BioEnergetics

BIOENERGETICSBIOENERGETICS
Reading:
Harper’s Biochemistry pp. 123-129
Lehninger Principles of Biochemistry
3rd Ed. pp. 485-522

OBJECTIVESOBJECTIVES
To gain an understanding of concepts used to deal
with energy flow in living organisms.
To understand the following terms and concepts
1. Enthalpy
2. Entropy
3. Free Energy
4. Bioenergetic coupling of chemical reactions
5. Additivity of free energy changes
6. Relationship between standard free energy and
equilibrium constant
7. Role of ATP as energy currency of cell

Bioenergetics- BiochemicalBioenergetics- Biochemical
ThermodynamicsThermodynamics
Quantitative study of the energy transductions that
occur in living cells, and of the nature and function
of the chemical processes underlying these
transductions
Provides underlying principles to explain why some
reactions may occur while other do not
Non-biological systems may use heat energy to
perform work, whereas biological systems are
essentially isothermic and use chemical energy to
power living processes

Biomedical Importance of BioenergeticsBiomedical Importance of Bioenergetics
Fuel is required to provide energy for normal
processes, so understanding energy production and
utilization is fundamental to understanding normal
nutrition and metabolism
Starvation - occurs when available energy reserves
are depleted
Certain forms of malnutrition are associated with
energy imbalance e.g. marasmus- wasting disease
due to insufficient energy and protein intake
Excess storage of surplus energy results in obesity
which can have negative effects on health

Gibbs Free Energy Change (Gibbs Free Energy Change (∆∆G)G)
∆G is that portion of the total energy change in a
system that is available for doing work - it is the
useful energy
When a reaction proceeds with a release of free
energy (i.e. the system changes so as to possess
less free energy), the free energy change, ∆G, has
a negative value and the reaction is said to be
exergonic
In endogonic reactions, the system gains energy
and ∆G is positive
Units of ∆G - joules/mole (J/mol)
- calories/mole (cal/mol)

Enthalpy, HEnthalpy, H
Enthalpy is the heat content of the reacting system
It reflects the number and kinds of chemical bonds
in the reactants and products
When a chemical reaction releases heat, it is said
to be exothermic - the heat content of the products
is less than that of the reactants and by convention,
∆H has a negative value
Reacting systems that take up heat from their
surroundings are endothermic and have positive
value of ∆H
Units of ∆H - joules/mole (J/mol)
- calories/mole (cal/mol)

Entropy, SEntropy, S
Entropy is a quantitative expression for the
randomness or disorder of a system
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
Units of ∆S - J/mol·K
- cal/mol·K
K= units of absolute temperature (25°C = 298K)

Entropy, SEntropy, S
Example
Oxidation of glucose C6H12O6+6 O2→6 CO2+ 6 H2O
Increase in number of molecules, or when a solid is
converted to liquid or gas, generates molecular
disorder. Entropy increases.

Relationship betweenRelationship between ∆∆G,G, ∆∆H, andH, and ∆∆SS
Under conditions existing in biological systems
(constant temperature and pressure), changes in
free energy (∆G), enthalpy (∆H), and entropy (∆S)
are related to each other quantitatively:
∆G = ∆H - T ∆S, where T= absolute temp (K)
∆H has a negative sign when heat is released by
the system to the surroundings
∆S has a positive sign when entropy increases

In a favorable exergonic process which releases
heat and increases entropy:
e.g. Oxidation of glucose
∆G = (negative value of ∆H) - (T ⋅ positive value ∆S)
∆G = negative value
For favorable or spontaneous processes, ∆G has a
negative value
Relationship betweenRelationship between ∆∆G,G, ∆∆H, andH, and ∆∆SS

Free EnergyFree Energy
We must subtract the energy lost to increasing
entropy of the system from the total enthalpy
change to figure the amount of energy left over
available for useful work:
∆G = ∆H - T∆S
At equilibrium in a closed system no net change in
free energy can occur, ∆G = 0 and ∆H = T∆S

Example:
Heat water in tea kettle - steam is produced and
potentially capable of doing work
Allow to cool, no work is done, temp of
surroundings increases by infinitesimal amount until
equilibrium is reached. Kettle and surroundings are
at the same temp, the free energy that was once in
the kettle has disappeared.
∆H (change in heat) = T∆S (change in entropy)
∆G = ∆H - T∆S, ∆G = 0, no free energy
available to do work
Irreversible
Free EnergyFree Energy

For general reaction: aA+bB cC+dD, where
a,b,c,d = number of molecules of A,B,C,D
Equilibrium constant, Keq = [C]c
[D]d
[A]a
[B]b
When a reacting system is not at equilibrium, the
tendency to move toward equilibrium represents a
driving force, the magnitude of which can be
expressed as the free energy change for the
reaction, ∆G
At 25°C (298K) and at [1M] for participants, driving
force = ∆G ° = standard free energy change
If [H+] involved, 1M H+ = pH 0

For biochemical reactions, at pH 7, define:
∆Go´
= standard transformed free energy change
∆Go´
= - RT ln K´
eq = -2.303 RT log K´
eq

Indicates how much free energy is available from
the indicated reaction under standard conditions

Actual free energy changes depend onActual free energy changes depend on
reactant and product concentrationsreactant and product concentrations
The standard free-energy change tells us which
direction and how far a given reaction will go to
reach equilibrium when the initial concentration of
each component is 1.0M, the pH is 7.0, the temp is
25°C, and the pressure is 1 atm.

However, actual free-energy change, ∆G, is a
function of the reactant and product concentrations
and the prevailing conditions:
∆G and ∆Go´
are related for A+B C+D by
∆G = ∆Go´
+ RT ln [C][D]
[A][B]
At equilibrium, ∆G = 0
0 = ∆Go´
+ RT ln [C][D]
[A][B]
∆Go´
= -RT ln Keq
Actual free energy changes depend onActual free energy changes depend on
reactant and product concentrationsreactant and product concentrations

ExampleExample
Oxaloacetate + acetyl-CoA + H2O→citrate + CoA + H+
At pH 7 and 25°C in rat heart mitochondria - oxaloacetate = 1µM; acetyl-
CoA = 1µM; citrate = 220 µM; CoA = 65 µM
∆G°´= -32.2 kJ/mol
RT = 2.48 kJ/mol
What is direction of metabolite flow?
Solution - calculate ∆G, positive or negative?
∆G = ∆G°´+ RT ln [P][P]
[R][R]
∆G = -32.2 + 2.48 ln [220][65]x10-12
[1][1] x10-12
∆G = -32.2 + 2.48 x ln[14300]
= -32.2 + 23.7
= -8.5 kJ/mol

For sequential reactions,
A B and B C, the
∆Go´
values are additive:
∆Go´
Total = ∆G1
o´
+ ∆G2
o´
A thermodynamically
unfavorable reaction
(endergonic) can be
driven in the forward
direction by coupling it to
an exergonic reaction
Standard Free-Energy Changes are AdditiveStandard Free-Energy Changes are Additive

Vital processes- e.g.
synthetic reactions,
muscle contractions,
active transport, obtain
energy by chemical
linkage, or coupling, to
oxidative reactions
One way of coupling an
exergonic to an
endogonic process is to
synthesize a compound
of high energy potential
in the exergonic reaction
and to incorporate the
new compound into the
endergonic reaction

~ Can
theoretically
serve as a
transducer of
energy for a
wide range of
reactions
E

Example
Synthesis of glucose 6-phosphate
Glucose + Pi→glucose 6-phosphate + H2O
∆Go ´
= 13.8 kJ/mol
(will not proceed spontaneously in this direction)
ATP + H2O→ADP + Pi; ∆Go´
= - 30.5 kJ/mol
These reactions share the common intermediates Pi and H2O
and may be expressed as:
(1) Glucose + Pi→glucose 6-phosphate + H2O
(2) ATP + H2O→ADP + Pi
Glucose + ATP→ADP + glucose 6-phosphate
∆Go ´
= 13.8 kJ/mol + (-30.5 kJ/mol) = -16.7 kJ/mol
Overall reaction is exergonic
The actual pathway of glucose 6P formation is different from

ATP is the shared
chemical intermediate
linking energy-releasing
to energy-requiring cell
processes. Its role in
the cell is analogous to
that of money in an
economy: It is
“earned/produced” in
exergonic reactions
and “spent/consumed”
in endergonic
reactions.
ATP has a special role as energy currencyATP has a special role as energy currency

Hydrolysis causes charge
separation, relieving
electrostatic repulsion
among the four negative
charges on ATP
Inorganic phosphate
released is stabilized by
formation of a resonance
hybrid
ADP2-
produced ionizes
Greater degree of
solvation of ADP and Pi
than ATP
Chemical basis for the large free-energyChemical basis for the large free-energy
change associated with ATP hydrolysischange associated with ATP hydrolysis

ATP has two “high-energy” phosphateATP has two “high-energy” phosphate
groupsgroups
Standard free-energy of hydrolysis of ATP is
intermediate in list of organophosphates

ATP can act as a donor
of high-energy
phosphate to
compounds below it in
the table
ADP can accept high-
energy phosphate to
form ATP from those
compounds above it in
the table
This forms ATP/ADP
cycle

Adenylyl Kinase Interconverts AdenineAdenylyl Kinase Interconverts Adenine
NucleotidesNucleotides
Adenylyl Kinase (or myokinase) is present in most
cells and catalyzes the interconversion of ATP and
AMP to ADP and vice versa
ATP + AMP 2 ADP
Allows high-energy phosphate in ADP to be used in
synthesis of ATP
Allows AMP (formed as a consequence of several
activating reactions involving ATP) to be recovered
by rephosphorylation to ADP

The transfer of
these groups
couples the
energy of ATP
breakdown to
endergonic
transformation
of substrates
ATP can Donate Phosphoryl,ATP can Donate Phosphoryl,
Pyrophosphoryl, or Adenylyl GroupsPyrophosphoryl, or Adenylyl Groups

Involves attachment of
the carrier coenzyme A
Direct condensation of a
fatty acid with coenzyme
A is endergonic, but
process is made
exergonic by stepwise
removal of two
phosphoryl groups from
ATP
Hydrolysis of PPi to 2Pi
by inorganic
pyrophosphatase
releases additional
energy
Activation of a fatty acidActivation of a fatty acid

Other Nucleoside Triphosphates ParticipateOther Nucleoside Triphosphates Participate
in the Transfer of High-Energy Phosphatein the Transfer of High-Energy Phosphate
By means of the enzyme nucleoside diphosphate
kinase (NDK), nucleoside triphosphates similar to
ATP but containing different bases (U,G,C) can be
synthesized
ATP + UDP ADP + UTP
ATP + GDP ADP + GTP
ATP + CDP ADP + CTP
Similarly, specific nucleoside monophosphate
kinases (NMK) exist:
ATP + nucleoside  ADP + nucleoside ~
NDK

SummarySummary
Biological systems are isothermic and use chemical
energy to power living processes
Chemical reactions are influenced by two forces:
(1) The tendency to achieve the most stable
bonding state (enthalpy, H)
(2) The tendency to achieve the highest degree of
randomness (entropy, S)
The net driving force of a reaction, ∆G, the free-
energy charge, represents the net effect of those
two factors: ∆G = ∆H - T∆S.

SummarySummary
The standard free-energy change, ∆Go´
, is a physical
constant for a given reaction and is related to the
equilibrium constant:
∆Go´
= - RT ln K´
eq
The actual free energy change, ∆G, is a variable which
depends on ∆Go ´
and the actual conditions:
∆G = ∆Go ´
+ RT ln [products]
[reactants]
∆G large, negative - reactions go in forward direction
∆G large, positive - reactions go in reverse
∆G is zero - system is at equilibrium

SummarySummary
Endergonic processes occur only when coupled to
exergonic processes. Free-energy changes are
additive for successive reactions sharing a common
intermediate.
ATP acts as the energy currency of the cell and is
the chemical link between catabolism and
anabolism. Its exergonic conversion to ADP and Pi
or to AMP and PPi is coupled to a large number of
endergonic reactions.

BioEnergetics

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