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BIOENERGETICS Bart Dzudzor

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Page 1: BIOENERGETICS (2)

BIOENERGETICS

Bart Dzudzor

Page 2: BIOENERGETICS (2)

Lecture Objectives:

• Be able to identify energy sources and there utilization.• Be able to understand how Endergonic processes

proceed by Coupling to Exergonic processes.• Be able to predict the spontaneity of biochemical

reactions based on the signs of ΔG, ΔH and ΔS.• Be able to calculate standard free energy changes of

biochemical reactions given the equilibrium constant and vice versa.

• How high-energy phosphates (ATP) play a central role in energy capture and transfer and also acts as the “energy currency” of the cell.

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Some Biomedical Importance of Bioenergetics

* Bioenergetics, or biochemical thermodynamics, is the study of the energy changes accompanying biochemical reactions. * Death from starvation occurs when available energy

reserves are depleted.* Certain forms of malnutrition are associated with energy imbalance (marasmus). * Thyroid hormones control the rate of energy release

(metabolic rate), and disease results when they malfunction.

* Excess storage of surplus energy causes obesity, one of the most common diseases of Western society.

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Introduction to Metabolism

Metabolism = the sum of all chemical reactions that take place in a cell or organism.

Bioenergetics = the quantitative study ofthe energy transductions that occur inin living cells and the nature and functionof the chemical processes underlying thesetransductions.

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What do cells and organism use their energy to do?

• Cells need energy for several cellular processes including;

• motility • pumping of molecules across membranes• maintaining temperature• cell growth and division• secreting molecules• sending signals to one another• synthesizing or degrading macromolecules• energy to do work, etc

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Sources of energy

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• Catabolic pathways generally converge• Anabolic pathways diverge• Some pathways are cyclic

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Overview of catabolism Complex

Monomeric unit

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Glycolysis

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Acetyl-Coenzyme A

Learn to

recognize this

molecule

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Principles of Metabolic Regulation A B C DE1 E2 E3

G1 G2 G3General features of metabolic pathways:• Individual reaction steps may be reversible, but the overall

pathway is irreversible.

• The irreversible committed step is usually an early reaction step. (G1 << 0)

• The rate of metabolism (flux) through a pathway is regulated at the committed reaction step (rate-limiting step).

• Opposing degradation/biosynthetic pathways that proceed through the same metabolites must differ in at least one enzyme to allow for independent regulation in each direction. (Often referred to as “coordinated control” or “reciprocal regulation”.)

• The enzymes that mediate a given reaction pathway are found in the same cellular compartment or organelle.

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Glycolysis

The irreversible committed step is usually an early reaction step. (G1 << 0)

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The enzymes that mediate a given reaction pathway are found in the same cellular compartment or organelle

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II. How are reaction steps regulated?A. Enzyme amount (esp. procaryotes)

1. Transcription2. Translation3. mRNA stability4. Protein stability (degradation)

B. Enzyme activity (positive or negative effectors)1. Metabolites2. Active site inhibitors3. Allosteric modulators4. [ATP]/[ADP] energy status5. [NADH]/[NAD+] reduction status6. Covalent modification of enzyme

(phosphorylation)

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III. Regulation by metabolites within the same pathway

A. Simple feedback inhibition

B.Complex feedback inhibition

1. Cumulative2. Concerted3. Sequential4. Isoenzymes for multiple effectors

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Laws of Thermodynamics• The first law of thermodynamics states that the total

energy of a system, including its surroundings, remains constant.

• It implies that within the total system, energy is neither lost nor gained during any change.

• However, energy may be transferred from one part of the system to another or may be transformed into another form of energy.

• In living systems, chemical energy may be transformed into heat or into electrical, radiant, or mechanical energy.

• Energy-utilizing reactions perform various necessary and, in many instances, tissue-specific, cellular functions, for example, nerve impulse conduction, muscle contraction, growth, and cell division.

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Laws of Thermodynamics cont’d.

• The second law of thermodynamics states that the total entropy of a system must increase if a process is to occur spontaneously.

• Entropy is the extent of disorder or randomness of the system and becomes

maximum as equilibrium is approached.

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Biological Systems Obey the PhysicalLaws of Thermodynamics

H = U + PV

G = H - TS

G = Gibbs Free EnergyH = EnthalpyS = EntropyU = Internal Energy

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Free Energy• Hypothetical quantity - allows chemists to asses whether reactions

will occur

• The Gibbs free energy: G = H - TS

• For any process at constant P and T:

G = H - TSIf G = 0, reaction is at equilibriumIf G < 0, reaction proceeds as written

Exergonic: Spontaneous processes with negative (-) G

Endergonic:Processes that are not spontaneous with positive (+) G

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Variation of Reaction Spontaneity (Sign of G) with the signs of H and S

Pag

e 56

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Free Energy Calculations:

G < 0, product formation is favoredG = 0, reaction at equilibriumG > 0, substrate formation favored

At equilibrium,

If all [i] = 1 M,

c d

a bG'o = -RT ln eq eq

eq eq

[C] [D] [A] [B]

aA + bB cC + dD , Grxn

enzyme

G = G'o

Standard free E change at pH 7Biochemists use this

Standard free E change under standard conditionspH 0, 1 atm pressure &When recatants & products are present initially at 1 M

G = G'o + RT ln[C] [D] [A] [B]

c d

a b

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Living systems can couple energy requiring reactions tothose which are spontaneous (exergonic)

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A chemical example of reaction coupling

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Exergonic reaction drives the endergonic processes

Some coupled reactions involving ATP

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Reasons why Certain compounds or Bonding arrangements are energy rich

1. Products of hydrolysis of an energy-rich bond may exist in more resonance forms than the precursor molecule.

• The more possible resonance forms in which a molecule can exist stabilize that molecule. Eg. Fewer resonance forms can be written for ATP or PPi than for Pi.

2. Many high-energy bonding arrangements have groups of similar electrostatic charges located in close proximity to each other in such compounds.

• Because like charges repel one another, hydrolysis of energy-rich bonds alleviates this situation lending stability to products of hydrolysis.

3. Hydrolysis of certain high-energy bonds results in formation of an unstable compound which may isomerize spontaneously to form a more stable compound. ΔGo’ for isomerization is considerable, and the final product is much more stable. e.g. PEP enolpyruvate to pyruvate.

4. If a product of hydrolysis of a high-energy bond is an undissociated acid, dissociation of the proton and its subsequent buffering may contribute to the overall ΔGo’ of the hydrolytic reaction.

In general, any property or process that lends stability to products of hydrolysis tends to confer a high-energy character to that compound.

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Adenosine triphosphate (ATP)

High-energy bonds

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Why is ATP a “high-energy” compound ?

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Garrett & GrishamFig. 3.8

ATP iskinetically

verystable

ATP in water is not readily convertedto ADP, but needs enzymes to mediate hydrolysis.

ATP + H2O Slow reaction

ADP + HPO42-

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ATP the universal energy carrier

It releases a significant amount of free energy upon hydrolysis. But not too much so it can be a between “high energy”phosphate donors and low energy acceptors.

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Why are the hydrolysis products (ADP+Pi) more stable than the reactant (ATP)?

(1) More resonance forms per phosphate in hydrolysis products

ATP

ADP Pi

(2) Charge separation in products

* More possible forms a molecule can exist stabilizes the molecule* Pi has more resonance forms than ATP

Similar electrostatic charges close together in ATP compared to products

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Creatine

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Hydrolysis of energy rich compound can form an unstable compound but can isomerize spontaneously to form stable compund

Stable form

Spontaneous isomerization

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Consider a reaction: A + B C Then:

G = Go + RT ln ([C]/[A][B])Go = free energy change of the reaction under standard conditions

Go used in physical chemistryGo' used in biochemistry

In "Chemistry":The standard state convention defines the standard state of solute as that with unit activity at 25oC and 1 atm. (So if H+ is a reactant or product, pH = 0.)

In "Biochemistry":The standard state convention is modified because most reactions occur in dilute solutions near pH 7 with activities of water and proton at unity

Remember that G = G'

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Free energy change is dependent on concentration

G Actual free energy change under specified conditions, including concentration of reactants and products

G0 Standard Free energy change, all reactants and products in their standard states, i.e.

1 M concentration.

Unfortunately, if [H+] = 1 M, pH = 0, which is not consistent with biochemical processes.

G0' Standard Free Energy change for the biochemical standard state, all reactants and products at 1 M

except

[H+] = [OH-] = 10-7 M, which allows pH = 7.standard transformed constants

Page 40: BIOENERGETICS (2)

Free Energy Calculations:

G < 0, product formation is favoredG = 0, reaction at equilibriumG > 0, substrate formation favored

At equilibrium,

If all [i] = 1 M,

c d

a bG'o = -RT ln eq eq

eq eq

[C] [D] [A] [B]

aA + bB cC + dD , Grxn

enzyme

G = G'o

Standard free E change at pH 7Biochemists use this

Standard free E change under standard conditionspH 0, 1 atm pressure &When reactants & products are present initially at 1 M

G = G'o + RT ln[C] [D] [A] [B]

c d

a b

Page 41: BIOENERGETICS (2)

1) Calculate the equilibrium constant for the hydrolysis of glucose-1-phosphate at 37oC.

From tableGo’ = - 20.9 kJ.mol-1

@ equilibrium G = 0

Answer: K = 3300

c d

a bG'o = -RT ln eq eq

eq eq

[C] [D] [A] [B]

KG = G'o + RT ln

[C] [D] [A] [B]

c d

a b

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Consider a reaction: A + B C

Go = free energy change of the reaction under standard conditions

G = Go + RT ln Keq

G = Go + RT ln([C]/[A][B])

At equilibrium, G = 0

Therefore, Go = - RT ln Keq = - RT ln ([C]/[A][B])

Thus the equilibrium constant can be calculated from standard free energy data and vice versa.

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Phosphocreatine + H2O Creatine + Pi

Go' at 37oC = - 43.1 kJ/mol

Physiological concentrations of phosphocreatine, creatine and Pi are between 1 and 10

mM.

Assuming 1 mM concentration and using equation

G = Go' + RT ln ([C][D]/[A][B])

G = - 43.1 kJ/mol + (8.314 J/mol. K) (310 K) ln ([0.001][0.001]/[0.001])

G = - 60.9 kJ/mol

Difference between standard state and 1 mM concentration for the above reaction is -17.8 kJ/mol

Free energy change can be very different from standard state if concentration of reactant and product are

different from unit activity (1 M)

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Variation of Keq with G° at 25°C

Pag

e 57

-

-

-

-

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Free Energies of Formation of Some Compounds of Biochemical Interest

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Pag

e 58

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Pag

e 58

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Consider a reaction: A + B C

Go = free energy change of the reaction under standard conditions

G = Go + RT ln Keq

G = Go + RT ln([C]/[A][B])

At equilibrium, G = 0

Therefore, Go = - RT ln Keq = - RT ln ([C]/[A][B])

Thus the equilibrium constant can be calculated from standard free energy data and vice versa.

Page 49: BIOENERGETICS (2)

Linking of individual amino acids

Protein Cells cope this situation by coupling this reaction to a highly negative G reaction

Many biological reactions lead to an increase in order (decrease in entropy)

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An unfavorable reaction can proceed spontaneously if it is coupled to an

energetically favorable reaction

A B + X G = + 50 kJ/mol

X Y + Z G = -100 kJ/mol

Overall reaction: A B + Y + Z

G = - 50 kJ/mol

Page 51: BIOENERGETICS (2)

PEP + H2O Pyruvate + Pi G = -78 kJ/mol

ADP + Pi ATP + H2O G = 55 kJ/mol

PEP + ADP Pyruvate + ATP Total G = -23 kJ/mol

Values from human erythrocytes