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CHEM 1125Q 1st Edition Lecture 4Outline of Last Lecture (Ch. 10)I.Outline of Current Lecture (Ch. 10- Ch. 11)II. Ch. 10 Cont. Lattice Energy and the Stability of Ionic CompoundsA. Born-Haber CycleIII. Ch. 11: GasesA. CharacteristicsIV. Kinetic Molecular TheoryV. Molecular SpeedA. Root-mean-square SpeedB. Diffusion Vs. EffusionVI. Graham’s LawVII. Gas PressureCurrent LectureII. Ch. 10 Cont. Lattice Energy and the Stability of Ionic CompoundsA. Born-Haber Cyclea. Cycle that analyzes reaction energies in order to determine lattice energies for ionic compoundsb. Combines rules of ΔH, Ionization Energy, and Electron Affinityc. Uses a method similar to Hess’s Law in format, in order to cancel out certain properties to obtain the final equation. The resulting energy value (kJ/mol) is the lattice energy of the compoundIII. Ch. 11: GasesA. Characteristicsa. Gas takes the shape and volume of any container which holds itb. The density of a gas is much lower than those of its corresponding liquids and solidsc. Density is also highly variable based on its environment’s temperatureand pressured. Gas is easily compressed into smaller volumes because of its lower densitye. It can form solutions and mixtures with other gasesIV. Kinetic Molecular TheoryThese notes represent a detailed interpretation of the professor’s lecture. GradeBuddy is best used as a supplement to your own notes, not as a substitute.A. Kinetic Molecular Theory is how a gas’s molecular nature corresponds to its macroscopic propertiesa. When particles are separated by long distances, the volume that individual molecules have is insignificantb. Molecules are constantly moving at random, whether in a straight path or colliding into other molecules with perfectly elastic collisionsc. Molecules neither attract nor repel each otherd. Average Kinetic Energy is proportional to absolute temperaturei. EK ∝ Te. Pressure is caused by gas molecules colliding with its container’s wallsi. When volume decreases, Density increases, and molecules collide more oftenii. Pressure increases as molecule collision frequency increasesf. Average Kinetic Energy (KE) of gas can be increased by heating iti. Molecules move faster when heatedii. When they move faster, they collide more often and at greaterspeedsg. The Total KE of one mole of a gas= (3/2)RTi. R= Universal Gas Constanth. The Average KE of one molecule= ½mu2i. m=massii. u2= average square speedi. The Average KE for one mole= NA(½mu2)=(3/2)RTi. NA= Avogadro’s NumberV. Molecular SpeedA. Root-mean-square Speeda. The root-mean-square speed (urms) is the speed of a molecule with an average KEb. urms = √((3RT)/ Μ)c. Μ = molar massd. urmsis inversely proportional to the square root of the molar mass and directly proportional to temperature e. When 2 gases are at the same temperature, urms is comparablef. (urms1 / urms2) = √(Μ2/Μ1)B. Diffusion Vs. Effusiona. Diffusion is when different gases mix as a result of frequent collisions and random motionb. Effusion is when gas molecules escape their container into a region of vacuumVI. Graham’s LawA. Graham’s Law states that the rates of diffusion and effusion are inversely proportional to the square root of molar massa. The heavier a gas is, the slower it movesVII. Gas PressureA. Pressure is force applied per unit of areaa. SI unit for Force = Newton (N)b. 1N = 1kg x m/s2c. SI unit for Pressure = Pascal (Pa)d. 1Pa = 1 N/m2e. Pascals are small units of pressuref. Pressure can also be measured in atm, mmHg, torr, and barg. 1 atm = 101.325 Pa (atm=Pressure at sea level)h. 1 mmHg = 133.322 Pa (mmHg = Barometer measure)i. 1 torr = 133.322 Pa (torr = name given to mmHg in honor of the Barometer creator)j. 1 bar = 1x105 Pa (same order of magnitude as atm)B. Pressure is measured by a Barometera. Standard Atmospheric Pressure = atmi. The pressure that would support a column of Hg 760mm


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