Transcript Ch04_Outline

Chapter 4
Thermochemistry

Thermodynamics
Dr.Imededdine Arbi Nehdi
Chemistry Department, Science
College, King Saud University
Thermochemistry
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• The study of energy and its transformations
is known as thermodynamics
• The relationships between chemical
reactions and energy changes is an aspect
of thermodynamics called thermochemistry
Thermochemistry
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Energy
• Energy is the ability to do work or
transfer heat.
– Energy used to cause an object that has
mass to move is called work.
– Energy used to cause the temperature of
an object to rise is called heat.
Thermochemistry
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Potential Energy
Potential energy is energy an object
possesses by virtue of its position or chemical
composition.
Fig. 4.1 A bicycle at the top of a hill (left) has a high potential
energy. As the bicycle proceeds down the hill (right), the potential
energy is converted into kinetic energy; the potential energy is lower
at the bottom than at the top.
Thermochemistry
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Kinetic Energy
Kinetic energy is energy an object possesses
by virtue of its motion.
1
Ek =  mv2
2
where m is mass, v is the speed,
and Ek is the kinetic energy.
Thermochemistry
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Units of Energy
• The SI unit of energy is the joule (J).
kg m2
1 J = 1 
s2
• An older, non-SI unit is still in
widespread use: the calorie (cal).
1 cal = 4.184 J
Thermochemistry
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Definitions:
System and Surroundings
• The system includes the
molecules we want to
study (here, the hydrogen
and oxygen molecules).
• The surroundings are
everything else (here, the
cylinder and piston).
If the hydrogen and oxygen react to form water, energy is liberated.
2 H2 (g) + O2(g)
2 H2O (l) + energy
The system has not lost or gained mass; it undergoes no exchange
of matter with its surrounding. However, it does exchange energy with its
surroundings in the form of work and heat.
Thermochemistry
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Definitions: Work
• Energy used to
move an object over
some distance is
work.
• w=Fd
where w is work, F
is the force, and d is
the distance over
which the force is
exerted.
Thermochemistry
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Heat
• Energy can also be
transferred as heat.
• Heat flows from
warmer objects to
cooler objects.
Thermochemistry
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We can now provide a more precise definition of energy:
Energy is the capacity to do work or transfer heat.
Thermochemistry
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Conversion of Energy
• Energy can be converted from one type to
another.
• For example, the cyclist above has potential
energy as she sits on top of the hill.
Thermochemistry
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Conversion of Energy
• As she coasts down the hill, her potential
energy is converted to kinetic energy.
• At the bottom, all the potential energy she had
at the top of the hill is now kinetic energy.
Thermochemistry
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First Law of Thermodynamics
• Energy is neither created nor destroyed.
• In other words, the total energy of the universe is
a constant; if the system loses energy, it must be
gained by the surroundings, and vice versa.
Thermochemistry
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Internal Energy
The internal energy of a system is the sum of all
kinetic and potential energies of all components
of the system; we call it E.
Thermochemistry
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Internal Energy
By definition, the change in internal energy, E,
is the final energy of the system minus the initial
energy of the system:
E = Efinal − Einitial
Thermochemistry
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Changes in Internal Energy
• If E > 0, Efinal > Einitial
– Therefore, the system
absorbed energy from
the surroundings.
– This energy change is
called endergonic.
Thermochemistry
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Changes in Internal Energy
• If E < 0, Efinal < Einitial
– Therefore, the system
released energy to the
surroundings.
– This energy change is
called exergonic.
Thermochemistry
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Changes in Internal Energy
• When energy is
exchanged between
the system and the
surroundings, it is
exchanged as either
heat (q) or work (w).
• That is, E = q + w.
Thermochemistry
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E, q, w, and Their Signs
Thermochemistry
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Exchange of Heat between
System and Surroundings
• When heat is absorbed by the system from
the surroundings, the process is endothermic.
Thermochemistry
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Exchange of Heat between
System and Surroundings
• When heat is absorbed by the system from
the surroundings, the process is endothermic.
• When heat is released by the system into the
surroundings, the process is exothermic.
Thermochemistry
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State Functions
Usually we have no way of knowing the
internal energy of a system; finding that value
is simply too complex problem.
Thermochemistry
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State Functions
• However, we do know that the internal energy
of a system is independent of the path by
which the system achieved that state.
– In the system below, the water could have reached
room temperature from either direction.
Thermochemistry
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State Functions
• Therefore, internal energy is a state function.
• It depends only on the present state of the system, not on
the path by which the system arrived at that state.
• And so, E depends only on Einitial and Efinal.
• The value of state function does not depend on the
particular history of the sample, only on its present
condition.
Thermochemistry
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State Functions
• However, q and w are
not state functions.
• Whether the battery is
shorted out or is
discharged by running
the fan, its E is the
same.
– But q and w are different
in the two cases.
a) A battery shorted out by a wire loses energy
to the surroundings as heat; no work is performed.
b) A battery discharged through a motor loses energy as work
(to make the fan turn|). Some heat will be released to the surroundings.
Thermochemistry
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Work
Usually in an open
container the only work
done is by a gas
pushing on the
surroundings (or by
the surroundings
pushing on the gas).
Thermochemistry
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Work
We can measure the work done by the gas if
the reaction is done in a vessel that has been
fitted with a piston.
w = -PV
Thermochemistry
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Enthalpy
• If a process takes place at constant
pressure (as the majority of processes we
study do) and the only work done is this
pressure-volume work, we can account for
heat flow during the process by measuring
the enthalpy of the system.
• Enthalpy is the internal energy plus the
product of pressure and volume:
H = E + PV
Enthalpy is a state function
Thermochemistry
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Enthalpy
• When the system changes at constant
pressure, the change in enthalpy, H, is
H = (E + PV)
• This can be written
H = E + PV
Thermochemistry
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Enthalpy
• Since E = q + w and w = -PV, we can
substitute these into the enthalpy expression:
H = E + PV
H = (q+w) − w
H = q
• So, at constant pressure, the change in
enthalpy is the heat gained or lost by the
system when the process occurs under
constant pressure.
H = Hfinal – Hinitial = qp
Thermochemistry
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Endothermicity and
Exothermicity
• A process is
endothermic when
H is positive.
Thermochemistry
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Endothermicity and
Exothermicity
• A process is
endothermic when
H is positive.
• A process is
exothermic when
H is negative.
Thermochemistry
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Enthalpy of Reaction
The change in
enthalpy, H, is the
enthalpy of the
products minus the
enthalpy of the
reactants:
H = Hproducts − Hreactants
- It is found experimentally that 890 kJ of heat is produced when
1 mol of CH4 is burned in a constant-pressure system.
-The combustion of 2 mol of CH4 with 4 mol of O2
-releases twice as much heat, 1780 kJ.
Thermochemistry
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Enthalpy of Reaction
This quantity, H, is called the enthalpy of
reaction, or the heat of reaction.
(a)
(b)
(c)
Fig. a)A candle is held near a balloon filled with H2 gas and O2 gas.
b) The H2 (g) ignites, reacting with O2 (g) to form H2O (g).
The resultant explosion produces the yellow ball of flame.
The system gives off heat to its surroundings.
c) The enthalpy diagram of this reaction.
Thermochemistry
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The Truth about Enthalpy
1. Enthalpy is an extensive property.
2. H for a reaction in the forward
direction is equal in size, but opposite
in sign, to H for the reverse reaction.
3. H for a reaction depends on the state
of the products and the state of the
reactants.
Thermochemistry
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Calorimetry
Calorimeter
Since we cannot know
the exact enthalpy of
the reactants and
products, we measure
H through
calorimetry, the
measurement of heat
flow by measuring the
temperature change it
produces.
Thermochemistry
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Heat Capacity and Specific Heat
The amount of energy required to raise the
temperature of a substance by 1 K (1C) is its
heat capacity.
Thermochemistry
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Heat Capacity and Specific Heat
We define specific heat capacity (or simply
specific heat) as the amount of energy
required to raise the temperature of 1 g of a
substance by 1 K.
Thermochemistry
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Heat Capacity and Specific Heat
The amount of heat (q) transferred (gained or loosed) by
a masse (m) of substance with a specific heat capacity
(Cs) and temperature change of T is:
q = Cs x m x T
So if we known (q) we can calculate specific heat:
Cs =
q
m  T
The molar heat capacity (Cm) of a substance is:
Cm = Cs x M
Thermochemistry
M: molar mass of the substance (g/mol)
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Constant Pressure Calorimetry
-By carrying out a reaction in
aqueous solution in a simple
calorimeter such as this one, one
can indirectly measure the heat
change for the system by measuring
the heat change for the water in the
calorimeter.
• -This calorimeter is used for the
reactions occurring in solution.
- Because the calorimeter is not
sealed, the reaction occurs under
atmospheric pressure.
Fig. Coffee-cup calorimeter,
in which reactions
occur at constant pressure.
Thermochemistry
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Constant Pressure Calorimetry
Because the specific heat for water is well
known (4.184 J/g-K), we can measure H
for the reaction with this equation:
qsoln = m  Cs  T
qsoln = - qrxn
qsoln : the heat gained (loosed) by the solution
qrxn : the heat gained (loosed) by the reaction
m : masse of solution
Cs : specific heat of solution (  4.184 J/g-K)
Thermochemistry
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Bomb Calorimetry
• Reactions can be carried out in a
sealed “bomb” such as this one.
• The heat absorbed (or released)
by the water is a very good
approximation of the enthalpy
change for the reaction.
• Combustion reactions are the
most conveniently studied using
bomb calorimeter.
• After the sample (organic
compound) has been placed in
the bomb, the bomb is sealed Fig. Cutaway view of a bomb
and pressurized with oxygen. calorimeter, in witch reactions Thermochemistry
occur at constant volume.
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Bomb Calorimetry
• Combustion reactions are the most
conveniently studied using bomb
calorimeter.
• After the sample (organic compound)
has been placed in the bomb, the bomb
is sealed and pressurized with oxygen.
• It is then placed in the calorimeter,
which is essentially an insulated
container, and covered with an
accurately measured quantity of water.
• When all the components within the
calorimeter have come to the same
temperature, the combustion reaction is
initiated by passing an electrical current
through a fine wire that is in contact
with the sample.
Fig. Cutaway view of a bomb
• When the wire gets sufficiently hot, calorimeter, in witch reactions Thermochemistry
occur at constant volume.
the sample ignites.
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Bomb Calorimetry
• Because the volume in the bomb
calorimeter is constant, what is
measured is really the change in
internal energy, E, not H.
• For most reactions, the difference is
very small.
• So the heat of the reaction (q rxn) is :
q rxn = - Ccal x T
• Ccal : Heat capacity of the calorimeter
• T:the temperature change of
calorimeter
Thermochemistry
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Hess’s Law
 H is well known for many reactions,
and it is inconvenient to measure H
for every reaction in which we are
interested.
• However, we can estimate H using
published H values and the
properties of enthalpy.
Thermochemistry
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Hess’s Law
Hess’s law states that
“if a reaction is carried
out in a series of
steps, H for the
overall reaction will be
equal to the sum of
the enthalpy changes
for the individual
steps.”
Thermochemistry
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Hess’s Law
Because H is a state
function, the total
enthalpy change
depends only on the
initial state of the
reactants and the final
state of the products.
Thermochemistry
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Enthalpies of Formation
An enthalpy of formation, Hf, is defined
as the enthalpy change for the reaction
in which a compound is made from its
constituent elements in their elemental
forms.
2C (graphite) + 3H2(g) + ½ O2 (g)
C2H5OH (l)
fH = -277.7 kJ
Thermochemistry
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Thermochemistry
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Standard Enthalpies of Formation
Standard enthalpies of formation, Hf°, are
measured under standard conditions (25 °C
and 1.00 atm pressure).
By definition, the standard enthalpy of formation of the most stable form of
any element is zero (Hf (O2g) = 0, …)
Thermochemistry
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Calculation of H of reaction using
enthalpies of formation
C3H8 (g) + 5 O2 (g)  3 CO2 (g) + 4 H2O (l)
• Imagine this as occurring
in three steps:
C3H8 (g)  3 C (graphite) + 4 H2 (g)
3 C (graphite) + 3 O2 (g)  3 CO2 (g)
4 H2 (g) + 2 O2 (g)  4 H2O (l)
Thermochemistry
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Calculation of H
C3H8 (g) + 5 O2 (g)  3 CO2 (g) + 4 H2O (l)
• Imagine this as occurring
in three steps:
C3H8 (g)  3 C (graphite) + 4 H2 (g)
3 C (graphite) + 3 O2 (g)  3 CO2 (g)
4 H2 (g) + 2 O2 (g)  4 H2O (l)
Thermochemistry
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Calculation of H
C3H8 (g) + 5 O2 (g)  3 CO2 (g) + 4 H2O (l)
• Imagine this as occurring
in three steps:
C3H8 (g)  3 C (graphite) + 4 H2 (g)
3 C (graphite) + 3 O2 (g)  3 CO2 (g)
4 H2 (g) + 2 O2 (g)  4 H2O (l)
Thermochemistry
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Calculation of H
C3H8 (g) + 5 O2 (g)  3 CO2 (g) + 4 H2O (l)
• The sum of these
equations is:
C3H8 (g)  3 C (graphite) + 4 H2 (g)
3 C (graphite) + 3 O2 (g)  3 CO2 (g)
4 H2 (g) + 2 O2 (g)  4 H2O (l)
C3H8 (g) + 5 O2 (g)  3 CO2 (g) + 4 H2O (l)
H1 = - Hf°(C3H8(g)), H2 = 3 Hf°(CO2(g)), H3 = 4 Hf°(H2O(l)),
Hrxn = H1 +H2 + H3 = 2220 kJ
Thermochemistry
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Calculation of H
We can use Hess’s law in this way:
H =  nHf°products –  mHf° reactants
where n and m are the stoichiometric
coefficients.
Thermochemistry
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Calculation of H
C3H8 (g) + 5 O2 (g)  3 CO2 (g) + 4 H2O (l)
H = [3(-393.5 kJ) + 4(-285.8 kJ)] – [1(-103.85 kJ) + 5(0 kJ)]
= [(-1180.5 kJ) + (-1143.2 kJ)] – [(-103.85 kJ) + (0 kJ)]
= (-2323.7 kJ) – (-103.85 kJ) = -2219.9 kJ
Thermochemistry
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