GR

RSPT 1201 INTRO TO RESPIRATORY CARE UNIT II Lecture notes

ENERGY

  • Energy cannot be destroyed; it can only be stored or transferred.
  • All matter possesses internal energy:
    • Potential Energy: The energy of position (Stored energy)
      • The result of strong attractive forces between molecules
      • Is not relative to the environment
      • Examples: height/distance or mass (distance from earth or stretched)
    • Kinetic Energy: The energy of motion
      • All matter has some kinetic energy
      • Attractive forces in gases are weak, so energy in gases is mostly kinetic energy
        • Space among particles allows them to move freely
      • Is relative to the environment to other moving and stationary objects in its immediate environment.
      • Examples: speed/velocity & mass

STATES OF MATTER

  • There are 3 primary states of matter: solids, liquids, and gases.
    • Solids
      • Maintain fixed shape and volume
      • Have strong mutual attractive forces
      • Have the shortest distance to travel until they collide (packed tightly together)
      • Very low kinetic energy
      • Increasing pressure will not compress the solid to a smaller volume
    • Liquids
      • Molecules have less mutual attractive forces than solids
      • Molecules can move more freely (not enough room to flow around each other)
      • Quite dense
      • Cannot be easily compressed
      • More kinetic energy
      • Take shape of container
      • Viscosity: Thickness; opposition to flow
      • Cohesiveness: unified; stick together as one body
      • Objects have a Buoyant force (forces of gravity around the object have weak intermolecular forces)
      • Density: heavier substance particles are packed more closely together.
    • Gases
      • Have weak attractive forces
      • Exhibit rapid, random motion, with frequent collisions
      • Have no fixed volume or shape
      • High Kinetic energy
      • Expand to fill container
      • No definite volume or shape
      • Increase in temperature leads to increase in pressure
      • Increase in pressure leads to increase in collisions
      • Velocity: How fast something moves in a particular direction

Kinetic Molecular Theory

  • Applies to gases
    • No energy is lost during molecular collisions
    • The volume of molecules is negligible (does not matter)
    • No forces of mutual attraction between molecules (bounce off each other)
      *Kinetic energy/activity Pg. 85, 96, 103
  • Brownian motion: The random motion of smaller particles (suspended matter <3um) deposit in respiratory region of the lung where bulk gas flow cease and most aerosol particles reach the alveoli by depositing on surface walls and diffuse into the lungs.
    • 1827, Robert brown- molecular motion. (See Pg. 835 for further info)
    • Plasma has been referred to as the fourth state.

Absolute Zero, Critical Temperature & Temperature Conversions

  • TEMPERATURE CONVERSION:

    • Absolute Zero: A temperature at which all molecular activity ceases, there is no kinetic energy.
      • A logical zero point to build a temperature scale.
      • -273 ^\circ C = 0 K
      • -460 ^\circ F = 0 K
      • Although researchers have come close to obtaining it; no one has actually achieved it:
        • Celsius & Fahrenheit =Properties of water;
        • Kelvin = Molecular motion
          *Example
          0 ^\circ C = 273 K
          20 ^\circ C = 293 K
          320 K = (320-273) = 47 ^\circ C
          -460 ^\circ F = 0 ^\circ K
          See Figure 6-2
      • Conversion Formulas:
        • ^\circ C= (^\circ F-32) ÷ 1.8
          • Note: 1.8 is derived from 180^\circ F ÷ 100 ^\circ C
        • ^\circ F=[^\circ C x 1.8] +32
        • R = F+ 460 or 0^\circ Rankin = -460^\circ F
    • Problems:
      1. Convert 90 ^\circ F to ^\circ C
      2. Convert 40 ^\circ C to ^\circ F
    • At really high temperatures molecular density comes into play Alters relationship between Pressure and volume

  • Critical Temperature:

    • Is the highest temperature at which a substance can exist as a liquid.
    • Kinetic activity is so great, attractive forces cannot be kept in a liquid state.
    • Critical temperature of water is 374^\circ C. At or above this temperature a vapor can no longer be liquefied no matter how much pressure is applied.
    • Gases compared to liquids, have much lower critical points. (Ref. Egan Table 6-5).
      Gas Critical Points:
GasDegree CDegree FAtmospheric Pressure
Helium-267.9-450.22.3
Oxygen-118.8-181.149.7
Carbon Dioxide31.187.973
Nitrous Oxide36.597.771.8
  • Correction factors (page 102). These critical points Table 6-5.
  • Critical Pressure is the lowest pressure necessary at the critical temperature of a substance to maintain it in a liquid state.
    • This critical pressure maintains equilibrium between liquid and gas form.
    • When we heat water (to 374^\circ C) it drives the pressure up, a pressure of 218 atm is needed to maintain equilibrium between the liquid and gaseous forms of water.
      • No pressure can return water vapor to its liquid form at a temperature greater than 374^\circ C
      • Critical point is the highest temp and lowest pressure to maintain this equilibrium

GAS PRESSURE-ATMOSPHERIC PRESSURE

  • Measuring Atmospheric Pressure or air pressure: the force exerted on a surface by the air above it as gravity pulls it to the earth.
    • Gravity increases gas density, molecular collision (kinetic energy) and gas tension; this explains why atmospheric pressure decreases with altitude.
    • Tension refers to pressure when dissolved in liquid (blood, fluid)
    • Pressure = Height x density. Pressure exerted by a liquid depends on its height & weight (density)
    • The average atmospheric pressure at sea level is:
      Unit Equivalents:
    • 1 Atmospheric Pr. = (Sea Level is 1atm)
      • 760 mm Hg (76 cm Hg) (76 x 13.6)=1034 cm H_2O
      • 29.9 in. (29.9 x .491) = 14.7 PSI
      • 33.9 feet salt water
      • 1 cm H_2O = 0.735 mm Hg = 0.0142 PSI
      • 1 mm Hg = 1.36 cm H_2O = 0.019 PSI
      • 1PSI (Pounds per second) = 51.7 mm Hg = 70.34 cm H_2O

GAS PRESSURE PROBLEMS

  • Convert 500 mm Hg = cm H_2O
  • Convert 29.4 PSI = _ mm Hg
  • Convert 3100 cm H2O = _ mm Hg
  • At Denver the atmospheric pressure is 500 mm Hg. What would the pressure of oxygen be in the atmosphere on a dry day?

Dalton’s Law (Partial Pressure)

  • The total pressure of a mixture of gases must equal the sum of the partial pressures of all component gases.
  • The partial pressure of a component gas must be proportional to its percentage in the mixture.
    • Air contains 21% oxygen and 79% nitrogen. Assuming the atmospheric pressure is 760 torr,
      PO2 = 760 torr x .21 = 159.6 torr PN2 = 760 torr x .79 = 600.4 torr
      *760 torr *Torr is short for Torricelli evented mercury barometer in 17th century (1 torr =1mmHg) *If the total pressure changes, the pressures of individual gasses will change accordingly. However, the concentration of each gas will not change. *Increase in Atmospheric pressure (atm) results in increase in partial pressure exerted
      *Problem: 1. A heliox gas cylinder contains 70% helium. If the pressure of the gas cylinder is 2200 PSI, what is the pressure of helium in that gas cylinder?

Dalton's Law Problems

  • At a depth of 33 feet under the sea, water exerts a pressure of 1520 torr. If the pressure exerted by nitrogen is 1064 torr what is the percentage of nitrogen at that level?
  • The pressure exerted by gas X in a mixture is 350 mm Hg, and the total pressure is 1050 mm Hg, calculate the percentage of gas X in the mixture.
  • A gas cylinder contains 4 gases named A, B, C, and D. The pressures of gas A are 175 mm Hg, gas B is 22.7 mm Hg, gas C is 113.5 mm Hg, and gad D is 342 mm Hg. What is the total pressure (P_{TOTAL}) of this gas mixture?

Avogadro’s Law

  • Equal volume of gases at same temperature and pressure must contain the same number of molecules.

    • One gram atomic weight of any substance contains exactly the same number of atoms, molecules, or ions.
      • This number, 6.023 x 10^{23}, is Avogadro’s constant. This is one mole.
      • Thus, one mole of a gas, at a constant temperature and pressure, should occupy the same volume as one mole of any gas.
      • This ideal volume is termed the molar volume. At STPD the ideal volume of any molar gas is 22.4L.
      • A helium balloon weighs much less than a balloon filled with oxygen. The Balloons contain same # of molecules, since the GMW of helium is lower than N2 or O2, the helium balloon is lighter.

  • Density: A ratio of a substance’s mass to volume.

    • Density of a gas = Molecular weight ÷ Universal molar volume (22.4)
    • Density of a gas mixture = the sum of the % of each gas density in the mixture.
      Density of air = [(GMW x %N2) + (GMW x %O2)] ÷ 22.4 L = (28 x .79) gm + (32 x .21) gm ÷ 22.4 L = 1.29 gm/L
      Example Problem:
      *The atomic weight of nitrogen is 14. The formula for nitrogen is N_2 and its gram molecular weight is 28 grams. At normal temperature and pressure what is the density of Nitrogen?

Density Problems

Given the gram molecular weight of gas X is 15, gas Y is 8, and gas Z is 21; and the gas X=20%, gas Y=30%, and gas Z=50% of the mixture. Calculate the density of a gas mixture, at normal temperature and pressure?

Diffusion

  • The process whereby molecules move from an area of higher concentration to areas of lower concentration.
    • Gases have high kinetic energy and diffuse more rapidly

Graham’s Law (Gas Diffusion):
The rate of diffusion of a gas (D) is inversely proportional to the square root of its gram molecular weight.
D_{gas} = 1 ÷ \sqrt{GMW}

Diffusion is based on Kinetic Activity, anything that increases molecular activity quickens diffusion.

Lighter gases diffuse rapidly where heavier gases diffuse more slowly.

Fick’s first law of diffusion formula

The rate of diffusion of a gas into another gas is proportional to its concentration The bulk movement of gas through a biologic membrane (Vgas) (A=Cross Sectional area; D=Diffusion coefficient; T=Thickness; P1-2= Partial pressure gradient)

V{gas} = \frac{A X D}{T} (P1 –P_2)

CO2 and O2 move between and through: alveoli, capillary blood, cells, tissues, pressure gradients of lungs to maintain cellular metabolism and gas exchange.

The way transport occurs depends on:

Surface area
Diffusion constant
Concentration (pressure) gradient
Moving out CO2 and O2 in

Solubility of Gases in a Liquid

Henry’s Law
As the kinetic energy of the gaseous solute increases, its molecules have a greater tendency to escape the attraction of the solvent molecules and return to the gas phase. Therefore, the solubility of a gas decreases as the temperature increases and vice versa.

  • Hyperbaric Oxygen therapy
  • Gases can dissolve in liquids. Henry’s law predicts how much of a given gas will dissolve in a liquid. At a given temperature, the volume of a gas that dissolves in a liquid (V) equals its solubility coefficient (α) times its partial pressure (P{GAS}). Formula: V = α x P{GAS}

Carbonated water and soda are good examples of gas (CO2) dissolved in water (H2O).

Temperature plays a major role in gas solubility.

  • Solubility Coefficient:
    • The solubility coefficient equals the volume of a gas that will dissolve in 1ml of a given liquid at standard pressure and specified temperature.
    • The Sol Coefficient of O_2 at 37 degree Celsius 760 torr is .023
    • The Sol Coefficient of CO_2 at 37 degree Celsius and 760 torr is .510
      Blood gas vs patient temp.

GAS LAWS

  • Boyle’s Law: at constant temperature, the volume of gas varies inversely with the pressure exerted on it.
    *Smaller container, particles travel faster, they hit the walls more often.
    *Kinetic energy increases, increase in frequency leads to increase in pressure.
    *Pressure becomes larger as gas (volume) becomes smaller.
    *See Body plethysmograph pg. 411-12, Hyperbaric Chamber 922

  • Charles’ Law: If the pressure and the mass of a gas remain constant, the volume of the gas varies directly with the absolute temperature.

    • Average kinetic energy in a gas is proportional to the temperature of a gas.
    • Mass is constant.
    • Particles move faster as the temperature gets warmer, increasing kinetic energy; As they move faster, force exerted on wall lead to increase in pressure.
    • If the container is flexible it will expand until the pressures of gas balances pressure of the atmosphere.
      Volume becomes larger as temperature increases.

  • Gay Lussac’s Law: If the volume and mass of a gas remain constant, the pressure of the gas varies directly with the absolute temperature.

    • Kinetic energy only increases if the average velocity of the particles increases; the faster particles hit the wall of the container, the greater force exerted.
    • As temperature increases, kinetic energy/activity increases, pressure increases.

  • Avogadro’s Law: As number of gas molecules increase, the frequency of collisions do too, this leads to an increase in the pressure of gas.

    • Flexible containers (like a balloon) will expand until the pressure of the gas inside once again balances with the pressure on the outside.
    • ↑ in Kinetic energy, ↑ gmw, ↑ in volume

  • Combined gas Law Combines Pressure, Volume and temperature

  • Ideal Gas Law For the gas law to be Ideal, it must include: Pressure, Volume, Temperature and Density
    *Atmospheric pressure
    *Liquid oxygen is produced by compressing and cooling air at a temperature below its boiling point (-183° C or -297° F), and then separating oxygen from liquefied air mixture. After we separate it from air oxygen is stored in an insulated container below its boiling point.
    *Oxygen will remain liquid at atmospheric pressure as long as the temperature does not exceed –183 ° C. If at any time the liquid oxygen exceeds its critical temperature of -118.8° C, it converts immediately into a gas.