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CHEM205 Thermodynamics.rtf

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CHEM 205
David Chen

THERMODYNAMICS Laws of Thermodynamics: 0th Law: absolute temperature (0K = no movement/vibrations) 1st Law: conservation of energy - keeping track of energy changes and allow us to calculate the heat a reaction produces 2nd Law: why some rxns occur but others don’t by keeping track of changes of both enthalpy and entropy (system and surroundings) 3rd Law: defines absolute entropy for a pure substance Ideal Gas (IG): PV = nRT Combined Gas Law: P V /1 1 P1V /T2 2 2 Dalton’s Law of Partial Pressure: P(total) = P1 + P2 + P3... Kinetic Model of Gas: - continuous/random motion - infinitesimally small points - move in straight lines until collision - do not interact each other except for collisions (i.e. no repulsive/attractive forces) E(kinetic) = ΔU = (f/2)RT where f is the degrees of freedom; 3 for monoatomic, 5 for diatomic Degrees of freedom are based on translational and rotational Van der Waals Equation: (P+an^2/V^2)(V-nb) = nRT 1. Repulsive interaction (Size effect): IG are assumed to have no volume but in reality gases take up some amount of volume. So, V is replaced by V - nb, where b is the volume/mole occupied by the gas. 2. Attractive forces: Molecules spend more time with each other during a collision and as a result, they collide less with the wall. Compressibility factor: Real Gas: z = pV/nRT Ideal Gas: z = 1 Virial Equation of State: z = PV/nRT = 1 + B2p*P + B3p*P + … where B2p, B3p... are Virial coefficients Terminology: - Irreversible: System =/= Surroundings - Reversible: System = Surroundings Thermodynamic Processes: - Isothermal: T = constant - Isobaric: P = constant - Isochoric: V = constant - Adiabatic: q = 0 (no heat enters/leaves) - Exothermic: q < 0 - Endothermic: q > 0 Properties: - System - part of universe we’re interested in - Surrounding - rest of universe - Open: allows exchange of matter and energy - Closed: allows exchange of energy - Isolated: does not allow exchange of matter or energy - State function: A => B depends only initial and final state; independent of path - Extensive: proportional amount of materials (i.e. m, V, E) - Intensive: depends on nature of materials (i.e. density, T, P) The First Law of Thermodynamics: Conservation of energy 1. Heat Transfer 2. Work Closed system: ΔU = UB - UA = q + w q - heat supplied TO the system w - work done ON the system Reversible and Irreversible Isothermal Expansions: Reversible: follows PV=nRT graph What is reversible? 1. IG Expansion/Compression: P = nRT/V 2. Heating/Cooling a stable substance 3. Chemical reaction at equilibrium 4. Phase transition along phase boundary Work: Irreverisible/constant external pressure: w = -PexΔV Reversible (Pex = P): dw = -PexdV => w = -nRT ln (V2/V1) Vacuum (Pex = 0): w = 0 ΔV = 0: w = 0, ΔU = Δqv Enthalpy H≡U + PV Isobaric: ΔH = qp Calorimetry: Measurement of Heat Heat Capacity C - the energy required to increase the temperature by 1o q = CΔT C≡q/ΔT (J K-1) ΔU = ∫ n*Cv,m dT = 3/2nRΔT Isobaric: Cp = n*Cp,m = (dq/dT)p ΔH = ∫ n*Cp,m dT = 5/2nRΔT Relationship between Cp and Cv: Cp - Cv = nR **Molar heat capacity does not exist for mixtures The Second Law of Thermodynamics - the entropy of an isolated system increases in the course of any spontaneous change - Spontaneous change - a change happens naturally without being driven by an external force - Nonspontaneous change - a change with no natural tendency Entropy: Entropy S - a measure of disorder Influenced by... - Thermal disorder - increase T = increase thermal motion - Positional disorder - spread into greater volume/mixing Reversible: dS = dq /Trev ΔS = ∫dq /T OR ΔSrevC ln (T /T ) 2 1 Irreversible: make hypothetical reversible path Isothermal: ΔS = q /Trev-w /T =revln (V /V ) O2 Δ1 = nRln (P /P ) 1 2 Isochoric: dq rev= n*CvdT = dU => ΔS = n*Cv ln (T /T ) 2 1 Isobaric: ΔS = n*Cp ln (T /2 )1 Adiabatic: ΔS = 0 Phase Chan
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