Transcript Figure 6.15 When a reaction

Chapter 6 Thermochemistry:
The Fire Within
Study of the heat released or required by chemical reactions.
This canopy walkway and net in the
Peruvian rainforest allow biomass
research to be conducted high above
the ground. The source of most of our
energy is the Sun. Through the process
of photosynthesis, solar radiation
causes chemical reactions in green
plants that store the energy for future
use. We make use of the energy
captured by plants when we burn fuels.
Fossil fuels contain energy that has
been stored for thousands of years, but
research into alternative fuels is
finding ways to make efficient use of
plants for fuel. This chapter presents
the basic concepts used in research
into the energy changes that
accompany all chemical reactions.
Assignment for Chapter 6
14, 23, 35, 47, 55, 61
Energy, Heat, Enthalpy
CH4(g)+2O2(g)CO2(g)+2H2O(l)+energy
C6H12O6(aq)+O2(g)6CO2(g)+6H2O(l)+energy
Energetics: How energy is transformed and used
Bioenergetics: Study of the use of energy in organisms
Conservation and transformation of energy
Heat
Enthalpy
Figure 6.1 The potential energy of a mass, m, is proportional
to its height, h, above a surface that is taken to correspond to
zero potential energy.
Figure 6.2 Kinetic energy (represented by the height of the
dark green bar) and potential energy (the light green bar) can
be converted into one another. However, their sum (the total
height of the bar) is a constant in the absence of external
influences, such as air resistance. A ball thrown up from the
ground loses kinetic energy as it slows but gains potential
energy. The reverse happens as it falls back to Earth.
Conservation of energy:
Ek+Ep=const. for an isolated body
Internal Energy
U=sum of the kinetic and
potential energies of all the
atoms and molecules in a
sample.
N
U   ( E k ,i  E p ,i )
i 1
Figure 6.3
In thermodynamics, the world is divided into a system (the
object of interest) and the surroundings (everything else). In
practice, the surroundings may be a constant-temperature
water bath. The arrows represent the energy being transferred
between the system and its surroundings.
Heat(energy) exchange(transfer)
Matter exchange(transfer)
Figure 6.4
We classify systems into one of three kinds, according to their
interactions with their surroundings. An open system can
exchange matter and energy with its surroundings. A closed
system can exchange energy but not matter. An isolated
system can exchange neither matter nor energy.
Figure 6.5
The booster rockets on the space shuttle form an open
system. The stream of gases produced by the chemical
reaction pours out of the engines and moves the rocket. (See
Applying Chemistry: Case Study 20.)
Figure 6.6
When we heat a system, we make use of a difference in
temperature between it and the surroundings to induce energy
to flow through the walls of the system. Heat flows from high
temperature to low.
Heat is a transfer of energy that occurs as
a result of a temperature difference.
1 cal = 4.184 J
Figure 6.7
When energy leaves a system as a result of a temperature
difference between the system and the surroundings, we say
that the system has lost energy as heat. This transfer of
energy stimulates the thermal motion of molecules in the
surroundings.
Thermal motion = random molecular
motion (in gases, liquids & solids)
Figure 6.8
A system does work when it expands against an external
pressure. Here we see a gas that pushes a piston out against a
pressure, P. We shall see shortly that the work done is
proportional to both the pressure and the change in volume
that the system undergoes.
Work is a transfer of energy that takes
place when an object is moved against
an opposing force.
Figure 6.9 When a system expands, it performs work on its
surroundings by forcing all the molecules in another object in
the surroundings to move in the same direction. Here we see
the expansion of a gas raise a weight. The expansion of gases
in the cylinders of automobiles do work by pushing on the
piston, which turns the gears that move the vehicle.
Three Ways for Changing the Internal
Energy of a System
• Adding matter
• Supplying energy as heat (thermal motion)
• Doing work (uniform motion)
Figure 6.10 When we wind a spring, the potential energy of
the atoms changes because they are squashed together and
repel one another. The internal energy of the spring rises as a
result of this increase in potential energy.
Another type of potential energy:
Interactions between atoms in a
system.
Figure 6.11 The internal energy of a system can be changed
either by doing work or by heating. The diagram shows that
the change in internal energy is positive (U increases) when
energy is supplied in either way. When energy leaves the
system as heat or work, the internal energy falls and DU is
negative.
Unit of Enery
1 J  1 kg  m 2 /s 2
DU  w  q
Figure 6.12 The inventor of this elaborate device, the Keely
motor, claimed that it could generate high pressures and
energies from a small amount of water. However, like other
perpetual motion machines, the Keely motor was found to be a
fraud. The work was accomplished by compressed air from a
hidden source.
The First Law of Thermodyna mics :
For an isolated system : DU  0
Figure 6.13 A reaction does work when it generates a gas. For
example, carbon dioxide is formed from the thermal
decomposition of calcium carbonate. As the gas is formed, it
drives back the surrounding atmosphere.
W=?
Figure 6.14 The gas in the piston expands a distance, l,
against a pressure, P. The volume increase is l  A, where A is
the cross-sectional area of the piston.
work done  distance moved  opposing force
opposing force  area  external pressure
work done  distance moved  area  external pressure
 change in volume  external pressure
w  -PextΔV
An Example for Calculating Work
A gas expands by 12.0 L against a pressure
of 2.0 atm. How much work is done?
w   Pex DV
 12.0L  2.0atm
 24.0L  atm
 24.0  10 3 m 3 1.01325  105 Pa
 2.43  103 J
Figure 6.15 When a reaction (such as the thermal
decomposition of calcium carbonate) takes place in a closed,
constant-volume container, the gas fills the container but
cannot expand against the surrounding atmosphere. As a
result, it does no work on the surroundings.
At constant v olume :
w0
DU  q
Figure 6.16 When a system is free to expand against an
external pressure, some of the energy supplied to it as heat
escapes back into the surroundings as work. As a result, the
change in internal energy is less than the energy supplied.
At constant pressure :
w   PDV
DU  q  PDV
Enthalpy
At constant pressure :
w   PDV
D U  q  PD V
DH  DU  PDV
q
Change in enthalpy of a system is equal to
the heat transferred to it at constant pressure
Figure 6.17 The thermite reaction is another highly exothermic
reaction—one that can melt the metal it produces. In this
reaction, aluminum metal is reacting with iron(III) oxide, Fe2O3,
causing a shower of molten iron sparks. In an exothermic
reaction, energy is lost as heat, the amount lost depending on
the amount of reactants available.
At constant pressure :
DH  0  exothermic process
Figure 6.18
The reaction between ammonium thiocyanate, NH4SCN, and
barium hydroxide octahydrate, Ba(OH)2•8H2O, absorbs a lot of
heat and can cause water vapor in the air to freeze on the
outside of the beaker. In an endothermic reaction, energy is
absorbed as heat.
2NH4SCN+ Ba(OH)2•8H2O
Ba(SCN)2+2NH3 (g)+10H2O (g)
At constant pressure :
DH  0  endothermic process
Classroom Exercise
• In an endothermic reaction, 5 kcal of heat is
absorbed and 1.2 L of gas is generated. How
much is the change of the internal energy?
At constant pressure:
w   PDV
DU  q  PDV
 5000 cal  4.184 J/cal  105  1.2J
 9800J  9.8 kJ
Figure 6.19
Heat capacity is an extensive property, so a large object (bottom) has a
larger heat capacity than a small object (top) made of the same material.
Heat capacity 
C
heat supplied
temperature rise
q
DT
An Ethanol Sample:
C
98 kJ (heat supplied)
2 0 C(temperature rise)
 49 kJ/ o C
Specific Heat Capacity
Specific Heat Capacity
(
heat supplied
temperature rise
Cs 
q
DT m
) / mass
Using Specific Heat Capacity
• 5.10 g sample of an alloy with specific heat capacity
of 0.124 J/C/g is heated from 24.2 C to 138.5 C.
How much energy is supplied?
Specific Heat Capacity
heat supplied
 ( temperatur
e rise ) / mass
Cs 
q
DT m
Heat supplied 
mass  specific heat capacity  temperatur e rise
 5.10 g  0.124
 72.3J
J
o
Cg
 (138.5  24.2) o C
Figure 6.20 A pyrotechnics expert attaches fuses containing
potassium chlorate to fireworks set up in mortars. The highly
exothermic reactions of fireworks do not begin until they are
intiated by a fuse. The fuse ignites a mixture of carbon and
other reducing agents that react with an oxidizer such as
potassium perchlorate.
Figure 6.21 The quantity of heat released or absorbed by a reaction can
be measured in this primitive version of a calorimeter. The outer
polystyrene cup acts as an extra layer of insulation to ensure that no
heat enters or leaves the inner cup. The quantity of heat released or
absorbed is proportional to the change in temperature of the
calorimeter.
Figure 6.22 A bomb calorimeter. The combustion is initiated
with an electrical fuse. Once the reaction has begun, energy is
released as heat that spreads through the walls of the bomb
into the water. The heat released is proportional to the
temperature change of the entire calorimeter assembly.
Calorimetry
• 50.0 g water at 20 C is mixed with 21 g of iron at
90.2 C. The equilibrium temperature is 23.2 C. Find
the specific heat capacity of iron.
The heat gain of water  The heat loss of iron
50 g  4.184J/ C/g  (23.2  20.0) C
o
o
 21 g  x J/ C/g  (90.2  23.2) C
o
x  0.48 J/ C/g
o
o
Enthalpy H = U + PV
Benoit Paul Émile Clapeyron
Rudolf Clausius
J. Willard Gibbs
Heike Kamerlingh Onnes
Thermodynamics related to eating/drinking, so much, more than entire physics…….
How does power come from?
H=U+PV
Ice cold water  hot water  vapor
Vaporization: enthalpy of
vaporization
DH vap  H vapor  H liquid
Figure 6.23
Melting (fusion) is an endothermic process. As molecules
acquire energy, they begin to struggle past their neighbors.
Finally the sample changes from a solid with ordered
molecules (left) to a liquid with disordered, mobile molecules
(right).
Melting/freezing: enthalpy of
fusion/freezing
DH fus  H liquid  H solid
Figure 6.24
The enthalpy change for the reverse of a process is the
negative of the enthalpy change for the forward process at the
same temperature.
Figure 6.25
The polar ice caps on Mars extend and recede with the seasons.
They are solid carbon dioxide and form by direct conversion of the
gas to a solid. They disappear by sublimation. Although some water
ice is also present in the polar caps, the temperature on Mars never
becomes high enough to melt it. On Mars, ice is just another rock.
Figure 6.26
Because enthalpy is a state property, the enthalpy of
sublimation at a given temperature can be expressed as the
sum of the enthalpies of fusion and vaporization measured at
the same temperature.
Sublimation: enthalpy of
sublimation
DH sub  H vapor  H solid
DH sub  DH fus  DH vap
Thermochemistry of Some Physical Changes
Vaporization: enthalpy of vaporization
DH vap  H vapor  H liquid
Melting/freezing: enthalpy of fusion/freezing
DH fus  H liquid  H solid
Sublimation: enthalpy of sublimation
DH sub  H vapor  H solid
DH sub  DH fus  DH vap
Figure 6.27 The heating curve of water. The temperature of the solid
rises as heat is supplied. At the melting point, the temperature remains
constant and the heat is used to melt the sample. When enough heat
has been supplied to melt all the solid, the temperature of the liquid
begins to rise again. A similar pause in the temperature rise occurs at
the boiling point.
DH  C s , solid mDT1  DH
fus
 C s ,liquidmDT 2 DH vap
 C s , vapor mDT 3
The Enthalpy of Chemical Change:
Reaction Enthalpies
• ReactantsPorducts+heat (exothermic)
• Reactants+heatProducts (endothermic)
CH4(g)+2O2(g)CO2(g)+2H2O(l)+heat
CH4(g)+2O2(g)CO2(g)+2H2O(l) ΔHr=-890.0 kJ/mol
(Thermochemical Equation)
Molar Reaction Enthalpy (CH4)
Figure 6.28
A biophysicist monitors an experimental fermentation
chamber in which fuel ethanol is being produced from waste
biomass by a genetically engineered strain of bacteria.
+
Many such chemical reactions
take place in your body:
Figure 6.29
This diagram shows how the value of the reaction enthalpy
depends on the physical states of a product. When water is
produced as a vapor rather than as a liquid in the combustion
of methane, 88 kJ remains stored in the system for every 2 mol
H2O produced.
CH4(g)+2O2(g)CO2(g)+2H2O(g) ΔHr=-802.0 kJ/mol
CH4(g)+2O2(g)CO2(g)+2H2O(l) ΔHr=-890.0 kJ/mol
Figure 6.30
The standard reaction enthalpy is the difference in enthalpy
between the pure products, each at 1 atm, and the pure
reactants at the same pressure and the specified temperature
(which is commonly but not necessarily 25°C). The scheme
here is for the combustion of methane.
Standard reaction enthalpy:
CH4(g)+2O2(g)CO2(g)+2H2O(l) ΔHo=-890.0 kJ/mol
Investigating Matter 6.1 (a)
Because fossil fuel reserves are limited, they must be
extracted wherever they are found. This platform is used to
pump petroleum from beneath the ocean; however, the natural
gas accompanying it cannot be easily transported and so is
burned off.
Investigating Matter 6.1 (b)
An agricultural researcher assesses the growth rate of a
seedling. Plant photosynthesis is only about 3% efficient, and
conditions that increase this efficiency are actively being
sought.
Reaction Enthalpy of A Reverse Reaction
P4(s)+6Cl2(g)4PCl3(l)
DH o  1279 kJ
4PCl3(l) P4(s)+6Cl2(g)
DH o  1279 kJ
C6H12O6(aq)+6O2(g)6CO2(g)+6H2O(l)
DH o  2808 kJ
6CO2(g)+6H2O(l) C6H12O6(aq)+6O2(g)
DH o  2808 kJ
Figure 6.31
If the overall reaction can be broken down into a series of
steps, then the corresponding overall reaction enthalpy is the
sum of the reaction enthalpies of the steps on the alternative
path. None of the steps need be a reaction that can actually be
carried out in the laboratory.
Hess’s Law
DH  DH  DH  DH    
o
Germain Henri Hess
(1802 - 1850)
o
1
o
2
o
3
Using Hess’s Law
C(s)+O2(g)CO2(g)
DH o  ?
C(s)+(1/2)O2(g)CO(g)
DH o  110.5 kJ
+ CO(g)+(1/2)O2(g)CO2(g)
DH o  283.0 kJ
C(s)+O2(g)CO2(g)
DH o  393.5 kJ
Using Hess’s Law
DH o  ?
3C(s)+4H2(g)C3H8(g) (X)
C3H8(g)+5O2(g)3CO2(g)+4H2O(l)
(A)
C(s)+O2(g)CO2(g) (B)
H2(g)+(1/2)O2(g)H2O(g) (C)
DH o  2220.0 kJ
DH o  394.0 kJ
DH o  286.0 kJ
X=3*B-A+4*C DH o  3  (-394) - (-2220)  4 * (-286) kJ
 106 kJ
Figure 6.32
The amount of heat produced or absorbed in a chemical
reaction can be determined from the reaction stoichiometry.
The Enthalpy of Chemical Change:
Reaction Enthalpies
Enthalpy of combustion
Enthalpy of formation
Enthalpy of Combustion
CH4(g)+2O2(g)CO2(g)+2H2O(g) ΔHc=-802.0 kJ/mol
How much heat is produced by burning 150.0 g of methane?
q  150.0gCH 4 
 - 8.32 10 kJ
3
1mol CH 4
16.04gCH4

890kJ
1molCH 4
Figure 6.33 During World War II, fuel was in short supply and all
manner of ingenious solutions were sought. However, as we can see
from this photograph of a vehicle powered by coal gas (a mixture of
carbon monoxide and hydrogen) in London, the low enthalpy density
of gases creates storage problems. A modern approach to using
gases to power a vehicle can be seen in Applying Chemistry: Case
Study 18.
Different Units for Enthalpy
Specific Enthalpy
The enthalpy of combustion per gram
DH /m Unit : kJ/g
o
Enthalpy Density
The enthalpy of combustion per liter
DH o /V Unit : kJ/L
Figure 6.34 The range and speed of this electric-powered car depend on
the type of battery it uses. For example, metal-hydride devices have a
longer range than lead-acid storage batteries. As this driver would
agree, recharging is generally a slow process. The cable attached to this
car may look like a gasoline hose, but it is actually delivering electricity
while its owner waits. The use of hydrogen as fuel could reduce the
number of refueling stops.
Standard Enthalpies of Formation
The standard enthalpy of formation of an element in
its most stable form is 0.
The stanard enthalpy of reaction
DH f  H
o
o
final
H
o
initial
  ni DH (products )   n j DH (reactants )
o
f
i
o
f
j
Figure 6.35
The reaction enthalpy can be constructed from enthalpies of
formation by imagining the formation of both the reactants
and the products from their respective elements. The reaction
enthalpy is the difference between the two.
Standard Enthalpies of Formation
4C(s)+6H2(g) +O2(g)2C2H5OH(l) ΔHc=-555.38 kJ
The stanard enthalpy of formation per mole
of ethanol
DH f (C 2 H 5OH, l) 
o
-555.38 kJ
2 mol C 2 H 5 OH
 -277.69 kJ/mol of C 2 H 5OH
Classroom Exercise
2C(s)+O2(g)2CO (g) ΔHc=-221.06 kJ
Standard enthalpies of formation of CO?
The stanard enthalpy of formation per mole
of CO
DH f o (CO,g) 
-221.06 kJ
2 mol CO
 -110.53 kJ/mol of CO
The standard enthalpy of formation of an element in
its most stable form is 0.
Using Standard Enthalpy of Formation
2C 2 H 2 ( g )  5O2 ( g )  4CO2  2 H 2O(l ), DH r  ?
The total enthalpy of formation of the reactants:
o
H initial
  n j DH of (reactants)  2DH of (C2 H 2 , g )  5DH of (O2 , g )
j
 {2  (226.73)  5  (0)} kJ  453.46 kJ
The total enthalpy of formation of the products:
H ofinal   ni DH of (products)
i
 4DH of (CO2 , g )  2DH of ( H 2O, l )  {4  (393.51)  2  (285.83)} kJ
 -2145.70 kJ
Enthalpy of Reaction:
DH ro  2599.16 kJ
Enthalpy of combustion per mole of ethyne:
DH co  1299.58 kJ/mol C2H 2
Fun Chemistry: Calorie of Food
• It is equal to specific combustion enthalpy: kcal/g
• It is in the sense of average and approximation,
depending on the location and growth conditions
of the original produces, the genetics of the
people who consumes the food…
Compound A + O2  Compound B + Compound C + …
ΔHc = calorie
It is assumed that the calorie of a compound (protein, carbonhydrate etc)
metabolized in body is the same as that when the compound is burnt
outside.
Measuring Calories
•
•
•
•
•
The caloric value of food is the energy produced by combustion of its proteins, carbohydrates
and fats. The amount of energy liberated by the catabolism of food in the body is almost the
same as the amount liberated when food is burnt outside the body. The energy liberated by
catabolic processes is used for maintaining body functions namely digestion, thermoregulation,
muscular contraction and nerve impulses conduction. The amount of energy liberated / unit
time is the metabolic rate. When food is burnt outside the body, all the energy is liberated as
heat.
The standard unit of heat energy is the calorie (cal), which is defined as the amount of heat
energy necessary to raise the temperature of 1 ml of water by one degree, from 15 to 16 celsius
at rest.
A slightly different calorie is used in engineering, the international calorie, which equals 1/860
international watt-hour (4.1868 J). A large calorie, or kilocalorie, usually referred to simply as a
calorie and sometimes as a kilogram calorie, equals 1,000 calories and is the unit used to express
the energy-producing value of food in the calculation of diets.
The energy released by combustion of foodstuffs outside the body can be measured directly and
indirectly. In direct calorimetry method the liberated energy can be measured using a bomb
calorimeter. It is a metal vessel surrounded by a water insulated container. The food is ignited by
an electric spark. The change in the temperature of water is a measure of the calories produced.
In indirect method, the energy production can also be calculated by
measuring the amount of oxygen consumed for combustion of food.
The amount of oxygen consumed / unit of time is proportionate to the
energy liberated. This method of energy estimation is called wet
combustion. The caloric value of carbohydrate is 4.1 Kcal /g,
protein is 5.65 Kcal/ g and fat is 9.4 Kcal/g.
Case Study 6 (a)
Regular exercise not only is good for the metabolism, it can be
fun, too, when we make it a part of daily life.
When you exercise, you burn the nutrients in your body
by speeding up metabolism and spend more calories.
Case Study 6 (b)
Energy consumed (in kilojoules per hour) in typical activities:
blue for a 70-kg male and pink for a 58-kg female.
Biological systems are wonderful
chemicals and wonderful chemical plants
Thermodynamics related to eating/drinking, so much, more than entire physics…….
Now, a six-table banquet is served.
When enjoying the delicacies,
refresh yourself of the meaning
of “calorie” Dr. Ding told you.
Assignment for Chapter 6
14, 23, 35, 47, 55, 61