Download Thermochemistry

Chapter 6
Thermochemistry
John A. Schreifels
Chemistry 211
1
Overview
• Understanding Heats of Reaction
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Energy and its Units
Heat of Reaction
Enthalpy and Enthalpy Change
Thermochemical Equations
Stoichiometry and Heats of Reaction
Measuring Heats of Reaction
• Uses of Heats of Reaction
– Hess’s Law
– Standard Enthalpies of Formation
– Fuels – Foods, etc.
John A. Schreifels
Chemistry 211
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Energy and Its Units
• Thermochemistry the study of the energy changes
that take place during a reaction.
– Reactions generally proceed in whichever direction that will
produce products with lower energy than the reactants.
Heat and Energy
• Heat: energy transferred from hotter to colder one.
• Kinetic energy: the energy of movement of matter
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EK  mv 2. Units: Joule = 1 kg*m /s .
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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.
John A. Schreifels
Chemistry 211
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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
John A. Schreifels
Chemistry 211
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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 focusing, 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.
– exothermic (exo “out of”). Heat flows from system to surroundings.
– endothermic (endo “into”). Heat flows from surroundings to system.
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C(gr) + O2(g)  CO2(g) + 393.5 kJ
CO2(g) + 393.5 kJ  C(gr) + O2(g)
John A. Schreifels
Chemistry 211
Exothermic
Endothermic
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Sign conventions
heat
work
Sign
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.
John A. Schreifels
Chemistry 211
Heat
Work
E
+
+
+
+

Depends

+
Depends



E = q + w
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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
– chemical identity of reactants and products
– 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.
John A. Schreifels
Chemistry 211
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Expansion Work
Work = 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.
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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.
John A. Schreifels
Chemistry 211
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Expansion Work2
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If work is performed at constant temperature, then the amount 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)
John A. Schreifels
Chemistry 211
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Energy and Enthalpy
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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.
John A. Schreifels
Chemistry 211
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Enthalpies of Physical and
Chemical Change
Heat Absorbed While Heating
Heat, kJ
Enthalpies of Physical Change:
One Mole of Water
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• Heating a substance increases the
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temperature; the amount of heat absorbed 50
Heat of Vaporization
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is proportional to the heat capacity of the
30 Heat of Fusion
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species being heated.
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• Amount of energy absorbed during phase
0
change is proportional to the heat of phase 150 200 250 300 350
Temperature, K
change.
• Sum the heats in each portion of the curve to determine overall
heat.
• The heat for converting a solid directly to a gas is called the heat
of sublimation and is equal to the sum of the heats of fusion and
vaporization at the same temperature.
John A. Schreifels
Chemistry 211
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400
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)
H =  890kJ
• 2CH4(g) + 4O2(g) 2CO2(g) + 4H2O(l)
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)
H = +890kJ.
John A. Schreifels
Chemistry 211
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Calorimetry and Heat Capacity
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Calorimeter = a device that measures the change in the heat content
or internal energy.
– Atmospheric pressure
– Bomb calorimeter
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Heat capacity = the amount of heat absorbed by a substance to raise
the temperature by a given amount.
Heat transferal to a substance like a solid or a liquid, causes a change
in temperature that is proportional to the amount of heat involved.
– H2O absorbs 4.18 J for every gram and °C
– Al absorbs 0.902 J for every gram and °C
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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.
John A. Schreifels
Chemistry 211
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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 calorimeter. H = 
qcalorimeter.
E.g.2 When 2.00 g of ethanol was burned, all of the reaction energy
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.
John A. Schreifels
Chemistry 211
14
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):
Hof = ?
½ N2(g) + O2(g)  NO2
• Forming NO2(g) from N2(g) can be thought of as 2
step process:
Formation of NO(g)
Oxidation of NO
Overall
John A. Schreifels
Chemistry 211
½ N2(g) + ½ O2(g)  NO(g)
NO(g) + ½ O2(g)  NO2(g)
½ N2(g) + O2(g)  NO2(g)
H° = +180 kJ
H° = 56 kJ
Hof = +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 for methanol decomposition to its elements
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
CH3OH(l)
 C(gr) + 2H2(g) + ½ O2(g)
H1° = ?
Form CO2
C(gr) + O2(g)
 CO2(g)
H2° = 393.51 kJ
Form H2O
2H2(g) + O2(g)
 2H2O(l)
H3° = 571.66 kJ
 CO2(g) + 2H2O(l)
H°= 726.4 kJ
Overall
CH3OH(g) + 3/2 O2(g)
H° = H1° + H2° + H3° or H1° = 726.4 + 393.51 +571.66 = +238.77 kJ
John A. Schreifels
Chemistry 211
16
Standard Heats of Formation
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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 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, H°
Most tables present enthalpy data in its standard state and as the heat of
formation.
E.g. (HCl) is 92.3 kJ and the reaction is:
½H2(g) + ½Cl2(g)  HCl(g)
= 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).
of elements in another form often given.
Na(s)  Na(g) H° = 107.8 kJ/mol. Also called the enthalpy of
sublimation.
John A. Schreifels
Chemistry 211
17
Calculations with Heat of
Formation
• H° of a reaction can be obtained from of all reactants and
products.
E.g. Determine the heat of combustion of ethanol, CH3CH2OH,
from heats of formation in the book.
Solution: CH3CH2OH + 3O2  2CO2(g) + 3H2O(l)
Hc°= ?
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
John A. Schreifels
Chemistry 211
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