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123 .201 (3)

# Chemical Thermodynamics (Unit 1).docx

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School
Massey University
Department
123
Course
123 .201
Professor
John Harrison
Semester
Spring

Description
Unit 1 Notes Chemical Thermodynamics It describes the physical properties of a system at equilibrium, independent of time and on a macroscopic scale, providing a relationship between these properties. Basic Terminology The System and its Surroundings:  The system in a collection of materials which is undergoing change  The surroundings is everything else The surroundings and system can interact using heat, work and matter. Radiation is a form of interaction that can be ignored. Types of Systems: 1. Open A permeable boundary that can exchange matter and energy 2. Closed No exchange of matter (only energy) between system and surroundings 3. Isolated No exchange of matter or energy The properties of a system can be described as either intensive (do not depend on size of system) or extensive (dependent on size). Intensive Extensive Temperature (T) Mass (m) Pressure (p) Volume (V) Concentration (c) Number of Particles (N) Internal Energy (U) Entropy (S) 1 Unit 1 Notes Different Types of Boundaries:  Adiabatic: Does not allow heat to be transferred between the system and the Opposites surroundings  Diathermic: Allows heat to be transferred between the system and the surroundings  Movable: Work (pressure volume) transfer Boundaries can be permeable (allows all kinds or matter to pass), semi-permeable (allows some kinds of matter to pass) or impermeable (no matter is allowed to pass through). Energy, Work and Heat Work: causes orderly motion of surrounding particles ∫ x = distance Energy can be exchanged as work but also heat. The disorder or surroundings increases when heat transfers from a system. Exothermic Releases energy from the system as heat Endothermic Heat is absorbed by the system First Law of Thermodynamics Zeroth Law of Thermodynamics Systems in thermal contact come to the same temperature First Law of Thermodynamics Energy cannot be created or destroyed 𝑈 the change in internal energy q is the energy gained by the system w is the work done on the system For infinitesimal change: NB: The value for absolute U cannot be measured, only the ΔU. 2 Unit 1 Notes Work  Expansion Work/Pressure Volume Work (w ) pV  Electrical Work (w elvoltage x charge)  Surface Expansion (w surface surface tension x area change) Reversible and Irreversible Processes Expansion will occur when and compression will occur when . If then the process is reversible. Characteristics of a Reversible Process:  The driving force is approximately equal to the opposing force  Can assume equilibrium conditions (i.e. pV=nRT is maintained at all times for an ideal gas) NB: Expansion and Compression are Irreversible Processes Expansion Work or Pressure-Volume Work This type of work arises from changes in volume ∫ if p is constant then, ext Expansion (dV is positive) wpVs negative Compression (dV is negative) wpVs positive Reversible Process w pV-p .ext = -p int. ΔV If the system gets bigger then work is being done by the system (hence –ve) Enthalpy Determining the heat output/input during a process. If the process is at constant pressure with only w pVvolved: NB: enthalpy is a state function qvis heat at constant volume qpis heat at constant pressure NB: both are constants Internal energy has a fixed volume and enthalpy has a fixed pressure. 3 Unit 1 Notes Heat Capacities (JK mol ) -1 At constant pressure At constant volume Relationship between heat capacities for a perfect gas: pV = nRT State Functions and State Equations  A state function only depends on start and finish only  Path functions depend on the pathway taken (change between states are always path functions Equations of State: equations relating some state properties to other state properties i.e pV = nRT Adiabatic Expansion of a Perfect Gas In an adiabatic process no heat is allowed to enter or leave the system, even if the surroundings and system are at different temperatures. i.e q = 0 The temperature will drop during adiabatic expansion. 4 Unit 1 Notes Reversible Adiabatic Change for a Perfect Gas m means constant moles is the heat capacity ratio. For a perfect monatomic gas, it is 5/3. For 1 mol of a perfect gas Pressure-Volume Indicator Diagram (PV Diagram) Irreversible Expansion and Compression Cycle 5 Unit 1 Notes Overall Heat is transferred from the system wtotal to the surroundings Reversible Expansion and Compression Cycle ( ) ( ) Since Carnot Cycle Step 1: Isothermal Expansion (ΔU=0) Step 2: Adiabatic Expansion (q=0) Step 3: Isothermal Compression (ΔU=0) Step 4: Adiabatic Compression (q=0) Efficiency = useful work in/heat in 6 Unit 1 Notes Thermochemistry Standard Enthalpy The standard enthalpy change ΔH° is the change in enthalpy for a process in which the initial and final substances are in their standard states. A specified temperature in its pure form at 1 bar Standard Transition Enthalpy Change The physical state of a substance is Standard Enthalpy of Fusionfus H°) changed Standard Enthalpy of Vaporisationvap H°) Standard Enthalpy of Sublimation (Δ H°) sub Standard Reaction Enthalpy Change Reactants in their standard states change to products in their standard states v = stoichiometric coefficient ° ∑ ° ∑ ° Hm° = standard molar enthalpy Standard Enthalpy of Formation ° ∑ ° ∑ ° The standard enthalpies of formation of all the elements in their standard states are zero. Hess’s Law An enthalpy change for a given reaction can be found using the same combination of enthalpy changes of the other reactions. Temperatu
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