1. The gas laws By Dr. Ahmed Mostafa Assist. Prof. of anesthesia & I.C.U

2. Condition of measuring the volume of gases Standard temperature and pressure: - Temperature = 0 °C. -Pressure = 760 mmHg. -Water vapour pressure = Zero.

3. Condition of measuring the volume of gases Body temperature and pressure: - Temperature = 37 °C. -Pressure = 760 mmHg. -Water vapour pressure = 47 mmHg.

4. Condition of measuring the volume of gases Ambient temperature and pressure: - Temperature = 20 °C. -Pressure = 760 mmHg. -Water vapour pressure = 47 mmHg.

5. 1 st gas law(Boyle’s law) At a constant temperature, the volume of a mass of gas is inversely proportional to the pressure i.e. -V α1/P -VP = Constant (k1)

6. 1 st gas law(Boyle’s law) Clinical applications: Calculation of the amount of O2 that will be available at atmospheric pressure form an O2 cylinder: V1X P1 = V2 X P2

7. 2 nd gas law(Charles’ law) At a constant pressure the volume of a gas is directly proportional to its absolute temperature i.e. -V αT -V/T = Constant(k2). When a given mass of gas is heated or cooled at a constant pressure, its volume increases or decreases by 1/273 of its original volume at 0°C for each degree rise or fall in temperature respectively.

8. 2 nd gas law(Charles’ law) Clinical applications: - Gases expand when they are heated and become less dense, thus hot air rises.

9. 2 nd gas law(Charles’ law)

10. The 3 rd gas law(Gay-Lussac’s law) - At constant volume the absolute pressure varies directly with absolute temperature (P/T = Constant). Pressure is proportional to temperature. -Clinical applications: An example is the hydrogen thermometer. A constant volume of hydrogen when heated produces a change in pressure.

11. The 3 rd gas law(Gay-Lussac’s law)

12. The universal gas equation (the ideal gas low) If the perfect gas laws and Avogadro’s hypothesis are combined PV/T = Constant. For one mole of gas, PV/T equals the universal gas constant R. The equation can be rearranged to PV = nRT (the universal gas equation) where n equals the number of moles present.

13. The universal gas equation (the ideal gas low) The practical application of this law is the use of pressure gauges to assess the contents of a cylinder. The volume, temperature and gas constant remain the same and pressure is therefore proportional to n, the number of moles.

14. Perfect gas It is the gas that completely obeys all three gas laws. Or a gas that contains molecules of infinitely small size, which, therefore, occupy no volume themselves and which have no force of attraction between them. It is important to realize that this is a theoretical concept and no such gas actually exists. Hydrogen comes the closest to being a perfect gas as it has the lowest molecular weight. In practice, most commonly used anesthetic gases obey the gas laws reasonably well.

15. Dalton’s law of partial pressures In a mixture of gases, the pressure each gas exerts is the same as that which it would exert if it alone occupied the container. Clinical applications: 1- Calculation of partial pressure of gas in a mixture.

16. Dalton’s law of partial pressures In a mixture of gases, the pressure each gas exerts is the same as that which it would exert if it alone occupied the container. Clinical applications: 2- Manufacturing a cylinder producing 10% CO2 in O2 mixture.

17. Avogadro’s hypothesis Equal volumes of gases at the same temperature and pressure contain equal numbers of molecules. A mole is the quantity of a substance containing the same number of particles as there are atoms in 0.012 kg of carbon 12. The number of particles is: 6.022 x 10 23 (Avogadro’s number)

18. Avogadro’s hypothesis One mole of gas at standard temperature and pressure is contained in a volume of 22.4 liters. One mole of gas at temperature (20°C) is contained in a volume of 24 liters.

19. Avogadro’s hypothesis Clinical applications: to calculate the volume of nitrous oxide in a cylinder. A nitrous oxide cylinder contains 2.2 kg of nitrous oxide. The molecular weight of nitrous oxide is 44. One mole is 44 g. At STP we know that 44 g occupies 22.4 liters, therefore 2200 g occupies 22.4 x 2200/44 = 1120 liters.

20. Temperature scales Different thermometers use particular thermometric properties. For example, a mercury-in-glass thermometer uses the change in length of a column of mercury confined to a capillary tube of uniform bore; a platinum thermometer uses the increase in resistance with increasing temperature.

21. Temperature scales To establish a temperature scale it is necessary to make use of fixed points: a fixed point is the single temperature at which it can be confidently expected that a particular physical event always takes place.

22. Temperature scales The ice point: is the temperature at which pure ice can exist in equilibrium with water at standard atmospheric pressure. The steam point: is the temperature at which pure water is in equilibrium with its vapor at standard atmospheric pressure.

23. Temperature scales The triple point of water: Is that unique temperature at which pure ice, pure water and pure water vapor can exist together at equilibrium. The triple point is particularly useful, because there is only one pressure at which all three phases (solid, liquid and gas) can be in equilibrium with each other.

24. Temperature scales Critical temperature: is the temperature above which a gas cannot be liquefied however much pressure is applied (for CO2 Tc = 31.1 0 C). Critical pressure: is the minimum pressure that causes liquefaction of a gas at its critical temperature (for CO2pc = 73 atm). Specific critical volume: is the volume occupied by 1 kg of a gas at its critical temperature and pressure.

25. Temperature scales Therefore one can define a gas as a substance in the gaseous phase above its critical temperature. Vapor is the term applied to a substance in the gaseous phase below its critical temperature. Thus, simply increasing the pressure can liquefy a vapor, but not a gas. The relationship between pressure, volume and temperature is displayed as a family of isotherms.

26. Temperature scales Isotherms: (خط التحاور)

27. Temperature scales Oxygen, nitrogen and hydrogen are traditionally called permanent gases, because it was thought they could not be liquefied. This is because each of these gases has a critical temperature below room temperature (oxygen –118 o C, nitrogen – 146 o C, hydrogen –240 o C).

28. Temperature scales Poynting effect (overpressure effect): the critical temperature and pressure of a gas may be affected when it is mixed with another gas. For example, in a cylinder of Entonox, the new critical temperature of nitrous oxide (known as the pseudo-critical temperature) changes to –6 o C. Therefore, precautions regarding the cooling of cylinders should be taken into account.

29. Adiabatic compression or expansion of gases Adiabatic, when applied to the expansion or compression of a gas, means that heat energy is not added or removed when the changes occur. Thus when compression of a volume of gas occurs, it is accompanied by a temperature rise, and similarly expansion of a volume of gas will produce a temperature fall.

30. Adiabatic compression or expansion of gases Practical consequences of this are that compression of gases will require added cooling to avoid unwanted heating of the system. Alternatively, expansion of gases in the airway during jet ventilation can produce localised cooling, which in turn can reduce the humidity of injected gases.

31. Adiabatic compression or expansion of gases A practical application of the adiabatic expansion of gases lies in the cryoprobe. Here expansion of gas in the probe is used to produce low temperatures in the tip for cryotherapy.

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33. Thank you Dr. Ahmed Mostafa