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THERMOCHEMISTRY Thermodynamics The study of Heat and Work and State Functions

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Thermodynamics The study of Heat and Work and State Functions Energy & Chemistry ENERGY is the capacity to do work or transfer heat. HEAT is the form of energy that ... – PowerPoint PPT presentation

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Title: THERMOCHEMISTRY Thermodynamics The study of Heat and Work and State Functions


1
THERMOCHEMISTRY Thermodynamics The study of Heat
and Work and State Functions
2
Energy Chemistry
  • ENERGY is the capacity to do work or transfer
    heat.
  • HEAT is the form of energy that flows between 2
    objects because of their difference in
    temperature.
  • Other forms of energy
  • light
  • electrical
  • kinetic and potential

3
Energy Chemistry
  • Burning peanuts supply sufficient energy to boil
    a cup of water.
  • Burning sugar (sugar reacts with KClO3, a strong
    oxidizing agent)

4
Energy Chemistry
  • These reactions are PRODUCT FAVORED
  • They proceed almost completely from reactants to
    products, perhaps with some outside assistance.

5
Energy Chemistry
  • 2 H2(g) O2(g) --gt
  • 2 H2O(g) heat and light
  • This can be set up to provide ELECTRIC ENERGY in
    a fuel cell.
  • Oxidation
  • 2 H2 ---gt 4 H 4 e-
  • Reduction
  • 4 e- O2 2 H2O ---gt 4 OH-

CCR, page 845
6
Potential Kinetic Energy
  • Potential energy energy a motionless body has
    by virtue of its position.

7
Potential Energy on the Atomic Scale
  • Positive and negative particles (ions) attract
    one another.
  • Two atoms can bond
  • As the particles attract they have a lower
    potential energy

NaCl composed of Na and Cl- ions.
8
Potential Energy on the Atomic Scale
  • Positive and negative particles (ions) attract
    one another.
  • Two atoms can bond
  • As the particles attract they have a lower
    potential energy

9
Potential Kinetic Energy
Kinetic energy energy of motion Translation
10
Potential Kinetic Energy
Kinetic energy energy of motion.
11
Internal Energy (E)
  • PE KE Internal energy (E or U)
  • Int. E of a chemical system depends on
  • number of particles
  • type of particles
  • temperature

12
Internal Energy (E)
  • PE KE Internal energy (E or U)

13
Internal Energy (E)
  • The higher the T the higher the internal energy
  • So, use changes in T (?T) to monitor changes in E
    (?E).

14
Thermodynamics
  • Thermodynamics is the science of heat (energy)
    transfer.

Heat energy is associated with molecular motions.
Heat transfers until thermal equilibrium is
established.
15
Directionality of Heat Transfer
  • Heat always transfer from hotter object to cooler
    one.
  • EXOthermic heat transfers from SYSTEM to
    SURROUNDINGS.

16
Directionality of Heat Transfer
  • Heat always transfer from hotter object to cooler
    one.
  • ENDOthermic heat transfers from SURROUNDINGS to
    the SYSTEM.

17
Energy Chemistry
  • All of thermodynamics depends on the law of
  • CONSERVATION OF ENERGY.
  • The total energy is unchanged in a chemical
    reaction.
  • If PE of products is less than reactants, the
    difference must be released as KE.

18
Energy Change in Chemical Processes
PE of system dropped. KE increased. Therefore,
you often feel a T increase.
19
UNITS OF ENERGY
  • 1 calorie heat required to raise temp. of 1.00
    g of H2O by 1.0 oC.
  • 1000 cal 1 kilocalorie 1 kcal
  • 1 kcal 1 Calorie (a food calorie)
  • But we use the unit called the JOULE
  • 1 cal 4.184 joules

20
HEAT CAPACITY
  • The heat required to raise an objects T by 1 C.

21
Specific Heat Capacity
  • How much energy is transferred due to T
    difference?
  • The heat (q) lost or gained is related to
  • a) sample mass
  • b) change in T and
  • c) specific heat capacity

22
Specific Heat Capacity
  • Substance Spec. Heat (J/gK)
  • H2O 4.184
  • Ethylene glycol 2.39
  • Al 0.897
  • glass 0.84

Aluminum
23
Specific Heat Capacity
  • If 25.0 g of Al cool from 310 oC to 37 oC, how
    many joules of heat energy are lost by the Al?

24
Specific Heat Capacity
  • If 25.0 g of Al cool from 310 oC to 37 oC, how
    many joules of heat energy are lost by the Al?

where ?T Tfinal - Tinitial q (0.897
J/gK)(25.0 g)(37 - 310)K q - 6120 J
Notice that the negative sign on q signals heat
lost by or transferred OUT of Al.
25
Heat Transfer No Change in State
  • q transferred (sp. ht.)(mass)(?T)

26
Heat Transfer with Change of State
  • Changes of state involve energy (at constant T)
  • Ice 333 J/g (heat of fusion) -----gt Liquid
    water
  • q (heat of fusion)(mass)

27
Heat Transfer and Changes of State
Liquid ---gt Vapor
  • Requires energy (heat).
  • This is the reason
  • a) you cool down after swimming
  • you use water to put out a fire.

energy
28
Heating/Cooling Curve for Water
Evaporate water
Heat water
Note that T is constant as ice melts
Melt ice
29
Heat Changes of State
  • What quantity of heat is required to melt 500. g
    of ice and heat the water to steam at 100 oC?

Heat of fusion of ice 333 J/g Specific heat of
water 4.2 J/gK Heat of vaporization 2260 J/g
30
Heat Changes of State
  • How much heat is required to melt 500. g of ice
    and heat the water to steam at 100 oC?
  • 1. To melt ice
  • q (500. g)(333 J/g) 1.67 x 105 J
  • 2. To raise water from 0 oC to 100 oC
  • q (500. g)(4.2 J/gK)(100 - 0)K 2.1 x 105
    J
  • 3. To evaporate water at 100 oC
  • q (500. g)(2260 J/g) 1.13 x 106 J
  • 4. Total heat energy 1.51 x 106 J 1510 kJ

31
Chemical Reactivity
  • What drives chemical reactions? How do they
    occur?
  • The first is answered by THERMODYNAMICS and the
    second by KINETICS.
  • Have already seen a number of driving forces
    for reactions that are PRODUCT-FAVORED.
  • formation of a precipitate
  • gas formation
  • H2O formation (acid-base reaction)
  • electron transfer in a battery

32
Chemical Reactivity
  • But energy transfer also allows us to predict
    reactivity.
  • In general, reactions that transfer energy to
    their surroundings are product-favored.

So, let us consider heat transfer in chemical
processes.
33
Heat Energy Transfer in a Physical Process
  • CO2 (s, -78 oC) ---gt CO2 (g, -78 oC)

Heat transfers from surroundings to system in
endothermic process.
34
Heat Energy Transfer in a Physical Process
  • CO2 (s, -78 oC) ---gt CO2 (g, -78 oC)
  • A regular array of molecules in a solid -----gt
    gas phase molecules.
  • Gas molecules have higher kinetic energy.

35
Energy Level Diagram for Heat Energy Transfer
36
Heat Energy Transfer in Physical Change
CO2 (s, -78 oC) ---gt CO2 (g, -78 oC) Two things
have happened!
  • Gas molecules have higher kinetic energy.
  • Also, WORK is done by the system in pushing aside
    the atmosphere.

37
FIRST LAW OF THERMODYNAMICS
  • ?E q w

Energy is conserved!
38
SYSTEM
?E q w
39
ENTHALPY
  • Most chemical reactions occur at constant P, so

Heat transferred at constant P qp qp ?H
where H enthalpy
and so ?E ?H w (and w is usually small) ?H
heat transferred at constant P ?E ?H change
in heat content of the system ?H Hfinal -
Hinitial
40
ENTHALPY
  • ?H Hfinal - Hinitial

If Hfinal gt Hinitial then ?H is positive Process
is ENDOTHERMIC
If Hfinal lt Hinitial then ?H is negative Process
is EXOTHERMIC
41
USING ENTHALPY
  • Consider the formation of water
  • H2(g) 1/2 O2(g) --gt H2O(g) 241.8 kJ

Exothermic reaction heat is a product and ?H
241.8 kJ
42
USING ENTHALPY
  • Making liquid H2O from H2 O2 involves two
    exothermic steps.

H2 O2 gas
Liquid H2O
H2O vapor
43
USING ENTHALPY
  • Making H2O from H2 involves two steps.
  • H2(g) 1/2 O2(g) ---gt H2O(g) 242 kJ
  • H2O(g) ---gt H2O(liq) 44 kJ
  • --------------------------------------------------
    ---------------------
  • H2(g) 1/2 O2(g) --gt H2O(liq) 286 kJ
  • Example of HESSS LAW
  • If a rxn. is the sum of 2 or more others, the net
    ?H is the sum of the ?Hs of the other rxns.

44
Hesss Law Energy Level Diagrams
Forming H2O can occur in a single step or in a
two steps. ?Htotal is the same no matter which
path is followed.
45
Hesss Law Energy Level Diagrams
  • Forming CO2 can occur in a single step or in a
    two steps.
  • ?Htotal is the same no matter which path is
    followed.

46
  • This equation is valid because ?H is a STATE
    FUNCTION
  • These depend only on the state of the system and
    not on how the system got there.
  • V, T, P, energy and your bank account!
  • Unlike V, T, and P, one cannot measure absolute
    H. Can only measure ?H.

47
Standard Enthalpy Values
  • Most ?H values are labeled ?Ho
  • Measured under standard conditions
  • P 1 bar 105 Pa 1 atm /1.01325
    Concentration 1 mol/L
  • T usually 25 oC
  • with all species in standard states
  • e.g., C graphite and O2 gas

48
Enthalpy Values
Depend on how the reaction is written and on
phases of reactants and products
  • H2(g) 1/2 O2(g) --gt H2O(g)
  • ?H -242 kJ
  • 2 H2(g) O2(g) --gt 2 H2O(g)
  • ?H -484 kJ
  • H2O(g) ---gt H2(g) 1/2 O2(g)
  • ?H 242 kJ
  • H2(g) 1/2 O2(g) --gt H2O(liquid)
  • ?H -286 kJ

49
Standard Enthalpy Values
  • NIST (Natl Institute for Standards and
    Technology) gives values of
  • ?Hfo standard molar enthalpy of formation
  • the enthalpy change when 1 mol of compound is
    formed from elements under standard conditions.
  • See Table 6.2

50
?Hfo, standard molar enthalpy of formation
  • Enthalpy change when 1 mol of compound is formed
    from the corresponding elements under standard
    conditions
  • H2(g) 1/2 O2(g) --gt H2O(g)
  • ?Hfo (H2O, g) -241.8 kJ/mol
  • By definition,
  • ?Hfo 0 for elements in their standard states.

51
Using Standard Enthalpy Values
  • Use ?Hs to calculate enthalpy change for
  • H2O(g) C(graphite) --gt H2(g) CO(g)

52
Using Standard Enthalpy Values
  • H2O(g) C(graphite) --gt H2(g) CO(g)
  • From reference books we find
  • H2(g) 1/2 O2(g) --gt H2O(g) ?Hf - 242
    kJ/mol
  • C(s) 1/2 O2(g) --gt CO(g) ?Hf - 111
    kJ/mol

53
Using Standard Enthalpy Values
  • H2O(g) --gt H2(g) 1/2 O2(g) ?Ho 242 kJ
  • C(s) 1/2 O2(g) --gt CO(g) ?Ho -111 kJ
  • --------------------------------------------------
    ------------------------------

H2O(g) C(graphite) --gt H2(g) CO(g)
?Honet 131 kJ
To convert 1 mol of water to 1 mol each of H2 and
CO requires 131 kJ of energy. The water gas
reaction is ENDOthermic.
54
Using Standard Enthalpy Values
Calculate ?H of reaction?
  • In general, when ALL enthalpies of formation are
    known

?Horxn ? ?Hfo (products) - ? ?Hfo (reactants)
Remember that ? always final initial
55
Using Standard Enthalpy Values
  • Calculate the heat of combustion of methanol,
    i.e., ?Horxn for
  • CH3OH(g) 3/2 O2(g) --gt CO2(g) 2 H2O(g)
  • ?Horxn ? ?Hfo (prod) - ? ?Hfo (react)

56
Using Standard Enthalpy Values
CH3OH(g) 3/2 O2(g) --gt CO2(g) 2 H2O(g)
?Horxn ? ?Hfo (prod) - ? ?Hfo (react)
  • ?Horxn ?Hfo (CO2) 2 ?Hfo (H2O)
  • - 3/2 ?Hfo (O2) ?Hfo (CH3OH)
  • (-393.5 kJ) 2 (-241.8 kJ)
  • - 0 (-201.5 kJ)
  • ?Horxn -675.6 kJ per mol of methanol

57
Measuring Heats of Reaction
CALORIMETRY
  • Constant Volume Bomb Calorimeter
  • Burn combustible sample.
  • Measure heat evolved in a reaction.
  • Derive ?E for reaction.

58
Calorimetry
Total heat evolved qtotal qwater qbomb
59
Measuring Heats of Reaction CALORIMETRY
  • Calculate heat of combustion of octane. C8H18
    25/2 O2 --gt 8 CO2 9 H2O
  • Burn 1.00 g of octane
  • Temp rises from 25.00 to 33.20 oC
  • Calorimeter contains 1200 g water
  • Heat capacity of bomb 837 J/K

60
Measuring Heats of Reaction CALORIMETRY
  • Step 1 Calc. heat transferred from reaction to
    water.
  • q (4.184 J/gK)(1200 g)(8.20 K) 41,170 J
  • Step 2 Calc. heat transferred from reaction to
    bomb.
  • q (bomb heat capacity)(?T)
  • (837 J/K)(8.20 K) 6860 J
  • Step 3 Total heat evolved
  • 41,170 J 6860 J 48,030 J
  • Heat of combustion of 1.00 g of octane -
    48.0 kJ
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