Chapter 14.docx

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Department
Chemistry & Biochemistry
Course
CHM 142
Professor
Dr.James D.Garrity
Semester
Summer

Description
First law of thermodynamics—energy is neither created nor destroyed Second law of thermodynamics  The amount of energy in the universe available to do work is constantly decreasing Spontaneous processes—can proceed without any outside intervention  Processes that are spontaneous in one direction are nonspontaneous in the reverse direction  No correlation between spontaneity and rate or exothermic vs. endothermic  A process that is spontaneous at one temperature, may not be spontaneous at another temperature  Spontaneity depends on dispersion of energy that occurs during a process Entropy (S)—multi-faceted concept  dispersion of energy at a given temperature  The extent of ―randomness‖ in a system  q/T units. J/K  Energy distribution is affected by molecular motion and volume  State funciton The motion of molecules is quantitized  Different molecular states related to molecular motion are separated by specific energies (don’t exist between energy levels/states) Energy State or Energy level—allowed value of energy Microstate—a unique distribution of molecules of a sample among energy levels (like a factorial) Three types of motions  Translational—movement through space  Rotational—spinning motion around axis perpendicular to bond  Vibrational—movement of atoms towards/away from each other The range of speeds of a molecule follows a Boltzmann distribution  The range of speeds increases as temperature increases The number of times a molecule occupies an accessible energy level also follows the Boltzmann solution  Number of accessible energy level increases as volume increases  Increasing volume stretches the energy of distribution o Accessible energy levels move closer together S=k*ln(W)  W=# of microstates ΔS = S fina– Sinitial Isothermal process—a process that takes place at a constant temp (ex. melting) Reversible process—a process that can be run in the reverse direction with no heat flow into or out of the system (idealized) For an isothermal process, ΔS sy=q re/T At 0C, ice and water coexist. No net change in composition of the mixture  To melt 1 mole of ice at 0C requires 6.01 kJ of energy absorbed from surroundings o ΔS sy=q re/T = 6.01*10^-3 J/ 273K = 22 J/K  The entropy of the universe increases in this spontaneous process Endothermic, ΔS = positive due to increased freedom of movement in dissolved ions ΔS univer= ΔS sys+ ΔS surr Third law of thermodynamics—entropy of a perfect crystal is 0 at absolute 0. Entropy increases as temperature increases  Increased KE increases the number of accessible microstate
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