RSPT 1060 MODULE C Lesson #3 Properties of Gases
OBJECTIVES • At the end of this module, the student will: • Define terms associated with the properties of gases and gas mixtures. • Identify abbreviations used during expressions of the properties of gases or gas mixtures. • State the six principle assumptions that explain the unique properties of gases. • State the effect of gas particle size on gas density. • State the effect of the distance between particles of a gas on its compressibility. • State the effect of kinetic energy of a gas on diffusion of that gas through an environment. • State the effect of molecular attraction of a gas on its physical properties.
OBJECTIVES • At the end of this module, the student will: • State the effect of kinetic energy of a gas on the pressure it exerts. • State the relationship between the kinetic activity of a gas, the temperature of the gas and the pressure it exerts. • List the major gases that comprise the atmosphere and their fractional concentrations. • Describe the clinical significance of hyperbaric & hypobaric conditions. • State Dalton’s Law. • Given the barometric pressure, calculate the partial pressure of a gas. • Explain how changes in altitude will affect partial pressure and fractional concentration of the gas. • Describe why gases need to be converted to Standard Conditions.
OBJECTIVES • At the end of this module, the student will: • Given a table of conversion factors, convert between the following standard conditions: • ATPS • STPD • BTPS • Describe how gases in the atmosphere differ from gases in the lungs and blood. • State the partial pressure of water vapor at body temperature (37 C) with a relative humidity of 100%. • List the normal values for various partial pressures found in the lungs and blood • State how pressure is measured. • Express one atmosphere in units commonly encountered in Respiratory Therapy. • Convert between units of pressure commonly encountered in Respiratory Therapy.
OBJECTIVES • At the end of this module, the student will: • Explain the significance of Avogadro's Law and number. • Calculate the gram-molecular weight of a gas. • Calculate the density of gases and gas mixtures commonly found in Respiratory Therapy. • Explain the relationship between density, mass and volume.
Kinetic-Molecular Theory Six principle assumptions that explain the unique properties of gases.
Assumption 1: Gas Particle Size Theory: Gas consists of tiny particles called molecules. Gas Property: Gas density is very low. (mass/volume)
Assumption 2: Distance Between Particles Gas Property:Gases are easily compressed. Theory: The distance between molecules of gas is very great compared to the size of the molecules. The volume occupied by a gas mainly consists of the empty space between molecules. Clinical example: Gas compressed into cylinders
Assumption 3: Gas Particle Movement Theory: Gas molecules are in rapid motion and move in straight lines, frequently colliding with each other and with the walls of the container. Gas Property: Gases diffuse. One gas can move through the empty space between another gas. Clinical example: odors moving across a room
Assumption 4: Molecular Attraction Theory: Gas molecules do not attract each other. Gas Property: Gases can fill any size container so they have no definite volume or shape. Gases are easily expanded. Gases flow. Clinical example: gas expands as it leaves a compressed gas cylinder.
Assumption 5: Collisions and Energy Loss Theory: When molecules of a gas collide with each other or with the walls of the container, they bounce back with no loss of energy. Such collisions are said to be perfectly elastic. Gas Property: Gas exerts pressure equally in all directions. Gases do not settle out and stop moving. Clinical examples: Gases may layer somewhat in the atmosphere (heavy near bottom) but molecular activity and gas flow limits this to some extent. In a gas cylinder the pressure is the same everywhere inside.
Assumption 6: Kinetic Energy and Temperature Theory: The average kinetic energy of the molecule is the same for all gases at the same temperature. The average kinetic energy increases as the temperature increases and decreases as the temperature decreases. Gas Property: Heating a gas increases the pressure it exerts on the walls of it’s container and cooling a gas decreases the pressure exerted on the walls of it’s container. Clinical examples: Gas cylinder in a fire can rupture. Gay-Lussac’s Law.
Question: • They have large spaced between molecules & atoms. • The have low densities • They have high densities • I ONLY • II ONLY • I and II ONLY • II and III ONLY • I, II, and III Why can gases be compressed easily?
Question: • Molecular particles of gases are in constant motion • Molecules of one gas cannot move in the spaces between other gas molecules • Gas molecules are highly attracted to each other • The distance between gas molecules is very small Which of the following is true regarding why a gas with an odor can travel across a room?
Question: • True • False All gas molecules eventually run out of energy and settle to the lowest surface.
Question: A. True B. False The pressure inside a compressed gas cylinder is greater at the bottom than it is at the top.
Question: • Increase the energy & increase the activity • Increase the energy but decrease the activity • Decrease the energy but increase the activity • Decrease the energy and decrease the activity • It has no affect on energy or activity What will heating a gas do to its kinetic energy & molecular activity?
Gas in the Atmosphere • Definitions • Relationship between: • Altitude • Pressure • Gas Concentrations
Definitions: • Pressure – (P)The force per unit of surface area (pounds per square inch or psi). Colliding molecules exert pressure. • Tension – Pressure of gas exerted in a liquid. • Fractional Concentration – Percent expressed as a decimal. Percent is parts per 100 parts. • Partial Pressure – Portion of the total pressure being exerted by one gas in a gas mixture.
Definitions: • Barometric Pressure – (PBARO) Pressure exerted by gases in the atmosphere also known as atmospheric pressure or ambient pressure. • One Atmosphere – (PATM) 760 mmHg pressure • Hypobaric – Pressure below 760 mmHg • Hyperbaric – Pressure above 760 mmHg • Water Vapor Pressure – (PH2O) Pressure exerted by water in the gas (vapor) form. 1034 cm H2O…14.7 psi…760 torr…29.92 in Hg
Relationships Relationship between: • Altitude • Pressure • Gas Concentrations
PO2 mmHg 18 29 47 73 110 159
Altitude • At Sea Level • Normobaric • One atmosphere (760 mmHg) • Above Sea Level • Hypobaric condition • < 760 mmHg • Below Sea Level • Hyperbaric condition • >760 mmHg
Sea Level • 760 mmHg • One atmosphere • Composition of air: • FO2 = .2095 (21%) • FN2 = .7809 (78%) • FCO2 = 0.0003 (0.03%) • Traces of Argon, Neon, Krypton, Hydrogen, Xenon. Ozone, Radon
Altitude, Pressure & Gas Concentrations Calculating Partial Pressure Pgas = Fractional Concentration x Barometric Pressure • Example: FO2 x PBaro = PO2 .2095 (21%) x 760 = 159.22 or159 mmHg
Above Sea Level Less then 760 mmHg • Hypobaric • Composition of air: • FO2 = .2095 (21%) • FN2 = .7809 (78%) • FCO2 = 0.0003 (0.03%) • Traces of Argon, Neon, Krypton, Hydrogen, Xenon. Ozone, Radon
Altitude and Pressure Changes • As altitude increases, gravitational pull decreases. • Gas molecules move apart. • Density decreases. • Molecule collisions decrease. • Pressure decreases.
Summary: Hypobarism and Pressure Changes • As the altitude increases the barometric pressure ___________. • As the altitude increases the gas composition (%) _______________. • As the altitude increases, the partial pressure exerted by each gas ________. • _____________ occurs at high altitudes because the oxygen pressure is too low.
Why are oxygen masks needed on airplanes? People who have traveled by air are familiar with the safety instructions given by the crew before flight. Instructions are included on how to use the oxygen masks. When and why are these masks needed? Cruising altitude 30,000 ft PBaro outside plane = 226 mmHg FO2 = _____________ PO2 = _____________
Why are oxygen masks needed on airplanes? Oxygen masks provide FO2 of 0.7 (70%) Cruising altitude 30,000 ft PB outside plane = 226 mmHg What PO2 will this provide? _______ Is this adequate to support life? ____
Below Sea Level Above 760 mmHg • Hyperbaric • Composition of air: • FO2 = .2095 (21%) • FN2 = .7809 (78%) • FCO2 = 0.0003 (0.03%) • Traces of Argon, Neon, Krypton, Hydrogen, Xenon. Ozone, Radon
Summary: Hyperbarism and Pressure Changes • As the altitude decreases the barometric pressure ___________. • As the altitude decreases the gas composition (%) ________________. • As the altitude decreases, the partial pressure exerted by each gas ________. • Hypoxemia __________ below sea level because the oxygen pressure is _____.
Hyperbaric Medicine • 2 – 3 Atmospheres - used to increase oxygen in the blood • Up to 6 atmospheres - used for high pressure therapy.
Egan Fig. 38-25 A, Fixed hyperbaric chamber. B, Monoplace chamber.
Hyperbarism: Indications • Indications for 2 – 3 Atmospheres • Used to increase partial pressure of oxygen in the blood • CO Poisoning • Wound care (kills anaerobic bacteria) • Gas Gangrene • Cyanide poisoning • Extreme blood loss (anemia)
Hyperbarism: Indications • Indications for Up to 6 atmospheres used for high pressure therapy. • Treatment for nitrogen bubbles in the blood stream. • The bends (decompression sickness) as a result of deep sea diving accidents. • Air embolism
Hyperbarism: Complications • Barotrauma of ears, gas in pleural space, gas bubbles in blood stream. • Oxygen toxicity • Fire hazards • Rapid decompression problems • Cataract worsening • Claustrophobia
Hyperbaric Oxygen Therapy • For additional information read: Egan - Hyperbaric Oxygen, pages 891 – 895
Dalton’s Law • “The total pressure of a gas mixture is equal to the sum of the partial pressures of its individual gases.” • PTOTAL = P1 + P2 + P3 + P4… • Each gas exerts a pressure proportional to its concentration.
Dalton’s Law • Partial Pressure - Partial pressure is the pressure exerted by a single gas. • PPartial = PTotal x Fractional Concentration • Example: 760 x 0.21 = 160 mmHg 760 x 0.4 = 304 mmHg • Example: • PAtm = PN2 + PO2 + PCO2 + PAr…..
Dalton’s Law - Practice • What is the total pressure in a dry gas mixture given the following information: • PO2 150 mmHg • PN2 500 mmHg • Anesthetic gas 8 mmHg Answer: 150 + 500 + 8 = 658 mmHg NOTE: Dry means there is no effect from water vapor pressure….
Water Vapor Pressure • Pressure exerted by water in the vapor (gaseous) form. • PH2O • In the lungs: • All gas is warmed to body temperature (37° C) and humidified to a saturation of 100% (fully saturated). • This results in a water vapor pressure (PH2O) of 47 mmHg. • It also results in a content (density) of 44 mg/L.
Water Vapor in Calculations • Before working with gas mixtures containing water vapor – remove the water vapor by subtracting the proper amount. • Inspired (I) (remove 47 mm Hg) • Dry(subtract nothing) • In between, the amount of water vapor will vary based on temperature.
Water Vapor - Practice • If the total pressure in a gas mixture is 730 mmHg, what is the partial pressure of the anesthetic gas if the PO2 is 150 mmHg, PH2O is 47 mmHg and PN2 is 480 mmHg?
Water Vapor Pressure • Question: • How do I know how much pressure the water vapor is exerting? • Answer: • It is based on two things: • Temperature of the gas • % saturation • It’s easy - Look at the chart!
Water Vapor Pressures & Contents at Selected Temperatures TABLE 6-3 Egan page 103
Dalton’s Law - Practice Gas fully saturated with water Temperature 37°C PBaro 760 mmHg FiO2 0.21 What is PO2? Gas fully saturated with water Temperature 32°C PBaro 760 mmHg FiO2 0.21 What is PO2?