Download THERMODYNAMICS

THERMODYNAMICS
SPECIFIC LEARNING OBJECTIVE
At the end of the session the student should be
able to explain:
•Energy
•The first law of thermodynamics
•Entropy
•Free energy
Thermodynamic
In Chemistry
System
In Living
System
THERMODYNAMICS
Thermodynamic is the law that
formulated from observation on
conversion of energy from one form
to the other. i.e. transduction
What is the energy?
1. Energy is a much used term, but it
represents a rather abstract
concept.
2. Energy is usually defined as the
capacity to do work.
3. Chemist define work as directed
energy change resulting from a
process
The type of Energy
1. Kinetic energy
2. Radiant energy
3. Thermal energy
4. Chemical energy
5. Potential energy
Definition of type energy
1. Kinetic energy – the energy
produced by a moving object
2. Radiant energy : comes from the
sun (solar energy) and is Earth’s
primary energy source. Solar
energy heats the atmosphere and
Earth’s surface, stimulates the
growth of vegetation through the
process known as photosynthesis,
and influences global climate
patterns.
continued
The Activated Complex form the reaction:
A+B
AB#
P
Definition of type energy
3.Thermal energy is the energy
associated with the random
motion of atoms and molecules.
4. Chemical energy is stored within
the structural units of chemical
substances; its quantity is
determined by the type and
arrangement of atoms in the
substance being considered.
5. Potential energy is energy that is
also available by virtue of an
object’s position.
Conclusion of Energy
• All forms of energy can be
interconverted (at least in principle)
from one form to another
Scientists have concluded that energy
can be neither destroyed nor created.
• Thermodynamic Law
THERMODYNAMICS
Thermodynamic I is the law
of conservation of energy
Thermodynamic I
Work
Thermodynamic II
Work and heat are
not state functions
Thermodynamic III
Heat
• The relationship between chemical energy and
other forms of energy, with examples.
Energy change in chemical reactions
Almost all chemical reactions absorb
or produce (release) energy, generally
In the form of heat.
• Heat is the transfer of thermal energy
between two bodies that are at different
temperatures.
Although ”Heat” itself implies the
transfer of energy, we customarily talk of
”heat absorbed” or ”heat released” when
describing the energy changes that occur
during a process.
Energy changes associated with
chemical reactions
System
Surroundings
SYSTEM AND SURROUNDING
System we mean that the part of
the world we are investigating.
Surrounding we
Three type of systems
be
Open
Close
mean everything else
Surrounding
Isolated
system
System be open
• Two of this examples are the
examples of open system:
1.e.g. in the living organism, which
takes up nutrients, releases the
waste products, and generates work
and heat.
2.An example in body, the body takes
up nutrient, and then release urine
which contains toxin, carbon
dioxide, and so on.
System be closed
Example of close system:
• An example of close system is living
of an microorganism, it was sealed
inside a perfectly insulated box, it
will, together with the box, constitute
a closed system.
HEAT
q, to be the manner of energy transfer that
results from a temperature difference
between the system and its surrounding
Positive and negative
sign of heat
Heat input to a
system is
considered a
positive quantity
Heat evolved by a
system is
considered a
negative quantity.
WORK
w, to be the transfer of energy between the
system of interest and its surroundings as
a result of existence of unbalanced forces
between two.
Positive and negative sign of work
If the energy of the system is
increased by the work, we say
that work is done on the
system by surroundings, and
we take it to be a positive
quantity
if the energy of the system is
decreased by the work, or the
system does work on the
surroundings, or that work is
done by the system, and we
take it to be a negative
quantity
• The effect of work is equivalent to the
raising or lowering of mass in the
surroundings.
• Work is done by the system because the
mass is raised
work is done on the system because the
mass is lowered.
ENERGY
Energy is a state function
It is a property that depends only upon the
state of the system, and not upon how the
system was brought to that state, or upon
the history of the system.
Thermodynamic I study
of conservation of energy
The first law of thermodynamic
•∆U=q+w
which is essentially a statement of the law
of conservation of energy.
Where :
1. The term ∆ U represents the change of internal
energy of the system,
2. q is the thermal energy (heat) added to the
system, and w is the work done on the
system.
The chemical reactions that need energy
1. The Photoelectric Effect  this is mystery
in physics.
Experiments had already demonstrated
that electrons were ejected from the
surface of certain metals exposed to light
of at least a certain minimum frequency.
Einstein suggested that a beam of light is a
stream of particles. These particles of light
are called photons. Using Planck’s
quantum theory of radiation as a starting
point, Einstein deduced that each photon
must possess energy E, given by the
equation :
E = hv
In which v is the frequentcy of light and h is
Planck’s constant
The equation of E = hv
E = hv
E = KE + BE  hv = KE + BE
in which :
* KE is the kinetic energy of the
ejected electron and
* BE is the binding energy of the
electron in the metal
The energies that the electron
in the hydrogen atom
• En = – RH
1
n2
• In which RH, the Rydberg constant,
has the value 2.18 x 10-18 J
• The number n is an integer called
the principal quantum number; it
has the value n = 1, 2, 3,….
Strength of Covalent Bond
The Strength of Covalent Bond is
defined by the amount of energy
needed to break it.
A quantitative measure of stability
of a molecule is its bond
dissociation energy (or bond
energy).
For example:
H2(g)  H(g) + H(g)
H = 436.4 kJ
HCl(g)  H9g) + Cl(g) H = 431.9 kJ
Covalent Bond in Organic Compounds
STRUCTURE
SATURATED:
 Bonding : are
formed by overlap
of two atomic
orbitals, each of
which contains one
electron
UNSATURATED :
 and  bonding.
 bonding


 (pi) bond
ENTHALPY
• The enthalpy of a system, which has
the symbol H, is that of Heat content
(heat of reaction) and is measure of
the change in total bonding energy
during a reaction. It is defined
mathematically as :
H = U + PV; H is a function of state
The standard enthalpy change for any
reaction (∆H0rxn)can determine by
using standard enthalpies of
formation (∆H0f) and Hess’s Law
CONSTANT PRESSURE PROCESSES
• Most processes occur in the open
at one atmosphere pressure.
In these cases, P1 = P2 = P, say,
and
•
∆H = ∆U + P ∆V
Positive and negative sign of Enthalpy
• ∆ H has a negative sign for an
exothermic change (heat is
released), is mean the bonds in
the products are stronger (more
stable) than the bonds in the
reactants.
• ∆ H has a positive sign for an
endothermic change (heat is
absorbed), is mean the bonds in
the products are weaker (less
stable) than the bonds in the
reactants.
HESS’S LAW
The principle of constant heat
summation, often known as Hess’s
Law, is thus seen to lead directly from
the fact that H is a function of state.
Hess’s Law be valid for : ∆r H˚ or ∆f H˚,
˚ = all reactants and products are in
• their standard states.
f = formation standard enthalpies of
formation
Pº = 1 atmosphere, and temperature
25ºC or 298.15°K
•
This idea is immensely
powerful, because it enables
Hº298 values to be determined
for any reaction, as long as
the H of formation are known
for each reactant and product.
• ∆r H = H prod – H react
Example No. 1 :
• Consider the following two chemical
equations.
• 1. C(s) + ½ O2 (g)
CO (g)
∆r H (1) = -110.5 kJ
2. CO (g) + ½ O2 (g)
CO2 (g)
∆r H (2) = -283.0 kJ
How many Joule ∆r H (3) = ….? For below
equation
C (s) + O2 (g)
CO2 (g)
∆r H (3) = ...?
Example No. 2 :
2 P(s) + 3 Cl2(g)
2 P(s) + 5 Cl2(g)
2 PCl3(l) ∆rH (1)= -640 kJ
2 PCl5(s) ∆rH (2)= -887 kJ
Please calculate the value of ∆r H for below
equation
PCl3(l) + Cl2(g)
PCl5(s) ∆r H (3) = .....?
• Please you make the application of Hess’s
Law, consider the use of
solution No. 2:
2 P(s) + 3 Cl2(g)
2 P(s) + 5 Cl2(g)
2 PCl3(l) ∆rH (1)= -640 kJ
2 PCl5(s) ∆rH (2)= -887 kJ
Please calculate the value of ∆r H for below
equation
PCl3(l) + Cl2(g)
PCl5(s) ∆r H (3) = .....?
• Please you make the application of Hess’s
Law, consider the use of
A. – 247 kJ
B. + 247 kJ
C. – 124 kJ
D. + 124 kJ
E. – 1527 kJ.
Spontaneous Changes
The process tends to occur or not
Two driving forces in nature
1. The towards
minimization of
energy is one such
directing influence,
but there is also a
tendency for material
to become more
physically
disorganized.
2. The tendency
for entropy to
increase is
nature’s
second driving
force.
ENTROPY
• The symbol of entropy = S
a thermodynamic function of state
NATURAL OR IRREVERSIBLE PROCESS
• The entropy of system and surroundings
together increases during all natural or
irreversible process;
• ∆ Ssystem + ∆ Ssurrounding = ∆ Suniverse > 0
REVERSIBLE PROCESS
• For reversible process, the total
entropy is unchanged;
• ∆ Ssys + ∆ Ssur = ∆ Suniverse = 0
CYCLIC PROCESSES
• For a cyclic process, a process in which the
final state is the same as the initial state,
∆S = 0
Changes of entropy
with temperature
• ∆ S = S2 – S1 = CP ln T2/T1 (P
constant)
• ∆ S = S2 – S1 = CV ln T2/T1 (V
constant)
Absolute entropy
The third Law of Thermodynamics
All truly perfect crystals at absolute
zero temperature have zero entropy.
FREE ENERGY
Gibbs Free Energy
The Gibbs energy determines the
direction of a Spontaneous
Process for a System at Constant
Pressure and Temperature
G is function of state
Gibbs free energy, G,. It is a function
of state which provides possible or
not a change of any kind will tend to
occur.
The value of ∆ G
• For a favorable reaction, ∆G has a
negative value, meaning that
energy is released to the
surroundings  Exergonic
• For a unfavorable reaction, ∆G has
a positive value, meaning that
energy is absorbed from the
surroundings  Endergonic
REACTION AT CONSTANT
TEMPERATURE & PRESSURE
dG ≤ 0 (constant T and P)
The quantity G is called the Gibbs energy
Value of G in a system
at constant T and P
• The Gibbs energy will decrease as
the result of any spontaneous
processes until the system reaches
equilibrium, where d G = 0.
• The Gibbs free energy is
defined as:
• G = H - TS
RELATIONSHIP BETWEEN THE PROCESSES
WITH GIBB’S FREE ENERGY
Spontaneous processes, that is, those with
negative ∆ G values, are said to be
exergonic; they can be utilized to do work.
Processes that are not spontaneous, those
with positive ∆ G values, are termed
endogonic; they must be driven by the input
of free energy.
Processes at equilibrium, those in which the
forward and backward reactions are exactly
balance, are characterized by ∆ G = 0.
Thermodynamic
In Chemistry
System
In Living
System
What was The Thermodynamic
Studied in Living System?
•
•
•
•
•
•
•
•
Thermodynamic In Living System
1. ∆ H (heat)
2. ∆ S (the extent of disorder of the
system)
3. ∆ G (Gibbs change in free energy that
proportion of the total energy change
in a system, that is available for doing
work)
Thermodynamic In Living System
• Under the conditions of biochemical
reactions,
• 1. ∆ H (heat) is approximately equal
•
to ∆ E, the total change in internal
•
energy of the reaction,
•
∆ G = ∆ H – T ∆S, become:
•
∆ G = ∆ E – T ∆S
What is the difference between chemical
reaction in nonbiologic systems and in
biologic systems?
Nonbiologic systems may utilize heat
energy to perform work, but biologic
systems are essentially isothermic and
use chemical energy to power living
processes.
ATP (Adenosine Triphosphate):
The Primary Energy Carrier)
• Certain bonds in ATP save the energy
released during the oxidation of
carbohydrates, lipids, and proteins.
• The ATP molecules act as energy
carries, and deliver the energy to the
parts of the cell where energy is
needed to power muscle contraction,
biosynthesis, and other cellular work.
Energy source in the body
ATP plays a central
role in the
transference of free
energy from the
exergonic to the
Endergonic
processes.
It serves as a carrier
of chemical energy
between high energy
phosphate donors
and low energy
phosphate acceptors
Structure of ATP
Energy source in the body
ATP consists of
adenine (a purine),
ribose and three
phosphate groups,
out of which the
two terminal
phosphate groups
being anhydride
bonds are the high
energy groups
Structure of ATP
Energy change in chemical reactions
Almost all chemical reactions
absorb or produce (release)
Energy.
The standard free energy, i.e.
∆ G0’ of hydrolysis of ATP
1. ATP + H20  ADP + Pi
∆ G0’ = -7.3 kcal/mol 
used for doing work
2. ATP + H20  AMP + PPi
∆ G0’ = -7.7 kcal/mol
Chemical reaction can use up,
or produce useful energy.
Exergonic reactions produce an energy
output (ΔG = -  means that the process is
not favorable)
Endergonic reactions require an energy
Input (ΔG = +  the criterion for a favorable
process in a nonisolated system, at
constant temperature and pressure)
Biochemical system couple these energy
yielding (exergonic: unstable to stable)
with energy requiring (endergonic: stable
to unstable) to make cellular metabolism
work.
Energy transfer in the body
The working cell
• The chemical reactions within cells
are accompanied by changes in
energy.
• Cells accomplish their tasks by
coupling energy-requiring reactions
with energy-producing reactions
Protein + ATP  Pro-phosphate complex
+ ADP (protein is phosphorylated)
What is the reaction called, if a
reaction between solute and solvent
needs heat?
•
•
•
•
•
A. exergonic reaction
B. endergonic reaction
C. exothermic reaction
D. endothermic reaction
E. kinetic reaction
• The reaction of glucose become to
glucose-6-phosphate as follows:
•
Pi + glucose
glucose-6-P + H2O
•
ΔG0 = +13.8 (kJ.mol-1)
•
ATP + H2O
ADP + Pi
•
ΔG0 = -30.5 (kJ.mol-1)
ATP + glucose
ADP + glucose-6-P
•
ΔG0 = -16.7 (kJ.mol-1)
What is the reaction above called?
A. exergonic reaction
B. endergonic reaction
C. exothermic reaction
D. endothermic reaction
E. kinetic reaction
Energy transfer in the body
• Phosphoryl-transfer Reactions
• R1-O-PO32- + R2-OH
R1-OH + R2-O-PO32• Are of enormous metabolic significance.
Some of the most important reactions of
this type involve the synthesis and
hydrolysis of ATP:
•
ATP + H2O
ADP + Pi
•
ATP + H2O
AMP + PPi
• For examples: next slide
Continuation: Energy transfer in the body
• The metabolism of glucose is its
conversion to glucose-6-phosphate :
• Endergonic half-reaction 1:
• Pi + glucose
glucose-6-P + H2O
ΔG0 = +13.8 (kJ.mol-1)
• Exergonic half reaction 2:
• ATP + H2O
ADP + Pi
•
ΔG0 = - 30.5 (kJ.mol-1)
ATP + glucose
ADP +glucose-6-P
• ΔG0 = -16.7 (kJ.mol-1)  Exergonic
Active transport: An energy-requiring
process involving the movement of
substances across a membrane
•
High muscle activity:
1. relaxed muscle + ATP
Contracted muscle +
ADP + Pi
2. ADP + phosphocreatinine
ATP + creatine
•
Low muscle activity:
1. Catabolic energy + ADP +Pi
ATP
2. ATP + Creatine
ADP + Phosphocreatine
THERMODYNAMICS OF LIVE
1. Living organism are open system and
therefore can never be at equilibrium.
2. The free energy from this process is
used to do work and to produce the
high degree of organization
characteristic of life.
3. Living system must maintain a
nonequilibrium state for several
reasons.
For example: the ATP-generating
consumption of glucose.
Alterations in Body Temperature
• FEVER AND HYPERTHERMIA:
• Fever:
• Is an elevation of body temperature
above the normal circadian range as
the result of a change in the
thermoregulatory center located in
the anterior hypothalamus.
• A normal body temperature is
ordinarily maintained, despite
environmental variations, through
the ability of the thermoregulatory
center to balance heat oproduction
by tissues (notably, muscles and the
liver) with heat dissipation.
Continuation:
• With fever, the balance is shifted to
increase the core temperature.
• Hyperthermia:
• Is an elevation of body temperature
above the hypothalamic set point due to
insufficient heat dissipation (e.g. in
association with exercise perspirationinhibiting drugs, or a hot environment) 
the topic in Lab activity (salicylat
poisoning)
Summary References:
1. Warn, J.R.W., 1999, Concise Chemical
Thermodynamics, Second Edition,
Stanley Thornes Ltd., United Kingdom.
2. McQuarrie, D.A., Simon, J.D., 1997,
Physical Chemistry a Molecular
Approach, University Science Books,
Sausalito.