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:

• 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.

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.