BIOENERGETICS.pptx

BIOCHEMISTRY
The study of Structure and Functions of
Biomolecules with the Reactions and
Processes they undergo

Carbohydrate
Fatty Acids
Proteins
Nucleic Acids
Enzymes and
Coenzymes;
Energy;
Conditions for
the Reaction
METABOLISM

BIOENERGETICS
SoM&D - UDOM
Medical Biochemistry

• Order is expensive
–Its existence is an active process
–It is an unnatural state of affairs
–If left to its devices nature will favor disorder
• An unattended environment will overtime become messy but it
wont suddenly become neat

• Equilibrium is not a prerequisite for homeostasis
• There are many examples of in living organisms of homeostasis existing when
state of affairs is far from equilibrium
• Energy is spent to pay for the decrease in entropy of the system

Bioenergetics
• Is a field of biochemistry dealing with the transfer
and utilization of energy in living systems

Bioenergetics
• The study of bioenergetics concerns the initial and the final energy
states of a reaction irrespective of the time it takes to go from
reactants to products
• Free energy AKA Gibbs free energy (G): That portion of the energy of a
system available to do work as the system proceeds toward
equilibrium under conditions of constant temperature and pressure
• Change in Gibbs free energy (ΔG) is the main measure in bioenergetics
– It is a measure of the energetic feasibility of a reaction and allows a prediction
of whether or not the reaction will take place

Why study Bioenergetics
• Living organisms must perform work to stay alive, to grow and to
reproduce. The metabolic processes through which living organisms
accomplish these tasks utilize energy/involve energy transactions
• Hence the rationale for the ability of living organisms to harness energy
and to channel it into biological work (metabolism) as a fundamental
requirement for life
• Metabolism (vital for life) entails the chemical transformations taking place
in a cell or an organism
– Occurs through series/cascades of enzyme catalyzed reactions, referred to as
metabolic pathways
• Interconversion of metabolites, Catabolism and Anabolism

Why study Bioenergetics!
• Living organisms carry out a remarkable variety of
energy conversions from one form to another
• The chemical energy in fuels is used to bring about the
synthesis of complex macromolecules and structures
from simple precursors through making and breaking of
bonds

• The energy is also used to maintain concentration
gradients across membranes, heat production, for
motion, thinking, growth, etc.
• Energy manipulation for Biological Work for Life
• We have established that living systems transact energy;
–Do these transaction conform to the physical principals that
guide energy dynamics? Yes, How?

Bioenergetics Obey General Thermodynamics
• Living organisms conform to the general laws of
thermodynamics in their manipulation of energy
–Energy is transferred from one part of the system to another or
transformed from one form to another but it is not lost or created
–In these energy transformations entropy of the universe must
increase
–Living cells are isothermal, source of energy is usually not heat flow
from their sorroundings, they instead depend on free energy

Illustration by: Felipe Barros

• Energy manipulation occurs in living organisms
–Living organisms can neither create, from nothing nor destroy it to
nothing
–Autotrophs transform light energy into chemical energy
–Chemical energy is transformed into heat, electrical, mechanical or
chemical energy conserved by the formation of compounds
–Energy is extracted (catabolism) in a stepwise ordered fashion
• Fuel oxidation – glycolysis, fatty acid oxidation, TCA cycle

• Energy manipulation occurs in living organisms cont..
–Conserved for subsequent use in anabolism,
mechanical/electrical activity and/or lost (as heat)
• Conserved through the synthesis of ATP, high energy compounds
or the reduction of electron carriers NAD, NADP, FAD
–Order and organization in living systems must be
compensated/paid for by release of energy that will increase
entropy of surroundings

Bioenergetics in Biochemistry
• Explores the energy transformations which take place in the
reactions of the ATP-ADP cycle
• The ATP-ADP cycle refers to the utilization of ATP chemical
bond energy to do work required for life, and the continuous
oxidation of fuels (carbohydrate, lipid, protein) to regenerate
this ATP
–Mechanical work, transport work, biosynthetic work, etc.

Thermodynamics and Energy Transformation in
Living Systems
• Biological energy transactions obey the laws of thermodynamics
• First law is the principle of conservation of energy. Energy may
change form or it may be transported from one region to another,
but it cannot be created nor destroyed
– although energy can be converted into different forms, the total amount
of energy in a system must remain constant
– There seems to be a preferred direction (predictability) of energy flow

Living Systems
• Second law says that the universe always tends toward
increasing disorder. In all natural processes, the entropy of the
universe increases.
–Spontaneous processes occur in directions that increase the overall
disorder of the universe (system and surroundings)
–Sum of the entropies of a system and its surroundings must always
increase

Living Systems
• Second law
–All processes whether chemical or biological progress toward a
situation of maximum randomness or disorder (entropy, S)
–Equilibrium in a system is achieved when the entropy is at a
maximum
• Energy changes in biological systems are governed by the first
and second law of thermodynamics

Living Systems
• Reacting system: a collection of matter that is undergoing a
particular chemical or physical process.
– An organism, a cell, two or more reacting compounds
• Surroundings: everything else apart from the “reacting
system”
• Universe: the reacting system and surroundings

Living Systems
• Living cells and organisms are open systems, they exchange
materials and energy with their environment, they are never at
equilibrium with their surroundings

Living Systems
–Energy is spent to create and maintain the order in the cell or
organism and surplus energy is released to the surroundings to
create disorder such that the resultant of the decrease of entropy in
the system and increase of entropy in the surroundings is a net
increase of entropy of the universe
–This is how they are able to maintain a higher degree of order than
their surroundings without violating the 2nd law of thermodynamics

Thermodynamic Quantities
• These describe energy changes that occur in a chemical
reaction
–Free energy/Gibb’s free energy
–Enthalpy
–Entropy

Gibb’s free energy
• The amount of energy available to do work during a reaction,
at constant temperature and pressure, is expressed as free-
energy or Gibb’s free energy (G)
–A state function that relates enthalpy and entropy at a particular
temperature
–Expressed in joules/mole (J/mol) or calories/mole

• The Gibbs free energy (G) of a system is a measure of the
amount of usable energy (energy that can do work) in that
system
• The change in Gibbs free energy during a reaction provides
useful information about the reaction's energetics and
spontaneity (whether it can happen without added energy)

• Each compound involved in a chemical reaction contains a
certain amount of chemical potential energy, related to the
kind and number of its bonds

Gibb’s free energy change
• The energy that cells use to perform metabolic work is free
energy, described by the Gibb’s free energy, G
• Energy changes associated with reactions taking place in
biological systems are therefore reported as changes in free
energy, ΔG

• Heterotrophic cells acquire free energy from nutrient
molecules and photosynthetic cells acquire it from absorbed
solar radiation
• Both kinds of cells transform this free energy into ATP and
other energy-rich compounds capable of providing energy for
biological work at constant temperature and pressure

• Initially reactants posses higher G than the products and the
forward reaction proceeds with net loss of G,
• The reaction proceeds until a point of equilibrium where the
energy content of products and reactants balances, ΔG is zero
–The tendency of reacting systems to move towards equilibrium
represents a driving force, its magnitude is ΔG
• Because of the associated release of energy
–Nb: There is an association between ΔG of a reacting system and
reaction equilibrium, represented by equilibrium constant (Keq)

• There is release of free energy, G, when chemical reaction
systems are moving towards equilibrium
–Change in free energy can be calculated
–It is expressed as Gibb’s free energy change or simply as ΔG

• The equilibrium constant, Keq, defines this equilibrium
–There is a relationship between Keq and change in standard free
energy of a reaction in a particular direction
• The position of the equilibrium (the extent of converting
reactants to products) depends on
–Minimizing the difference in Gibb’s free energy between reactants
and products, ΔG
–Maximizing entropy, ΔS

Utility of Gibb’s free energy change
• The Gibbs free energy change of a reaction allows for prediction of:
– The direction where a chemical reaction is spontaneous,
– Their exact equilibrium position and
– The amount of work they can perform at constant temperature and pressure
• Reactions that result in a decrease in free energy, (i.e the products
have less free energy than reactants) are said to be exergonic. Such
reactions have a negative free energy change, -ΔG
• Exergonic reactions tend to occur spontaneously

Utility of Gibb’s free energy change
• Reactions that result in an increase in free energy, (i.e the
products have more free energy than reactants) are said to be
endergonic. Such reactions have a positive free-energy
change, +ΔG.
• Endergonic reactions are not spontaneous, require net energy
input to occur

Entropy
• The randomness or disorder of the components of a chemical
system is expressed as entropy, S
• Entropy is expressed in joules/mole. Kelvin (J/mol.K)
• Entropy is related to the number of accessible microstates
(degree of freedom) a substance can assume

Entropy
• Any change in randomness of the system is expressed as
entropy change, ΔS
• ΔS is positive if randomness increases
• When the products of a reaction are less complex and more disordered than
the reactants, the reaction is said to proceed with a gain in entropy
• The thermodynamically favored reaction direction must take
into account the degree of randomness or disorder of the
chemical system

Order to Disorder
• The spontaneous mixing of gases is driven by an increase in
entropy
• Ordered arrangement is replaced by a random mixture

Order to Disorder
• ΔS is positive
–Ice melting, vapourization
• ΔS is negative
–Freeze, condensation, crystallization

Entropy – alternative explanations
• How dispersed the energy is within a system amongst
the ways that that system can contain energy
• Energy that is distributed or dispersed among the
various motions of molecules of the system
–Movement of heat energy from hot to cold substance
(dispersal is favored)

Entropy – alternative explanations
• “Entropy is the tendency for heat energy to become
evenly distributed over time
• Energy is often defined in physics as the ability to do work
• If energy is evenly distributed, it cannot move anymore and no
work gets done. So there may be some energy in the system
but with no differential (distribution) that energy is not
available to do any work” by J. Werbock – extracted from researchgate

How is Order Possible in Biological Systems?
• Living cell exhibit ordered processes eg. Growth and differentiation,
and build very ordered substances eg. DNA, RNA, proteins
• Does this defy 2nd law of thermodynamics?
• Cell is not an isolated system, it can take energy from the
environment – food and light
• This energy is used to generate order in the cell but, through heat and
other simpler byproducts (CO2, H2O) released into the environment,
the surroundings are disordered, therefore total entropy of
“universe” is increasing (system + surroundings)

How is Order Possible in Biological Systems?
• The 2nd law of thermodynamics states that the entropy of the
universe increases during all chemical and physical processes,
but it does not require that the entropy increase must take
place in the reacting system itself
• The order produced within cells as they grow and divide is
compensated for by the disorder they create in their
surroundings through the numerous reactions in the course of
their growth and division

Enthalpy
• Enthalpy, H: Is the heat content of the reacting system
• It reflects the number and kinds of chemical bonds in the
reactants and products

Enthalpy
• Heat given up from or absorbed into a reaction
–In exothermic reactions the heat content of the products is
less than that of the reactants
• ΔH has a negative value
–In endothermic reactions the heat content of the reactants
is less than that of the products, i.e the reaction system
takes up heat from its surroundings
• ΔH has a positive value

Enthalpy
• Is a (convenient) way of looking at the energy of a system
• Enthalpy is expressed in joules/mole (J/mol) or calories/mole

The Gibb’s Equation
• The Gibb’s equation describes the relationship between the
three thermodynamic variables/quantities
ΔG= ΔH – TΔS
(T= absolute temperature in Kelvin, 25°C= 298K)
• The free energy change, which is a combination of enthalpy
and entropy changes, can be used to predict the direction in
which a reaction will spontaneously proceed.

• ΔS has a positive sign when entropy increases
• ΔH has a negative sign when heat is released from the
system to its surroundings
• Both conditions are typical of energetically favorable
processes and tend to make ΔG negative
• ΔG of a spontaneously reacting system is always negative

• The Gibb’s free energy change depends on:
–The difference in chemical bond energy between the products
and the reactants(enthalpy change or ∆H).
• ∆H = the chemical bond energy of products – the chemical bond
energy of the reactants
–The amount of energy unavailable for work because it has gone
into an increased disorder of the system
• The increase in entropy or ∆S
–The initial concentration of substrates and products

Favorable reactions
• If in a particular reaction direction of a reacting system:
– The H is negative, heat is released during reaction. Reaction is exothermic
in that direction
– The  S is positive, randomness increases
– Then  G < 0, the reaction is "spontaneous“, “
(thermodynamically/energetically favourable) in that direction
• Such a reaction release energy and is termed exergonic

Unfavorable reactions
• If in a particular reaction direction of a reacting system:
– The H is positive, heat is consumed during reaction. The reaction is
endothermic in that direction
– The  S is negative, randomness decreases
– Then  G > 0, the reaction is “not spontaneous“ (thermodynamically
unfavourable) in that direction
– Such a reaction consumes energy and they are termed endergonic.

.
Exergonic and Endergonic Reactions
Exergonic Reaction
Products -less energy than reactants
Energy released
Usually Entropy increases
spontaneous
Endergonic Reaction
Products -more energy than reactants
Energy required
Usually Entropy decreases
Not spontaneous 68

Sign of ΔG predicts the direction of a Reaction
• The change in free energy, ΔG can be used to predict the
direction of a reaction at constant temperature and pressure.
• Consider the reaction:
A → B

1. Negative ΔG:
• If ΔG is a negative number, there is a net loss of
energy, and the reaction goes spontaneously as
written,
• that is, A is converted into B.
• The reaction is said to be Exergonic.

2. Positive ΔG:
• If ΔG is a positive number, there is a net gain of
energy,
• and the reaction does not go spontaneously from
B to A .
• Energy must be added to the system to make the
reaction go from B to A,
• and reaction is said to be Endergonic.

3. ΔG is Zero:
• If ΔG = 0, the reactants and products are in
equilibrium,
• Note, when a reaction is proceeding
spontaneously, that is, free energy is being lost,
then the reaction continues until ΔG reaches zero
and equilibrium is established.
• No energy gain or lost

ΔG of the forward and back reactions
• The free energy of the forward reaction A → B is
equal in magnitude but opposite in sign to that of
the back reaction, B → A.
• E.g, if ΔG of the forward reaction is -5 J/mol, then
that of the back reaction is +5 J/mol.
• Note: ΔG can also be expressed in calories per
mole or cal/mol

Reaction Spontaneity
ΔG= ΔH – TΔS

< ΔH > ΔH
> ΔS Spontaneous Likely spontaneous
at high temperatures
< ΔS Likely spontaneous at
low temperatures
Never spontaneous
ΔG= ΔH – TΔS

ΔG= ΔH – TΔS
• ∆H < 0 & ∆S > 0: enthalpically favored (exothermic) and entropically
favored. Spontaneous reaction (exergonic) at all temperatures
• ∆H < 0 & ∆S < 0: enthalpically favored but entropically opposed.
Spontaneous reaction only at temperatures below T=∆H/∆S
• ∆H > 0 & ∆S > 0: enthalpically opposed (endothermic) but entropically
favored. Spontaneous reaction only at temperatures above T=∆H/∆S
• ∆H > 0 & ∆S < 0: enthalpically and entropically opposed.
Unspontaneous reaction (endergonic) at all temperatures

• ∆G is not an indicator of the velocity of the reaction or the
rate at which equilibrium is reached
• The velocity of the reaction depends on the amount of
enzyme available and the energy of the activation of the
reaction

Standard vs Non-standard Free Energy Change
• The change in free energy can be designated as standard or
actual, symbolized as ΔGo and ΔG, respectively
• Standard free energy change, ΔGo (with the superscript “o”),
is the free energy change at standard temperature (298K) and
pressure (1 atm) when reactants and products are at a
concentrations of 1 mol/L
–ΔGo is a constant for a given reaction

• Biochemical Standard free energy change, ΔGoʹ (transformed
standard free energy change) Is defined as a standard change
in free energy (i.e at 298K, 1atm and 1 mol/L reactants and
products) at pH 7

• Actual change in free energy, ΔG (without the superscript
“o”), is the more general/actual free energy change because
• It expresses the change in free energy and, thus, the direction
of a reaction at any specified concentration of products and
reactants
• In the cell, reactants and products are almost never at 1 mol/L
(1M) concentrations

Standard vs Actual Free Energy Change
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Standard vs Actual Free Energy Change
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Actual Free Energy Change, ΔG
ΔG depends on the concentration of Reactants and
Products
• The ΔG of the reaction A → B depends on the concentration of the
reactant and product.
• At constant Temperature and Pressure, the following relationship
can be derived:
ΔG = ΔGo + RT ln [B]/ [A]

ΔG = ΔGo + 2.303 RT Log [B]
[A]
• Where
–ΔG = Change in Free Energy
–ΔGo = The Standard Free Energy Change
–R = The Gas Constant (8.3 J/mol.K)
–T = The Absolute Temperature (K)
–[A] = The Actual Reactant Concentration
–[B] = The Actual Product Concentration

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Tractor capability coupled with soil condition (dry vs wet)
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are-the-advantages-of-ploughing-
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• A reaction with a positive ΔGo can proceed in the
forward direction (have a negative overall ΔG) if the
ratio of products to reactants ([B]/[A]) is sufficiently
small
–That is there is a large amount of reactants and very little
products

• ∆G depends on [reactant] and [products]. Therefore, it
can be negative even if ∆G°’ is positive, if the
[reactants] is much much larger than [products]
–Or if products are immediately removed from the system
• So a positive ∆G°’ doesn’t always mean that the
reaction won’t go forward in vivo
–It depends on how the levels of reactants and products
are controlled

• ∆G of ATP hydrolysis in vivo (in the cellular
environment) is larger negative number than the ∆G°’
of ATP hydrolysis in vitro (in experimental setup)
• At a normal physiological condition in cells, the
[ADP].[Pi]/[ATP] ratio is maintained at a low fraction value

Standard free energy change, ΔGo
• The standard free energy change, ΔGo, is so called because it
is equal to the free energy change, ΔG, under standard
conditions,
• that is, when Reactants and Products are kept at 1 mol/L
concentrations
• Under these conditions, the natural logarithm of the ratio of
products to reactants is zero
–Ln 1=0 or Log 1 =0

Therefore, ΔG = ΔGo + RT ln[B]/[A]
ΔG = ΔGo + 2.303 RT Log [B]
[A]
Becomes,
ΔG = ΔGo + 0
ΔG = ΔGo

Relationship Between ΔGo and Keq
• The ∆G° (or ∆G°’ ) is related to the equilibrium constant,
Keq, for a reaction

• In a reaction A → B, a point of equilibrium is reached at which
no further net chemical change takes place,
• that is, when A is being converted to B as fast as B is being
converted to A.
• In this state, the ratio of [B] to [A] is constant, regardless of the
actual concentrations of the two compounds:

• Keq = [B]eq
[A] eq
Where
• Keq is the equilibrium constant
• [A]eq and [B]eq are the concentrations of A and B at
equilibrium

• If the reaction A → B is allowed to go to equilibrium at
constant Temperature and Pressure,
• then at equilibrium the overall free energy change ΔG is zero
• Therefore,
ΔG = ΔGo + 2.303 RT Log[B]eq/[A]eq = 0

• Where the actual concentrations of A and B are equal to the
equilibrium concentrations of reactant and product, [A]eq and
[B]eq, and their ratio as shown above is equal to the Keq.
• Thus,
ΔGo = - 2.303 RT Log Keq
ΔGo = -RT lnKeq

• ΔGo can be obtained from the equilibrium constant of
a reaction system
•

This reaction allows some simple predictions:
• If Keq = 1, Then ΔGo = 0 A ↔ B
• If Keq > 1, Then ΔGo < 0 A → B
• If Keq < 1, Then ΔGo > 0 A ← B

ΔGo is predictive only under standard conditions:
• Under standard conditions, ΔGo can be used to predict the
direction a reaction proceeds because, under these
conditions, ΔGo is equal to ΔG.
• However, ΔGo cannot predict the direction of a reaction
under physiologic conditions
• If the standard free energy change (∆Go) can be determined,
then the actual free energy change of a reaction (∆G) can be
calculated.

• Although ΔGo represents energy changes at these non-
physiologic concentrations of reactants and products,
• It is nonetheless useful in comparing the energy changes of
different reactions.
–Free energy change of different reactions have been determined
experimentally at standard conditions
• Furthermore, ΔGo can readily be determined from
measurement of the equilibrium constant

• In cells metabolites do not exist in molar concentrations
• We talk about ΔG, rather than ∆G°’

Reaction Coupling
• Free energy changes in a (metabolic) pathway are
additive and those with positive values can be driven by
others with negative ∆G values.
• The biochemical standard free energy changes ∆G°’ are
additive in any sequence of consecutive reactions, as are
the free energy changes, ΔG

• This additive property of free energy changes is very
important in biochemical pathways through which substrates
must pass in a particular direction ( for example A →B → C →
D → E…)
• As long as the sum of the ΔGs of the individual reactions in a
pathway is negative, the pathway can potentially proceed as
written
–even if some of the individual reactions of the pathway have a
positive ΔG
• The actual rate of the reactions does, of course, depend on
the activity of the enzymes that catalyze the reactions.

Biochemical Reactions and their ΔGs
• Hydrolysis reactions – favourable, smaller –ve ΔG
• Isomerizations – occur are near equilibrium
• Oxidation of reduced fuels – favourable, large –ve ΔG

ATP and Other High Energy metabolites
• Some nutrient molecules have fuel qualities
–High in chemical potential energy
• Catabolism of such nutrient molecules by heterotrophs
releases this chemical free energy
• Some of this nutrient derived free energy is used (conserved
in) to synthesize ATP from ADP and Pi, directly or indirectly

• Catabolism of such nutrient molecules by heterotrophs
releases this chemical free energy

• ATP donates some of this conserved (or stored) energy in
anabolic processes to make them thermodynamically feasible,
yielding ADP + Pi or AMP + PPi
–It does this as it participates covalently in those processes (most
cases)
–In other cases through straight forward hydrolysis of ATP without
covalent association with the reactants (mechanical motion/trans
membrane transportation processes)

• ATP functions as an intermediate energetic compound that
metabolic processes can use
–Other polyphosphorylated nucleotides are also energy rich
compounds and are used similar to or interchangeably with ATP
–GTP, UTP, etc.
• There are other ‘energy rich’ metabolites that provide free
energy for metabolic reactions
–Can be synthesized directly from catabolism(oxidation) of fuels
–Can be synthesized from hydrolysis of ATP

• Some of these high-energy compounds can transfer a phosphate
group to make ATP
• The other high energy intermediates/metabolites referred to are:
– Phosphorylated compounds, some with mixed anhydride bond
– Thioesters where sulfur atom replace the Oxygen in the ester bond
• ATP and these other energy rich intermediates/metabolites are called
so because of the large negative free energy change values of their
hydrolysis
– They include PEP, Creatine phosphate, 1,3-Bisphosphoglycerate, Acyl-CoA

• Some of these high-energy compounds can transfer a phosphate
group to make ATP

• The designation as high-energy compounds implies that the
products of their hydrolytic cleavage are more stable forms
than the original compound
– it is the thermodynamic stability/favourableness of the hydrolytic
products of the high energy compounds that is the source of the
free energy released when their hydrolyzed,

ATP as the Energy Currency
• ATP is short for adenosine triphosphate
• ATP has a central role in the transfer of energy in biological systems.
(the energy currency)
• It links catabolism and anabolism
• Through catabolism of nutrient molecules, heterotrophs obtain energy
and use it to make ATP

• ATP in turn is used as a source of energy for anabolic
reactions.
–i.e Endergonic processes such as:
• The synthesis of metabolic intermediates and macromolecules
from smaller precursors
• Mechanical motion
• Transmembrane transportation of molecules, etc.

Why is ATP an Energy-Rich Molecule
I. The hydrolytic cleavage of the terminal phosphoanhydride
bond of ATP relieves some of the intramolecular
electrostatic repulsion existing in ATP
–Repulsion between the negatively charged Oxygen atoms of
phosphates and between the positively charged Phosphorus atoms
II. Formation of several resonance forms, by delocalization of
the pi electrons, stabilizes free inorganic phosphate (Pi)
–The Pi can not enjoy the freedom of resonance when it is bonded to
ADP

Why is ATP an Energy-Rich Molecule
III. The entropy increases. There is a greater stability in the
products because there exists a greater entropy; i.e. more
randomness. 1 mole of reactants has a higher energy than 2
moles of products. Disorder is favored over order according
to the 2nd law of thermodynamics
• The three factors make ATP hydrolysis relatively highly
exergonic

Resonance Stabilization
• Resonance allows for electron delocalization, in which the
overall energy of a molecule is lowered since its electrons
occupy a greater volume
• Molecules that experience resonance are more stable than
those that do not. These molecules are said to be resonance
stabilized

• The biochemical free energy change, ∆G°’, of ATP
hydrolysis is very different from the actual free energy
change, ∆G, in cellular conditions
–The [ATP], [ADP], [Pi] are far from the standard values used
to obtain ∆G°’
–Regulatory mechanism of ATP synthesis and breakdown hold
[ATP] far above the equilibrium concentration with its
hydrolytic products, ADP and Pi

–This makes ∆G a larger negative number than ∆G°’
–When cellular [ATP] decreases not only is the cell losing
energy molecules (losing its energy charge) but the potency
of the available (remaining) ATPs diminishes as well – i.e
their hydrolysis yields less free energy

• With the exception of usage of ATP in processes involving:
–Mechanical motion by the change of protein conformation (muscle
contraction, receptor activation)
–Active transmembrane transport
• Most other reactions in which ATP is the source of free energy
–ATP provides the energy by participating in the reaction covalently
–Groups Pi, PPi or AMP, are transiently transferred (covalently bound)
to the reactant (or enzyme) and then later displaced as the products
are released

• ATP donates phosphoryl, pyrophosphoryl or adenylyl groups in its
covalent involvement in metabolic reactions that use ATP to become
feasible
• Which group is donated (gets covalently attached to the reactant) and
which is displaced depends on which of the α, β and δ phosphates of
ATP has been attacked
– Hydrolysis of the α-β phosphoanhydride bond releases more free energy than
hydrolysis of the β-γ phosphoanhydride bond
• Reactions where AMP is donated are called adenylylation
– Inorganic pyrophosphatase hydrolyzes the PPi to 2Pi

ATP – ADP+Pi vs ATP – AMP+PPi

• Adenylylation is more thermodynamically favourable than transfer of
phosphory group
• Adenylylation is usually the mechanism of energy coupling in the very
unfavourable metabolic reactions that use ATP as the source of free
energy
– Higher ∆G°’ of the α-β phosphoanhydride bond hydrolysis
– Extra free energy from PPi hydrolysis

• Group transfer

Other High Energy Metabolites
• 1,3-Bisphosphoglycerate
– A glycolytic intermediate
– Hydrolysis of its mixed anhydride has a large negative ∆G°’ which is used to
transfer a phosphate to ADP producing ATP
– The product 3-phosphoglycerate enjoys resonance stabilization as it
interconverts between the two possible resonance forms
– For the above reasons 1,3-BPG high energy metabolite and it hydrolysis is
exergonic

• Phosphoenolpyruvate
– Another glycolytic intermediate, has a phosphate ester bond
– Hydrolysis of phosphoester bond has a large negative ∆G°’ which is used to
transfer a phosphate to ADP producing ATP
– Pyruvate produced by the hydrolysis of PEP has possibility for tautomerization
to a more stable keto form of the compound
• This keto – enol tautomerization is not possible in PEP, the possibility to form a more
stable intermediate makes this dephosphorylation favourable and therefore exergonic
– That free energy released does the work of phosphorylating ADP into ATP

• Phosphocreatine
– Is a form of energy storage in muscle, a quick supply of ATP during muscular
activity (recycling ATP from ADP)
– Hydrolysis of P-N bond has a large negative ∆G°’ which is used to transfer a
phosphate to ADP producing ATP
– creatine produced enjoys resonance stabilization, this favors its
dephosphorylation
– That free energy released does the work of phosphorylating ADP into ATP
• When there is ample ATPs some can hydrolyze and power creatine phosphorylation and
therefore storing/conserving that energy in the high energy molecule phsophocreatine

• Thioesters
– AcylCoA such as acetylCoA are common and important thioesters in
metabolism
– Thioesterification of an acyl group to CoA activates/energizes the acyl ready
for further reactions such as transacylation/condensation/oxidation-reduction
reactions
– Acetate (or other carboxylate) has possibility for resonance stabilization which
it looses when in a thioester bond with CoA (hence becoming less stable)
– This chemical fact (state of affairs) makes the thioesters like acetylCoA high
energy compounds
– The energy that holds this compound together is released as free energy
upon the thioester bond hydrolysis and it can be used to power reactions

• These hydrolysis reactions with large -ve ∆G°’, produce products more
stable than reactants for the following reasons:
– Relieving bond strain and electrostatic repulsions by charge separation
– Products can ionize and achieve better stability through resonance
delocalization
– Product stabilization through isomerization
– Products are ions which can undergo resonance stabilization

ATP as the energy currency
• The energy from ATP is obtained from the energy that is released
when ATP is hydrolyzed to ADP + Pi or AMP +2Pi.
• The hydrolysis of ATP and other high energy phosphate compounds
is accompanied by a large negative free energy change
• This energy is coupled to drive other unfavorable reactions

Energy coupling
• The central issue in bioenergetics is the means by which energy from
fuel metabolism (or light capture) is coupled to a cell’s energy
requiring reactions.
• Because cell function depends on macromolecules such as DNA,
proteins whose free energy of formation is positive, cells couple
these energy-requiring (endergonic) reactions to other reactions that
release energy (exergonic) so that the overall process is exergonic
• The usual source of free energy in coupled biological reactions is the
energy released by the breakage of phosphoanhydride bonds such
as those of ATP.

Simple examples as analogy of reaction coupling
• An object at the top of an inclined plane has a certain amount of potential
energy.
• Through appropriate string and pulley device this object can be coupled to
another smaller object such that when this elevated object slides down the
spontaneous downward motion will lift the smaller object. This amount of
energy available to do work is the free-energy change, G

• The energetically
unfavourable reaction Y → X
is driven by the energetically
favourable reaction, C → D,
because the free energy
change for the coupled
reactions is overall negative
Coupled Reaction

Coupled Reaction
• The third reaction is the sum of the reactions 1 and 2, and the G₃
is the arithmetic sum of G₁ and G₂. Because G₃ is negative, the
overall reaction is exergonic and proceeds spontaneously.

Example of a Coupled Reaction
• Example: synthesis of glucose-6-phosphate from glucose
1. Glucose + Pi → G-6-P + H2O ∆G°´= 13.8 kJ/mol
2. ATP + H2O → ADP + Pi ∆G°´= -30.5 kJ/mol
• The first reaction will not proceed alone as it has a
positive ∆G°´
• H2O and Pi are the common intermediates in these
reactions
• Overall reaction is:
ATP + glucose → ADP + G-6-P
• Overall ∆G°´= 13.8 - 30.5 = -16.7 kJ/mol
• So overall reaction is exergonic
• Energy stored in the bonds of ATP is used to drive the
synthesis of G-6-P

Biological Oxidation-Reduction Reactions
• The transfer of electrons in redox reactions is a common feature in
metabolism
• One chemical species looses electrons, it is oxidized, another one
gains electrons, it is reduced
• This flow of electrons is directly or indirectly responsible for all work
done by living organisms
• The ultimate source of electrons in heterotrophs are reduced
compounds, usually food substances (fuel)

Biological Oxidation-Reduction Reactions
• Electrons harvested as the fuels are oxidized and converted into
series of metabolic intermediates are carried by specialized carriers
to electron acceptors
• This train of transfer of electrons to an acceptor with higher
reduction potential releases energy that is channeled (transduced) to
useful cellular work - they flow down the potential energy slope
• Various enzymes and other proteins can use this energy to do
biological work
– Proton pumping by mitochondrial membrane enzyme complexes
– which provides the energy that later powers ATP synthesis

Types of Biological Redox Reactions
• Transfer as electrons
• Transfer as hydrogen atom(s), e- and H+ or hydride ion :H-
– FAD accepts electrons as the hydrogen atom (H) which is equivalent to two H+
and two electrons
– NAD+ accept electrons as the hydride ion (H:-) which is equivalent to one H+
and 2 electrons

Types of Biological Redox Reactions
• Direct combination with O2
– The hydrocarbon is the electron donor and the O2 the electron acceptor

Oxidation-Reduction Potential
• Reduction potential, E, is the measure of the affinity of a
chemical species for electrons
–Reduction potentials determined at standard conditions E°
• Electrons flow from the species (half cell) with lower E° to that
with higher/more positive E°
–Strength of this tendency to flow is proportional to ΔE°

Oxidation-Reduction Potential
• Reduction potential is affected not just by the nature of the
chemical species but also buy their activity, which is a function
of their concentration
–There is standard reduction potential-E°, and actual reduction
potential-E
• The energy associated with this flow of electrons is a form of
free energy
–Its magnitude, ΔG° is proportional to ΔE°

BIOENERGETICS.pptx

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