1. Thermodynamics & ATP Review thermodynamics, energetics, chemical sense, and role of ATP

2. Lecture 24 Thermodynamics in Biology

3. A Simple Thought Experiment

4. Driving Forces for Natural Processes Enthalpy –Tendency toward lowest energy state Form stablest bonds Entropy –Tendency to maximize randomness

5. Enthalpy and Bond Strength Enthalpy = ∆H = heat change at constant pressure Units –cal/mole or joule/mole 1 cal = 4.18 joule Sign –∆H is negative for a reaction that liberates heat

6. Entropy and Randomness

7. Entropy = S = measure of randomness –cal/deg·mole T∆S = change of randomness For increased randomness, sign is “+”

8. “System” Definition

11. Cells and Organisms: Open Systems Material exchange with surroundings –Fuels and nutrients in (glucose) –By-products out (CO 2 ) Energy exchange –Heat release (fermentation) –Light release (fireflies) –Light absorption (plants)

12. 1 st Law of Thermodynamics Energy is conserved, but transduction is allowed Transduction

13. 2 nd Law of Thermodynamics In all spontaneous processes, total entropy of the universe increases

14. 2 nd Law of Thermodynamics ∆S system + ∆S surroundings = ∆S universe > 0 A cell (system) can decrease in entropy only if a greater increase in entropy occurs in surroundings C 6 H 12 O 6 + 6O 2  6CO 2 + 6H 2 O complex simple

15. Entropy: A More Rigorous Definition From statistical mechanics: –S = k lnW k = Boltzmann constant = 1.38  10 –23 J/K W = number of ways to arrange the system S = 0 at absolute zero (-273ºC)

16. Gibbs Free Energy Unifies 1 st and 2 nd laws ∆G –Gibbs free energy –Useful work available in a process ∆G = ∆H – T∆S –∆H from 1 st law Kind and number of bonds –T∆S from 2 nd law Order of the system

17. ∆G Driving force on a reaction Work available  distance from equilibrium ∆G = ∆H – T∆S –State functions Particular reaction T P Concentration (activity) of reactants and products

18. Equilibrium ∆G = ∆H – T∆S = 0 So ∆H = T∆S –∆H is measurement of enthalpy –T∆S is measurement of entropy Enthalpy and entropy are exactly balanced at equilibrium

19. Effects of ∆H and ∆S on ∆G Voet, Voet, and Pratt. Fundamentals of Biochemistry. 1999.

20. Standard State and ∆Gº Arbitrary definition, like sea level [Reactants] and [Products] –1 M or 1 atmos (activity) T = 25ºC = 298K P = 1 atmosphere Standard free energy change = ∆Gº

21. Biochemical Conventions: ∆Gº Most reactions at pH 7 in H 2 O Simplify ∆Gº and K eq by defining [H + ] = 10 –7 M [H 2 O] = unity Biochemists use ∆Gº and K eq

22. Relationship of ∆G to ∆Gº ∆G is real and ∆Gº is standard For A in solution –G A = G A + RT ln[A] For reaction aA + bB  cC + dD –∆G = ∆Gº + RT ln –Constant Variable (from table) º [C] c [D] d [A] a [B] b }

23. Relationship Between ∆Gº and K eq ∆G = ∆Gº + RT ln At equilibrium, ∆G = 0, so –∆Gº = –RT ln –∆Gº = –RT ln K eq [C] c [D] d [A] a [B] b [C] c [D] d [A] a [B] b

24. Relationship Between K eq and ∆Gº

25. Will Reaction Occur Spontaneously? When: –∆G is negative, forward reaction tends to occur –∆G is positive, back reaction tends to occur –∆G is zero, system is at equilibrium ∆G = ∆Gº + RT ln [C] c [D] d [A] a [B] b

26. A Caution About ∆Gº Even when a reaction has a large, negative ∆Gº, it may not occur at a measurable rate Thermodynamics –Where is the equilibrium point? Kinetics –How fast is equilibrium approached? Enzymes change rate of reactions, but do not change K eq

27. ∆Gº is Additive (State Function) Reaction A  B B  C Sum: A  C Also: B  A Free energy change ∆G 1 º ∆G 2 º ∆G 1 º + ∆G 2 º – ∆G 1 º

28. Coupling Reactions Glucose + HPO 4 2–  Glucose-6-P ATP  ADP + HPO 4 2– ATP + Glucose  ADP + Glucose-6-P ∆Gº kcal/mol kJ/mol +3.3 +13.8 –7.3 – 30.5 –4.0 – 16.7

29. Resonance Forms of P i –– –– –– ––

30. Phosphate Esters and Anhydrides

31. Hydrolysis of Glucose-6-Phosphate ∆Gº = –3.3 kcal/mol = –13.8 kJ/mol

32. High ∆Gº Hydrolysis Compounds ∆Gº = –14.8 kcal/mol = –61.9 kJ/mol

33. High ∆Gº Hydrolysis Compounds ∆Gº = –11.8 kcal/mol = –49.3 kJ/mol

34. High ∆Gº Hydrolysis Compounds ∆Gº = –10.3 kcal/mol = –43 kJ/mol

35. Phosphate Anhydrides (Pyrophosphates) ∆Gº = –7.3 kcal/mol = –30.5 kJ/mol

36. Thiol Esters ∆Gº = –7.5 kcal/mol = –31.4 kJ/mol

37. Thiol Esters Thiol ester less resonance-stabilized

38. “High-Energy” Compounds Large ∆Gº hydrolysis –Bond strain (electrostatic repulsion) in reactant ATP –Products stabilized by ionization Acyl-P –Products stabilized by isomerization PEP –Products stabilized by resonance Creatine-P

39. “High-Energy” Compounds “High-energy” compound is one with a ∆Gº below –6 kcal/mol (–25 kJ/mol)

40. High-Energy Compounds

41. Group Transfer Potential

60. Lecture 25 Chemical Sense in Metabolism

61. Making and Breaking C–C Bonds Homolytic reactions Heterolytic reactions

62. Making and Breaking C–C Bonds Nucleophilic substitutions

63. Nucleophilic Substitution Reactions S N 1

64. Carbocation

65. Common Biological Nucleophiles

66. S N 2 Nucleophilic Substitution –– ––

67. Reactivity is S N 2 Reactions

68. Leaving Group Must accommodate a pair of electrons –And sometimes a negative charge

69. Major Role of Phosphorylation Converts a poor leaving group ( – OH) into a good one (P i, PP i )

70. Acid Catalysis of Substitution Reactions This H is often donated by an acidic sidechain of enzyme

71. Central Importance of Carbonyls 1. Can produce a carbocation 2. Can stabilize a carbanion

72. Biological Carbonyls

73. Aldol Condensation

76. Aldolase Reaction Glycolysis and gluconeogenesis

77. Claisen Condensation

79. Thioesters in Biology In thioesters, the carbonyl carbon has more positive character than carbonyl carbon in oxygen ester.

80. “High-Energy” Thioester Compounds

81. Coenzyme A

82. Fatty Acid Metabolism Uses Claisen condensation Thiolase acts in fatty acid oxidation for energy production

83. Thiolase: Role of Cys-SH

85. Energy Diagram for Reaction ‡ is the transition state –Pentacovalent carbon, for example

86. Functional Groups on Enzymes Amino acid side chains – –Imidazole –

87. Functional Groups on Enzymes Coenzymes/cofactors –Pyridoxal phosphate Metal ions and complexes – Mg 2+, Mn 2+, Co 2+, Fe 2+, Zn 2+, Cu 2+, Mo 3+

88. Enzyme Inhibitors and Poisons Chelating agents –EDTA (divalent cations) –CN – (Fe 2+ ) Cofactor analogs –Warfarin Suicide substrates

89. Lecture 26 ATP and Phosphoryl Group Transfers

90. Phosphate Esters and Anhydrides

91. Phosphoryl Group Transfers

92. Phosphoryl (Not Phosphate) Transfers

93. Nucleophilic Displacements

94. ATP as a Phophoryl Donor 2 roles for ATP –Thermodynamic Drive unfavorable reactions –Mechanistic Offer 3 electrophilic phosphorous atoms for nucleophilic attack

95. ATP as Phosphoryl Donor 3 points of nucleophilic attack

96. Adenylyation: Attack on  -P

97. Adenylation: Attack on  -P

98. Pyrophosphorylation: Attack on  -P

99. Phosphorylation: Attack on  -P

100. Amino Acid Sidechains as Nucleophiles

101. Enzymatic Phosphoryl Transfers Four classes –Phosphatases Water is acceptor/nucleophile –Phosphodiesterases Water is acceptor/nucleophile –Kinases Nucleophile is not water –Phosphorylases Phosphate is nucleophile

102. Phosphatases: Glucose-6- Phosphatase

103. Phosphatases: Glucose-6- Phosphate

104. Phosphodiesterases: RNAase

106. Kinases:  -Phosphoryl Transfer Transfer from ATP

107. Kinases: P-Enzyme Intermediates

109. Kinases

110. Pyruvate Kinase Makes ATP (∆Gº= –31 kJ/mol) from PEP ∆Gº= –62 kJ/mol

111. Phosphoryl-Group Transfer Potential

112. Significance of “High-Energy” P Compounds Drive synthesis of compounds below Phosphated compounds are more reactive –Thermodynamically –Kinetically If organism has ATP (etc…), it can do work and resist entropy  Cells must get ATP