Download H = 890kJ - George Mason University

Chapter 5 THERMOCHEMISTRY
 The Nature of Energy
 The First Law of Thermodynamics
 Enthalpy
 Enthalpies of Reaction
 Calorimetry
 Hess’ Law
 Enthalpies of Formation
 Foods and Fuels
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THERMOCHEMISTRY
Thermochemistry the study of the energy
changes that take place during a reaction.
 Reactions generally proceed in whichever
direction produces products that have lower
energy than the reactants.
Heat and Energy
 Heat: energy transferred from hotter to colder
one.
 Kinetic energy: the energy of movement of
1
matter EK  mv 2 . Units: Joule = 1 kgm2/s2.
2
E.g. what is the kinetic energy of 50.0 kg person
running at a velocity of 20 m/s.
 Potential energy: stored energy. E.g. water at
the top of a mountain, a compressed spring, a
chemical bond.
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Energy Changes and Energy
Conservation
First law of Thermodynamics: Energy is neither
created nor destroyed but may be converted from
one form to another.
Energy forms:
 Thermal energy a form of kinetic energy;
energy transfer results in a temperature change.
 Chemical energy a form of potential energy.
Energy is stored in chemical bonds and released
when a compound reacts.
 During reaction, energy is usually transformed
from chemical to thermal energy.
 First law can be written as:
E=q+w
where q = heat involved in the process and
w = work done by or to the system.
 Work can be electrical or pressure –volume
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Internal Energy and The First Law of
Thermodynamics
Internal Energy, E, is the sum of the potential
and kinetic energy of a system.
System - that part of the universe upon which
we are focussing, e.g. reactions.
Surroundings - eveything else in the universe
which is not the system.
State function - property depending only upon
initial and final states and not upon path.
 Extent of transfer of energy is: E = Efinal  Einitial.
 System is the reference point and a negative
sign indicates that energy is flowing from the
system to the surrounding.
(1) exothermic (exo “out of”). Heat flows
from system to surroundings.
(2) endothermic (endo “into”). Heat flows
from surroundings to system.
C(gr) + O2(g)  CO2(g) + 393.5 kJ Exothermic
CO2(g) + 393.5 kJ  C(gr) + O2(g)
Endothermic
E.g. what is the change in energy when a system
provides 225 J if heat to surroundings while doing
185 J of work on the surrounding.
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Sign conventions
Sign
heat
+

work
+

When
heat transferred from surroundings
to system (temperature of system
often increases.
heat transferred from system to
surroundings (temperature of
surroundings often increases).
work is done on system
work is done by system
 Sign of E will depend upon the sign of q and
w.
heat
+
+


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work
+



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E
+
depends
depends

Internal Energy and The First Law of
Thermodynamics2
 The conditions of measurement must be
included when discussing the total internal
energy since it is related to
a) chemical identity of reactants and products
b) their temperature, pressure, and physical
state.
 Internal energy of a system is a state function.
State Function a property of the system which
depends only in the initial and final states and is
independent of the history of the system.
 Several energy functions to be discussed have
this property.
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Expansion Work
Work is defined as force acting over some
distance: w =  d x F (referenced to the
system).
 During reactions often there is an expansion of
gases against some pressure where pressure is
equal to the force per unit area:
F
P  or F  PxA .
A
 Work is obtained by substitution:
w =  d x F =  d x (PxA) or
w =  PV.
 The first law can be restated as E = q  PV.
 This equation indicates that the amount of heat
involved in a reaction will be reduced by the
amount of work being done for a given change in
the internal energy.
E.g. Calculate the work done when during a
reaction the gaseous products cause the volume
to change from 22.4 L to 44.8 L against a constant
pressure of 1.00 atm.
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Expansion Work2
 If work is performed at constant temperature,
then the anmount of work performed will depend
upon the change in the number of moles (n)
 Modifications of the ideal gas law (PV = nRT
where n = # mol and R = 8.3145 J/molK) lead to
an alternative way of determining work. PV =
nRT
 The presence of solids and liquids need not be
considered since the molar volume of either a
solid or liquid is about 1000x smaller than the
molar volume of a gas.
E.g. determine the work performed during the
combustion of methane at 1.00 atm and 298.15 K.:
CH4(g) + 2O2(g)  CO2(g) +2H2O(l)
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Energy and Enthalpy
 From the first law: q = E + PV.
 With no change in volume the equation
simplifies to qV = E.
 At constant pressure: qP = E + PV.
 There are times when both volume and pressure
can change; the heat involved in the reaction is
then a more complicated function of E.
Enthalpy: the heat output at constant pressure.
H = E + PV.
 In general, H = E + PV + VP.
 At constant pressure, a change in enthalpy is
given by: H = E + PV = qP.
 Normally, H and E are fairly close to each
other in magnitude. In the combustion of
propane (see book), E = 2043 kJ, H = 2041
kJ and w = PV = 2kJ.
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Enthalpies of Physical and Chemical
Change
Enthalpies of Physical Change:
Heat, kJ
Heat Absorbed While Heating
 Heating a substance
One Mole of Water
70
increases the
60
50
Heat of Vaporization
temperature; the
40
amount of heat
30 Heat of Fusion
20
absorbed is
10
0
proportional to the
150
200
250
300
350
400
Temperature, K
heat capacity of the
species being heated.
 Amount of energy absorbed during phase
change is proportional to the heat of phase
change.
 Sum the heats during each portion of the curve
to determine overall heat.
 It is also possible for a substance to be
converted from a solid directly to a gas; the heat
of this process is called the heat of sublimation
and is equal to the sum of the heats of fusion
and vaporization at the same temperature.
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Enthalpies of Chemical Change
 H is an extensive property – its value depends
upon the amount of reactants.
 H is attached to the chemical equation to
indicate the amount of heat involved in the
reaction.
E.g. the combustion of methane:
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l)
2CH4(g) + 4O2(g)  2CO2(g) + 4H2O(l)
H =  890kJ
H =  1780kJ
E.g.2 determine the amount of heat that would be
evolved when 150 g of methane is burnt.
 Reversing reaction changes the sign of the heat.
CO2(g) + 2H2O(l)  CH4(g) + 2O2(g)
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H = +890kJ.
Calorimetry and Heat Capacity
Calorimeter = a device used to measure the
heat involved during a physical or chemical
change.
Heat capacity = the amount of heat absorbed
by a substance to raise the temperature by a
given amount.
 Calorimeters a device that measures the change
in the heat content or internal energy.
a) Atmospheric pressure
b) Bomb calorimeter
 Heat transferal to a substance like a solid or a
liquid, causes a proportional change in
temperature increases
1. H2O absorbs 4.18 J for every gram and °C
2. Al absorbs 0.902 J for every gram and °C
3. The amount of heat absorbed is directly
proportional to amount of absorbing species:
q  C  n  T
 s  m  T
where s = specific heat capacity, C = molar heat
capacity and T = Tfinal  Tinitial.
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Calorimetry and Heat Capacity2
 Energy change from any source such as
reactions or phase change can be measured
with heat capacity..
E.g. How much heat is required to heat 500.0 g of
water from 20.0°C to 100.0°C.
 The enthalpy change in the system is the
negative of the heat of the calorimeter. E.g.
exothermic reactions gives off heat to the
calorimeter. H =  qcalorimeter.
E.g.2 When 2.00 g of ethanol was burned, all of
the energy of reaction was used to heat water in a
calorimeter. Determine H for the reaction if the
temperature of 200.0 g of water increased from
25.0°C to 89.0°C.
 Heat capacity of a whole calorimeter is used for
complicated calorimeters such as the bomb
calorimeter.
E.g. 800.0 J of heat caused the temperature of a
calorimeter was found to increase by 2.0 K. In
some other reaction, the temperature of the
calorimeter was found to increase by 5.0 K.
Calculate the heat of the reaction.
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Hess’s Law
Hess’s law when a reaction at constant temperature and pressure
can be written as the summation of a series of reactions, the
enthalpy change, H, of the reaction is equal to the summation of
the H’s of the individual reactions.
E.g. determine the heat of formation of NO2(g):
½ N2(g) + O2(g)  NO2 Hof = ?
 Forming NO2(g) from N2(g) can be thought of as 2 step process:
Formation of NO(g)
Oxidation of NO
Overall
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½ N2(g) + ½ O2(g)  NO(g) H° = +180 kJ
NO(g) + ½ O2(g)  NO2(g) H° = 56 kJ
½ N2(g) + O2(g)  NO2(g) Ho = +124 kJ
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Hess’s Law2
Missing steps in a sequence can be determined using Hess’s law.
E.g. determine the heat of formation of methanol from the heat of
combustion and the other given reactions. Heat of combustion is
CH3OH(g) + O2(g)  CO2(g) + H2O(l) H = 726.4 kJ.
Decomp. CH3OH
Form CO2
Form H2O
Overall
CH3OH(l)
C(gr) + O2(g)
2H2(g) + O2(g)
CH3OH(g) + 3/2 O2(g)
 C(gr) + 2H2(g) + ½ O2(g)
 CO2(g)
 2H2O(l)
 CO2(g) + 2H2O(l)
H° = ?
H° = 393.51 kJ
H° = 571.66 kJ
H°= 726.4 kJ
Solution: Since the reaction sequence adds to give the overall reaction
we can add the heats.
726.4 = x + (393.51) + (571.66)
x = 238.77 kJ;
Hof = x = 238.77 kJ.
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Hess’s Law3
 H° of a reaction can be obtained from Hof of all reactants and
products. (See last example).
E.g. Determine the heat of combustion of ethanol, CH3CH2OH, from
heats of formation in the book.
Solution: CH3CH2OH + 3O2  2CO2(g) + 3H2O(l)
Hocomb = ?
Hocomb  3Hof ,H O  2Hof ,CO  Hof ,CH CH OH
2
2
3
2
 (3  ( 285 .63 )  2  ( 393 .51)  ( 1368 )) kJ
 276 .51 kJ
For any general reaction such as: aA + bB  cC + dD,
H  c  Hof ,C  d  Hof ,D  a  Hof ,A  b  Hof ,B
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Standard Heats of Formation
 Standard state the pure form of a substance at 1 atm usually at 25°C.
 Standard reaction enthalpies, H°, difference in enthalpy between
products and reactants of a reaction each in their standard states.
 Standard heat (enthalpy) of formation Hof the standard reaction
enthalpy per mol for the synthesis of a compound from its elements.
 Since reaction enthalpy depends upon conditions of experiment, it is usually
reported at some reference condition, Ho
 Most tables present enthalpy data in its standard state and as the heat of
formation.
E.g. Hof (HCl) is 92.3 kJ and the reaction is:
½H2(g) + ½Cl2(g)  HCl(g) Hof = 92.3 kJ
 H of pure elements in their most stable form under standard conditions is
defined as zero. E.g. Na(g), Na(s); C(g), C(gr), C(d).
 Hof of elements in another form often given.
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Na(s)  Na(g) H° = 107.8 kJ/mol. Also called the enthalpy of sublimation.
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