CHM 110 CHAPTER 6 THERMOCHEMISTRY: Energy Flow and Chemical Change

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CHM 110CHAPTER 6THERMOCHEMISTRY Energy Flow
and Chemical Change

• Dr. Floyd Beckford
• Lyon College

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ENERGY

• Energy the capacity to do work or supply heat
• ENERGY Work heat
• Energy is classified into two types
• 1. Kinetic energy, EK energy of motion
• EK ½ mv2
• where m mass of object
• v velocity of object

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• 2. Potential energy, EP energy a system
• possesses by virtue of its composition or
• position
• The SI unit of energy is the Joule, J
• 1 J 1 kgm2/s2
• An older unit, the Calorie, is still used
• particularly in nutrition
• 1 cal 4.184 J

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ENERGY CHANGES

• Energy changes are dictated by the Law of
• Conservation of Energy - energy cannot be
• created or destroyed it can only be converted
  from
• one form into another
• HEAT the energy transferred from one object
• to another as a result of as temperature
  difference
• between them

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INTERNAL ENERGY

• The substances involved in the process under
• investigation makes up the system
• Everything in the systems environment
• constitutes its surroundings
• The internal energy, E, of the system represents
  
• all the energy contained within the system the
• sum of all types of energies present in the system

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• The First Law of Thermodynamics states the
• total energy of an isolated system is constant
• It is important to be able to measure energy
• changes which occur during reactions
• With respect to internal energy
• ?E Efinal Einitial
• Energy changes are always measured with
• respect to the system

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• Energy flow from the system ? ?E is negative
• - heat is lost and Efinal lt Einitial
• When Efinal gt Einitial heat flows into the
  system
• from the surroundings ?E is positive

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• The thermodynamic state of a system is defined
• by a set of conditions that completely specifies
  all
• the properties of the system
• These properties are called state functions
• The value of a state function depends ONLY on
• the state of the system and NOT on the way in
• which the system came to be in that state
• e.g. altitude

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• A change in a state function describes a
• difference between two states it is independent
• of the process or path way by which the change
• takes place
• It is the changes relating to state functions
• which usually interests us
• For any state function, X
• ?X Xfinal - Xinitial

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WORK

• All energy transfer to and from the system
• occur as heat or work
• In chemical systems, the most common type of
• work observed is pressure-volume or PV work
• - work done as a result of volume changes
• Volume changes may be expansion or
• contraction

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w -P ?V
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• During expansion, work is done BY the system
• ON the surroundings
• - sign of work is negative
• V2 gt V1 so ?V V2 V1 gt 0
• So w -P ?V lt 0
• This work can be due to the number of moles

of gas CO2(s) ? CO2(g)
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• The reverse process is called contraction work
  
• is done ON the system BY the surroundings
• V1 gt V2 so ?V V2 V1 lt 0
• And
• w -P ?V positive
• This change can be due to a decrease in the
• number of moles of gas
• When ?V 0 no PV work is done

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Calculate the work (in kilojoules) done
during the synthesis of ammonia in which the
volume contracts from 8.6 L to 4.3 L at a
constant external pressure of 44 atm. In which
direction does the work energy flow? What is the
sign of the energy change? 1 Latm 101 J
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ENERGY AND ENTHALPY

• The change in internal energy is ?E
• ?E q w q heat transferred
• System gains heat q is positive
• System loses heat q is negative
• ?E q - P ?V and so
• q ?E P ?V
• At constant volume qv ?E

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• If the reaction is carried out at constant
  pressure
• the heat transferred is called the heat of
  reaction
• or the enthalpy change, ?H of the reaction
• qp ?E P ?V ?H
• Enthalpy is a state function. So
• ?H Hproducts Hreactants
• For ordinary reactions ?E and ?H are roughly
• the same value

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THERMODYNAMIC STANDARD STATE

• Consider the following equation

• This equation is referred to as a thermochemical
• equation
• Note that the physical states of the substances
• MUST be specified when enthalpy changes are
• reported

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• C3H8(g) 5O2(g) ? 3CO2(g) 3H2O(g)
• ?H -2043 kJ
• C3H8(g) 5O2(g) ? 3CO2(g) 3H2O(l)
• ?H -2219 kJ
• It is also necessary to specify the conditions
• under which the reaction was run
• The thermodynamic standard state for a system
• must be defined

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• 1. Most stable form of a substance at 1 atm
• pressure and at a specified temperature, usually
• 25 C
• 2. 1 M concentration for all substances in
  solution
• Measurements done under standard conditions
• are are indicated with a superscript
• Reactants ? Products ?Hrxn
• ?Hrxn standard enthalpy change for reaction

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• H2O(s) ? H2O(g) ?Hsubl
• H2O(s) ? H2O(l) ?Hfusion
• H2O(l) ? H2O(g) ?Hvap
• For the above processes the enthalpy changes
• are referred to as enthalpy of sublimation,
  fusion,
• and vaporization respectively
• Note that during a phase change the temperature
• remains constant

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• 1. Ba(OH)28H2O(s) 2NH4Cl(s) ? BaCl2(aq)
• 2NH3(aq) 10H2O(l) ?H 80.3 kJ
• 2. 2H2(g) O2(g) ? 2H2O(g) ?H -484 kJ
• Reaction 1 absorbs heat from the surroundings
• the reaction is said to be endothermic
• When heat flows OUT of the system the reaction
• is exothermic reaction 2

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• Points to note about ?H
• 1. The value of ?H is for the thermochemical
• equation
• - this is because enthalpy is an extensive
• property
• 2. ?H values refer to the reaction proceeding in
• the direction as written
• - the sign of ?H must be changed for the
• reverse reaction

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How much heat (in kilojoules) is evolved or
absorbed in the burning 15.5 g of propane
as given in the following equation C3H8(g)
5O2(g) ? 3CO2(g) 3H2O(l) ?H -2219 kJ
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CALORIMETRY

• The amount of heat transferred during a
• reaction can be measured by calorimetry
• Based on observing the temperature change
• when a system absorbs or releases heat
• The reaction is carried out in a calorimeter
• The calorimeter measures ?H or ?E depending
• on the setup

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• Various heat transfer processes occur during a
• calorimetry experiment
• The heat released warms both the surrounding
• water as well as heat the calorimeter
• The amount of heat absorb by the calorimeter
• depends on its heat capacity
• - the amount of heat required to raise the
• temperature of an object by a certain amount

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• A more useful quantity is the specific heat
  the
• amount of heat necessary to raise the temperature
• of exactly 1 g of a substance by exactly 1 K
• In general
• Heat transferred heat gained by water heat
• gained by calorimeter
• For any substance
• q mc?T

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Assuming that Coca Cola has the same
specific heat as water (4.18 J/gC), calculate
the amount of heat (in kilojoules) transferred
when one can (about 350 g) is cooled from 25 C
to 3 C. What is the specific heat of lead if it
takes 96 J to raise the temperature of a 75 g
body by 10.0 C?
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HESSS LAW

• Hesss Law the enthalpy change for a reaction
• is the same whether it occurs by one step or by
• any series of steps
• Another way to state this the overall enthalpy
• change for a reaction is equal to the sum of the
• enthalpy changes for the individual steps in the
• reaction

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• ?Hrxn ?H1 ?H2 ?H3
• 1, 2, 3, refers to balanced thermochemical
• equations
• Consider the following equation
• C(graphite) ½ O2(g) ? CO(g) ?Hrxn ?
• The following reactions are known
• C(graphite) O2(g) ? CO2(g) ?Hrxn -393.5
• CO(g) ½ O2(g) ? CO2(g) ?Hrxn -283.0

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• C(graphite) O2(g) ? CO2(g) ?Hrxn -393.5
  (1)
• CO2(g) ? CO(g) ½ O2(g) ?Hrxn 283.0 (2)
• _________________________________________
• C(graphite) ½ O2(g) ? CO(g) ?Hrxn ?
• From Hesss Law
• ?Hrxn ?Hrxn (1) ?Hrxn (2)
• - 393.5 283.0
• - 110.5 kJ/mol

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Determine the heat of formation of liquid
hydrogen peroxide at 25 C from the following
thermochemical data. H2(g) ½ O2(g) ? H2O(g)
?H -241.82 kJ/mol 2H(g) O(g) ? H2O(g)
?H -926.92 kJ/mol 2H(g) 2O(g) ? H2O2(g) ?H
-1070.60 kJ/mol 2O(g) ? O2(g) ?H
-498.34 kJ/mol H2O2(l) ? H2O2(g) ?H
51.46 kJ/mol
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HEAT OF FORMATION

• Consider again the final reaction we just saw
• The 110.5 kJ is the amount of energy liberated
• when CO is formed from carbon and oxygen
• This is the standard heat of formation, ?Hf,
  for
• CO the enthalpy change for the formation of 1
• mole of a substance in its standard state from
  its
• constituents in their standard states

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• In using this definition the ?Hf for any
  element
• in its standard state is zero
• ?H ?Hf(products) - ?Hf(reactants)
• For any reaction
• aA bB .. ? cC dD
• ?H c ?Hf (C) d ?Hf (D)
• a ?Hf (A) b ?Hf (B)

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e.g. SiH4(g) 2O2(g) ? SiO2(s) 2H2O(l) ?Hrxn
?Hf (SiO2) 2 ?Hf (H2O) ?Hf
(SiH4) 2 ?Hf (O2) But ?Hf (O2) 0 So
?Hrxn ?Hf (SiO2) 2 ?Hf (H2O)
?Hf (SiH4)
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Acetic acid (CH3COOH) is made by the reaction of
C2H5OH with oxygen C2H5OH(l) O2(g) ?
CH3COOH(l) H2O(l) Use the following data to
calculate ?H (in kJ) for the reaction ?Hf
(C2H5OH) -277.7 kJ/mol ?Hf (CH3COOH) -484.5
kJ/mol ?Hf (H2O) -284.8 kJ/mol
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Calculate ?Hf (in kilojoules per mol) for
benzene, C6H6, from the following data 2C6H6(l)
15O2(g) ? 12CO2(g) 6H2O(l) ?H -6534
kJ ?Hf (CO2) - 393.5 kJ/mol ?Hf (H2O)
285.8 kJ/mol
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BOND ENERGIES

• Bond energy, D the amount of energy needed
• to break one mole of bonds to form products
• X-Y ? X Y ?H D
• Hesss law can be applied to bond energies
• ?H D(bonds broken) D(bonds formed)
• e.g. CH4(g) 3Cl2(g) ? CHCl3(g) 3HCl(g)
• ?H (3DCl-Cl 3DC-H) (3DC-Cl 3DH-Cl)