Chapter 6
THERMOCHEMISTRY
• Energy: The capacity to do work or to
produce heat
• Law of Conservation of Energy: Energy can
be converted from one form to another but
can be neither created nor destroyed. The
energy of the universe is constant (Euniv =
constant)
The Two Types of Energy:
Potential Energy: Energy due to position or
composition - can be converted to work
Kinetic Energy: Energy due to the motion of
the object and depends on the mass of the
object m and its velocity , K.E =1/2m2
• Work: Defined as force acting over a distance.
• Temperature V. Heat: Temperature reflects
random motions of particles, therefore related to
kinetic energy of the system. Heat involves a
transfer of energy between two objects due to
temperature difference.
• State Function (Property): A property of the
system that depends only on its present state. A
state function does not depend on the system’s
past or future. It does not depend on how the
system arrived at the present state (independent
of pathway).
• System: The part of the universe on which we
wish to focus attention.
• Surroundings: Include everything else in the
universe
Universe = System + Surroundings
• Exothermic Reaction: A reaction results in the
evolution of heat. Energy flows out of the system.
A reaction gives off heat, the surrounding become
warmer.
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) + energy (heat)
The Combustion of Methane
Endothermic Reaction: Reactions that absorb
energy (heat) from the surroundings. Heat flow is
into a system, the surroundings become cooler.
N2(g) + O2(g) + energy (heat)  2NO(g)
• For exothermic process – energy gained by the
surroundings must be equal to the energy lost by
the system. For endothermic process the
situation is reversed.
• Thermodynamics: The study of energy and its
interconversions.
The Energy Diagram for the Reaction of
Nitrogen and Oxygen to Form Nitric Oxide
First Law of thermodynamics: The energy of
the universe is constant.
• Internal Energy (E): The internal energy (E) of
a system can be defined as the sum of the kinetic
and potential energies of all the particles in the
system. The internal energy of a system can be
changed by a flow of work, heat or both.
E = q + w
E = change in system’s internal energy
q = heat
w = work
• Sign of heat flow: ‘q’ is positive when heat
flows into the system from the surroundings.
When q is positive, the process is called
endothermic.
‘q’ is negative when heat flows out of the
system to the surroundings. When q is negative
the process is called exothermic.
• Pressure is defined as force per unit area,
P = F/A  F = P x A
Work = force x distance
Work = F x h (h = distance)
= P x A x h
Change in volume, V = final volume – initial volume
= A x h
work = P x A x h = P x V
For an expanding gas, V is a positive quantity because
the volume is increasing. V and w have opposite signs,
which leads the equation, (work flows into the
surrounding, w is negative)
W = -PV
Work = pressure x volume
The Volume of a Cylinder
Enthalpy and Calorimetry
• Enthalpy (H): Defined as
H = E + PV where, E = Internal energy,
P = Pressure, V = Volume
At constant pressure,
E = qp + w
E = qp - PV
qp = E + PV
where, qp = heat at constant pressure
H = E + (PV) = E + PV
Therefore, H = qp
At constant pressure, the change in enthalpy H
of the system is equal to the energy flow as heat,
i.e, flow of heat is a measure of the change in
enthalpy.
Enthalpy change,
H = Hproducts – Hreactants
exothermic reactions: H < 0 (negative)
endothermic reactions: H > 0 (positive)
Rules of Thermochemistry
1. The magnitude of H (heat flow) is directly
proportional to the amount of reactant or
product (extensive property).
2. H for a reaction is equal in magnitude but
opposite in sign to H for the reverse reaction.
3. If a reaction is multiplied or divided by a
number the H of the reaction is also multiplied
or divided by that number.
4. The values of H for a reaction is the same
whether it occurs directly or in a series of steps.
5. The H of a reaction depends on the physical
state of the reactants and products.
Example: When 1 mole of methane (CH4) is
burned at constant pressure, 890 kJ of energy is
released as heat. Calculate H for a process in
which a 5.8 g sample of methane is burned at
constant pressure.
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) H = -890 kJ
1molCH 4  890kJ
58
. gCH 4 

  320kJ
16.0 gCH 4 molCH 4
• If we have 1.00 g of O2 then H = ?
1molO2  890KJ
100
. gO2 

  13.9 kJ
32.0 gO2 2molO2
Calorimetry
• Calorimeter: The device used experimentally to
determine the heat associated with a chemical
reaction.
• Calorimetry: The science of measuring heat, is
based on observing temperature change when a
body absorbs or discharges energy as heat.
• Heat Capacity: The amount of energy required to
raise the temperature of a system by1oC (J/oC)
C = [heat absorbed(q)] / [increase in
temperature(T)]
q = c x T
A Coffee-Cup Calorimeter Made
of Two Styrofoam Cups
A Bomb Calorimeter
Specific Heat Capacity: The energy required to
raise the temperature of one gram of a substance
by one degree celsius (J/oC.g).
• Molar Heat Capacity: The energy required to
raise the temperature of one mole of a substance
by one degree celsius (J/oC.mole).
• Energy released by the reaction
= energy absorbed by the solution
= specific heat capacity x mass of solution
x increase in temperature
= s x m x T
Example: If 100 J of energy is added to 50 g of
Cu, initially at 25.0oC, what will be the final
temperature? (Specific heat capacity = 0.382
J/oC.g )
q  m s T
100 J  50 g  0.382 J / C. g   T
o
100 J
T 
 5.2C
50 g  0.382 J / C. g
o
o
o
 T  Tf  Ti
Tf  Ti   T  25C  5.2C  30.2C
o
o
o
• Example: How much energy does it take to
heat 120 g of water from 20oC to 45oC?
q = m x s x T
= 120 x 4.18 J / oC .g x (45oC - 20oC)
= 12540 J
= 12.540 kJ
• Example: A 55.0 g piece of metal was heated to 99.8oC
and dropped into a calorimeter which contains 225 mL
of water (density = 1.00g / mL) and is at 21.0oC. The
final temperature of the water and the metal is 23.1oC.
What is the specific heat of the metal?
Energy released by the metal (qmetal)
= Energy absorbed by water (qwater)
Qwater = m x s x T
= 225 g x 4.18 J/oC.g x (23.1oC -21.0oC)
= 1975 J
Qmetal = -1975 J = m x s x T
= 55.0 g x s x (23.1oC –99.8oC)
s = -1975 J / -4218.5 goC = 0.468 J/oC.g
Hess’s Law
• In going from a particular set of reactants to a
particular set of products, the change in enthalpy
is the same whether the reaction takes place in
one step or in a series of steps.
Example:
Reactants
N2(g) + 2O2(g)
Products
2NO2(g) H1 = 68 kJ
…Hess’s Law continued…
This reaction also can be carried out in two
distinct steps,
N2(g) + O2(g)
2NO(g)
H2 = 180 kJ
2NO(g) + O2(g)
2NO2(g) H3 = -112 kJ
Net reaction:
N2(g) + 2O2(g) 2NO2(g) H2+ H3 = 68 kJ
H1 = H2+ H3 = 68 kJ
The Principle of Hess’s Law
Characteristics of Enthalpy Changes
• If a reaction is reversed, the sign of H
is also reversed
Xe(g) + 2F2(g) XeF4(s)
H = -251 kJ
XeF4(s)
Xe(g) + 2F2(g) H = +251 kJ
…Characteristics of Enthalpy Changes continued…
• The magnitude of H is directly proportional to
the quantities of reactants and products in a
reaction. If the co-efficients in a balanced
reaction are multiplied by an integer, the value of
H is multiplied by the same integer.
H is an extensive property, depends on the
amount of substance reacting.
Xe(g) + 2F2(g) XeF4(s) H = -251 kJ
2Xe(g) + 4F2(g) 2XeF4(s) H = 2(-251 kJ)
= -502 kJ
Standard Enthalpies of Formation
• The standard enthalpy of formation (Hof)
of a compound is defined as the change in
enthalpy that accompanies the formation of
one mole of a compound from its elements
with all substances in their standard states.
Standard States
•
•
•
•
For a compound:
The standard state of a gaseous substance is a
pressure of exactly 1 atmosphere.
For a pure substance in a condensed state (liquid
or solid), the standard state is the pure liquid or
solid.
For a substance present in a solution, the standard
state is a concentration of exactly 1M.
For an element:
The standard state of an element is the form in
which the element exists under conditions of 1
atmosphere and 25oC.
Pathway for the Combustion of Methane
A Schematic Diagram of the Energy Changes for the
Reaction CH4(g) + 2O2(g)  CO2(g) + 2H2O(l)
Change in Enthalpy
• The enthalpy change for a given reaction can be
calculated by subtracting the enthalpies of
formation of the reactants from the enthalpies of
formation of the products.
Horeaction = np Hof (products) -  nr Hof (reactants)
• Example:
4NH3(g) + 5O2(g)
4NO(g) + 6H2O(l)
4(-46) 5(0)
4(+90) 6(-242)
Hrxn = 360 + (-1452) – (-184)
= -1092 + 184 = -908 kJ
Present Sources of Energy
Energy Sources Used in the United States
• Petroleum is a thick, dark liquid composed
mostly of compounds called hydrocarbons that
contain carbon and hydrogen. Example, Gasoline
(C5 – C10), Kerosene and Jet fuel (C10 – C18),
Diesel fuel, Heating oil, Lubricating oil (C15 –
C25), Asphalt (>C25 ).
• Natural gas consists mostly of methane (CH4),
but it also contains ethane (C2H6), propane
(C3H8) and butane (C4H10).
• Coal was formed from the remains of plants that
were buried and subject to high pressure and heat
over long periods of time.
Effects of Carbon dioxide on Climate
The earth receives a tremendous quantity of radiant
energy from the sun. Some of this energy is
absorbed by plants for photosynthesis and some by
the oceans to evaporate water but most of it is
absorbed by soil, rocks, and water, increasing the
temperature of the earth’s surface. Molecules in the
atmosphere, principally H2O and CO2, strongly
absorb infrared radiation and radiate it back toward
the earth so a net amount of thermal energy is
retained by the earth’s atmosphere, causing the
earth to be much warmer than it would be without
its atmosphere.
The Earth’s Atmosphere
Atmospheric CO2 Concentration
New Energy Sources
• As we search for the energy sources of the future,
we need to consider economic, climatic, and
supply factors.
• There are several potential energy sources: the
sun (solar), nuclear processes (fission and fusion),
biomass (plants), and synthetic fuels.
• Direct use of the sun’s radiant energy to heat our
homes and run our factories and transportation
systems seems a sensible long-term goal.
Coal Conversion
One alternative energy source involves using a
traditional fuel – coal – in new ways.
• Since transportation costs for solid coal are
high, more energy-efficient fuels are being
developed from coal.
• One possibility is to produce a gaseous fuel.
• To convert coal from a solid to a gas requires
reducing the size of the molecules; the coal
structure must be broken down in a process
called coal gasification.
Hydrogen as a Fuel
• The combustion reaction is
H2(g) + 1/2O2(g)  H2O(l) Ho = -286 kJ
The heat of combustion of per gram is
approximately 2.5 times that of natural gas.
• In addition, hydrogen has a real advantage over
fossil fuels in that only product of hydrogen
combustion is water; Fossil fuels also produce
carbon dioxide.
• However, even though it appears that hydrogen
is a very logical choice as a major fuel for the
future, there are three main problems: the cost of
production, storage, and transport.
…hydrogen as a Fuel continued…
• Although hydrogen is very abundant on earth,
virtually none of it exits as the free gas. The main
source of hydrogen gas is from the treatment of
natural gas with steam:
CH4(g) + H2O(g)  3H2(g) + CO(g)
This reaction is highly endothermic; treating
methane with steam is not an efficient way to
obtain hydrogen for fuel. It would be much more
economical to burn the methane directly.
A virtually inexhaustible supply of hydrogen exists
in the waters of the world’s oceans. However, the
reaction H2O(l)  H2(g) + 1/2O2(g), requires 286
kJ of energy per mole of liquid water, large-scale
production of hydrogen from water is not
economically feasible. However, several methods
for such production are currently being studied:
electrolysis of water, thermal decomposition of
water, thermochemical decomposition of water and
biological decomposition of water.
• A current research goal is to find a system for
which the required temperatures are low enough
that sunlight can be used as the energy source.
Summary
• Energy: Potential Energy and Kinetic Energy
• Law of Conservation of Energy: energy can’t be
created or destroyed
• Heat: A means for the transfer of energy
• Universe = System + Surroundings
• Exothermic Reaction and Endothermic Reaction
• First Law of Thermodynamics: Euniv = Constant
• Internal Energy, E = q + w
• Enthalpy H = E + PV
• Rules of Thermochemistry
…Summary continued…
• Calorimetry: Science of measuring heat
• Heat Capacity, specific Heat Capacity, and
Molar Heat Capacity:
• Energy changed = m x s x T
• Energy Released = Energy Absorbed
• Hess’s Law: One step or series of step H same
• Change in Enthalpy:
Horeaction = np Hof (products) -  nr Hof (reactants)
• Energy Sources