pressure are kept constant, then the volume of the gas is directly proportional to the number of molecules of gas 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

37. if the number of gas molecules and the temperature remain constant, then the pressure is inversely proportional to the volume i f 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

38. 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

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

40. 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

41. t he total pressure of a mixture of gases equals the sum of the partial pressures of the individual gases

42. 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

43. Clinical Implication c ompression of gases will require added cooling in cyroprobe , expansion of gas in the probe – low temperature in probe tip c ompression of air rapidly in compressor >> ↑ temp >> need of coolant c ylinder connected to an anesthetic machine rapidly turned on >> ↑↑ temperature in gauges & pipelines >> fire or explosion

44. Cryoprobe rapidly expanding gas through a capillary tube causes cooling N 2 O , He, Argon, N 2 c ooling causes degeneration, necrosis wart/mole removal, nerve degeneration for pain

45. 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 - 183 C and through fractional distillation, liquid oxygen collected in the lower part is separated from nitrogen with a boiling point of - 197 o C which collects at the top of the container

46. 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

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

48. Clinical Implication explains the second gas effect when using nitrous and a volatile anaesthetic in o xygen 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.

49. 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)

50. Clinical Implication a naesthetic vapour diffusing into breathing circuits and later acting as vaporizers at the time of discontinuation of anaesthetic agents N 2 O diffusion into cuff of ETT diffusion of N 2 O into air filled cavities

51. Critical temperature temperature above which a gas cannot be liquefied No matter how much pressure! For N 2 O 36.5 o C , -119 o C for O 2 for CO 2 = 31.1 o C

52. Critical Pressure m inimum pressure that causes liquefaction of a gas at its critical temperature (for CO 2 pc = 73 atmospheres) s o CO 2 liquefies ↓ 73 atm at 31.1 C

53. 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

54. is the temperature of a gas mixture at which the gas mixture may separate out into constituents gases Entonox N 2 O 50% / O2 50% = - 5.5°C for cylinders (most likely at 117 bar) N 2 O 50% / O2 50% = - 30° C for piped gas

55. Thank you!!!

56. venturi effect coanada effect