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Chemistry (213)
Felix Lee (19)
Chapter 4

Discovering Chemistry Chapter 4.docx Discovering Chemistry Chapter 4.docx

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School
Western University
Department
Chemistry
Course
Chemistry 1027A/B
Professor
Felix Lee
Semester
Fall

Description
Discovering Chemistry Chapter 4 November 7, 2011  Systems and Surroundings o System: the part of the universe we are studying; usually the chemical reaction o Surroundings: the rest of the universe o System plus Surroundings equals Universe o System exchanges heat with Surroundings o Heat always flows from the warmer object to the colder object  What is Energy? o Energy is the capacity to do work o Heat energy; transferred due to temperature changes; symbol used is q o Kinetic Energy: due to motion o Potential Energy: Due to position o Chemical Energy: a form of potential energy o Thermodynamics focuses on:  Heat  Pressure-Volume (PV) or expansion work  Heat Capacity and Specific Heat o How much energy is required to raise the temperature of something by one degree o Heat capacity, C, is for objects  C has units J/degrees C or J/K o Specific heat capacity, c, is for substances  C has units j/g x degrees C or J/g x K o Molar heat capacity, C, is for substances using moles  C has units J/mol x degrees C or J/mol x K o Used to determine the amount of heat transferred  Q = C (delta) T or q = mc (delta) T or q = nC (delta) T o (delta)T =finalTinitial o (delta)T is the same in degrees C or K o See 4.1.1 o See 4.1.2  Sign Conventions o q always indicates the amount of heat transferred o We need to know if this heat is being lost of gained o If heat is lost, the sign of q is negative (-) and this is said to be an exothermic process o If heat is gained, the sign of q is positive (+) and this is said to be an endothermic process  Thermal Equilibrium o Two objects of different temperatures come into contact with one another, “thermal equilibrium” is when the two objects come to one final, common temperatufe, T o Heat is lost by the warmer object and gained by the cooler object:  - Heat lost = + Heat gained  -qwarmer +qcooler o See Example 4.1.3  The First Law of Thermodynamics o The total energy of a system and surroundings is conserved o Energy can be converted from one form to another o Cannot be created or destroyed  Intensive and Extensive Properties o Intensive Property: a physical property that does not depend on the amount of material in the system  Example: temperature, boiling point, melting point, density o Extensive Property: a physical property whose value is proportional to the size of the system (the amount of material present in the system)  Example: mass, volume, energy  State Functions o State functions depend only on the present state of the system; it doesn’t matter how it got there  Example: temperature, internal energy, enthalpy  Pressure-Volume Work o Work, w, is defined as the product of the force acting on an object and the distance, d, it moves in response to the force  W = Fd  Units are J, or Nm o Chemists are interested in pressure-volume (PV) work, the work involved with the expansion or compression of gases  W = -P(delta)V o If system does work on surroundings, w is negative (-) – lots of gas produced o If surroundings do work on system, w is positive (+) o Units of work are J; (1 Latm = 101.325 J or 1 LkPa = 1 J) o If there is no volume change, no work is done  Internal Energy o Heat and Work are equivalent ways of transferring energy in or out of a system o Internal Energy, E, is the sum of all the kinetic and potential energies of all the chemical species in the system; an extensive property.  E is a State Function o Absolute internal energy, E cannot be determined, but changes in internal energy, (delta)E, can be measured or calculated o If (delta)E > 0, the surroundings do work or supplies heat to the system o If (delta)E < 0, the system does work or supplies heat to the surroundings o (Delta)E is a State Function  Determination of Internal Energy o (Delta)E is the sum of the heat (q) and work done (w), on or by the system  (Delta)E = q + w o At constant volume: (delta)E = q  (delta)Ev= q o The internal energy change of a system is equal to the heat transferred when the volume does not change  Energy and Enthalpy o Enthalpy, H, is the energy content of a system at constant pressure: o Enthalpy is an extensive property (proportional to the amount of substance) and a state function  H = E + PV (by definition) o Changes in enthalpy can be determined:  (delta)H = (delta)E + (delta) PV o At constant pressure: (delta)H = (delta)E + P(delta)V  Heat Flow Expressions o Q is heat flow associated with the same process o (delta)H is heat flow for a process at constant pressure o (delta)H is heat flow for a process at standard conditions (1.00 atm or 101.325 kJ and 298.15 K) o At constant pressure, P =ext o So (delta)H = (delta)E extdelta)V o (delta)E = q + w and w = -P(delta)V ext o (delta)H = (q+w) + extdelta)V  Substitute for w o (delta)H = (q p  Physical Processes: Changes of State o Solid  liquid (melting, or fusion) o Liquid  gas (boiling, or vaporization) o Temperature of substance remains constant during a change of state or “phase” o Energy involved is known as latent heat; usually given on a molar basis  H o (s)  H O (l) q = 6.01 kJ/mol 2 2 o If process is carried out at constant pressure then q = (delta)H; if at standard conditions, (delta)H o See Example 4.2.1  Enthalpy Change of Chemical Reactions o (delta)H degreereaction{enthalpy (products) – enthalpy (reactangts) o When heat is released, system has lost energy and enthalpy has decreased (delta)H<0  Reaction is exothermic o When heat is absorbed, system has gained energy and enthalpy has increase (delta)H>0  Reaction is endothermic o If a reaction is reversed the sign of (delta)H changes  Measuring (Delta)H; Calorimetry o Calorimeter is a container used to measure heat transfer in chemical processes o Usually calibrated so that the amount of heat that is absorbed (the heat capacical C ) is known o Heat (q) flows between system and surroundings o Either sysr qsurrs assigned a negative sign  Q sys -Qsurrregardless of heat flow direction) o If reaction is exothermic, heat flows from system to surroundings  Q sys Q rxn -Qsurr  Simple Calorimeter o System is the reactants in solution o Surroundings is everything else (calorimeter and solution) o Heat absorbed by calorimeter is negligible o Solution absorbs heat o qsys -qsurrhich becomes:  qrxn= -qsolution  qsurr qsolutionmc(delta)T o See Example 4.2.2  Calorimeter Calibration o Calorimeters absorb or release heat and are part of the surroundings; therefore must be calibrated o Q rxn -Q surrnd –Q surr -(Qcalq soln o Q rxn -(QcalQ soln o The “calorimeter constant”, C cal units of J/degrees C or J/K o 1) Add hot water to calorimeter: q hot water-cal  (m hot waterchot waterdelta)Thot water -(cal (delta)Tcal o 2) Reaction with known (delta)H (or q): qrxn= -qcal  Q = -(C x (delta)T ) rxn cal cal o See Example 4.2.3  Bomb Calorimeter o Constant volume apparatus used to measure heat liberated during combustion reactions  q = q = -q = -(q + q ) sys rxn surr cal soln o See Example 4.2.4 o See Example
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