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QUIZ

QUIZ. Please put away everything except a pencil or pen and calculator. RT Physics Week 2. Objectives. Atomic Particles , Number and Mass (weight) Henry’s and Grahams Law Avogadro’s Law Mass and Weight Density and Specific Gravity. Atoms.

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QUIZ

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  1. QUIZ • Please put away everything except a pencil or pen and calculator

  2. RT Physics Week 2

  3. Objectives • Atomic Particles , Number and Mass (weight) • Henry’s and Grahams Law • Avogadro’s Law • Mass and Weight • Density and Specific Gravity

  4. Atoms • Basic unit of matter that consists of a dense central nucleus surrounded by a cloud of negatively charged electrons. The atomic nucleus contains a mix of positively charged protons and electrically neutral neutrons • The electrons of an atom are bound to the nucleus by the electromagnetic force. • Likewise, a group of atoms can remain bound to each other by chemical bonds based on the same force, forming a molecule. An atom containing an equal number of protons and electrons is electrically neutral, otherwise it is positively or negatively charged and is known as an ion.

  5. Atoms • An atom is classified according to the number of protons and neutrons in its nucleus: the number of protons determines the chemical element, and the number of neutrons determines the isotope of the element

  6. Atomic Particles • subatomic particles are particles smaller than atoms • Basic structures include protons, neutrons and electrons

  7. Proton • Proton is a subatomic particle with a positive charge and a mass of 1 atomic mass unit and is located within the nucleus. When the number of protons in the nucleus is equal to the number of electrons orbiting the nucleus the atom is electrically neutral. • The number of protons contained in the nucleus = Atomic Number • Each element has a different atomic number • Ex: Oxygen atomic number is 8 • Hydrogen 1 • Helium 2

  8. Neutron • Neutron is also in the atoms nucleus • It is a subatomic particle with no electric charge and a mass of 1 amu • The sum of the neutrons inside the nucleus he number of the protons = atomic mass • Most atoms of elements can accommodate additional neutrons in their nucleus • Ex: Oxygen, three variations with varying atomic mass

  9. Electron • Electron is the lightest subatomic particle. It is a negatively charged particle. • Its mass is only 1/1,840 the mass of a proton. An electron is therefore considered to be mass-less in comparison with proton and neutron and is not included in calculating atomic mass of an atom.. Under ordinary conditions, electrons are bound to the positively charged nucleus by the attraction created from opposite electric charges. • An atom may have more or fewer electrons than the positive charge of its nucleus and thus be negatively or positively charged as a whole; a charged atom is called ion. In plasma, electrons exist in free state along with ions.

  10. Ions • An ion is a charged atom or molecule. It is charged because the number of electrons do not equal the number of protons in the atom or molecule. An atom can acquire a positive charge or a negative charge depending on whether the number of electrons in an atom is greater or less then the number of protons in the atom. • When an atom is attracted to another atom because it has an unequal number of electrons and protons, the atom is called an ION. • If the atom has more electrons than protons, it is a negative ion, or ANION (a – will be applied). If it has more protons than electrons, it is a positive ion or CATION (a + will be applied) Ex: Na+ Cl-

  11. What is the difference between a compound and a molecule? • A molecule is formed when two or more atoms join together chemically. • A compound is a molecule that contains at least two different elements. All compounds are molecules but not all molecules are compounds. • Molecular hydrogen (H2), molecular oxygen (O2) and molecular nitrogen (N2) are not compounds because each is composed of a single element. • Water (H2O), carbon dioxide (CO2) and methane (CH4) are compounds because each is made from more than one element.

  12. Chemical Bonding Electrovalent Bonding Covalent Bonding

  13. Chemical Bonding (type I) • Electrovalent (Ionic) Bonding • Two or more elements combine with each other by transferring electrons • One electron leaves outermost electron shell of one atom and enters outermost shell of other atom • Tendency in nature is for all atoms to seek eight electrons in their outermost electron shell (RULE OF EIGHT) • Ex: Na and Cl, Cl has 7 electrons in its outermost shell and Na has one. When the two elements react NA transfers its lone electron to Cl making Cl negatively charged, and sodium becomes positively charged since it lost an electron. Creating ions of Na and Cl • The bond that holds Na and Cl together is called electrovalent • The electrostatic attraction between these opposite charges maintains molecular integrity (NaCl)

  14. Chemical Bonding (type II) • Covalent Bonding • Results from the sharing of one or more pairs of electrons to form covalent molecules • Example: two Oxygen Molecules • Oxygen has two unpaired electrons in its outermost shell (total of 6), this combines with two from another Oxygen molecule to form O2

  15. Chemical Bonding • Some compounds are held together from a combination of both types of bonds • As elements react to form compounds, electrons are either shared or transferred or both • The electrons involved in these processes are called valence electrons • The valence of an atom is the number of electrons transferred (donated or received) in forming ions • Ex: Na has a valence of 1 because it donates one electron and forms the Na+ cation. The Cl- anion also has a valence of 1, it receives one electron and becomes Cl-. • Oxygen has a valence of 2, because two electrons are involved in the covalent bond formation • Some have multiple valences such as Fe, either Fe2+ or Fe3+

  16. Henry’s Law

  17. Henry’s Law • Describes the diffusion rate, or dissolving of a gas molecules into liquid. • It states “For a given temperature, the rate of a gas’s diffusion into a liquid is proportional to the partial pressure of that gas and its solubility coefficient”. • Does not apply to gases that chemically react with the solvent

  18. Henry’s Law • Increased pressure will increase diffusion. • Ex: In the Lung PAO2 represents the partial pressure of alveolar oxygen. The lung has other pressures such as Nitrogen, Co2 and Water Vapor, hence why PAO2 is called partial pressure • Increasing PAO2 with inspired Oxygen above room air (21%) will increase this partial pressure and allow more to enter the blood to attach to Hemoglobin and dissolve in plasma • In the blood the partial pressure of oxygen dissolved is PaO2 • PAO2 = Partial pressure of Alveolar oxygen (NORMAL 100 mmHg) • PaO2 = Partial pressure of Arterial oxygen (NORMAL 80-100 mmHg) • *Pressure alone does not determine diffusion, it also has to do with temperature and solubility of that gas. For example CO2 although far less pressure in the lung diffuses faster than Oxygen.

  19. Review of PAO2 (100 mmHg), PaO2 (90 mmHg), PvO2 (40mmHg) PACO2 (40 mmHg), PaCO2 (40 mmHg), PvCO2 (47 mmHg) PAGE 208

  20. Respiration • External respiration or breathing is a process of inhaling the air into the lungs and expelling the air that contains more carbon-di-oxide from the lungs to the outer environment. Exchange of gases in and out of the blood also takes place simultaneously. External respiration is a physical process in which oxygen is taken up by capillaries of lung alveoli and carbon-dioxide is released from blood. • Internal respiration or cellular respiration/tissue respiration is a metabolic process during which oxygen is released to tissues and carbon dioxide is absorbed by the blood. Once inside the cell the oxygen is utilized for production of energy in the form of ATP (adenosine triphosphate)

  21. Henry’s Law Continued Solubility Coefficient of Oxygen (only O2 and Co2 are capable of diffusing into the blood under normal circumstances, Nitrogen does not diffuse due to its large size) • Determined by: • Plasma as the solvent • 760 torr (sea level barometric pressure) 1 atm • Body Temperature Under these circumstances oxygen’s solubility coefficient has a value of 0.023. Meaning: • 0.023 ml of Oxygen can be dissolved by 1 ml of plasma • 0.023 ml O2/ml plasma / 760 torr PO2

  22. Henry’s Law Continued Solubility Coefficient of Oxygen Conversion Since 0.023 ml of Oxygen dissolves in 1 ml of plasma under a pressure of 760 torr of Oxygen, the amount of oxygen able to be dissolved in 1 ml of plasma for each torr of oxygen is determined by: 0.023 ml O2 /ml Plasma = 0.00003 ml O2/ml plasma/torr PO2 --------------------------------- 760 torr Further broken down to: 0.00003 mlO2/ml plasma/torr Po2 x 100 ml plasma= 0.003 ml O2/ 100 ml plasma/torr PO2

  23. Henry’s Law Continued Solubility Coefficient of Oxygen Conversion • Physiologic gas volumes are expressed in terms of volumes percent (vol%). • 0.003 ml O2/100 ml plasma/ torr becomes: • 0.003 vol%/torr PO2 • USED in the following formula: • Total Oxygen Content (formula to determine tissue oxygenation) • CaO2 = (Hb x 1.34 x SaO2) + (0.003 x PaO2) • Amount of O2 combined with Hb + Amount dissolved in Plasma

  24. CaO2 Formula • Total Oxygen Content (formula to determine tissue oxygenation) • CaO2 = (Hb x 1.34 x SaO2) + (0.003 x PaO2) NORMAL 17-21vol% • Amount of O2 combined with Hb + Amount dissolved in Plasma • Hb = Hemoglobin • Normal males: 13.5 to 17.5 grams per deciliter • Normal femailes: 12.0 to 15.5 grams per deciliter • 1.34 = number of millimeters of Oxygen that combine with each gram of hemoglobin • SaO2 = Percentage of saturated Hemoglobin with Oxygen • Normal 92-99% • 0.003= Solubility coefficient of O2 • PaO2 = Partial pressure of oxygen in the blood • Normal value 80-100 mmHg • EX: CaO2 = (15 x 1.34 x .95) + (0.003 x 90) • 19.1 + 0.27 = 19.37 vol%

  25. Solubility of CO2 • CO2 diffuses about 22 times faster than O2 into the blood • Based on: • Plasma being the solvent • Normal body temperature • 760 mmHg atmospheric blood pressure • The solubility coefficient for CO2 at 37C and 760 torr is 0.510 (A lot higher than O2) • 0.510 ml CO2/ml plasma/760 torr • Reduced to • 0.510 ml CO2/ml plasma = 0.00067 ml CO2/ml plasma/torr PCo2 • --------------------------------- 760 torr Further broken down to: 0.067 vol%

  26. PAO2 formula • Called the Alveolar air equation • PAO2 = FIO2 (PB- PH2O) – PaCO2 /RQ • PAO2 = Alveolar air equation • Normal 100 mmHg • FIO2 = Fractional inspired Oxygen • Room air = 21%, with supplemental Oxygen can increase to 100% • PB = Barometric pressure • At sea level = 760 mmHg • PH2O = Water vapor pressure • At normal BTPS = 47 mmHg • PaCO2= Partial pressure of Co2 in the blood • Normally 35-45 mmHg, varies with ventilation status • RQ= Respiratory quotient • Default number 0.8, this represents cellular uptake and elimination of O2 and CO2

  27. PAO2 formula • Inputting normal values: • PAO2= .21(760 mmHg-47mmHg)- 40 / 0.8 0.21 (713) – 50 149.73 – 50 = 99.73 (rounded to 100) This varies with changes in all the variables in the formula. Under normal lung conditions, a PAO2 of 100 results in a PaO2 of 80-100 mmHg

  28. Graham’s Law • States “the rate of diffusion of a gas through a liquid is directly proportional to its solubility coefficient and inversely proportional to the square root of its density”. • Describes the diffusion rate of one gas into another gas. • Gram molecular weight equals the number of particles in a given amount of matter. • Molecular weight of CO2 = 44.01 • Molecular weight of Oxygen = 31.99

  29. Grahams Law

  30. Graham’s Law • DENSITY: the mass per unit volume of a substance. The standard units for density are grams per cubic centimeter (gm/cm3), but the density of medical gases is usually expressed in grams per liter (g/l). • The medical community prefers numbers that do not have a zero to the left of the decimal point. You will see this later in the reporting of blood electrolytes and hydrogen ion concentration • The density of oxygen can therefore be expressed as 0.00143 gm/cm3 or 1.43 g/l. You will usually see the later unit set used in medical applications.

  31. Weight and Mass • 1. WEIGHT: measurement of the force of gravity acting upon an object (usually the earth's gravity). • 2. MASS: measurement of the amount of matter present in an object independent from the force of gravity acting on it.

  32. Weight and Mass • GRAM MOLECULAR WEIGHTS( GMW): The molecular weight of a substance in grams. To find the GMW of a medical gas we must know the atomic weights of several common chemical elements. • Substance SymbolAtomic Weight • A) Hydrogen H1 • B) Helium He4 • C) Carbon C12 • D) Nitrogen N14 • E) Oxygen O16 • F) Room Air 28.8 • NOTE: Nitrogen and Oxygen are found in the atmosphere in gaseous form as diatomic elements. So oxygen gas will have an atomic weight of 16 X 2 or 32, and nitrogen gas will have an atomic weight of 14 X 2 or 28.

  33. Room Air • Room air is not a pure substance; it is a mixture of gases. It contains about 79% nitrogen (N2) and 21% oxygen (O2) and small amounts of other gases. We can determine the relative GMW for room air by multiplying the fractional concentration of each gas by its molecular weight and adding the results. The GMW of room air can also be used to find the specific gravity of other medical gases because air is the usual standard for specific gravity of gases. • NitrogenOxygen • GMW air = (.79 x 28) + (.21 x 32) • =( 22.1 ) + ( 6.7 ) • GMW air = 28.8 • NOTE: The above method can also be used to find the relative GMW of any mixture of gases, ie: 60% He and 40% O2 or 95% O2 and 5% CO2.

  34. Avogadro's Law • "Equal volumes of gases at the same temperature and pressure contain the same number of molecules." • Therefore, the gram molecular weight (GMW) of a substance at STPD contains 6.023 X 1023 molecules and occupies a volume of 22.4 liters. • Avogadro's law can be used to find the density of medical gases.

  35. Density of Gases • Gases are influenced by changes in temperature and pressure • Calculates under STP conditions • Calculated by dividing volume occupied by 1 mole of gas at STP, that is 22.4 liters, into the gram of molecular weight of that gas • Density = gram molecular weight / 22.4 liters • Example: • Density of O2 = Weight of O2 32g /22.4 liters = 1.43g/L • Gases such as Helium have far less density • Oxygen has higher density than air and tends to accumulate at the lowest point (Ex: oxygen enclosure)

  36. Specific Gravity • The density of a material divided by or compared with the density of a standard. • In physics, the density of water is most often used as a standard. • When finding the specific gravity of a medical gas, air is used as the standard. • This is done because it is useful to know how the density of a medical gas compares with room air. • When the specific gravity of a substance is calculated, the standard must be indicated. This is done by indicating (Air = 1) for medical gases or (Water = 1) for liquids.

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