THERMOCHEMISTRY OR THERMODYNAMICS

1
THERMOCHEMISTRY OR THERMODYNAMICS

• Chapter 6

2
Energy and Chemistry

• ENERGY is the capacity to do work or transfer
  heat.
• HEAT is the form of energy that flows between 2
  samples because of their difference in
  temperature.
• Other forms of energy
• light electrical nuclear
• kinetic potential

3
Temperature v. Heat

• Temperature reflects random motions of
  particles, therefore related to kinetic energy of
  the system.
• Heat involves a transfer of energy between 2
  objects due to a temperature difference

4
Law of Conservation of Energy

• Energy can be converted from one form to another
  but can neither be created nor destroyed.
• (Euniverse is constant)

5
Kinetic and Potential Energy

• Potential energy energy a motionless body has
  by virtue of its position.

6
Kinetic and Potential Energy
Kinetic energy energy of motion.
7
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

James Joule 1818-1889
8
Extensive Intensive Properties

• Extensive properties depends directly on the
  amount of substance present.
• mass
• volume
• heat
• heat capacity (C)

• Intensive properties is not related to the amount
  of substance.
• temperature
• concentration
• pressure
• specific heat (s)

9
System and Surroundings

• System That on which we focus attention
• Surroundings Everything else in the universe
• Universe System Surroundings

10
Exo and Endothermic

• Heat exchange accompanies chemical reactions.
• Exothermic Heat flows out of the system (to
  the surroundings).
• Endothermic Heat flows into the system (from
  the surroundings).

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Endo- and Exothermic
Surroundings
Surroundings
System
System
heat
heat
E(system) goes down
E(system) goes up
ENDOTHERMIC
EXOTHERMIC
12
Enthalpy

• DH Hfinal - Hinitial
• If Hfinal gt Hinitial then DH is positive
• Process is ENDOTHERMIC
• If Hfinal lt Hinitial then DH is negative
• Process is EXOTHERMIC

13
Upon adding potassium hydroxide to water the
following reaction takes place NaOH(S)
NaOH(aq) for this reaction at constant pressure,
?H -43 kj/mol1 .Is the reaction exo- or
endothermic2. Does the water get warmer?3. What
is the enthalpy change for the solution if 14 g
of NaOH is added?
14
Exercise

• Consider the combustion of propane
• C3H8(g) 5O2(g) ? 3CO2(g) 4H2O(l)
• ?H 2221 kJ
• Assume that all of the heat comes from the
  combustion of propane. Calculate ?H in which 5.00
  g of propane is burned in excess oxygen at
  constant pressure.
• 252 kJ

15

• Consider the reaction
• H2(g) O2(g) H2O(l) DH
  286 kJ
• Which of the following is true?
• a) The reaction is exothermic.
• b) The reaction is endothermic.
• c) The enthalpy of the products is less than that
  of the reactants.
• d) Heat is absorbed by the system.
• e)         Both A and C are true. AC

16

• Consider the reaction
• When a 24.8-g sample of ethyl alcohol (molar
  mass 46.07 g/mol) is burned, how much energy is
  released as heat?
• c

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• The total volume of hydrogen gas needed to fill
  the Hindenburg was 2.01x108 L at 1.00 atm and
  24.7C.
• How much energy was evolved when it burned?
• 2H2(g) O2(g) 2H2O(l) DH 286 kJ

18

  2.35 109 kJ
7.37 102 kJ
19
q msDt
q heat (J) m mass (g) s specific
heat (j/gCo) Dt change in temperature (Co)
Simple Calorimeter
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Some Heat Exchange Terms

• specific heat capacity (s)
• heat capacity per gram J/C g or J/K g
• molar heat capacity (s)
• heat capacity per mole J/C mol or J/K mol

21
Heat Capacity
22
Specific Heat Capacity

• Substance Spec. Heat (J/gK)
• H2O 4.184
• Al 0.902
• 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?
• Spec of Al0.902
• where DT Tfinal - Tinitial

heat gain/lost q m s DT
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 DT Tfinal - Tinitial
• q (0.902 J/gK)(25.0 g)(37 - 310)
• q - 6160 J

heat gain/lost q m s DT
25
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?
• q - 6160 J
• Notice that the negative sign on q signals heat
  lost by or transferred out of Al.bv 233hhbn

26

• Copper has a specific hear of .382j/goC. If 2.51
  g of cooper absorbs 2.75 j of heat , what is the
  change in temp ?

27

• Cooper has a specific heat of .382j/g/ oC. the
  temperature of an unknown mass of cooper
  increases by 4.50 oC when it absorbs 3.97J of
  heat. What is the mass of the copper?

28

• Heating curves

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REMEMBER!!!

• In regular calorimetry pressure is constant, but
  the volume will change.
• In bomb calorimetry, volume is constant.

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Calorimetry

• Constant volume calorimeter is called a bomb
  calorimeter.
• Material is put in a container with pure oxygen.
  Wires are used to start the combustion. The
  container is put into a container of water.
• The heat capacity of the calorimeter is known and
  tested.

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Bomb Calorimeter

• thermometer
• stirrer
• full of water
• ignition wire
• Steel bomb
• sample

32

• Suppose we wish to measure the energy of
  combustion of octane (C8H18), a component of
  gasoline. A 0.5269-g sample of octane is placed
  in a bomb calorimeter known to have a heat
  capacity of 11.3 kJ/ºC.
• This means that 11.3 kJ of energy is required to
  raise the temperature of the water and other
  parts of the calorimeter by 1ºC. The octane is
  ignited in the presence of excess oxygen, and the
  temperature increase of the calorimeter is
  2.25ºC.
•  

33

• The amount of energy released is calculated as
  follows
• Energy released by the reaction
• ?T x heat capacity of calorimeter

34

• A bomb calorimeter has a heat capacity of 9.47
  kJ/K. When a 2.01-g sample of (C3H6) was burned
  in this calorimeter, the temperature increased by
  4.26 K. Calculate the energy of combustion for
  the sample.

35

• A bomb calorimeter has a heat capacity of 9.47
  kJ/K. When a 2.01-g sample of (C3H6) was burned
  in this calorimeter, the temperature increased by
  4.26 K. Calculate the energy of combustion for
  the sample.

36

• A 2.200-g sample of quinine (C6H4O2) is burned in
  a bomb calorimeter whose total heat capacity is
  7.854kj/ºC. the temperature of the calorimeter
  plus the contents increased from 23.44ºC to
  30.57ºC.
• What is the heat of combustion?
• per gram of quinine?
• Per mole of Quinine?

37

• .5865g sample of lactic acid HC3H5O3 is burned in
  a calorimeter whose heat capacity is 4.812kj/ºC.
  The temperature increases from 23.10ºC to 24.95ºC
  . Calculate the heat of combustion of lactic acid
  per gram ?

38

• The heat of combustion of pentane, is -131.64
  kJ/g. Combustion of 4.50 g of pentane causes a
  temperature rise of 2.00C in a certain bomb
  calorimeter. What is the heat capacity of this
  bomb calorimeter?

39

• The combustion of 0.1584g benzoic acid increase
  the Temperature of a bomb calorimeter by 2.54C.
  The energy released by the combustion is
  26.42kj/g.Calculate the heat capacity of the bomb
  Calorimeter .

40
Standard States

• Compound
• For a gas, pressure is exactly 1 atmosphere.
• For a solution, concentration is exactly 1 molar.
• Pure substance (liquid or solid), it is the pure
  liquid or solid.
• Element
• The form N2(g), K(s) in which it exists at
  1 atm and 25C.

41
Hesss Law

• Reactants ? Products
• The change in enthalpy is the same whether the
  reaction takes place in one step or a series of
  steps.

42
Using Enthalpy

• Consider the decomposition of water
• H2O(g) 286 kJ ---gt H2(g) 1/2 O2(g)
• Endothermic reaction heat is a reactant
• DH 286 kJ

43
Using Enthalpy

• Making H2 from H2O involves two steps.
• H2O(l) 44 kJ ---gt H2O(g)
• H2O(g) 242 kJ ---gt H2(g) 1/2 O2(g)
• --------------------------------------------------
  ---------------
• H2O(l) 286 kJ --gt H2(g) 1/2 O2(g)
• Example of HESSS LAW
• If a rxn. is the sum of 2 or more others, the net
  DH is the sum of the DHs of the all rxns.

44
Calculations via Hesss Law

• 1. If a reaction is reversed, ?H is also
  reversed.
• N2(g) O2(g) ? 2NO(g) ?H 180 kJ
• 2NO(g) ? N2(g) O2(g) ?H ?180 kJ
• 2. If the coefficients of a reaction are
  multiplied by an integer, ?H is multiplied by
  that same integer.
• 6NO(g) ? 3N2(g) 3O2(g) ?H ?540 kJ

45
Using Enthalpy

• Calc. DH for
• S(s) 3/2 O2(g) --gt SO3(g)
• S(s) O2(g) --gt SO2(g) -320.5 kJ
• SO3(g) --gt SO2(g) 1/2 O2(g) 75.2 kJ
• _______________________________________
• S(s) 3/2 O2(g) --gt SO3(g) -395.7 kJ

46
energy
S solid
direct path
DH

1
O
2

-320.5 kJ
3/2 O
2
DH
SO
gas
2
-395.7 kJ
1/2 O
2
SO
gas
3
DH
-75.2 kJ
2
? DH along one path
? DH along another path
47

• Determine the heat of reaction for the
  decomposition of one mole of benzene to acetylene
• C6H6(l) 3C2H2(g)
• given the following thermo chemical equations
• 2C6H6(l) 15O2(g) 12CO2(g) 6H2O(g) DH
  -6271 kJ
• 2C2H2(g) 5O2(g) 4CO2(g) 2H2O(g) DH -2511
  kJ

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• See smart Hess Law

49
Using Standard Enthalpy Values

• H2O(g) C(graphite) --gt H2(g) CO(g)
• (product is called water gas)

50

• This equation is valid because DH is a STATE
  FUNCTION
• These depend only on the state of the system and
  not how it got there.

51
Change in Enthalpy

• Can be calculated from enthalpies of formation
  of reactants and products.
• ?Hrxn ?np?Hf?(products) ?
  ?nr?Hf?(reactants)
• ? ? H is an extensive property--kJ/mol
• For the reaction 2H2 (g) O2 (g) ---gt 2H2O(g)

52
Standard Enthalpy Values

• NIST (Natl Institute for Standards and
  Technology) gives values of
• DHof standard molar enthalpy of formation
• This is the enthalpy change when 1 mol of
  compound is formed from elements under standard
  conditions. DHof is always stated in terms of
  moles of product formed.

53
DHof, standard molar enthalpy of formation

• H2(g) 1/2 O2(g) --gt H2O(g)
• DHof -241.8 kJ/mol
• By definition, DHof 0 for elements in their
  standard states.

54
Using Standard Enthalpy Values

• Use DHfs to calculate enthalpy change DH for
• H2O(g) C(graphite) --gt H2(g) CO(g)
• From reference books we find
• DHf of H2O vapor - 242 kJ/mol
• H2(g) 1/2 O2(g) --gt H2O(g)
• DH f of CO - 111 kJ/mol
• C(s) 1/2 O2(g) --gt CO(g)

55
Using Standard Enthalpy Values

• Calculate the heat of combustion of methanol,
  i.e., DHorxn for
• CH3OH(g) 3/2 O2(g) --gtCO2(g) 2 H2O(g)
• DHof (-201.5kj) (-393.5kj)
  (-241.8kj)
• DHorxn ? DHof (prod) - ? DHof (react)

56
Using Standard Enthalpy Values

• CH3OH(g) 3/2 O2(g) --gt CO2(g) 2 H2O(g)
• DHorxn ? DHof (prod) - ? DHof (react)
• DHorxn (-393.5 kJ) 2 (-241.8 kJ)
• - 0 (-201.5 kJ)
• DHorxn -675.6 kJ per mol of methanol

57
h

• 2Al(s) Fe2O3(s) Al2O3(s)Fe(s)

Al2O3 DHof 646kj/mol
Fe2O3(s) DHof -826kj/mol
58

• 2KIO3 12HCl 2ICl KCL6H204Cl2

KIO3 -501
6H20 -286
HCl -92
ICl -24
KCL -435
59

• The standard enthalpy of combustion of ethane
  gas C2H4 is 1411.1 kj/mol a 298k
• Given the following enthalpy of formation
  calculate the enthalpy of formation of ethane gas
  
• CO2DHof -393.5kj/mol
• H20DHof -285.8kj/mol

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Pathway for the Combustion of Methane
61
Schematic diagram of the energy changes for the
combustion of methane.
62
Greenhouse Effect
-- a warming effect exerted by the earths
atmosphere due to thermal energy retained by
absorption of infrared radiation.
Greenhouse Gases CO2 H2O CH4 N2O