Chapter 7 Thermochemistry

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Chapter 7Thermochemistry
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The Nature of Energy

• The capacity to do work or to produce heat.

Burning sugar (sugar reacts with KClO3, a strong
oxidizing agent)
Burning peanuts supply sufficient energy to boil
a cup of water.
CHEMICAL ENERGY
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The Nature of Energy

• Kinetic Energy and Potential Energy
• Kinetic energy is the energy of motion
• motion - translation, rotation, vibration
• Potential energy is the energy an object
  possesses by virtue of its position, condition,
  or composition.
• gravity, electrostatic, chemical
• can be converted into kinetic energy

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The Nature of Energy

• Units of Energy
• Unit for energy is the joule, J
• We sometimes use the calorie instead of the
  joule
• 1 cal 4.184 J (exactly)
• A nutritional Calorie
• 1 Cal 1000 cal 1 kcal

James Joule 1818-1889
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Some Terminology in Thermochemistry

• Systems and Surroundings
• System part of the universe we are interested
  in.
• Surroundings the rest of the universe.

Universe System Surroundings
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Some Terminology in Thermochemistry

• Extensive Property a property that depends on
  the amount of substance in the system
• it is divisible
• e.g. V, n, mass
• Intensive Property a property that does NOT
  depend on the amount of substance in the system
• it is indivisible
• e.g. P, T

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Some Terminology in Thermochemistry
ENERGY is the capacity to do work or transfer
heat. WORK movement against a force w force x
distance The total energy of the system is called
its internal energy, U. Internal energy can be
increased by doing work on the system or
decreased by the system doing work on the
surroundings. Energy is stored in molecular
potential and kinetic energy.
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Some Terminology in Thermochemistry

• HEAT is the form of energy that flows between two
  samples because of their difference in
  temperature.
• during heat flow, Temperature may change or Phase
  may change (an isothermal process).
• The internal energy of the system can be changed
  by transferring heat
• DU q
• q positive - heat enters system
• q negative - heat leaves system
• Heat transfer changes molecular potential and
  kinetic energy.

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Conservation of Energy

• In interactions between a system and its
  surroundings the total energy remains constant
  energy is neither created nor destroyed.

qsystem -qsurroundings
qsystem qsurroundings 0
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Heat Capacity

• The quantity of heat required to change the
  temperature of a system by one degree is called
  the heat capacity of the system.
• If the system is one gram of material this is the
  specific heat C (J/gC)
• Heat change is given by

DT is POSITIVE heat is gained by the system, q
is positive DT is NEGATIVE heat lost by the
system, q is negative
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Specific Heat Capacity
Substance Spec. Heat (J/gK) H2O 4.184 Aluminiu
m 0.902 Glass 0.84
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Using the Heat Capacity
If 25.0 g of Al cools from 310 oC to 37 oC, how
many joules of heat energy are lost by Al?
heat gained/lost q (specific heat C)(mass)(DT)
q (0.902 J/gK)(25.0 g)(37 - 310)K q - 6160 J
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Experimental Determination of Specific Heat
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Example
Determining Specific Heat from Experimental
Data. Use the data presented on the last slide to
calculate the specific heat of lead.
qlead -qwater
qwater mc?T (50.0 g)(4.184 J/g C)(28.8 -
22.0)C
qwater 1.4x103 J
qlead -1.4x103 J mc?T (150.0 g)(c)(28.8 -
100.0)C
clead 0.13 Jg-1C-1
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Heats of Reaction and Calorimetry

• Chemical energy.
• Contributes to the internal energy of a system.
• When a reaction occurs in the system,heat can be
  released or absorbed by the system.
• Heat of reaction, qrxn.
• The quantity of heat exchanged between a system
  and its surroundings when a chemical reaction
  occurs within the system, at constant
  temperature.

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Heats of Reaction

• If the reaction produces heat
• qrxn lt 0 Exothermic reactions.
• If the reaction consumes heat
• qrxn gt 0 Endothermic reactions.
• Calorimeter
• A device for measuring quantities
• of heat.

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Bomb Calorimeter
Calorimetry is the method used to measure heats
of reaction.
qrxn -qcal
qcal qbomb qwater qwires
Define the heat capacity of the calorimeter
qcal Smici?T C?T
i
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Example

• Using Bomb Calorimetry Data to Determine a Heat
  of Reaction.
• The combustion of 1.010 g sucrose, in a bomb
  calorimeter, causes the temperature to rise from
  24.92 to 28.33C. The heat capacity of the
  calorimeter assembly is 4.90 kJ/C.
• What is the heat of combustion of sucrose,
  expressed in kJ/mol C12H22O11
• Verify the claim of sugar producers that one
  teaspoon of sugar (about 4.8 g) contains only 19
  calories.

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Example
Calculate qcalorimeter
qcal C?T (4.90 kJ/C)(28.33-24.92)C
(4.90)(3.41) kJ 16.7 kJ
per 1.010 g
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Example
Calculate qrxn in the required units
-16.7 kJ
qrxn -qcal
-16.5 kJ/g
1.010 g
Calculate qrxn for one teaspoon
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Coffee Cup Calorimeter

• A simple calorimeter.
• Well insulated and therefore isolated.
• Measure temperature change.

qrxn -qcal
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Work
In addition to heat effects, chemical reactions
may also do work.

• work against the surroundings by the system when
  the system expands its volume PV work
• Gas formed pushes against the atmosphere.
• Volume changes.

• work force distance
• P force/area
• work pressure area distance
• work pressure volume
• wsystem -PDV

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The First Law of Thermodynamics

• Internal Energy, U.
• Total energy (potential and kinetic) in a system.

• Translational kinetic energy.
• Molecular rotation.
• Bond vibration.
• Intermolecular attractions.
• Chemical bonds.
• Electrons.

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The First Law of Thermodynamics

• A system contains only internal energy.
• A system does not contain heat or work.
• These quantities are important during a change in
  the system.
• Law of Conservation of Energy
• The energy of an isolated (cannot exchange q or w
  with surroundings) system is constant.
• Energy can be converted from one form to another
  but can neither be created nor destroyed.

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The First Law of Thermodynamics
Now if both heat and work can be transferred into
and out of the system
DU q w
Energy is conserved!
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The First Law of Thermodynamics
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More Terminology

• State of the system - a condition with determined
  by the values of its properties
• an example is the equilibrium state
• The state can be described by an equation of
  state
• PV nRT
• State Function (state property) a property
  whose value of the system that depends on the
  present state.
• D represents change DV, DP, DE
• They are independent of pathway It doesnt
  matter how this state was achieved.
• The internal energy (U) of a system is a state
  function

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Functions of State

• U is a function of state.
• Not easily measured.
• ?U has a unique value between two states
  (independent of path).
• Is easily measured.

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More Terminology
The internal energy (U) of a system is a state
function In practice we do not determine this
value rather we work with changes that occur with
heat (q) or work (w) q and w are NOT state
functions their values depend on the path
followed when the system undergoes
change! However, together they do constitute a
state function U q w
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Pressure Volume Work
w F ? d (m ? g)
? ?h
(m ? g)
? ?h
? A

A
P?V w -Pext?V
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(No Transcript)
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Work not a state function
If we do the same expansion in two steps
1st step
P 2.40 atm 1.80 atm and V 1.02L
1.36L
w1 -PDV -1.80(1.36-1.02) -0.61
2nd step
P 1.80 atm 1.20 atm and V 1.36L
2.04
w2 -PDV -1.2(2.04-1.36) -0.82
Two-steps w1 w2 -1.43L/atm -1.44 x 102 J
one-step w -1.43L/atm -1.24 x 102 J
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m g h
Pi
Pf
1
Vi
Vf
gas
Vi Vf P constant Pi Pf
V constant
2
1 2
Thus W depends on the path
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And the heat???
Same thing!!
Thus The amount of transferred heat depends on
the path