Download Chapter 5 Thermochemistry

Spring 2016
Chapter 5
Thermochemistry
Thermochemistry
Nature of Energy
 Thermochemistry is the study of the energy released or
absorbed during a chemical reaction
It is an an aspect of thermodynamics.
Thermodynamics is the study of energy and its
transformations.
Thermochemistry
1
Spring 2016
Energy
 The ability to do work or transfer heat.
Thermochemistry
Energy
 The ability to do work or transfer heat.
Thermochemistry
2
Spring 2016
States and forms of Energy
 Energy can be in one of two states:
potential energy or kinetic energy.
Thermochemistry
Potential Energy
Stored energy. Energy ready to go.
Gravitational (potential) energy:
Energy an object possesses by virtue of its
position.
Thermochemistry
3
Spring 2016
Kinetic Energy
Energy an object possesses by virtue of its
motion.
1
2
KE =  mv2
Thermochemistry
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). which is defined as the amount of heat
required to raise the temperature of one gram of water
from 14.5oC to 15.5oC.
1 cal = 4.184 J
Thermochemistry
4
Spring 2016
System and Surroundings
 The system is the small portion of the universe in
which we are interested in, such as the water in a
beaker as shown in the figures . The
surroundings are everything else.
The system and its surroundings are separated by a
boundary. Heat is transferred across the boundary
between a system and its surroundings
Thermochemistry
Work
 Energy used to move
an object over some
distance.
 w = F x d,
where w is work, F is
the force, and d is the
distance over which
the force is exerted.
Thermochemistry
5
Spring 2016
Example
Calculate the amount of work that has to be done
to lift a 5-kilo bag of groceries a distance of 1.5
meter from the floor to the top of the kitchen
counter.
Thermochemistry
Heat
 Heat: Energy used to cause the
temperature of an object to rise
 Heat flows from warmer objects to
cooler objects.
 Energy can also be transferred as
heat.
Thermochemistry
6
Spring 2016
Transfer of Energy
a) The potential energy of this ball is
increased when it is moved from the
ground to the top of the wall.
Thermochemistry
Transfer of Energy
a) The potential energy of this ball is
increased when it is moved from the
ground to the top of the wall.
b) As the ball falls, its potential energy is
converted to kinetic energy.
Thermochemistry
7
Spring 2016
Transfer of Energy
a) The potential energy of this ball is
increased when it is moved from the
ground to the top of the wall.
b) As the ball falls, its potential energy is
converted to kinetic energy.
c) When it hits the ground, its kinetic
energy falls to zero (since it is no
longer moving); some of the energy
does work on the ball, the rest is
dissipated as heat.
Thermochemistry
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
8
Spring 2016
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
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
Use Fig. 5.5
Thermochemistry
9
Spring 2016
Changes in Internal Energy
 If E > 0, Efinal > Einitial
Therefore, the system
absorbed energy from
the surroundings.
This energy change is
called endergonic.
Thermochemistry
Changes in Internal Energy
 If E > 0, Efinal > Einitial
Therefore, the system
absorbed energy from
the surroundings.
This energy change is
called endergonic.
Thermochemistry
10
Spring 2016
Changes in Internal Energy
 If E < 0, Efinal < Einitial
Therefore, the system
released energy to the
surroundings.
This energy change is
called exergonic.
Thermochemistry
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
11
Spring 2016
E, q, w, and Their Signs
Thermochemistry
Exchange of Heat between System
and Surroundings
 When heat is absorbed by the system from the
surroundings, the process is endothermic.
Thermochemistry
12
Spring 2016
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 to the
surroundings, the process is exothermic.
Thermochemistry
State Functions
 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.
The properties of the system depends only on its
present state and not on how the state was reached
Thermochemistry
13
Spring 2016
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.
Thermochemistry
State Functions
 However, q and w are
not state functions.
Thermochemistry
14
Spring 2016
pressure-volume 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
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
15
Spring 2016
Enthalpy
 Enthalpy is a measure of the total energy of a system.
 Enthalpy is the internal energy plus the product of
pressure and volume:
H = E + PV
 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.
Thermochemistry
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
16
Spring 2016
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.
Thermochemistry
Enthalpy of Reaction
In a chemical reaction, the
quantity, H, is called the
enthalpy of reaction, or the
heat of reaction.
Thermochemistry
17
Spring 2016
Endothermicity and Exothermicity
 A process is endothermic,
when H is positive.
Thermochemistry
Endothermicity and Exothermicity
 A process is endothermic
when H is positive.
 A process is exothermic
when H is negative.
Thermochemistry
18
Spring 2016
The Truth about Enthalpy
1.
2.
3.
Enthalpy is an extensive property.
H for a reaction in the forward direction is equal in
size, but opposite in sign, to H for the reverse
reaction.
H for a reaction depends on the state of the products
and the state of the reactants.
Thermochemistry
Heat Capacity and Specific Heat
 The amount of energy required to raise the
temperature of a substance by 1C is its heat
capacity.
 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C .
Thermochemistry
19
Spring 2016
Heat Capacity and Specific Heat
Specific heat, then, is
heat transferred
Specific heat =
mass  temperature change
q
s=
m  T
q = m  s  T
Thermochemistry
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 H values that
are published and the properties of enthalpy.
Thermochemistry
20
Spring 2016
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
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
21
Spring 2016
Enthalpies of Formation
The enthalpy of formation or "heat of formation" of a
compound is the change of enthalpy that accompanies
the formation of 1 mole of a substance from its elements
in their elemental form
 For example, enthalpy of formation of carbon dioxide
would be the enthalpy of the following reaction:
 C(s, graphite) + O2 (g) → CO2 (g)
Thermochemistry
Standard Enthalpies of Formation
Standard enthalpies of formation, Hf
, are measured under standard
conditions (25°C and 1.00 atm pressure).
The standard enthalpy of formation for an element in its standard state is
ZERO!!!! So, H°f for C (s, graphite) is zero, but the
H°f for C (s, diamond) is 2 kJ/mol. That is because graphite is the
standard state for carbon, not diamond.
Thermochemistry
22
Spring 2016
Calculation of H°rxn
ΔH°rxn = ΣmΔH°f (Products) – ΣnΔH°f (Reactants)
where n and m are the stoichiometric coefficients.
Thermochemistry
Calculation of H°rxn
C3H8 (g) + 5 O2 (g)  3 CO2 (g) + 4 H2O (l)
Thermochemistry
23
Spring 2016
Calculation of H°rxn
C3H8 (g) + 5 O2 (g)  3 CO2 (g) + 4 H2O (l)
H°rxn= [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
Energy in foods and fuels
Most of the fuel in the food we eat comes from carbohydrates
and fats.
Fuel value = energy released when one gram of material is
combusted
Most of fuel energy comes from coal, natural gases, charcoal,
etc.
Thermochemistry
24
Spring 2016
END
Thermochemistry
Practice exercise
SAMPLE EXERCISE 5.1 Describing and calculating energy changes
1. A bowler lifts a 5.4-kg (12-lb) bowling ball from ground level to a
height of 1.6 m (5.2 feet) and then drops the ball back to the
ground.
(a)
What happens to the potential energy of the bowling ball as it is
raised from the ground?
(b)
What quantity of work, in J, is used to raise the ball?
(c)
After the ball is dropped, it gains kinetic energy. If we assume that
all of the work done in part (b) has been converted to kinetic
energy by the time the ball strikes the ground, what is the speed
of the ball at the instant just before it hits the ground?
(Note: The force due to gravity is F = m  g, where m is the mass
of the object and g is the gravitational constant; g = 9.8 m/s2.)
Thermochemistry
25
Spring 2016
(a) Because the bowling ball is raised to a greater height above the
ground, its potential energy increases.
(b) The ball has a mass of 5.4 kg, and it is lifted a distance of 1.6 m.
To calculate the work performed to raise the ball,
Thus, the bowler has done 85 J of work to lift the ball to a height of
1.6 m.
(c) When the ball is dropped, its potential energy is converted to kinetic
energy. At the instant just before the ball hits the ground, we assume
that the kinetic energy is equal to the work done in part (b), 85 J:
v
2  Ek
2  85

 5.6 m / s
m
5.4
Thermochemistry
SAMPLE EXERCISE 5.2 Relating heat and work to changes of internal energy
Two gases, A(g) and B(g), are confined in a cylinder-and-piston arrangement
like that in Figure 5.3.
Substances A and B react to form a solid product:
As the reactions occurs, the system loses 1150 J of heat to the surrounding.
The piston moves downward as the gases react to form a solid. As the volume
of the gas decreases under the constant pressure of the atmosphere, the
surroundings do 480 J of work on the system. What is the change in the
internal energy of the system?
Solution
Heat is transferred from the system to the surroundings, and work is done on
the system by the surroundings, so q is negative and w is positive: q =–1150 J
and w = 480 kJ. Thus, E is
The negative value of E tells us that a net quantity of 670 J of energy has
been transferred from the system to the surroundings.
Thermochemistry
26
Spring 2016
SAMPLE EXERCISE 5.2 Relating heat and work to changes of internal energy
Two gases, A(g) and B(g), are confined in a cylinder-and-piston arrangement
like that in Figure 5.3.
Substances A and B react to form a solid product:
As the reactions occurs, the system loses 1150 J of heat to the surrounding.
The piston moves downward as the gases react to form a solid. As the volume
of the gas decreases under the constant pressure of the atmosphere, the
surroundings do 480 J of work on the system. What is the change in the
internal energy of the system?
Solution
Heat is transferred from the system to the surroundings, and work is done on
the system by the surroundings, so q is negative and w is positive: q =–1150 J
and w = 480 kJ. Thus, E is
The negative value of E tells us that a net quantity of 670 J of energy has
been transferred from the system to the surroundings.
Thermochemistry
SAMPLE EXERCISE 5.3 Determining the Sign of H
Indicate the sign of the enthalpy change, H, in each of the following processes
carried out under atmospheric pressure, and indicate whether the process is
endothermic or exothermic:
(a) An ice cube melts
(b) 1 g of butane (C4H10) is combusted in sufficient oxygen to give complete
combustion to CO2 and H2O.
Solution
In (a) the water that makes up the ice cube is the system. The ice cube absorbs
heat from the surroundings as it melts, so H is positive and the process is
endothermic. In (b) the system is the 1 g of butane and the oxygen required to
combust it. The combustion of butane in oxygen gives off heat, so H is
negative and the process is exothermic.
Thermochemistry
27
Spring 2016
SAMPLE EXERCISE 5.3 Determining the Sign of H
Indicate the sign of the enthalpy change, H, in each of the following processes
carried out under atmospheric pressure, and indicate whether the process is
endothermic or exothermic:
(a) An ice cube melts
(b) 1 g of butane (C4H10) is combusted in sufficient oxygen to give complete
combustion to CO2 and H2O.
Solution
In (a) the water that makes up the ice cube is the system. The ice cube absorbs
heat from the surroundings as it melts, so H is positive and the process is
endothermic. In (b) the system is the 1 g of butane and the oxygen required to
combust it. The combustion of butane in oxygen gives off heat, so H is
negative and the process is exothermic.
Thermochemistry
PRACTICE EXERCISE
Hydrogen peroxide can decompose to water and oxygen by the following
reaction:
Calculate the value of q when 5.00 g of H2O2(l) decomposes at constant
pressure.
Thermochemistry
28
Spring 2016
SAMPLE EXERCISE 5.4
How much heat is released when 4.50 g of methane gas is burned in a
constant-pressure system? (Use the information given in Equation 5.18.)
Solution
By adding the atomic weights of C and 4 H, we have 1 mol CH4 = 16.0 g CH4.
Thus, we can use the appropriate conversion factors to convert grams of CH 4
to moles of CH4 to kilojoules:
The negative sign indicates that 250 kJ is released by the system into the
surroundings.
Thermochemistry
SAMPLE EXERCISE 5.4
How much heat is released when 4.50 g of methane gas is burned in a
constant-pressure system? (Use the information given in Equation 5.18.)
Solution
By adding the atomic weights of C and 4 H, we have 1 mol CH4 = 16.0 g CH4.
Thus, we can use the appropriate conversion factors to convert grams of CH 4
to moles of CH4 to kilojoules:
The negative sign indicates that 250 kJ is released by the system into the
surroundings.
Thermochemistry
29
Spring 2016
SAMPLE EXERCISE 5.3 Determining the Sign of H
Indicate the sign of the enthalpy change, H, in each of the following processes
carried out under atmospheric pressure, and indicate whether the process is
endothermic or exothermic:
(a) An ice cube melts
(b) 1 g of butane (C4H10) is combusted in sufficient oxygen to give complete
combustion to CO2 and H2O.
Solution
In (a) the water that makes up the ice cube is the system. The ice cube absorbs
heat from the surroundings as it melts, so H is positive and the process is
endothermic. In (b) the system is the 1 g of butane and the oxygen required to
combust it. The combustion of butane in oxygen gives off heat, so H is
negative and the process is exothermic.
Thermochemistry
SAMPLE EXERCISE 5.3 Determining the Sign of H
Indicate the sign of the enthalpy change, H, in each of the following processes
carried out under atmospheric pressure, and indicate whether the process is
endothermic or exothermic:
(a) An ice cube melts
(b) 1 g of butane (C4H10) is combusted in sufficient oxygen to give complete
combustion to CO2 and H2O.
Solution
In (a) the water that makes up the ice cube is the system. The ice cube absorbs
heat from the surroundings as it melts, so H is positive and the process is
endothermic. In (b) the system is the 1 g of butane and the oxygen required to
combust it. The combustion of butane in oxygen gives off heat, so H is
negative and the process is exothermic.
Thermochemistry
30
Spring 2016
SAMPLE EXERCISE 7 Relating Heat, Temperature Change, and Heat
Capacity
(a) How much heat is needed to warm 250 g of water (about 1 cup) from 22 oC
(about room temperature) to near its boiling point, 98 oC? The specific heat of
water is 4.18 J/g-K. (b) What is the molar heat capacity of water?
Solution
(a) The water undergoes a temperature change of
T = 98 ºC – 22 ºC = 76 ºC = 76 K
q = s x m  T
= (4.18 J/g-K)(250 g)(76 K) = 7.9  104 J
(b) The molar heat capacity is the heat capacity of one mole of substance.
Using the atomic weights of hydrogen and oxygen, we have
1 mol H2O = 18.0 g H2O
From the specific heat given in part (a), we have
Thermochemistry
SAMPLE EXERCISE 7 Relating Heat, Temperature Change, and Heat
Capacity
(a) How much heat is needed to warm 250 g of water (about 1 cup) from 22 oC
(about room temperature) to near its boiling point, 98 oC? The specific heat of
water is 4.18 J/g-K. (b) What is the molar heat capacity of water?
Solution
(a) The water undergoes a temperature change of
T = 98 ºC – 22 ºC = 76 ºC = 76 K
q = s x m  T
= (4.18 J/g-K)(250 g)(76 K) = 7.9  104 J
(b) The molar heat capacity is the heat capacity of one mole of substance.
Using the atomic weights of hydrogen and oxygen, we have
1 mol H2O = 18.0 g H2O
From the specific heat given in part (a), we have
Thermochemistry
31
Spring 2016
SAMPLE EXERCISE 7
The enthalpy of reaction for the combustion of C to CO 2 is – 393.5 kJ/mol C,
and the enthalpy for the combustion of CO to CO2 is – 283.0 kJ/mol CO:
Using these data, calculate the enthalpy for the combustion of C to CO:
Thermochemistry
SAMPLE EXERCISE 8
Calculate H for the reaction
given the following chemical equations and their respective enthalpy changes:
Thermochemistry
32
Spring 2016
SAMPLE EXERCISE 5.4
How much heat is released when 4.50 g of methane gas is burned in a
constant-pressure system? (Use the information given in Equation 5.18.)
Solution
By adding the atomic weights of C and 4 H, we have 1 mol CH4 = 16.0 g CH4.
Thus, we can use the appropriate conversion factors to convert grams of CH 4
to moles of CH4 to kilojoules:
The negative sign indicates that 250 kJ is released by the system into the
surroundings.
Thermochemistry
SAMPLE EXERCISE 5.4
How much heat is released when 4.50 g of methane gas is burned in a
constant-pressure system? (Use the information given in Equation 5.18.)
Solution
By adding the atomic weights of C and 4 H, we have 1 mol CH4 = 16.0 g CH4.
Thus, we can use the appropriate conversion factors to convert grams of CH 4
to moles of CH4 to kilojoules:
The negative sign indicates that 250 kJ is released by the system into the
surroundings.
Thermochemistry
33
Spring 2016
SAMPLE EXERCISE 7 Relating Heat, Temperature Change, and Heat
Capacity
(a) How much heat is needed to warm 250 g of water (about 1 cup) from 22 oC
(about room temperature) to near its boiling point, 98 oC? The specific heat of
water is 4.18 J/g-K. (b) What is the molar heat capacity of water?
Solution
(a) The water undergoes a temperature change of
T = 98 ºC – 22 ºC = 76 ºC = 76 K
q = s x m  T
= (4.18 J/g-K)(250 g)(76 K) = 7.9  104 J
(b) The molar heat capacity is the heat capacity of one mole of substance.
Using the atomic weights of hydrogen and oxygen, we have
1 mol H2O = 18.0 g H2O
From the specific heat given in part (a), we have
Thermochemistry
SAMPLE EXERCISE 7 Relating Heat, Temperature Change, and Heat
Capacity
(a) How much heat is needed to warm 250 g of water (about 1 cup) from 22 oC
(about room temperature) to near its boiling point, 98 oC? The specific heat of
water is 4.18 J/g-K. (b) What is the molar heat capacity of water?
Solution
(a) The water undergoes a temperature change of
T = 98 ºC – 22 ºC = 76 ºC = 76 K
q = s x m  T
= (4.18 J/g-K)(250 g)(76 K) = 7.9  104 J
(b) The molar heat capacity is the heat capacity of one mole of substance.
Using the atomic weights of hydrogen and oxygen, we have
1 mol H2O = 18.0 g H2O
From the specific heat given in part (a), we have
Thermochemistry
34