7.1 Origins of Thermodynamics Developed in 19 th century to answer question about how to build a better steam engine – Driving force of industrial revolution Terms used in equation relate more to steam engines than biochemical reactions – heat and work

7.2 First Law of Thermodynamics First law of thermodynamics is the conservation of energy System – portion of the universe we wish to study Surroundings – everything else Exchanges can occur between system and surroundings

7.2 First Law of Thermodynamics Exchanges are heat and work – Heat Abbreviation is q Flow of energy between objects of different temperature Change is called internal energy – Work Abbreviation is w Refers to translational, vibrational, or rotational energy Change is called internal energy

7.2 First Law of Thermodynamics Equation for first law of thermodynamics – –  E = q + w  E = change in internal energy – as a change from one state to another Each state is characterized by a set of state variables – Temperature (T) – Pressure (P) – Volume (V) – Number of moles (n)

7.2 First Law of Thermodynamics Enthalpy – changes in chemical reactions involve alterations in translation, vibration and rotation These contribute to bond energy Enthalpy is an additive property – usually defined as  H – Under constant pressure our equation becomes  H =  E + P  V

7.2 First Law of Thermodynamics Exothermic – gives off heat Endothermic – absorbs heat Exothermic reactions are ones that are often favored – However we cannot use it to predict the direction of reactions 교과서 94 쪽 참조

7.3 – Entropy and the Second Law of Thermodynamics Entropy is the missing factor to determine direction of a reaction Second law states that the entropy change in the universe for any process is greater than zero

7.4 Free Energy Free energy (G) is the combination of first and second law of thermodynamics –  G =  H - T  S – Applies to system under consideration Three possible values for free energy – Positive – Negative – zero

7.4 Free Energy  G =  H - T  S Reaction is favorable when  G is less than zero – Called exergonic – Contributions to this are  H – energy of bond formation or bond breakage  S – probability or arrangement factor – When reaction is exergonic – heat term is dominant Reaction is enthalpy driven Binding is favorable and more important than entropy – If T  S is the more dominant, the reaction is entropy driven

7.4 Free Energy Endergonic reactions –  G is greater than zero – Not possibly thermodynamically Equilibrium defined by Free energy –  G = 0 – Forward and reverse reactions are not favored

7.5 Standard Free Energy Standard free energy change can be related to equilibrium constant. –  G ° = -RTlnK eq Direct relations means that substrate and product concentration and free energy for a reaction are directly related. Standard conditions – 0 C and 1 atm – Means to compare the laboratory measures of different observers

7.6 Nonstandard Free Energy Changes Under cellular conditions, reactions are at steady state rather than at equilibrium The standard free energy equation than needs to add Q – the mass action ratio –  G = - RTlnK eq + RTlnQ

7.6 Nonstandard Free Energy Changes  G = - RTlnK eq + RTlnQ In the cells  G has to be less than zero for the reaction to proceed. – Q < K eq If the reaction is endergonic – the reaction will not proceed – Q > K eq A relationship exists between equilibrium, standard free energy change and actual free energy change

7.7 Near Equilibrium and Metabolically Irreversible Reactions Reactions in cells are part of sequences called pathway In a linear pathway all reactions take place at the same rate – They are all exergonic Can be divided into two groups – Near Equilibrium – Metabolically Irreversible Reactions

7.7 Near Equilibrium and Metabolically Irreversible Reactions Near equilibrium reactions – Q is close to K eq – Typical case for cellular reactions – Two consequences of this reactions Reaction is sensitive to changes in substrate or product Steps are rarely sites of control – Enzymes that catalyzes are in larger amounts

7.7 Near Equilibrium and Metabolically Irreversible Reactions Metabolically Irreversible Reactions – Q is much less than K eq – Greatly displaced from their equilibrium positions – Cannot run in reverse directions under cellular conditions – Sites of metabolic control Allosteric or covalent – Cellular content of enzymes is generally low

7.8 ATP Adenosine 5’triphosphate (ATP) is central energy intermediate – Hydrolysis reactions of ATP

7.8 ATP Smaller change in standard free energy for hydrolysis of AMP to adenosine than for hydrolysis of ATP to adenosine diphosphate – Increase in entropy occurs – Repulsion interactions from negative charges are relieved – Individual phosphates have more resonance possibilities

7.8 ATP Most of ADP and AMP that exist in the cell are bound to proteins – Must measure indirectly –  G is about – 59 kJ/mol in cytosol of muscle ATP hydrolysis does not occur physiologically – Phosphoryl group is transferred

7.9 Energy Coupling with ATP Energy coupling is joining of a process is endergonic in isolation with one that is exergonic to produce a process that is exergonic Asparagine synthetase reaction example – Formation of aspartate (reverse reaction)  G ° is kg/mol – ATP hydrolysis – 30kJ/mol – Standard Free Energy for entire reaction kJ/mol

7.9 Energy Coupling with ATP Actual reaction mechanism is more complicated – Two nucleophilic substitutions occur

7.9 Energy Coupling with ATP Maintenance reactions for ATP – Creatine Phosphokinase – NDP kinase – Adenylate Kinase

7.9 Energy Coupling with ATP Creatine Phosphokinase – Catalyzes the reaction of creatine with ATP – Phosphoryl group of ATP is transferred to a nitrogen of creatine – Creatine phosphokinase enzyme is highly expressed in muscle cells

7.9 Energy Coupling with ATP Creatine Phosphokinase cont. – Allows creatine phosphate to serve as very rapid energy store – Allows ATP formation to remain effectively constant despite rapid changes in energy demand – Near – Equilibrium reaction

7.9 Energy Coupling with ATP NDP kinase – Allows nucleotides other than ATP to serve as energy donors – Likely to be a metabolically irreversible reactions – Overall standard free energy is near zero – Relaxed specificities as it accepts all nucleotides

7.9 Energy Coupling with ATP Adenylate Kinase – Equilibrium constant is approximately 1 – Near equilibrium reaction in the cell – Provides mean of rephosphorylated AMP – Results in two ADP product molecules

7.10 NADH Electron carriers can also be energy intermediates between catabolism and anabolism Principle electron carrier is NADH

7.10 NADH Change in redox states involves two connected processes – Oxidation is loss of one or more electrons – Reduction is gain of one or more electrons – Two processes must occur together It is possible to focus on just half of the redox reaction at a time – called a half-reaction

7.10 NADH NADH is not the only electron carrier – NADPH is a hydride carrier Has one extra phosphoryl group Most enzymes use one or the other – Transhydrogenase can convert NADH to NADPH Metabolically irreversible reaction

7.11 Mobile Cofactors and the Pathway View Mobile cofactors – links between or within metabolic sequences – ATP, ADP, NADH and NAD are examples Connections between reactions occur in pathway view Pathway is the product of one reaction becomes the substrate of the next

7.11 Mobile Cofactors and the Pathway View Different pathways can be connected in parallel Prosthetic groups – bound cofactors – Must be regenerated with whole rest of enzyme First substrate in the pathway is the pathway substrate Final product in the pathway is the pathway product

7.11 Mobile Cofactors and the Pathway View If pathway is a linear sequence than there is one rate for the entire pathway - called the pathway flux Mobile cofactors are limited and regenerated by an external pathway