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Chapter 5:. Thermochemistry. Thermochemistry:. Energy Kinetic & Potential First Law of Thermo internal energy, heat & work endothermic & exothermic processes state functions Enthalpy Enthalpies of Reaction. Calorimetry heat capacity and specific heat constant-pressure calorimetry

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Chapter 5:

Thermochemistry


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Thermochemistry:

  • Energy

    • Kinetic & Potential

  • First Law of Thermo

    • internal energy, heat & work

    • endothermic & exothermic processes

    • state functions

  • Enthalpy

  • Enthalpies of Reaction


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  • Calorimetry

    • heat capacity and specific heat

    • constant-pressure calorimetry

    • bomb calorimetry (constant-volume calorimetry)

  • Hess’s Law

  • Enthalpies of Formation

    • for calculation of enthalpies of reaction

  • Foods and Fuels


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Energy

  • work is a form of energy w = F x d

  • energy is the capacity to do work or transfer heat

  • Kinetic Energy

    • energy of motion E = ½ mv2

  • potential energy

    • energy of position

    • applies to electrostatic energy

    • applies to chemical energy (energy of bonds)


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  • energy units

    • one joule = energy of a 2 kg mass moving at 1 m/s

    • E = ½ mv2 (½)(2 kg) (m/s)2 = kg m2/s2 = 1 J

    • 1 cal = 4.184 J 1 kcal = 1 food calorie (Cal)

      Systems & Surroundings

  • system -- chemicals in the reaction

  • surroundings -- container & all outside environment

  • closed system can exchange energy (but not matter) with its surroundings


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Closed System

energy(as heat or work)

2H2(g) + O2(g)

¯

2H2O(l)+ energy

(system)

no exchg of matterwith surroundings


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First Law of Thermo.

  • Energy is always conserved

    • any energy lost by system, must be gained by surroundings

  • Internal Energy -- total energy of system

    • combination of all potential and kinetic energy of system

    • incl. motions & interactions of of all components

    • we measure the changes in energy E = Efinal - Einitial


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  • + D E = Efinal > Einitial system has gained E from surroundings

  • - D E = Efinal < Einitial system has lost E to surroundings

  • Relating D E to heat and work

    • D E = q + w q is positive if heat goes from surroundings to system w is positive if work is done on system by surroundings


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+q +w

- q - w

surroundings

surroundings

system

system

heat

heat

work

work


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Endothermic

  • system absorbs heat or heat flows into the system

    Exothermic

  • system gives off heat or heat flows out of the system

    State Function

  • a property of a system that is determined by specifying its condition or state (T, P, etc.)

  • internal energy is a state function, \DE depends only on Efinal & Einitial


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Enthalpy

  • for most reactions, most of the energy exchanged is in the form of heat, that heat transfer is called enthalpy, H

  • enthalpy is a state function

  • like internal energy, we can only measure the change in enthalpy, DH

    • DH = qp when the process occurs under constant pressureDH = Hfinal - Hinitial = qp

  • - DH Þ exothermic process

  • +DH Þ endothermic process


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system

DH > 0

surroundings

system

surroundings

DH < 0


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Enthalpies of Reaction

  • D Hrxn = Hprod - Hreact

  • enthalpy is an extensive property

    • magnitude of D H depends directly on the amount of reactant

    • C(s) + 2H2(g)® CH4(g)D H = -74.8 kJ/mol

    • 2C(s) + 4H2(g)® 2CH4(g)D H = -149.6 kJ/2mol


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    • enthalpy change for forward rxn is equal in magnitude but opposite in sign for the reverse rxn

      • CH4(g) ® C(s) + 2H2(g)D H = +74.8 kJ/mol

      • C(s) + 2H2(g)® CH4(g) D H = - 74.8 kJ/mol

    • enthalpy change for a reaction depends on the state of the reactants and products

      • C(g) + 2H2(g)® CH4(g)D H = -793.2 kJ/mol

      • 2H2(g) + O2(g)® 2H2O(g)D H = -486.6 kJ/mol

      • 2H2(g)+ O2(g)® 2H2O(l)D H = -571.7 kJ/mol


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    DH = Hfinal - Hinitial

    H2O(g)

    -241.8 kJ

    44 kJ

    -

    +

    Enthalpy

    H2O(l)

    -285.8 kJ


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    Practice Ex. 5.2:

    • Hydrogen peroxide can decompose to water and oxygen . Calculate the value of q when 5.00 g of H2O2(l) decomposes at constant pressure.

      • 2H2O2(l)® 2H2O(l) + O2(g)D H = -196 kJ

      • 5.00 g H2O2(l) x 1 mol = 0.147 mol H2O2(l) 34.0 g H2O2(l)

      • 0.147 mol H2O2(l) x -196 kJ H2O2(l) = -14.4 kJ 2 mol


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    Calorimetry

    • experimental determination of D H using heat flow

      heat capacity

    • measures the energy absorbed using temperature change

    • the amount of heat required to raise its temp. by 1 K

    • molar heat capacity -- heat capacity of 1 mol of substance


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    specific heat

    • heat energy required to raise some mass of a substance to some different temp.

    • specific heat = quantity of heat trans. (g substance) (temp. change)

    • = q . m DT

    • S.H. = joule g K

    • q = (S.H.) (g substance) (D T)

    remember: this is change in temp.


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    Practice Ex. 5.3:

    • Calculate the quantity of heat absorbed by 50.0 kg of rocks if their temp. increases by 12.0 °C if the specific heat of the rocks is 0.82 J/gK.

      • S.H. x g x DT = joules

      • What unit should be in the solution?

      • joules -- quantity of heat

      • 0.82 J x 50.0 x 103 g x 12.0 K = 4.9 x 105 Jg K


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    Constant-Pressure Calorimetry

    • D H = qp at constant pressure as in coffee cup calorimeter

      • heat gained by solution = qsoln

      • \qsoln = (S.H.soln)(gsoln)(DT)

      • heat gained by solution must that which is given off by reaction

      • \qrxn = - qsoln = - (S.H.soln)(gsoln)(DT)

    must be opposite in sign

    if DT is positive then qrxn is exothermic


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    Practice Ex. 5.4:

    • When 50.0 mL of 0.100 M AgNO3 and 50.0 mL of 0.100 M HCl are mixed in a c.p. calorimeter, the temp. of the mixture increases from 22.30°C to 23.11°C. Calculate D H for this reaction, assuming that the combined solution has a mass of 100.0 g and a S.H. = 4.18 J/g °C.

      • AgNO3(aq) + HCl(aq) ® AgCl(s) + HNO3(aq)

      • qsoln = 4.18 J x 100.0 g soln x 0.81°C = 3.39 x 102 J g °C

      • qrxn = - qsoln = - 3.39 x 102 J = - 68,000 J or 0.00500 mol - 68 kJ/mol


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    insulating cup

    rxn

    q

    soln


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    Bomb Calorimetry (Constant-Volume)

    • bomb calorimeter has a pre-determined heat capacity

    • sample is combusted in the calorimeter and D T is used to determine the heat change of the reaction

    • qrxn = - Ccalorimeter x D T

    heat capacity of calorimeter

    because rxn is exothermic


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    thermometer

    insulation

    water

    rxn


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    Practice Ex. 5.5:

    • A 0.5865 g sample of lactic acid, HC3H5O3, is burned in a calorimeter with C = 4.812 kJ/°C. Temp. increases from 23.10°C to 24.95°C. Calculate heat of combustion per gram and per mole. D T = +1.85°C

      • qrxn = - (4.812 kJ/°C) (1.85°C) = - 8.90 kJ per 0.5865 g lactic acid-8.90 kJ = - 15.2 kJ/g 0.5865 g

      • - 15.2 kJ x 90 .1 g = - 1370 kJ/mol 1 g 1 mol


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    Hess’s Law

    • rxns in one step or multiple steps are additive because they are state functions

      • eg.

      • CH4(g) + 2O2(g)® CO2(g) + 2H2O(g)D H = - 802 kJ

      • CH4(g) + 2O2(g)® CO2(g) + 2H2O(l)D H = - 890 kJ

    2H2O(g)® 2H2O(l)D H = - 88 kJ


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    Practice Ex. 5.6:

    • Calculate D H for the conversion of graphite to diamond:

      • Cgraphite® Cdiamond

      • Cgraphite + O2(g)® CO2(g)D H = -393.5 kJ

      • Cdiamond + O2(g)® CO2(g)D H = -395.4 kJ

      • Cgraphite + O2(g)® CO2(g)D H = -393.5 kJ

      • CO2(g)® Cdiamond + O2(g)D H = 395.4 kJ

    Cgraphite® CdiamondD H = + 1.9 kJ


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    Enthalpies of Formation

    • enthalpies are tabulated for many processes

      • vaporization, fusion, formation, etc.

    • enthalpy of formation describes the change in heat when a compound is formed from its constituent elements, DHf

    • standard enthalpy of formation, DHfo, are values for a rxn that forms 1 mol of the compound from its elements under standard conditions, 298 K, 1 atm


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    • For elemental forms:

      • eg. C(s) graphite, Ag(s) , H2(g) , O2(g) , etc.

      • DHfo, for any element is = 0

    • used for calculation of enthalpies of reaction, DHrxn

    • DHrxn = S DHfo prod - S DHfo react


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    Practice Ex. 5.9:

    • Given this standard enthalpy of reaction, use the standard enthalpies of formation to calculate the standard enthalpy of formation of CuO(s)(

      • CuO(a) + H2(g)® Cu(s) + H2O(l)DHo = -130.6 kJ

      • DHrxn = S DH f o prod - S DH f o react

      • -130.6 kJ = [(0) + (-285.8)] - [(CuO) + (0)]

      • DHfo CuO = -155.2 kJ/mol


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