Transcript Process

LECTURE № 6
THEME: Theoretic bases of
bioenergetics. Electrochemistry.
associate. prof. Yevheniy. B. Dmukhalska
Plan
1.The basic concepts of
thermodynamics
2. First law of thermodynamics. Heat
(Q) and Work ( W)
3. Secohd law of thermodynamics.
Entropy (S)
4. Electrochemistry.
The branch of science which deals
with energy changes in physical and
chemical processes is called
thermodynamics
Some common terms which are
frequently used in the discussion of
thermodynamics are:
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Common terms of
thermodynamics
System
Parameter
Condition
(state)
Process
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System is a specified part of the
universe which is under observation
The remaining portion of the universe
which is not a part of the system is
called the surroundings
The system is separated by real or
imaginary boundaries.
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Classification of the thermodynamics
systems according to a structure
homogeneous
KNO3
heterogeneous
KNO3
PbI2↓
Types of Systems
ISOLATED
A system can neither
exchange matter nor
energy with the
surroundings
CLOSE
A system which can
exchange energy
but no mass with its
surroundings
OPEN
A
system
can
exchange both matter
and energy with the
surroundings.
Parameters
Extensive
(m, V, U, H, G, S, c)
The properties of the
system whose value
depends upon the
amount of
substance present
in the system
Intensive
(p, T, C, viscosity, surface
tension, vapour pressure)
The properties of the
system whose value does
not depend upon the
amount of substance
present in the system
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Process is the change of all or individual parameters
of the system during the length of time (the period of
time)
Classification
of a process according to the constant parameter of a
system are:
 Isothermic process – temperature is constant,
T=const
 Isochoric process – volume is constant V = const.
 Isobaric process – pressure of the system is
constant, p = const
 Adiabatic process – the system is completely
isolated from the surroundings. For an adiabatic
(Q=0) system of constant mass, ▲U=W
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Classification
of a process according to the releasing energy

Exothermic process is a process that releases
energy as heat into its surroundings. We say that in
an exothermic process energy is transferred ‘as heat’
to the surroundings. For example: a reaction of
neutralization (acid + basic).

Endothermic process is a process in which
energy is acquired from its surroundings as heat.
Energy is transferred ‘as heat’ from the
surroundings into the system. For example: the
vaporization of water
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Classification
of a process according to the direction of reaction
 Reversible
process is a process in
which the direction may be reversed at
any stage by merely a small change in a
variable like temperature, pressure, etc.
 Irreversible process is a process which
is not reversible. All natural process are
irreversible
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State of a system means the condition of the
system, which is described in terms of certain
observable (measurable) properties such as
temperature (T), pressure (p), volume (V)
State function (thermodynamic function)
 Internal energy U [J/mol]
 Enthalpy H [kJ/mol] or [kJ]
 Entropy S [J/mol K] or [J/K]
 Gibbs energy G [J/mol] or [J]
ΔU = U(products) – U(reactants)
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State function depends only upon the
initial and final state of the system
and not on the path by which the
change from initial to final state is
brought about.
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Internal energy U
It is the sum of different types of
energies associated with atoms and
molecules such as electronic
energy, nuclear energy, chemical
bond energy and all type of the
internal energy except potential
and kinetic energies.
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Heat (Q) is a form of energy which the
system can exchange with the
surroundings. If they are at different
temperatures, the heat flows from higher
temperature to lower temperature. Heat is
expressed as Q.
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Work (W) is said to be performed if the
point of application of force is displaced
in the direction of the force. It is equal to
the distance through which the force
acts.
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Enthalpy H
Chemical reactions are generally
carried out at constant pressure. ΔU
gives the change in internal energy at
constant volume. To express the energy
changes at constant pressure, a new
term called enthalpy was used.
Enthalpy cannot be directly measured,
but changes in it can be.
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Enthalpy H
A thermodynamic function of a
system, equivalent to the sum of
the internal energy of the system
plus the product of its volume
multiplied by the pressure
exerted on it by its surroundings.
▲H = ▲U + p▲V
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Heat absorbed by the system = H positive (Q negative
Heat evolved by the system = H negative (Q positive)
The signs of W or Q are related to the internal energy
change.
The meaning of the state functions in the
thermodynamic processes
Exothermic process
 Qv > 0, ▲U < 0
 Qp > 0, ▲H < 0
Endothermic process
 Qv < 0, ▲U > 0
 Qp < 0, ▲H > 0
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The first law of thermodynamics
Matter/energy
may be altered (converted), but not
created (from nothingness) nor destroyed (reduced to
nothingness).
The First Law teaches that matter/energy cannot
spring forth from nothing without cause, nor can it
simply vanish.
Energy can neither be created nor destroyed
although it may be converted from one form to
another.
The given heat for the system spends on the change
of the internal energy and producing the work:
Q = ▲U + W
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Bomb calorimeter for the determination of change in
internal energy
The process is carried out at constant volume, i.e., ΔV=0, then the product PΔV is also zero.
Thus, ΔU=Qv
The subscript v in Qv denotes that volume is kept constant.
Thus, the change in internal energy is equal to heat absorbed or evolved at constant
temperature and constant volume
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Thermochemistry
The study of the energy transferred as heat during
the course of chemical reactions.
Thermochemical reactions:
H2(g) + Cl2(g) = 2HCl; ▲ H = -184,6 kJ
1/2 H2(g) + 1/2 Cl2(g) = HCl; ▲ H = -92,3 kJ/mol
▲ H is calculated for 1 mole of product
▲H = ▲U + p▲V
▲H = ▲U + ▲nRT
Energy change at constant P = Energy change at
constant V + Change in the number of geseous
moles * RT
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The Hess’s law
Initial reactants
Н2
Н1
Н3
Н4
The products
of reaction
Н1 = Н2 + Н3 + Н4
If the volume or pressure are constant the
total amount of evolved or absorbed heat
depends only on the nature of the initial
reactants and the final products and doesn’t
depend on the passing way of reaction.
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Conclusions from the Hess law
1.Нc298(the standard enthalpy of combustion) =Нf298(the standard enthalpy of formation)
2.Н(formation)= ΣnНf298(products) - ΣnНf298(reactants)
3.Н(combustion) = ΣnНс298(reactants) - ΣnНс298(products)
4.For elementary substances Н0298 = 0
4.Н3=Н1-Н2
5.Н1=Н3-Н2
1
1
2
3
2
3
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Correlation U і Н:
H  U  pV  U  RT
If υ0, so НU:
СаО + СО2 → СаСО3
If υ0, so НU:
Na + H2O → NaOH + H2
If υ=0, so Н=U:
H2 + Cl2 → 2HCl
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Second law of thermodynamics

Second Law of Thermodynamics
(refrigerator): It is not possible for heat to
flow from a colder body to a warmer body
without any work having been done to
accomplish this flow.
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The amount of molecular randomness in a system
is called the system’s entropy (S).
Entropy is a measure of randomness or disorder
of the system
Free energy and free energy
change
The maximum amount of energy available to a
system during a process that can be converted
into useful work
It’s denoted by symbol G and is given by
▲G = ▲H - T ▲S
where ▲G is the change of Gibbs energy (free energy)
This equation is called Gibbs equation and is very useful
in predicting the spontaneity of a process.
N.B. Gibbs equation exists at constant temperature and
pressure
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1) Spontaneous (irreversible) process :
▲ G < 0, ▲S > 0, ▲H < 0
2) Unspontaneous (reversible) process :
▲ G > 0, ▲S < 0, ▲H > 0
3) Equilibrium state
▲G=0
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THIRD LAW OF THERMODYNAMICS:
The third law of thermodynamics, formulated
by Walter Nernst and also known as the
Nernst heat theorem, states that if one could
reach absolute zero, all bodies would have
the same entropy.
Temperature approaches absolute zero (0 K),
the entropy of a system approaches a
constant (and minimum) value. The entropy of
a perfect crystalline state is zero at 0 K.
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Electrochemistry

The branch of science, which deals
with the study oxidation-reduction
reaction
to
produce
the
interconversion of chemical and
electricl energy. of transition chemical
energy to electrical energy is known as
electrochemistry.
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Nernst’s equation
Мn++nе = М
Then the Nernst eqn. is applied as follows:
E = E0 – (RT/ nF) ln ([M]/ [Mn+])
where Е = electrode potential under given
concentration of Мn+ ions and temperature Т;
Е0 – standard electrode potential;
R – gas constant, R = 8.315 J/K .mol;
Т – temperature in К;
F – Faradays constant, F = 96,485 С /mol;
n –
number of electrons involved in the
electrode reaction.
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Standard (normal)
hydrogen electrode
Pt, Н2 (g)/Н+ (Concentration)
H2 = 2H+ + 2е2H+ + 2е- = H2
E = E0 – (RT/ 2F) ln (pH2/ [H+]2), E0H+/H2 = 0V.
In the standard hydrogen gas electrode, hydrogen at
atmospheric pressure is passed into 1 М НС1 in
which foil of the platinized platinum remains
immersed through which inflow or outflow of
electrons takes place.
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Since а cathode reaction is а reduction, the
potential produced at such an electrode is
called а reduction potential. Similarly, the
potential produced at an anode is called an
oxidation potential. These are known as
standard
reduction
potentials
or
standard electrode potentials. They are
usually tabulated for 25 С.
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Types of electrodes
1. Metal-metal
ion electrodes
2. Gas-ion electrodes
3. Metal-insoluble salt-anion electrodes
4. Inert "oxidation-reduction" electrodes
5. Membrane electrodes
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Electrochemical cells allow
measurement and control of a redox
reaction.
Electrodes of the first kind.
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An electrode of the first kind is а piece of pure
metal that is in direct equilibrium with the
cation of the metal. А single reaction is involved.
For example, the equilibrium between а metal Х
and its cation Х+n is:
Х+n + ne- = X (s)
for which
Еnd = Е0X+n
0.0592
1
0.0592
– -------- log ---- = Е0X+n + ---------- log aX+n
n
aX+n
n
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The metal - metal ion electrode consists of а
metal in contact with its ions in solution. An
example: silver metal immersed in а solution of
silver nitrate
As a cathode: the diagram: Ag+(aq)  Ag(s)
half-reaction equation is: Ag+ (aq) + e-Ag(s)
as an anode: the diagram: Ag(s)  Ag+(aq)
half-reaction equation is: Ag(s)  Ag+(aq) + еNernst’s equation:
E = E0 – (RT/ nF) ln ([Ag]/ [Agn+])
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Electrodes of the Second type.

Metals not only serve as indicator electrodes for
their own cations but also respond to the
concentration of anions that form sparingly soluble
precipitates or stable complexes with such cations.
AgCl + e- = Ag (s) + Cl

E0AgCl = 0.222 V
The Nernst expression for this process is:
EAgCl = E0AgCl – 0.0592 log [Cl-] = 0.222 + 0.0592
pCl
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In the metal-insoluble salt-anion
electrode, а metal is in contact with one of
its insoluble salts and also with а solution
containing the anion of the salt. An example
is the so-called silver - silver chloride
electrode, written as а cathode as:
Cl- (aq)  AgCl(s)  Ag(s)
for which the cathode half-reaction is:
AgCl (s) + е-  Ag(s) + Cl- (aq)
Nernst’s equation:
E = E0 – (RT/ 1F) ln ([Ag] [Cl-]/ [AgCl])
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An inert oxidation-reduction electrode consists
of а strip, wire, or rod of an inert materiel, say,
platinum, in contact with а solution, which contains
ions of а substance is two different oxidation states.
Thus, for the ferric - ferrous ion electrode
functioning as а cathode,
Fe3+, Fe 2+(aq)  Pt(s)
the iron(III), or ferric, ion, Fe+3(aq), is reduced to
the iron(II), or ferrous, ion, Fe+2(aq):
Fe+3(aq) + е- Fe+2(aq)
Nernst’s equation:
E = E0 – (RT/ 1F) ln ([Fe+2]/ [Fe+3])
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а membrane electrode - the glass
electrode.
 This can be depicted as:
 Pt(s) Ag(s)  AgC1(s)  HC1(aq,1M) 
glass 
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Cell can be depicted as:
reference electrode  salt bridge  analyte solution  indicator
electrode

Ecell = Eind + Eref + Ej
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1.Glass electrode – indicator electrode;
diagram which is:
Ag(s)  AgC1(s)  HC1(aq,1M)  glass 
 2.Bulb of glass electrode.
 3.Solution of unknown pH.
 4.Silver-silver
chloride electrode electrode;diagram which is:
 Cl- (aq)  AgCl(s)  Ag(s)
5.Amplifying potentiometer.

reference
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Cell potential or EMF of a cell.
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The difference between the electrode
potentials of the two half cell is known
as electromotive force (EMF) of the
cell or cell potential or cell voltage.
The EMF of the cell depends on the
nature of the reactants, concentration
of the solution in the two half cells, and
temperature.
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Reference electrode is electrode
potential which stabile
А hydrogen electrode is seldom used
as а reference electrode for day-today potentiometric measurements
because it is somewhat inconvenient
and is also а fire hazard.
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An ideal indicator electrode responds
rapidly and reproducibly to changes in the
concentration of an analyte ion (or group of
ions). Although no indicator electrode is
absolutely specific in its response, а few are
now available that are remarkably selective.
There are two types of indicator electrodes:
metallic and membrane.
Metallic indicator electrodes:
Electrodes of the first kind.
Electrodes of the Second Kind.
Membrane Electrodes
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The relationship between pH and the voltage of the
hydrogen elect
calomel electrode cell at 250С can be written as
Ecell
Ecalomel
1
pH = --------- - (---------- + ---- log pH2)
0.0592
0.0592
2
Ecell
pH = ---------- = constant
0.0592
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