Transcript 03. The Theoretic bases of bioenergetics

LECTURE № 3
THEME: Theoretic bases of
bioenergetics. Chemical kinetics and
biological processes.
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:
3
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
Open
Close
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.
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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.


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.
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A chemical kinetics of
biological processes
Definition

Chemical kinetics is that branch of
chemistry, which deals with the study
of the the rates of chemical reactions,
the factors affecting the rates of the
reactions and the mechanism by which
the reactions proceed.
Classification of chemical reactions
according to the quantity of stages (phases).
Simple
reactions
go in a one
elementary
chemical act
Complex(compound)
reactions go in several
stages
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Chain reactions
Primary process – chain initiating step (stage):
h
Cl2 === 2С1.
chlorine molecule absorbs one quantum of light (h)
and dissociates to give Cl atoms.
Secondary process – chain propagating step (stage):
1. Cl. + Н2 = HCl + H.
2. H. + Cl2 = HCl + Cl.
Third process – chain terminating step (stage):
Сl. + Cl. = Сl2
Parallel reactin

For example: Phenol with nitric acid, so have been
formed ortho-, pair- and meta-nitrophenol.
Series the reactions
are reaction which products firs
step (stage) are reactants for second step (stage):
A BCD….
C18H32O16 + HOH = C12H22O11
+ C6H12O6
Raffinose
disaccharide
monosaccharide
C12H22O11 + HOH = C6H12O6 +
C6H12O6
Monosaccharides
Reversible the reactions
reactions which are flowing past in
two parties: the forward reaction conducts to formation reaction product
and reverse reaction - decomposing
reaction
product
on
mother
substances.
k1
A + B + C = A1 + B1 + C1
k2

Classification of the chemical reaction
according to the quantity of the reacting
phases
• Homogeneous:
N2 (g) + H2 (g) → NH3 (g)
• Heterogeneous:
Mg (s) + HCl (l) → MgCl2 (l) + H2 g)
• Topochemical ( in the hard phase)
Classification of chemical reactions according
to the molecularity
The molecularity of an elementary reaction is the
number of molecules coming together to react
• Unimolecular (monomolecular):
Н2СО3 → Н2О + СО2
• Bimolecular:
CuO + CO → Cu + CO2
• Termolecular:
2 NO + O2= 2 NO2
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The rate of a reaction is the speed
at which a reaction happens.
The Kinetic curves of the rate
reaction’ mean value
Сonc
entra
tion
α
t
Veritable (true) rates:
dC = v·dt;
v = tgα
FACTORS WHICH INFLUENCE RATES OF
CHEMICAL REACTIONS
1. Concentration of the reacting
species.
2. Temperature of the system.
3. Nature of the reactants and
products.
4. Presence of a catalyst.
5. Surface area.
6. Exposure to radiation.
1.Concentration
of the reactants.
The rate of a reaction is directly proportional to the
concentration of the reactants.
Rate law expression
THE LAW OF MASSACTION
(The Rate laws)
The rate of reaction is proportional to the
concentrations of reactants raised to а power.
 A + B = C;  = k[A][B]
equation of the rate laws
 3H2+N2 =2NH3;  = k[H2]3[N2]
 The coefficient k is called the rate constant
for the reaction or velocity constant. The rate
constant is independent of the concentrations
but depends on the temperature.
 [A] = [B] = 1 mole/liter, then rate = k

3. Temperature of the system.
In general, an increase in the temperature increases the rate of almost all
chemical reactions. This effect is observed for exothermic as well as for
endothermic reactions.
A general approximate rule for the effect of temperature on
reaction rates is that the reaction rate for most of the chemical
reactions becomes almost double, for every 100 C rise in
temperature.
This is also called temperature coefficient. It is the ratio of rate
constants of the reaction at two temperature differing by 100 C
Thus
or Vant-Hoff’s rule :
Т
kT
time1



Т kТ time2
2
2
1
1
t
10
where T2T1
Temperature coefficient of reaction:
kt 10

kt



Arrhenius Equation
It is a well-known fact that raising the temperature
increases the reaction rate.
E a = activation energy
R = 8.314 [ J · mol -1 · K -1 ]
T = absolute temperature in degrees Kelvin
A = pre-exponential or frequency factor
A = p · Z, where Z is the collision rate and p is a
steric factor.
Z turns out to be only weakly dependant on
temperature.
Thus the frequency factor is a constant, specific for
each reaction.
4. Presence of a catalyst. A catalyst is a substance
which influences the rate of a reaction without
undergoing any chemical change itself. It has been
observed that many reactions are made to proceed
at an increased rate by the presence of certain
catalysts.
5. Surface area.
The large the surface area of the reactants, the faster is
rate of reaction. It has been observed that if one the
reactants is a solid, then the rate of the reaction
depends upon the state of sub-division of the solid.
6. Exposure to radiation.
In some cases, the rate is considerably increased by the
use of certain radiations. For example, reaction of
hydrogen and chloride takes place very slowly in the
absence of light. However, in the presence of light,
the reaction takes place very rapidly.
Catalysis.
A substance which changes the speed of a reaction
without being used itself is called a catalyst.
The phenomenon of increasing the rate of reaction by
the use of catalyst is called catalysis.
If а catalyst increases (accelerates) the speed of а
reaction, it is called а positive catalyst and the
phenomenon is called positive catalysis. On the other
hand, if а catalyst decreases (retards) the speed of а
reaction, it is called а negative catalyst and the
phenomenon is called negative catalysis.
1. Homogeneous catalysts.
If the catalyst is present in the same phase as the
reactants, it is called а homogeneous catalyst and this
type of catalysis is called homogeneous catalysis.
NO(g)
2 SO2(g) + О2(g) ===== SO3(g)
Н+ (aq)
CH3COOC2H5(l)+Н2О(l)=====СНЗСООН(l)+C2H5OH(1)
Н+ (aq)
С12Н22О11(aq)+Н2О(1)====С6Н12О6(aq)+С6Н12О6 (aq)
Sucrose
Glucose
Fructose
2. Heterogeneous catalysts.
If the catalyst is present in а different phase
than the reactants, it is called а
heterogeneous catalyst and this type of
catalysis is called heterogeneous catalysis.

Pt, 8000С
4NH3 + 5O2 ======== 4NO + 6Н2O
Types of catalysis
Positive
Negative
Autocatalysis
Homogeneous Heterogeneous
Acid-base specific
Enzyme
Acid-base
unspecific
Enzymes
Substance that acts as a catalyst in living organisms, regulating
the rate at which life's chemical reactions proceed without being
altered in the process.
Enzymes are classified by the type of reaction they catalyze:
1.Oxidation-reduction
2.Transfer
of a chemical group
3.Hydrolysis
4.Removal or addition of a chemical group
5.Isomerization
6.Polymerization
Influence on the activity of enzymes:
1. Enzyme activity can be affected by other molecules.
Inhibitors are molecules that decrease enzyme activity;
If a competing molecule blocks the active site or changes its
shape, the enzyme's activity is inhibited. If the enzyme's
configuration is destroyed (denaturated), its activity is lost.
Activators are molecules that increase activity.
Many drugs and poisons are enzyme inhibitors.
2. Activity is also affected by temperature
3. Chemical environment (pH).
4. The concentration of substrate.
The optimal meaning of рН
for enzymes
Enzyme
Substrate
рН
-fructofuranozydaza
Urease
Papain
Pepsin
Arginase
Saccharose
Urea
Protein
Protein
Arginine
4,5-6,6
6,7
5,0
1,5-2,0
9,5-9,9