Chapter 5: Thermochemistry

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Chapter 5 Thermochemistry
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Thermochemistry

• In most chemical reactions, energy is absorbed or
  released.
• Thermochemistry is the correlation of chemical
  processes and energy changes.

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Energy

• The capacity to do work or to produce heat.

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Law of Conservation of Energy

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

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Two Types of Energy

• Kinetic energy - The energy of a particle by
  virtue of its motion

EK ½ mv2

• Potential energy The energy of position
  relative to other objects.

PE mgh
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System and Surroundings

• The system is a well-defined part of the universe
  singled out for study.
• The surroundings is the remainder of the
  universe.
• In a closed system energy, but not matter,
  can be exchanged with the surroundings.

Closed System
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First Law of Thermodynamics

• There are essentially two ways to change the
  energy of a system - heat (q) and work (w).
• Our task is to understand how energy exchanges
  can occur between system and its surroundings

Delta E change of internal energy
Energy is conserved!
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Heat and work

• Heat (q) is a form of energy transfer
• Units 1 calorie (cal) 4.184 J
• Energy used to cause the temperature of an
  object to increase
• Example combustion releases the energy stored in
  molecules in the form of heat
• Work (w) is exertion of a force over a distance.
• We will only consider expansion work
• w - PDV where P ?
  pressure
• V
  ? volume

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Heat

• q gt 0 Heat is transferred from
  surroundings to system. Process is
  endothermic.
• q lt 0 Heat is transferred from system to
  surroundings. Process is exothermic. Å

Surroundings
System
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Work

• work force ? distance
• since pressure force / area,
• work pressure ? volume
• wsystem ?P?V

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Potential Energy due to gravity
Potential energy converted to kinetic energy
As ball strikes ground, kinetic energy used to
do work in squashing the ball the rest is given
off as heat
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Internal Energy

• Internal energy (E) - The combined kinetic and
  potential energies of all particles in a system.
• The internal energy of a system is usually not
  known. But energy changes can be measured.
• DE Efinal - Einitial

?E mg(hf - hi)
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Å
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Relating ?E to heat and work
?E q w
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Work

• w gt 0 Surroundings does work on system.
• w lt 0 System does work on surroundings. Å
• We will only consider expansion work
• w - PDV where P ?
  pressure
• V
  ? volume

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State Functions

• State function - A property of a system that
  depends only on its present state.
• The change in a state function does not depend on
  how the process is carried out.

?E is a State Function
The water could have reached 50C from either
direction
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Path Functions

• Work (w) and heat (q) are not state functions
  because they depend on how the process is carried
  out

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First Law Example
What is ?E if 100 kJ of heat is added to a rigid
container of gas with a pressure of 10 psi?
?E q w
q - P?V
Since the volume doesnt change, ?V 0
?E q
100 kJ
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First Law Example

• What is DE if an insulated system at 1.0 atm
  expands by a volume of 1.0 L?

V2 - V1 1.0 L
?E q w q - P?V
-P?V
Since no heat is transferred, q 0.
-1.0 L?atm
? 1.0 L
?E -1.0 atm
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Enthalpy

• Enthalpy H E PV
• ?E ?H ? P?V
• ?H ?E P?V
• At constant pressure,
• qP ?E P?V,
• where qP ?H at constant pressure
• ??H energy flow as heat (at constant
  pressure)

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The Enthalpy

• Enthalpy (H) is a state function defined as
• H E PV
• At constant pressure, the change in enthalpy is
• DH H2 - H1 E2 PV2 - (E1 PV1)
• DH DE PDV
• From the First Law DE q - PDV
• DH q (constant P)

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Chemical Energy

• The energy stored within substances and given off
  when they take part in chemical reactions.
• Can be converted to different forms of energy
• Burning fuel ? Heat ? electricity ? light
• Energy can neither be created or destroyed!

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Energy and Enthalpy

• Energy and enthalpy are numerically similar and
  have energy units. (If DV 0, DH DE.)
• Each is equal to q depending on how a process is
  carried out
• Constant volume qV DE
• Constant pressure qP DH
• Since most reactions are carried out at constant
  pressure, DH is the more useful quantity.

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Enthalpies of Reaction

• The enthalpy change for a chemical reaction is
  defined as
• DH Hproducts - Hreactants
• A thermochemical equation includes the DH value
  for stoichiometric quantities of reactants and
  products.
• CH4(g) 2 O2(g) CO2(g) 2 H2O(g)
  DH -802 kJ

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Examples

• What is the sign of ?H for
• SO3-2(aq) OCl?(aq) ? SO4-2(aq) Cl?(aq)
  ?H 0

lt
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Examples

• What is the sign of ?H for
• NH4NO3(s) ? NH4(aq) NO3?(aq) ?H 0

gt
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Enthalpy Diagrams

• An enthalpy diagram shows H for the initial and
  final states of a process.

• At constant pressure, q DH -802 kJ.

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Enthalpies of Reaction

• The enthalpy change depends on the states of
  reactants and products.
• CH4(g) 2 O2(g) CO2(g) 2 H2O(l) DH
  -890 kJ

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Enthalpies of Reaction

• DH for a reaction is opposite in sign to DH for
  the reverse reaction.

• CO2(g) 2 H2O(g) CH4(g) 2 O2(g) DH
  802 kJ

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Enthalpies of Reaction

• H is an extensive property, so DH depends on the
  amounts of reactants and products.
• What is DH for the combustion of 11.0 g of CH4 in
  excess oxygen?
• CH4(g) 2 O2(g) CO2(g) 2 H2O(g) DH
  -802 kJ

-550 kJ
11.0 g CH4
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Enthalpies of Reaction

• DH is a conversion factor between heat
  transferred and moles of substance.
• What mass of butane must react in order to
  produce 100 kJ of heat?
• 2 C4H10(g) 13 O2(g) 8 CO2(g) 10 H2O(g)
  DH -5317 kJ

2.19 g C4H10
100 kJ
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Heat and Temperature Change

• How is heat transferred related to the change in
  temperature of a system with mass m?
• q specific heat ? m ? ?T
• The specific heat of a substance is the amount of
  heat required to raise the temperature of 1 gram
  by 1 K.
• Hg 0.14 J/g?K
• Al 0.90 J/g?K
• H2O 4.18 J/g?K

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Heat and Temperature Change

• How much heat is required to raise the
  temp-erature of 12.5 g H2O from 20.0C to 30.0C?

q specific heat ? m ? ?T
?12.5 g
?10.0 K
522 J

• What is the specific heat of iron if 540 J of
  heat increases the temperature of 48.0 g by 25C?

0.45 J/g?K
specific heat
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Calorimetry

• A calorimeter measures the amount of heat
  transferred during a reaction.

A simple constant- Pressure calorimeter
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Calorimeter Constant

• Some of the heat from your reaction will be
  absorbed by the calorimeter
• The heat lost to the calorimeter must be
  accounted for
• To determine Ccal, use the equation
• (mh)(sp.ht)(?Th) -(mc)(sp.ht.)(? Tc) Ccal ?
  Tc
• Solve for Ccal

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Bomb CalorimetryConstant Volume Calorimetry

• Reaction carried out under constant volume.
• Use a bomb calorimeter.
• Usually study combustion.

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Hesss Law

• Consider a reaction carried out in two steps.
  What is ?H?
• C(s) ½ O2(g) CO(g) DH1 -110 kJ
• CO(g) ½ O2(g) CO2(g) DH2 -283 kJ
• ____________________ __________
• C(s) O2(g) CO2(g) DH
• For a reaction carried out in a series of steps,
  DH is the sum of enthalpy changes for the
  individual steps.

-393 kJ
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Hesss Law

• Hesss law is based on the concept of enthalpy as
  a state function.

Å
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?H1 ?H2 ?H3
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Formation Reactions

• Thermochemical data is provided in terms of
  formation reactions.
• Formation reaction The formation of one mole of
  a substance from its elements in their standard
  states (stable form at 25 and 1 atm)
• What is the formation reaction for CaCO3(s)?

CaCO3(s)
Ca(s)
C(graphite)
3/2 O2(g)
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Standard Enthalpies of Formation, ???Hf, at 298 K
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Standard Enthalpy of Formation

• The standard enthalpy of formation, DHfº, of a
  substance is the ?H? of its formation reaction.
• Ca(s) C(graphite) 3/2 O2(g) CaCO3(s)
• 
  DHf(CaCO3,s) -1207 kJ
• Ca(s) ½ O2(g) CaO(s) DHf(CaO,s) -636 kJ
• C(graphite) O2(g) CO2(g)
• 
  DHf(CO2,g) -394 kJ

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Standard Enthalpy of Formation

• What is DHf(O2,g)?
• O2(g) O2(g)

?H? 0
?Hf is zero for an element in its standard form.

• What is ?Hf(C,diamond)?
• C(graphite) ? C(diamond)

?H? 1.9 kJ

• What is ?Hf(Cl,g)?
• ½ Cl2(g) ? Cl(g)

?H? 122 kJ See Appendix
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Standard Enthalpy of Reaction

• Hesss Law can be used to find DHº for
• CaCO3(s) CaO(s) CO2(g)

DHº DHf(CaO,s) DHf(CO2,g) - DHf(CaCO3,s)
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Standard Enthalpy of Reaction

• In general, DH is given by
• DH å n DHf(products) - å n DHf(reactants)
• The ?H? value assumes reactants and products are
  at the same temperature.
• Since DHf values are tabulated at 25ºC, we can
  only calculate DH at this temperature.

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Standard Enthalpy of Reaction

• Calculate DH at 25 for
• H(aq) OH?(aq) H2O(l)
• DH DHf(H2O,l) - DHf(H,aq) - DHf(OH?,aq)
• -285.83 kJ

- (-230.0 kJ)
-55.8 kJ
- 0
Experimental value
DH -58 kJ
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Standard Enthalpy of Reaction

• What is ?H? for combustion of C4H10? You try
• C4H10(g) 6½ O2(g) 5 H2O(l) 4 CO2(g)

?H?
5 ?Hf?(H2O,l)
4 ?Hf?(CO2,g)
- ?Hf?(C4H10,g)
- 6½ ?Hf?(O2,g)
4(-393.5)
5(-285.83)
- (-124.73)
- 6½(0)
kJ
-2878 kJ
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Standard Enthalpy of Reaction

• DH for combustion of C6H12O6 (glucose) is
  -2816 kJ. What is DHf for glucose? You try
• C6H12O6(s) 6 O2(g) 6 CO2(g) 6 H2O(l)

6 ?Hf?(CO2,g)
6 ?Hf?(H2O,l)
?H?
- ?Hf?(C6H12O6,s)
- 6 ?Hf?(O2,g)
- ?H?
6?Hf?(CO2,g)
6?Hf?(H2O,l)
?Hf?(C6H12O6,s)
6(-285.8)
- 6(0)
6(-393.5)
- (-2816)
kJ
-1260 kJ
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Combustion Reactions

• Combustion reactions are exothermic because of
  the strong bonds formed in CO2 and H2O.
• Fuels such as coal, petroleum, and
  natural gas have high percentages
  of carbon.

Sources of Energy consumed in US
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Energy Content of Foods

• Chemical energy in animals is derived from
  carbohydrates, fats, and proteins.
• Fuel value (kJ/g)
• Carbohydrates 17
• Fat 38
• Protein 17
• Fuel value is usually expressed in kcal or Cal
  per serving.
• 1 Cal 1 kcal 4.184 kJ

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Energy Content of Foods

• If a person uses about 420 kJ/mi when running,
  how many candy bars are required to run three
  miles?
• 1 Butterfinger
• 42 g carbohydrates
• 11 g fat
• 3 g protein
• 56 g

710 kJ
420 kJ
50 kJ
1180 kJ