Outline

Introduction

> Bioenergetics

Energy, Heat, and Work

> > Isolated versus Open systems

> Internal Energy (E) and the State of a System

> > Internal Energy is a Function of the State of a System

> > > Amount of substance plus two of

> > > > 1) Temperature (T)

> > > > 2) Pressure (P)

> > > > 3) Volume (V)

> > Heat = q (+ = system absorbs heat)
> > Work = w (+ = work done by system)
> > Energy change = E

> First Law of Thermodynamics

> > Internal Energy Change (E = q-w) (Figure 3.1)
> > At constant P, w = PV, q = E + n*RT

> Enthalpy (H)

> > H = E +PV
> > At constant P, H = E + PV
> > At constant P, H = q

Entropy and the Second Law of Thermodynamics

> The Direction of Processes

> > Reversible = at equilibrium
> > Irreversible = far from equilibrium
> > Spontaneous/Favorable Processes

> Entropy (S) (Table 3.1)

> > Tendency of Systems of Molecules to Randomization
> > S = klnW (k = Boltzmann constant)

> Second Law - "The entropy of an isolated system will tend to increase to a maximum value"

Free Energy: The Second Law in Open Systems

> Gibbs Free Energy (Table 3.2)

> > G = H - TS
> > G = H - TS
> > G <0 means Exergonic, favorable process
> > G >0 means Endergonic, reverse process favored

> Example of Interplay of Enthalpy and Entropy (Figure 3.3)

> Summary of Interplay of Enthalpy and Entropy (Figure 3.4, Table 3.3)

> > 1. From G = H - TS, favorability of a reaction depends on temperature.
> > 2. Favorability of a reaction not related to its rate
> > 3. Entropy of an OPEN system can decrease - energy must be expended.
> > 4. Entropy of the whole universe must be increasing.

> Free Energy and Useful Work

> > 1. G is the part of the energy change (H) available to do useful work - work other than PV expansion.
> > 2. From G = H - TS, part of H is dissipated as heat (TS) and is not available for useful work.
> > 3. Efficiency = (Work accomplished)/(Theoretical work from free energy change).

Free Energy and Concentration

> Chemical Potential
> How Chemical Potential is Used: An Example (Figure 3.5)

Free Energy and Chemical Reactons: Chemical Equilibrium

> Free Energy Change and the Equibrium Constant

> > = Standard State Free Energy Change
> > = + RTln([products]/[reactants])
> > = + RTlnK, where K is equilbrium constant

> Free Energy Calculations: A Biochemical Example

> > 1. G6P <=> F6P ( = +1.7 kJ/mol means equilibrium concentration has more G6P than F6P)
> > 2. Plugging this into , one finds that (F6P)/(G6P) = 0.504
> > 3. Since G = + RTln([products]/[reactants]), displacements away from equilibrium will be moved towards equilibrium by corresponding force of free energy change brought about by the change. (Figure 3.6)

High Energy Phosphate Compounds: Free Energy Sources in Biological Systems

> Coupled Reactions

> > For
> > A <=> B ( = +10 kJ/mol) and
> > C <=> D ( = -30kJ/mol),
> > A + C <=> B + D ( = +10 - 30 = -20 kJ/mol)

> High-Energy Phosphate Compounds as Energy Shuttles (Figure 3.7)

> > ATP, PEP, Creatine Phosphate (CP) (Figure 3.8)

> > Resonance Statibilization of the Phosphate Products (Figure 3.9)

> > Additional Hydration of the Hydrolysis Products

> > Electrostatic Reulsion Between Charged Products

> > Enhanced Resonance Stabilization or Tautomerization of Product Molecules

> > > Pyruvate
> > > Creatine

> > Release of a Proton in Buffered Solutions

> Phosphate Transfer Potential (Figure 3.7)