Chemistry 100 Chapter 5
Energy Relationships in
Chemistry
Thermochemistry
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Thermodynamics – the study of energy
and its transformations.
Thermochemical changes – energy
changes associated with chemical
reactions.
System  that specific part of the
universe of interest to us.
Surroundings  the part of the universe
not contained in the system.
3 types of Systems
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open system  exchanges mass and
energy
closed system  exchanges energy
but no mass
isolated system  no exchange of
either mass or energy
Three Types of Systems
insulation
cork
Open system
Closed System
Isolated System
Different Types of Energy
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Energy – the ability to do work.
Thermal energy – associated with the
random motions of atoms and
molecules
Heat energy – transfer of thermal
energy between two objects at
different temperature.
Energy (cont’d)
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Chemical energy – energy stored
within the structural units of chemical
substance.
Potential energy – the ability of an
object to do work because of its
position in a field of force.
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Kinetic Energy – the work that can be
performed by a moving object.
The unit of energy
1 Joule (J)
=1 kg m2/s2
An older unit of energy
1 calorie (cal)
= 4.184 J exactly
The Law of Conservation of Energy
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The law of conservation of energy
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Energy is neither created nor destroyed
in ordinary chemical and physical
processes
Converted from one type into another.
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This is also stated in terms of the first
law of thermodynamics.
E  E f  E i
E = internal energy
change of the system
Ef and Ei  the energy
of the final and initial
states, respectively
First Law of Thermodynamics
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Chemical reactions either absorb or
release energy.
Two terms
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Exothermic reaction  heat is released
to the surroundings.
Endothermic reaction heat is supplied
to the system by the surroundings.
Exothermic
Heat
System
surroundings
An Exothermic Process
Endothermic
Heat
System
surroundings
An Endothermic Process
The First Law Restated
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chemical systems – examine the
conversion of heat energy into work.
E  E f  E i  q  w
Signs for Heat and Work
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Work done by system on surroundings
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Work done by surroundings on system
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w ‘+’
q < 0, heat flows to surroundings
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w ‘-’
Exothermic ‘-’
q > 0, heat flows to system
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Endothermic ‘+’
Pressure-volume Work
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Pressure – volume work
w = -Pop V = -Pop (Vf -Vi)
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This is the type of work done by the
pistons in our automobile engines!
The greater the magnitude of Pop, the
gas has to "work harder" to obtain the
same volume change.
Pressure-Volume Work
State and Path Functions
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E, H, V are examples of state
functions.
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State functions – numerical value doesn’t
depend on how the process is carried out.
Work (w) and q (heat) are path functions
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The amount of work done or heat released
depends on how the system changes states.
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Examine a chemical reaction.
C (s) + O2 (g)  CO2 (g)
E = E[CO2 (g)] – E[C(s)] – E[O2(g)]
This reaction has a negative enthalpy
change (H = -393.5 kJ).
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From the first law
surrE + sysE = 0
surrE = -sysE
The energy "lost" from the system is
"gained" in the surroundings.
Enthalpies of Formation – Standard
Reaction Enthalpies
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The enthalpy change for the reaction
rH = H(products) - H(reactants)
We cannot measure the absolute
values of the enthalpies!!
How do we ‘measure’ enthalpies (or
heat contents) of chemical species?
The Formation Reaction
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A "chemical thermodynamic reference
point."
For CO and CO2
C (s) + O2 (g)  CO2 (g)
C (s) + ½ O2 (g)  CO (g)
The "formation" of CO and CO2 from its
constituent elements in their standard
states under standard conditions.
The Formation Reaction
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The formation reaction
For the formation of 1.00 mole of
Na2SO3(s)
2 Na(s) + S(s) + 3/2 O2 (g)  Na2SO3 (s)
The ‘formation enthalpy of Na2SO3(s)’,
symbolised fH[Na2SO3 (s)]
Standard Conditions for Thermodynamic
Reactions
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The degree sign, either  or ,
indicates standard conditions
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P = 1.00 atm
[aqueous species] = 1.00 mol/L
T = temperature of interest (note 25C or
298 K is used in the tables in your text).
The Significance of the Formation
Enthalpy
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fH° is a measurable quantity!
Compare CO (g) with CO2 (g)
C (s) + 1/2 O2 (g)  CO (g)
fH° [CO(g)] = -110.5 kJ/mole
C (s) + O2 (g)  CO2 (g)
fH° [CO2(g)] = - 393.5 kJ/mole
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The formation enthalpy for CO2(g) is
larger than the formation enthalpy of CO
(g).
Reactions Enthalpies
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Formation enthalpies –
thermodynamic reference point,
Formation of the elements from
themselves is a null reaction – fH
(elements) = 0 kJ / mole.
The Combustion of Propane
The General Equation
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Calculate enthalpy changes from the
formation enthalpies as follows.
r H    n p f H  products    nr f H  reactants 
Reverse a reaction, the sign of the enthalpy
change for the reaction is reversed.
Multiply a reaction by an integer, the enthalpy
change is multiplied by the same integer.
The Measurement of Energy Changes –
Calorimetry
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Calorimetry – the measurement of heat
and energy changes in chemical and
physical processes.
Heat capacity (C) – the amount of heat
(energy) needed to raise the temperature
of a given mass of substance by 1°C.
Specific heat capacity (s) – the amount of
heat energy (in Joules, J) required to
raise 1 g of a substance by 1°C (units =
J/g °C).
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General expression for heat capacity
C =ms
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m is the mass of the substance (in
grams).
Molar heat capacity
Cm = M s
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M – molar mass of the substance
s – its specific heat capacity.
The Calorimeter
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A calorimeter – a device which
contains water and/or another
substance with a known capacity for
absorbing energy (heat).
Calorimeters are adiabatic systems.
All energy changes take place
within the calorimeter.
Adiabatic System
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Adiabatic system – thermally insulated
from the rest of the universe
No heat exchange between system
and surroundings!
For an adiabatic system,
qtot = qrxn + qH2O + qcal = 0
-qrxn = qH2O + qcal
The Constant Volume (Bomb)
Calorimeter
E = qv
The Constant Pressure Calorimeter
H = qp
Relating the Enthalpy to the
Internal Energy
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The enthalpy and the internal energy
both represent quantities of heat.
E = qv.
H = qp.
E and H are related as follows
H = E +Pop V
V = the volume change for the reaction.
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For reactions involving gases
V = ng /(RT Pop)
ng =  np (g) -  nr (g)
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For most reactions, ng is small.
The difference between the internal
energy change and the enthalpy
change is small.
Other important Enthalpy changes
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Many other important processes have
associated enthalpy changes.
The measurement of the heat changes
for these process can give us some
insight into the changes in
intermolecular forces that occur during
the transformation.
Heat of dilution and solution.
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solH = the heat absorbed or given off
when a quantity of solute is dissolved
in a solvent.
solH = H(sol’n) - H(component)
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H(component) = H (solid) + H(solvent)
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For the process,
HCl (aq, 6 M)  HCl (aq, 1 M).
A significant amount of heat is
released when the acid solution is
diluted.
This is the enthalpy of dilution of the
acid.
dilH = H(sol’n 2) – H(sol’n ,1)
Lattice Enthalpies
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Look at the following process.
NaCl (s)  Na+ (g) + Cl- (g)
H = latH = 788 kJ/mole  the lattice
enthalpy
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A very endothermic reaction!
Due to the strength of the ionic bond!
Latent Heats
Latent heats are the enthalpy changes
associated with phase transitions.
H2O (l)  H2O (g)
rH = vapH  the enthalpy of
vapourization.
H2O (s)  H2O (l)
rH = fusH  the enthalpy of fusion.
H2O (s)  H2O (g)
rH = subH  the enthalpy of sublimation.
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Foods and Fuels
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Most of the chemical reactions that
produce heat are combustion
reactions.
Note – all combustion reactions are
exothermic.
Fuel values are generally reported as
positive quantities.
Obtaining fuel values – calorimetry.
Calories, Food Calories, and
Kilojoules
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When we read our cereal boxes, we may
see the following
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1 bowl cereal = 30 g cereal = 132 Cal (490
kJ).
Isn’t 1 calorie = 4.184 J (not 4.184 kJ)?
The fuel values of foods are reported as
food calories (Cal).
1.00 food calorie (Cal) = 1000 thermal
calories (cal) = 4184 J = 4.184 kJ.
Combustion of Carbohydrates and
Fats
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Most of the energy our body needs
comes form the combustion of sugar and
fats.
For the glucose (blood sugar) combustion
C6H12O6 (s) + 6 O2 (g)  6 CO2 (g) + 6 H2O (l)
rH = -2816 kJ
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This energy is supplied quickly to the
body!
Average fuel value of carbohydrates = 17
kJ/g.
Fats
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The combustion (metabolism) of fats
also produces CO2 and H2O.
The combustion of tristearin
C57H110O6 (s) + 163/2 O2 (g)  57 CO2 (g)
+ 55 H2O (l).
rH = -37.8 x 104 kJ
Fuel Value of Fats
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Fats are the body’s ‘energy
stockpiles!’
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Insoluble in water.
Average fuel value = 38 kJ/g –
about twice that of the
carbohydrates.
Caloric Contents
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For proteins – average fuel value = 17
kJ/g, about the same value as for the
carbohydrates.
The relative amounts of proteins, fats,
and carbohydrates in foods determines
the caloric content.
Fossil Fuels
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Coal, petroleum, and natural gas are
known as fossil fuels. They are
collectively the major source of energy for
commercial and personal consumption.
Fossil fuels are mixtures of many
different kinds of organic compounds.
The fuel values of fossil fuels is directly
related to the amount of carbon and
hydrogen in the fuel.
Hydrogen As a Fuel
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Hydrogen has a huge fuel value (142
kJ/g).
The combustion product is innocuous –
water.
Obviously, there are problems!
Two major difficulties with H2 as a fuel
source.
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Where do we get the hydrogen?
How do we store the hydrogen?