Gen Chem Chapter 5 Reading Notes
Gen Chem Chapter 5 Reading Notes CHEM-111
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This 3 page Class Notes was uploaded by Madelyne Crawford on Wednesday September 28, 2016. The Class Notes belongs to CHEM-111 at Campbell University taught by Dr. Kesling in Fall 2016. Since its upload, it has received 5 views. For similar materials see General Chemistry in Chemistry at Campbell University.
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Date Created: 09/28/16
Chapter 5: Gases By Madelyne Crawford 5.1: Supersonic Skydiving and the Risk of Decompression Pressure- the force exerted per unit area by gas particles as they strike the surfaces around them o Constant force on all objects that are exposed to any sort of gas Total pressure exerted by gas depends on: o Concentration- the lower the concentration, the lower the pressure 5.2: Pressure- The Result of Molecular Collisions Constant collisions between atoms or molecules in a gas and surfaces around them= pressure Pressure = force / area Pressure decreases with increasing altitude Millimeter of mercury (mmHg)- common unit of measure for pressure Barometer- evacuated glass tube with tip submerged in mercury o When atmospheric pressure rises, mercury column rises too Torr- unit of millimeter of mercury o Named after Italian physicist Torricelli (1608-1647) o 1 mmHg = 1 torr Atmosphere (atm)- another unit of pressure; the average pressure at sea level o 1 atm= 760 mmHg pascal (Pa)- SI unit of pressure; 1 Pa = 1 N/m^2 o pascal is a lot smaller than the atm: 1 atm = 101,325 Pa 1 atm = 29.92 Hg and 1 atm = 14.7 psi (pounds per square inch) manometer- U-shaped tube containing dense liquid (usually mercury); measures pressure of a gas sample in a lab o if pressure of gas sample = atmospheric pressure, mercury levels on both sides will be the same o if pressure of gas sample > atmospheric pressure, Hg level on left is higher than on the right o if pressure of gas sample < atmospheric pressure, Hg level on right is higher than on the left o difference in height= difference between pressure and atmospheric pressure normal blood pressure: systolic= 100-119, diastolic= 60-79 5.3: The Simple Gas Laws- Boyle’s Law, Charles’s Law, and Avogadro’s Law simple gas laws- describe the relationship between pairs of these properties: pressure P, volume V, temp T, moles n Boyle’s Law- Boyle and Hooke used a J-tube to measure volume of a sample of gas at different pressures; discovered inverse relationship between volume and pressure Boyles Law = P1 x V1 = P2 x V2 o As pressure increases, volume decreases Most important rule of diving is to never hold breath because of Boyle’s law As long as temp and amount of gas remain constant, Boyle’s law can be used to calculate volume change following a pressure change (vice versa) Charles’s Law- V1 / T1 = V2 / T2 (1= initial 2= final) o J. A. C. Charles o When temp of gas sample increases, gas particles move faster o When air is heated, volume increases and makes less dense air Avogadro’s Law- when amount of gas increases at constant temp and pressure, volume increases in direct proportion- greater # gas particles fill more space o V1 / n1 = V2 / n2 5.4: The Ideal Gas Law Ideal gas law- PV = nRT ; ideal gas is a hypothetical gas that always follows this law o R (ideal gas constant) = 0.08206 (L X atm/mol X K) o Gay-Lussac’s law- relationship between pressure and temp; as temp increases, pressure increases 5.5: Applications of the Ideal Gas Law- Molar Volume, Density, and Molar Mas of Gas Molar volume- volume occupied by one mole of a substance Standard temp- T= 0 degrees C or 273 K Standard pressure- P=1.00 atm Density of gas= molar mass / molar volume ; density is directly proportional to molar mass o D= PM / RT 5.6: Mixtures of Gases and Partial Pressures Partial pressure- pressure due to any individual component in a gas mixture o P= n(RT / V) Dalton’s law of partial pressures- sum of partial pressures of components in gas mixture = total pressure Mole fraction- # moles of component in a mixture / total # moles Hypoxia- physiological condition produced by low oxygen levels Oxygen toxicity- caused by increased oxygen concentration in body tissues o Results in muscle twitching, tunnel vision, and convulsions Nitrogen narcosis- increase of nitrogen in the lungs; causes a drunk feeling in divers Vapor pressure- partial pressure of water in a mixture; depends on temp o Increases with increasing temp 5.7: Gases in Chemical Reactions- Stoichiometry Revised Calculation plan: P,V,T of gas A, then amount of A in mol, then amount of B in mol, PVT of gas B n= PV / RT V= nRT / P 5.8: Kinetic Molecular Theory- A Model for Gases Kinetic molecular theory- gas modeled as a collection of particles in constant motion; single particle moves in straight line until collision with another particle 3 Parts: o 1. Size of particle is very small (negligible) o 2. Avg kinetic energy of particle is proportional to temp (K) o 3. Collision of one particle with another is completely elastic (exchange of energy, but no overall loss of energy) o Kinetic energy = 0.5 X mv^2 In gas mixture, lighter particles travel faster than heavier ones 5.9: Mean Free Path, Diffusion, and Effusion of Gases Mean free path- average distance a molecule travels between collisions Diffusion- process of gas molecules spreading out in response to concentration gradient Effusion- process of gas escaping from a container into a vacuum through a small hole Graham’s Law of Effusion- rate A/rate B = square root(MA / MA) 5.10: Real Gases- The Effects of Size and Intermolecular Forces 1873- ideal gas equation changed to apply to all behaviors of real gases by Johannes van der Waals [P + a (n / v)^2] X [V – nb] = nRT Quizlet: https://quizlet.com/155631419/gen-chem-chapter-5-flash-cards/?new
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