1. Section 1 Introduction to Biochemical Principles

2. Chapter 4 Energy

3. Section 4.1: Thermodynamics  Energy is the basic constituent of the universe  Energy is the capacity to do work  In living organisms, work is powered with the energy provided by ATP (adenosine triphosphate)  Thermodynamics is the study of energy transformations that accompany physical and chemical changes in matter  Bioenergetics is the branch of TD that deals with living organisms

4. Section 4.1: Thermodynamics  Bioenergetics is especially important in understanding biochemical reactions  Reactions are affected by three factors:  Enthalpy (H)—total heat content  https://www.youtube.com/watch?v=SV7U4yAXL5I https://www.youtube.com/watch?v=SV7U4yAXL5I  Entropy (S)—state of disorder  https://www.youtube.com/watch?v=ZsY4WcQOrfk https://www.youtube.com/watch?v=ZsY4WcQOrfk  Free Energy(G)—energy available to do chemical work

5. Section 4.1: Thermodynamics  Three laws of thermodynamics:  First Law of Thermodynamics—Energy cannot be created nor destroyed, but can be transformed  Second Law of Thermodynamics—Disorder always increases  Third Law of Thermodynamics—As the temperature of a perfect crystalline solid approaches absolute zero, disorder approaches zero

6. Section 4.1: Thermodynamics  First two laws are powerful biochemical tools  Thermodynamic transformations occur in a universe composed of a system and its surroundings  Energy exchange between a system and its surroundings can happen in two ways: heat (q) or work (w)  Work is the displacement or movement of an object by force Figure 4.2 A Thermodynamic Universe

7. Section 4.1: Thermodynamics  First Law of Thermodynamics  Expresses the relationship between internal energy (E) in a closed system and heat (q) and work (w)  Total energy of a closed system (e.g., our universe) is constant   E = q + w  Unlike a human body, which is an open system  Enthalpy (H) is related to internal energy by the equation: H = E + PV  PV work is usually negligible in biochemical systems,  H is often equal to  E (  H =  E)

8. Section 4.1: Thermodynamics  First Law of Thermodynamics Continued  If  H is negative (  H <0) the reaction gives off heat: exothermic  If is  H positive (  H >0) the reaction takes in heat from its surroundings: endothermic  In isothermic reactions (  H =0) no heat is exchanged  Reaction enthalpy can also be calculated:   H reaction =  H products   H reactants  Standard enthalpy of formation per mole (25°C, 1 atm) is symbolized by  H f °

9. Section 4.1: Thermodynamics  First Law of Thermodynamics Continued   H reaction =  H products   H reactants C 2 H 5 OH (l) + (7/2) O 2 (g) ---> 2 CO 2 (g) + 3 H 2 O (l) Calculate the standard enthalpy of combustion for the following reaction: ΔH° f CO 2 = -393.5 kJ/mol ΔH° f H 2 O = -286.0 kJ/mol ΔH° f C 2 H 5 OH = -278.0 kJ/mol ΔH° f O 2 = 0 kJ/mol Answer? -1367 kJ/mol

10. Section 4.1: Thermodynamics  Second Law of Thermodynamics  Physical or chemical changes resulting in a release of energy are spontaneous  Nonspontaneous reactions require constant energy input Figure 4.3 A Living Cell as a Thermodynamic System

11. Section 4.1: Thermodynamics  As a result of spontaneous processes, matter and energy become more disorganized  Gasoline combustion  The degree of disorder is measured by the state function entropy (S) Figure 4.4 Gasoline Combustion

12. Section 4.1: Thermodynamics  Second Law of Thermodynamics Continued  Entropy change for the universe is positive for every spontaneous process   S univ =  S sys +  S surr  Living systems do not increase internal disorder; they increase the entropy of their surroundings  For example, food consumed by animals to provide energy and structural materials needed are converted to disordered waste products (i.e., CO 2, H 2 O and heat)  Organisms with a  S univ = 0 or equilibrium are dead

13. Section 4.2: Free Energy  Free energy is the most definitive way to predict spontaneity  Gibbs free energy change or  G  Negative  G indicates spontaneous and exergonic  Positive  G indicates nonspontaneous and endergonic  When  G is zero, it indicates a process at equilibrium Figure 4.5 The Gibbs Free Energy Equation

14. Section 4.2: Free Energy  Standard Free Energy Changes  Standard free energy,  G°, is defined for reactions at 25°C,1 atm, and 1.0 M concentration of solutes  Standard free energy change is related to the reactions equilibrium constant, K eq, at equilibrium (ΔG = 0)   G° = -RT ln K eq  Allows calculation of  G° if K eq is known  Because most biochemical reactions take place at or near pH 7.0 ([H + ] = 1.0  10 -7 M), this exception can be made in the 1.0 M solute rule in bioenergetics  The free energy change, relative to pH, is expressed as  G° ′, as follows:

15. Section 4.2: Free Energy  For the reaction HC 2 H 3 O 2 ↔ C 2 H 3 O 2 + H +, calculate the standard free energy change at equilibrium (ΔG ° ) and free energy change (ΔG ° f ). Assume that T = 25 ° C, P = 1 atm, R = 8.315 J/molK, and K eq = 1.80 ×10 -5. Is this reaction spontaneous? Not spontaneous. Rxn is spontaneous when pH drops below physiological pH.

16.  Coupled Reactions  Many reactions have a positive  G°′  Free energy values are additive in a reaction sequence  If a net  G°′ is sufficiently negative, forming the product(s) is an exergonic process Figure 4.6 A Coupled Reaction Section 4.2: Free Energy

17.  The Hydrophobic Effect Revisited  TD principles help us better understand spontaneous aggregation of nonpolar substances  NP substances disrupt water H-bond interactions (which are energetically favorable)  Entropy of water decreases when molecules “cage” the NP molecules  Aggregation of NP molecules decreases the surface area of their contact with water, increasing water’s entropy (i.e., free energy is negative, process is spontaneous)  Spontaneous exclusion of water contributes to membrane formation and protein folding

18.  Adenosine triphosphate is a nucleotide that plays an extraordinarily important role in living cells  Hydrolysis of ATP  ADP + P i (orthophosphate) or AMP + PP i (pyrophosphate) provides free energy through transfer of phosphoryl group Figure 4.7 Hydrolysis of ATP Section 4.3: The Role of ATP

19.  ATP produced from ADP + P i.  Drives reactions of several types: 1. Biosynthesis of biomolecules 2. Active transport across membranes 3. Mechanical work such as muscle contraction Figure 4.8 The Role of ATP Section 4.3: The Role of ATP

20.  Structure of ATP is ideally suited for its role as universal energy currency  Its two terminal phosphoryl groups are linked by phosphoanhydride bonds  Specific enzymes facilitate ATP hydrolysis Figure 4.9 Structure of ATP Section 4.3: The Role of ATP

21.  The tendency of ATP to undergo hydrolysis is an example of its phosphoryl group transfer potential  ATP acts as energy currency, because it can carry phosphoryl groups from high-energy compounds (e.g., PEP) to low-energy compounds (e.g., glucose, as part of glycolysis).  Process is described in detail in Chapter 8 of the textbook. Figure 4.10 Transfer of Phosphoryl Groups Section 4.3: The Role of ATP

23.  Several factors need to be considered to understand why ATP is so exergonic: 1. At physiological pH, ATP has multiple negative charges 2. Because of resonance stabilization, the products of ATP hydrolysis are more stable than resonance-restricted ATP  Resonance is when a molecule has two or more alternative structures that differ only in the position of their electrons 3. Hydrolysis products of ATP are more easily solvated 4. Increase in disorder with more molecules Figure 4.11 Contributing Structure of the Resonance Hybrid of Phosphate Section 4.3: The Role of ATP