Reginald H. Garrett Charles M. Grisham Reginald Garrett & Charles Grisham University of Virginia Chapter 3 Thermodynamics of Biological Systems

Essential Question What are the laws and principles of thermodynamics that allow us to describe the flows and interchanges of heat, energy, and matter in biochemical systems?

Outline What are the basic concepts of thermodynamics? What can thermodynamic parameters tell us about biochemical events? What is the effect of pH on standard-state free energies? What is the effect of concentration on net free energy changes? Why are coupled processes important to living things? What characteristics define high-energy biomolecules ? What are the complex equilibria involved in ATP hydrolysis? What is the daily human requirement for ATP? What are reduction potentials, and how are they used to calculate free energy changes in oxidation-reduction reactions?

3.1 What Are the Basic Concepts of Thermodynamics? Definitions for thermodynamics: The system: the portion of the universe with which we are concerned The surroundings: everything else Isolated system cannot exchange matter or energy Closed system can exchange energy Open system can exchange either or both

3.1 What Are the Basic Concepts of Thermodynamics?

The First Law – The Total Energy of an Isolated System is Conserved E (or U) is the internal energy - a function that keeps track of heat transfer and work expenditure in the system E is heat exchanged at constant volume E is independent of path E 2 - E 1 = ΔE = q + w q is heat absorbed BY the system w is work done ON the system Thus both q and w are positive when energy flows into a system

Enthalpy Enthalpy – a better function for constant pressure H = E + PV If P is constant, ΔH = q ΔH is the heat absorbed at constant P Volume is approximately constant for biochemical reactions (in solution) So ΔH is approximately the same as ΔE for biochemical reactions

The Second Law – Systems Tend Toward Disorder and Randomness Systems tend to proceed from ordered to disordered states The entropy change for (system + surroundings) is unchanged in reversible processes and positive for irreversible processes All processes proceed toward equilibrium - i.e., minimum potential energy

Entropy A measure of disorder An ordered state is low entropy A disordered state is high entropy dS reversible = dq/T

“What is Life?”, asked Erwin Schrödinger, in A disorganized array of letters possesses no information content and is a high-entropy state, compared to the systematic array of letters in a sentence. Erwin Schrödinger’s term “negentropy” describes the negative entropy changes that confer organization and information content to living organisms. Schrödinger pointed out that organisms must “acquire negentropy” to sustain life.

Energy dispersion Entropy can be defined as S = k ln W And ΔS = k ln W final – k ln W initial Where W final and W initial are the final and initial number of microstates of a system, and k is Boltzmann’s constant. Viewed in this way, entropy represents energy dispersion – the dispersion of energy among a large number of molecular motions relatable to quantized states (microstates). The definition of entropy above is engraved on the tombstone of Ludwig Boltzmann in Vienna, Austria

The Third Law – Why Is “Absolute Zero” So Important? The entropy of any crystalline, perfectly ordered substance must approach zero as the temperature approaches 0 K At T = 0 K, entropy is exactly zero For a constant pressure process: C p = dH/dT

Free Energy Hypothetical quantity - allows chemists to asses whether reactions will occur G = H - TS For any process at constant P and T: Δ G = Δ H - T Δ S If Δ G = 0, reaction is at equilibrium If Δ G < 0, reaction proceeds as written

 G and  G o ´ - The Effect of Concentration on ΔG How can we calculate the free energy change for reactions not at standard state? Consider a reaction: A + B  C + D Then: Thus concentrations at other than 1 M will change the value of Δ G

ΔG° Can Be Temperature Dependent

ΔS° Can Be Temperature Dependent

3.3 What is the Effect of pH on Standard State Free Energies? A standard state of 1 M for H + is not typical for biochemical reactions. It makes more sense to adopt a modified standard state – i.e., 1 M for all constituents except protons, for which the standard state is pH 7. This standard state is denoted with a superscript “°” For reactions in which H + is produced: ΔG°´= ΔG° + RT ln [H + ] And for reactions in which H + is consumed: ΔG°´ = ΔG° - RT ln [H + ]

3.4 What Can Thermodynamic Parameters Tell Us About Biochemical Events? A single thermodynamic parameter is not very useful Comparison of several thermodynamic parameters can provide meaningful insights about a process Heat capacity values can be useful A positive heat capacity change for a process indicates that molecules have acquired new ways to move (and thus to store heat energy) A negative heat capacity change means that the process has resulted in less freedom of motion for the molecules involved

3.4 What Can Thermodynamic Parameters Tell Us About Biochemical Events? Unfolding of a soluble protein exposes significant numbers of nonpolar groups to water, forcing order on the solvent and resulting in a negative entropy change.

3.4 What Can Thermodynamic Parameters Tell Us About Biochemical Events?

3.5 What are the Characteristics of High- Energy Biomolecules? Energy Transfer - A Biological Necessity Energy acquired from sunlight or food must be used to drive endergonic (energy-requiring) processes in the organism Two classes of biomolecules do this: Reduced coenzymes (NADH, FADH 2 ) High-energy phosphate compounds – with free energy of hydrolysis more negative than -25 kJ/mol

High-Energy Biomolecules Table 3.3 is important Note what's high - PEP and 1,3-BPG Note what's low - sugar phosphates, etc. Note what's in between - ATP Note difference (Figure 3.6) between overall free energy change - noted in Table and the energy of activation for phosphoryl-group transfer

3.5 What are the Characteristics of High- Energy Biomolecules?

3.5 What Are the Characteristics of High- Energy Biomolecules? The activation energies for phosphoryl group transfer reactions are substantially larger than the free energy of hydrolysis of ATP.

Group Transfer Potentials Quantify the Reactivity of Functional Groups Group transfer is analogous to ionization potential and reduction potential. All are specific instances of free energy changes.

ATP An Intermediate Energy Shuttle Device PEP and 1,3-BPG are created in the course of glucose breakdown Their energy (and phosphates) are transferred to ADP to form ATP But ATP is only a transient energy carrier - it quickly passes its energy to a host of energy- requiring processes

ATP Contains Two Pyrophosphate Linkages ATP contains two pyrophosphate linkages. The hydrolysis of phosphoric acid anhydrides is highly favorable.

Phosphoric Acid Anhydrides How ATP does what it does ADP and ATP are examples of phosphoric acid anhydrides Note the similarity to acyl anhydrides Large negative free energy change on hydrolysis is due to: electrostatic repulsion stabilization of products by ionization and resonance entropy factors

Ionization States of ATP ATP has four dissociable protons pK a values range from 0-1 to 6.95 Free energy of hydrolysis of ATP is relatively constant from pH 1 to 6, but rises steeply at high pH Since most biological reactions occur near pH 7, this variation is usually of little consequence

3.6 What Are the Complex Equilibria Involved in ATP Hydrolysis?

The Effect of Concentration Recall that free energy changes are concentration- dependent So the free energy available from ATP hydrolysis depends on concentration We will use the value of −30.5 kJ/mol for the standard free energy of hydrolysis of ATP At non-standard-state conditions (in a cell, for example), the ΔG is different Equation 3.13 allows the calculation of ΔG - be sure you can use it properly In typical cells, the free energy change for ATP hydrolysis is typically −50 kJ/mol

3.7 Why Are Coupled Processes Important to Living Things? Many reactions of cells and organisms run against their thermodynamic potential – that is, in the direction of positive ΔG Examples – synthesis of ATP, creation of ion gradients These processes are driven in the thermodynamically unfavorable direction via coupling with highly favorable processes

3.8 What is the Daily Human Requirement for ATP? The average adult human consumes approximately 11,700 kJ of food energy per day Assuming thermodynamic efficiency of 50%, about 5860 kJ of this energy ends up in form of ATP Assuming 50 kJ of energy required to synthesize one mole of ATP, the body must cycle through 5860/50 or 117 moles of ATP per day This is equivalent to 65 kg of ATP per day The typical adult human body contains 50 g of ATP/ADP Thus each ATP molecule must be recycled nearly 1300 times per day

ATP Changes K eq by 10 8 Consider a process: A → B Compare this to A + ATP → B + ADP + P i Assuming typical cellular concentrations of ATP, ADP and P i, and using the cellular free energy change for ATP hydrolysis, it can be shown that coupling ATP hydrolysis to the reaction of A → B changes the equilibrium ratio of B/A by more than 200 million-fold

3.9 What Are Reduction Potentials? How Are Reduction Potentials Used to Calculate Free Energy Changes for Oxidation-Reduction Reactions? High ℰ o ' indicates a strong tendency to be reduced Crucial equation: Δ G o ' = − n ℱΔ ℰ o ' Δ ℰ o ' = ℰ o '(acceptor) - ℰ o '(donor) Electrons are donated by the half reaction with the more negative reduction potential and are accepted by the reaction with the more positive reduction potential: Δ ℰ o ' positive, Δ G o ' negative If a given reaction is written so the reverse is true, then the Δ ℰ o ' will be a negative number and Δ G o ' will be positive

3.9 What Are Reduction Potentials? The standard reduction potential difference describing electron transfer between two species: Is related to the free energy change for the process by: ΔG o ' = −n ℱ Δ ℰ o ' Where n represents the number of electrons transferred, ℱ is Faraday’s constant, and Δ ℰ o ‘ is the difference in reduction potentials between the donor and acceptor.

Questions 1-6, 8,10-11, 18