Download THERMODYNAMICS

3/15/2015
Thermochemistry deals with changes in heat during
chemical reactions. A main goal of the study of
thermochemistry is to determine the quantity of heat
exchanged between a system and its surroundings. The
THERMODYNAMICS
system is the part of the universe being studied, while the
Thermochemistry
surroundings are the rest of the universe that interacts
with the system.
Definitions
2

System and surroundings
An open system is a system that freely
exchanges energy and matter with its
surroundings.
Open system
3
4
1
3/15/2015
Open systems

Closed system exchanges energy but not
matter with an outside system. Though it
is typically portion of larger system, it is
not in complete contact.
Closed system
WM_ THERMODYNAMICS
5
WM_ THERMODYNAMICS
6
Isolated system can exchange neither energy nor
matter with an outside system. While it may be
portion of larger system, it does not communicate
with the outside in any way.
The physical universe is an isolated system;
a closed thermos bottle is essentially an isolated
system (though its isolation is not perfect).
Isolated system
WM_ THERMODYNAMICS
7
WM_ THERMODYNAMICS
8
2
3/15/2015
9
Comparison of systems
10
Q. Identify system type (open, closed or isolated) from
descrition below and fill the enpty space in the table
below
System description
System type
Coffee in perfectly closed Thermos® flask
•Q,
•A closed system contains 2g of ice. Another 2g
of ice are added to the system. What is the final
mass of the system?
•Q
•An isolated system has an initial temperature of 30oC.
It is then placed on top of a bunsen burner for an hour.
What is the final temperature?
Combustion of gasoline in car engine
Mercury in thermometer
Living plant
Quiz
Electric battery
11
WM_ THERMODYNAMICS
12
3
3/15/2015
Q. Which type of thermodynamic system is:
1. an ocean?
2. an aquarium?
3. a pizza delivery bag?
4. a greenhouse?
Matter and
energy
ENERGY
13

WM_ THERMODYNAMICS
14
Matter is anything that has mass and
occupies space.

Mass is a measure of the quantity of
matter in a sample of any material.
 The more massive an object is, the more
force is required to put it in motion.
 All bodies consist of matter.

MATTER AND ENERGY
Energy is measure of the ability of a
body or system to do work or produce any
change. No activity is possible without
energy.
Energy
15
16
4
3/15/2015
Our senses of sight and touch usually tell
us that an object occupies space.
 In the case of colorless, odorless,
tasteless gases (such as air), our senses
may fail us.
 Additionally, we cannot say anything
about total energy which the object
possesses.

Chemical changes always involve energy
changes, some energy transformations do
not involve chemical changes at all.

For example, heat energy may be
converted into electrical energy or into
mechanical energy without any
simultaneous chemical changes.

17
18
Many experiments have demonstrated
that all of the energy involved in any
chemical or physical change appears in
some form after the change.
 These observations are summarized in the
Law of Conservation of Energy:
 Energy cannot be created or destroyed in
a chemical reaction or in a physical
change.
 It can only be converted from one form to
another.

Interceptor transforms energy of sun
into internal water energy
Warn water
Cold water
Warm air
Use of sun energy for
house heating
Law of Conservation of Energy
WM_ THERMODYNAMICS
19
20
5
3/15/2015



With the dawn of the nuclear age in the
1940s, scientists, and then the world,
became aware that matter can be converted
into energy.
In nuclear reactions, matter is transformed
into energy.
The relationship between matter and energy
is given by Albert Einstein’s now famous
equation:
At the present time, we have not
(knowingly) observed the transformation
of energy into matter on a large scale.
 It does, however, happen on an extremely
small scale in “atom smashers,” or particle
accelerators, used to induce nuclear
reactions (like in Large Hadron Collider in
CERN).

E = m c2
The Law of Conservation of Matter
and Energy
The Law of Conservation of Matter
and Energy
21

Now that the equivalence of matter and
energy is recognized, the Law of
Conservation of Matter and Energy
can be stated in a single sentence:
22
Energy is very important in every aspect
of our daily lives.
 The food we eat supplies the energy to
sustain life with all of its activities and
concerns.
 The availability of relatively inexpensive
energy is an important factor in our
technological society.

 The
combined amount of matter
and energy in the universe is
fixed.
MEANING OF THE ENERGY IN THE
LIFE
The Law of Conservation of Matter
and Energy
WM_ THERMODYNAMICS
WM_ THERMODYNAMICS
23
24
6
3/15/2015
•
•
•
•
The concept of energy is at the every
heart of science.
All physical and chemical processes are
accompanied by the transfer of energy.
Energy cannot be created or destroyed.
We must understand how to do the
“accounting” of energy transfers from one
body or one substance to another or from
one form of energy to another.
In thermodynamics we study the
energy changes that accompany
physical and chemical processes.
 Usually these energy changes involve
heat—hence the “thermo-” part of the
term.
 There are the two main aspects of
thermodynamics.

OBJECT OF THE
THERMODYNAMICS
„ACCOUNTING OF ENERGY’
29
30
•
The first aspect is thermochemistry.
This practical subject is concerned with how
we observe, measure, and predict energy
changes for both physical changes and
chemical reactions.

•
•
•
The second aspect is addressed to a more
fundamental aspect of thermodynamics.
How to use energy changes to tell us
whether or not a given process can occur
under specified conditions
to give predominantly products (or
reactants)
 how to make a process more (or less)
favourable.
FUNDAMMENTAL ASPECT OF
THERMODYNAMICS
THERMOCHEMISTRY
31
32
7
3/15/2015

We can define energy as follows.
 Energy
is the capacity to do
work or to transfer heat.
ENERGY - DEFINITION

Energy can take many forms:

electrical energy,

radiant energy (light),

nuclear energy,

chemical energy.
ENERGY FORMS
33
More common we classify energy into two
general types: kinetic and potential.


Q.
The kinetics energy of solid body with the
mass of 5 kg which moved with speed 8
m s-1 is equal:
a) 40 kg m s-1; b) 320 J; c) 160 J
d) 160 kg m2s-2
Kinetic energy is the energy of motion.
The kinetic energy of an object is equal to
one half its mass, m, times the square of
its velocity, v.

Ek 
34
1
 m  v2
2
The heavier a hammer is and the more rapidly it moves,
the greater its kinetic energy and the more work it can
accomplish.
KINETIC ENERGY
Kinetics energy in questions
35
WM_ THERMODYNAMICS
36
8
3/15/2015
Potential energy (EP) is the energy that
a system possesses by virtue of its
position or composition.
The work that we do to lift an object is
stored in the object as energy.


Q. When a bucket with 10 kg of water is
picked up at the height of 1 m the
potential energy is as follows:
a)
Ep= m g h
10 kg m; b) 100 J; c) 100 kg m2s-2;
d) 100 kg
Where: m – mass (kg); h – body movement (change of
height, m); g – gravitational acceleration, (10 m s-2 )
POTENTIAL ENERGY
Potential energy in question
37
38

Q. What type of energy does a stationary
pencil contain? falling pencil?
If we drop a hammer, its potential energy
is converted into kinetic energy as it falls,
and it could do work on something it
hits—for example, drive a nail or break a
piece of glass.
EXAMPLE: EpEk
39
WM_ THERMODYNAMICS
40
9
3/15/2015

Similarly, an electron in an atom has
potential energy because of the
electrostatic force on it that is due to the
positively charged nucleus and the other
electrons in that atom and surrounding
atoms.
Ep in ATOM

The atomic or molecular level, we can
think of each of these as either kinetic or
potential energy.

The chemical energy in a fuel or food
comes from potential energy stored in
atoms due to their arrangements in the
molecules.
ATOMIC LEVEL
41
Many forms of energy can be
interconverted and that in chemical
processes, chemical energy is converted
to heat energy or vice versa.
The amount of heat a process uses
(endothermic) or gives off (exothermic)
can tell us a great deal about that process.
 For this reason it is important for us
to be able to measure the intensity of
heat.
42

This stored chemical energy can be
released when compounds undergo
chemical changes, such as those that
occur in combustion and metabolism.
 Reactions that release energy in the form
of heat are called exothermic reactions.

INTENSITY OF HEAT
EXOTHERMIC REACTIONS
43
44
10
3/15/2015

Hydrocarbons—including methane, the
main component of natural gas, and
octane, a minor component of gasoline
undergo combustion with an excess of O2
to yield CO2 and H2O.

CH4 + 2O2

This reaction releases heat energy, Q.

Temperature measures the intensity of heat,
the “hotness” or “coldness” of a body.

A piece of metal at 100°C feels hot to the touch,
whereas an ice cube at 0°C feels cold.

Why? Because the temperature of the metal is
higher, and that of the ice cube lower, than body
temperature.

Heat is a form of energy that always flows
spontaneously from a hotter body to a colder
body—never in the reverse direction.
CO2 + 2H2O + Q
EXOTHERMIC REACTIONS
WM_ THERMODYNAMICS
TEMPERATURE vs. HEAT
45
EQUATION OF REACTION
WM_ THERMODYNAMICS

In such reactions, the total energy of the
products is lower than that of the
reactants by the amount of energy
released, most of which is heat.

Some initial activation (e.g., by heat)is
needed to get these reactions started.
This amount of energy is called activation
energy.
46
ACTIVATION Energy
52
WM_ THERMODYNAMICS
53
11
3/15/2015
Fig. 1

The amount of heat shown in such an
equation always refers to the reaction
for the number of moles of reactants and
products specified by the coefficients.
THERMOTERMIC REACTION
WM_ THERMODYNAMICS

54
The difference between the potential
energy of the reactants—one mole of
CH4(g) and two moles of O2(g)—and that
of the products—one mole of CO2(g) and
two moles of H2O(l)—is the amount of
heat evolved in this exothermic reaction
at constant pressure.
For this reaction, it is 890 kJ/mol.
WM_ THERMODYNAMICS
55
A process that absorbs energy from its
surroundings is called endothermic.
 For this type of reaction, the final level of
energy is higher than the initial level.

ENDOTHERMIC REACTION
56
WM_ THERMODYNAMICS
57
12
3/15/2015
When solid hydrated barium
hydroxide, Ba(OH)2 8H2O, and excess
solid ammonium nitrate, NH4NO3, are
mixed, a reaction occurs.
Fig. An endothermic process.
When solid hydrated barium hydroxide, Ba(OH)2 ∙ 8H2O, and excess
solid ammonium nitrate, NH4NO3, are mixed, a reaction occurs.
Profile of endothermic reaction
WM_ THERMODYNAMICS
Energy absorbed or
released
Relative Energy of
reactants & products
Sign of H
Exothermic
Reaction
Endothermic
Reaction
Energy is released.
It is a product
of the reaction.
Reaction vessel
becomes warmer.
Temperature inside
reaction vessel
increases.
Energy is absorbed.
It is a reactant of the
reaction.
Reaction vessel
becomes cooler.
Temperature inside
reaction vessel
decreases.
Energy of the
reactants is greater
than the energy of
the products
H(reactants) > H(products)
H = H(products) H(reactants)
= negative (-ve)
Energy of the
reactants is less than
the energy of the
products
H(reactants) < H(products)
H = H(products) H(reactants)
= positive (+ve)
The dissolution process is very
endothermic. If the flask is placed
on wet a wooden block, the water
freezes and attaches the block to
the flask
58
WM_ THERMODYNAMICS
Endothermic
Energy
Profile
60
59
Exothermic
Energy of reactants
Energy of reactants (NH3)
(N2 & H2) is greater
is less than the energy of
than the energy of
the products (N2 & H2).
the products (NH3).
Energy is absorbed
Energy is released
.
.
61
13
3/15/2015
Exothermic processes
Endothermic processes
making ice cubes
melting ice cubes
formation of snow in clouds
conversion of frost to water vapour
condensation of rain from water
vapor
evaporation of water
mixing sodium sulfite and bleach
baking bread
rusting iron
cooking an egg
burning sugar
producing sugar by photosynthesis
mixing water and strong acids
mixing water and ammonium nitrate
mixing water with an anhydrous salt
making an anhydrous salt from a
hydrate
crystallizing liquid salts (as in
sodium acetate in chemical
handwarmers)
melting solid salts
There are two ways of looking at what happens
to the enthalpy:
If the reaction is exothermic the products
have minimum enthalpy and the formation of
products (move toward the right) is favourable
If the reaction is endothermic the reactants
have minimum enthalpy and the formation of
products (move toward the right) is
unfavourable.
In this case the formation of reactants (move
toward the left) is favourable.
Author: Fred Senese [email protected];
http://antoine.frostburg.edu/chem/senese/101/thermo/faq/exothermic-endothermicexamples.shtml
62
63
Q. Which of the following is an endothermic
Q. Chemical reactions that absorb heat
energy are called __________ .
a. exothermic
b. eltothermic
c. endothermic
reaction?
a) Burning propane in a gas grill
b) Photosynthesis
c) Baking bread
Q. Electrolysis requires energy to make it
work. This means it is...
a) an endothermic reaction
b) an exothermic reaction
c) an eltothermic reaction
d) a chemical reaction
d) The mixture bubbled vigorously and the
temperature dropped by 40°C
e) A red glow spread through the mixture and the
temperature increased by 10°C
f) Electrolysis of water
WM_ THERMODYNAMICS
64
WM_ THERMODYNAMICS
65
14
3/15/2015


Some important ideas about energy are
summarized in the First Law of
Thermodynamics.
Energy is neither created nor destroyed in
ordinary chemical reactions and physical
changes.



The substances involved in the chemical
and physical changes that we are
studying are called the system.

Everything in the system’s
environment constitutes its
surroundings.

The universe is the system plus its
surroundings.
THE UNIVERSE, SYSTEM,
SURROUNDINGS
FIRST LAW OF THERMODYNAMICS
WM_ THERMODYNAMICS

66
The system may be thought of as the part
of the universe under investigation.
WM_ THERMODYNAMICS
67
Q. The first law of thermodynamics
states that energy is
The First Law of Thermodynamics tells us
that energy is neither created nor
destroyed.
a. increased during any process
b. decreased during any process
c. conserved during any process
Energy is only transferred between the
system and its surroundings.
FIRST LAW OF THERMODYNAMICS
WM_ THERMODYNAMICS
68
WM_ THERMODYNAMICS
69
15
3/15/2015
 The
thermodynamic state of a
system is defined by a set of
conditions that completely
specifies all the properties of the
system.
This set commonly includes:
the temperature,
 pressure,
 composition (identity and number of
moles of each component),
 physical state (gas, liquid, or solid) of
each part of the system.


THERMODYNAMIC STATE
WM_ THERMODYNAMICS
THERMODYNAMIC STATE
70
WM_ THERMODYNAMICS
71
The properties of a system—such as P, V,
T—are called state functions.
 The value of a state function depends
only on the state of the system and not
on the way in which the system came to
be in that state.
 A change in a state function describes a
difference between the two states.

Ice
liquid water
Changes in physical
state of water due
to temperature
changes

STATE FUNCTIONS
Steam
WM_ THERMODYNAMICS
72
73
16
3/15/2015
Indirect

It is independent of the process or
pathway by which the change occurs.
1
3
Indirect
Direct
2
74
75

For instance, consider a sample of one mole
of pure liquid water at 30°C and 1 atm
pressure.


It does not matter whether:
(1) the cooling took place directly (either
slowly or rapidly) from 30°C to 22°C,

If at some later time the temperature of the
sample is 22°C at the same pressure, then it
is in a different thermodynamic state.

or (2) the sample was first heated to 36°C,
then cooled to 10°C, and finally warmed to
30°C, then cooled to 22oC
We can tell that the net temperature change
is 8°C.


or (3) any other conceivable path was
followed from the initial state to the final
state.
STATE FUNCTIONS
STATE FUNCTION
WM_ THERMODYNAMICS
76
WM_ THERMODYNAMICS
77
17
3/15/2015


The change in other properties (e.g., the
pressure) of the sample is likewise
independent of path.

The most important use of state functions in
thermodynamics is to describe changes.
We describe the difference in any quantity, X,
as
X  X FINAL  X INITIAL


STATE FUNCTION
When X increases, the final value is greater
than the initial value, so ΔX is positive;
a decrease in X makes ΔX a negative value.
STATE FUNCTIONS
WM_ THERMODYNAMICS

You can consider a state function as
analogous to a bank account.

With a bank account, at any time you can
measure the amount of money in your
account (your balance) in convenient
terms—dollars and cents, euros and
eurocents, złoty and grosz.
78
COMPARISON
WM_ THERMODYNAMICS

Changes in this balance can occur for
several reasons, such as deposit of your
paycheck, using of VISA cart, or service
charges assessed by the bank.

In our analogy these transactions are not
state functions, but they do cause
changes in the state function (the balance
in the account).
79
ACCOUNT BALANCE
WM_ THERMODYNAMICS
80
WM_ THERMODYNAMICS
81
18
3/15/2015

- 25 $
-50 $
Δ = -150 $
Similarly, we shall see that the energy of
a system is a state function that can be
changed— for instance, by an energy
“deposit” of heat absorbed or work done
on the system, or by an energy
“withdrawal” of heat given off or work
done by the system.
-75 $
ENERGY OF THE SYSTEM
WM_ THERMODYNAMICS

We can describe differences between
levels of a state function, regardless of
where the zero level is located.

In the case of a bank balance, the
“natural” zero level is obviously the point
at which we open the account, before any
deposits or withdrawals.
82
ZERO LEVEL
WM_ THERMODYNAMICS

In contrast, the zero levels on most
temperature scales are set arbitrarily.

When we say that the temperature of an
ice–water mixture is “zero degrees
Celsius,” we are not saying that the
mixture contains no temperature!
85
ZERO LEVEL
WM_ THERMODYNAMICS
86
WM_ THERMODYNAMICS
87
19
3/15/2015

We have simply chosen to describe this
point on the temperature scale by the
number zero; conditions of higher
temperature are described by positive
temperature values, and those of lower
temperature have negative values, “below
zero.”
ARBITRARY SCALES
The phrase “15 degrees cooler” has the
same meaning anywhere on the scale.

Many of the scales that we use in
thermodynamics are arbitrarily defined in
this way.

Arbitrary scales are useful when we are
interested only in changes in the quantity
being described.
ARBITRARY SCALES
WM_ THERMODYNAMICS


88
Q. What is the change energy of the
WM_ THERMODYNAMICS

Any property of a system that depends only
on the values of its state functions is also a
state function.

For instance, the volume of a given sample of
water depends only on temperature,
pressure, and physical state; volume is a
state function.

We shall encounter other thermodynamic
state functions.
sausage after heating, if original energy is
4 kJ and 20 kJ is added to it?
a) 16 kJ; b) 4 kJ; c) 20 kJ; d) 24 kJ
89
STATE FUNCTION
WM_ THERMODYNAMICS
90
WM_ THERMODYNAMICS
91
20
3/15/2015
Most chemical reactions and physical
changes occur at constant (usually
atmospheric) pressure.
 The quantity of heat transferred into or
out of a system as it undergoes a
chemical or physical change at constant
pressure, qp, is defined as the enthalpy
change, H, of the process.

ENTHALPY

An enthalpy change is sometimes loosely
referred to as a heat change or a heat of
reaction.

The enthalpy change is equal to the
enthalpy or “heat content,” H, of the
substances produced minus the enthalpy
of the substances consumed.
ENTHALPY - HEAT
WM_ THERMODYNAMICS
92
WM_ THERMODYNAMICS
93
Calculate the reaction of decomposition of
ammonium nitrate. The reaction is
Δ H = Hfinal - Hinitial
and the enthalpies of the three compounds are given in
Table 1.
Δ H = Hsubstances produced – Hsubstance consumed
CHANGE OF ENTHALPY
WM_ THERMODYNAMICS
94
WM_ THERMODYNAMICS
95
21
3/15/2015
If you reverse the previous reaction,
the sign of the enthalpy of the reaction is reversed:
Δ H = +36 kJ
http://www.cliffsnotes.com/sciences/chemistry/chemistry/thermodynamics/e
nthalpy
96
97
WM_ THERMODYNAMICS
99
Standard enthalpies of formation
•Calculate the enthalpy change for the following
reaction and classify it as exothermic or
endothermic.
WM_ THERMODYNAMICS
WM_ THERMODYNAMICS
98
Compound
ΔH0
MgCl2 (S)
-642 kJ/mol
H2O (l)
-286 kJ/mol
MgO (S)
-602 kJ/mol
HCl (g)
-92 kJ/mol
22
3/15/2015

It is impossible to know the absolute
enthalpy (heat content) of a system.

Enthalpy is a state function, however, and
it is the change in enthalpy in which we
are interested; this can be measured for
many processes.
We can determine the energy change
associated with a chemical or physical
process by using an experimental
technique called calorimetry.
 This technique is based on observing
the temperature change when a system
absorbs or releases energy in the form of
heat.
 The temperature change is caused by the
absorption or release of heat by the
chemical or physical process under study.

CALORIMETRY
EHTHALPY AS STATE FUNCTION
WM_ THERMODYNAMICS


100
The experiment is carried out in a device
called a calorimeter, in which the
temperature change of a known amount
of substance (often water) of known
specific heat is measured.
The specific heat is the amount of heat
per unit mass required to raise the
temperature by one degree Celsius.
Specific heat
WM_ THERMODYNAMICS

Chemical reactions and physical changes occur
with either the simultaneous evolution of heat
(exothermic processes) or the absorption of
heat (endothermic processes).

The amount of heat transferred in a
process is usually expressed in joules or
in calories.

The SI unit of energy and work is the
joule (J), which is defined as 1 kg m2/s2.
101
SPECIFIC HEAT
WM_ THERMODYNAMICS
102
WM_ THERMODYNAMICS
103
23
3/15/2015

You may find it more convenient to think in
terms of the amount of heat required to raise the
temperature of one gram of water from 14.5°C
to 15.5°C, which is 4.184 J.

One calorie is defined as exactly 4.184 J.

The so-called “large calorie,” used to indicate the
energy content of foods, is really one kilocalorie,
that is, 1000 calories or 1 kcal.


The specific heat of a substance is the
amount of heat required to raise the
temperature of one gram of the substance
one degree Celsius (also one kelvin) with
no change in phase.

Changes in phase (physical state) absorb
or liberate relatively large amounts of
energy (see Figure –NEXT SLIDE).
We shall do most calculations in joules.
SPECIFIC HEAT
SPECIFIC HEAT
WM_ THERMODYNAMICS
104
WM_ THERMODYNAMICS

The specific heat of each substance,
a physical property, is different for the
solid, liquid, and gaseous phases of the
substance.

For example, the specific heat of:
ice is 2.09 J/g °C near 0°C;
liquid water it is 4.18 J/g °C;
steam it is 2.03 J/g °C near 100°C.
105
SPECIFIC HEAT
106
WM_ THERMODYNAMICS
107
24
3/15/2015
Specific heats and molar heat capacities for
various substances at 293 K (20oC)
SPECIFIC HEAT [c]
WM_ THERMODYNAMICS
108
Q. How many heat is needed to heat up 10 g of liquid
Substance
c in J/gm K
Aluminum
Copper
Gold
Lead
Silver
Zinc
Mercury
Alcohol(ethyl)
Water
Ice (-10 C)
0.900
0.386
0.126
0.128
0.233
0.387
0.140
2.4
4.186
2.05
Molar C
J/mol K
24.3
24.5
25.6
26.4
24.9
25.2
28.3
111
75.2
36.9
SPECIFIC HEAT
WM_ THERMODYNAMICS
109
Answer:
Q = m c ΔT
water from 10oC to 40oC.
[c of liquid water is 4.18 J g-1 oC-1].
Q = (10 g) . (4.18 J g-1 oC-1) ( 40oC – 10oC)
a) 418 J; b) 1254 J; c) 1672 J
Q = 10 . 4.18 . 30 = 1254 J
WM_ THERMODYNAMICS
110
WM_ THERMODYNAMICS
111
25
3/15/2015
Q.
A 385 grams chunk of iron is heated to
97.5oC. Then it is immersed in 247 gram of
water originally at 20.7oC. When thermal
equilibrium has been reached, the water and
iron are both at 31.6oC. Calculate the specific
heat of iron.
WM_ THERMODYNAMICS
112

A hot object, such as a heated piece of
metal (a), is placed into cooler water.

Heat is transferred from the hotter metal
bar to the cooler water until the two reach
the same temperature (b).

WM_ THERMODYNAMICS
113
Solution:
The number of heat gained by water from
temperature 20.7oC to 31.6o C =the amount of heat
which is lost by the iron.
Qwater = (247 g). (4.18 J g-1 oC-1) (31.6-20.7oC)
We say that they are then at thermal
equilibrium.
Qwater = (247) . (4.18). (10.9) = 11253.8 J
THERMAL EQUILIBRIUM
WM_ THERMODYNAMICS
114
WM_ THERMODYNAMICS
115
26
3/15/2015

End of part 1
WM_ THERMODYNAMICS
116
27