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CAS CH 109 (19)
Chapter 8

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CAS CH 109
Sean Elliott

Savan Shah CH 109 Chapter 8 Outline Chapter 8: Thermodynamics – The First Law I. Systems, States, and Energy - Systems • Thermodynamics – Study of transformations of energy from one form into another st • 1 Law – Concerned with keeping track of energy changes • 2 Law – Explains why some chemical reactions take place but others do not • System  Region we are interested in oOpen  Can exchange matter and energy with its surroundings oClosed  Can exchange energy but not matter oIsolated  Can exchange neither matter nor energy • Surrounding Where the system is immersed within • Universe = System + Surroundings • Work  Process of achieving motion against opposing force • Work = Opposing Force x Distance Moved 2 -2 • 1 J = 1 kg•m •s • 1 N = 1 kg•m•s -2 • 1 N x 1 m = 1 J • Energy  Capacity of system to do work oSystem does work on surroundings (Expansion)  Energy decreases oWork done onto system (Compression)  Energy increases oInternal Energy (U) • w = -P ex = External Pressure x Change in Volume (Pressure is constant) • ∆U = w (No type of transfer of energy is taking place) • Negative Energy  Energy leaving the system • Positive Energy  Energy entering the system • Expansion Work  Work arising from change in volume of system • Nonexpansion Work  Work that does not involve change in volume of system • 1 Pa • m = 1 J • If Pex 0, w = 0 (System does no expansion work when expanding into vacuum)  Free Expansion oReversible Process  One that can be reversed by infinitely small change in variable (infinitesimal); Ex: Pexncreased infinitesimally, piston moves inwards oIrreversible Process  Expansion against external pressure that differs by finite amount from pressure of system oIsothermal Expansion (Reversible Process)  Pressure of gas falls as it expands; To achieve reversible expansion, P mext decrease so that it equals pressure of gas V final  w = -nRT ln V initial -Heat  Energy transferred as result of temperature difference Savan Shah CH 109 Chapter 8 Outline • Flows from high-temps to low-temps • High-temp regions move more vigorously than low-temp regions • Heat = q • 1 cal = 4.184 J • Exothermic Process  Releases heat into surroundings (Negative E) • Endothermic Process Absorbs heat from surroundings (Positive E) • Adiabatic  Thermally insulating wall (∆U = w) • Diathermic  Wall that permits transfer of energy as heat (Typically raises temp of system) Heat supplied q • Heat Capacity (C)  Temperature RiseProduced  C = ∆T • q = C∆T • The larger the sample, the more heat is required to raise its temperature by given amount oSpecific Heat Capacity  C =sC/m oMolar Heat Capacity  C = m/n oSpecific Heat Capacity of water at room temp  4.18 J•°K •g-1 -1 oHeat supplied  q = mC ∆Ts oThe specific heat capacity of a dilute solution is normally taken to be same as that of pure solvent • Calorimeter  Device in which heat transfer is monitored by recording change in temperature produced by process taking place within it • ∆U = q + w oFirst Law  Internal energy of an isolate system in constant oState Function  Property that depends only current state of the system and is independent of how that state was prepared (Depend only on current state of the system, any change in its value is independent of how the change in state was brought about) oWork done by system is not state function  Depends on how the change is brought about oHeat is not a state function • System at high temperature has greater internal energy than the same system at lower temperature • Translational Energy  Energy of atom due to its motion through space • Rotational Energy  Energy of atom due to its rotational motion • Vibrational Energy  Sum of kinetic and potential contributions oKinetic contribution to vibration is due to relative motion of atoms in molecule oPotential contribution refle
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