Environmental Cycles of Metabolism

1. Environmental Cycles of Metabolism • Carbon is fixed (incorporated) by autotrophs (CO2) and heterotrophs (complex such as carbohydrates) • Nitrogen (N2) is solely introduced into biological systems through microbes • Also phosphate cycle, sulfur cycle, etc.

2. Modes of metabolism • Catabolism – nutrient breakdown • Anabolism – macromolecule synthesis • Both are linked via carriers of chemical energy NADH, ATP, NADPH, FADH2 • These sources of chemical energy allow cells to perform “work” (synthesis, etc…)

3. Consider the cell a “system” • Isolated system – cannot exchange energy or matter with its surroundings (not a cell) • Closed system – can exchange energy, but not matter with its surroundings (still not a cell) • Open system – can exchange energy and matter in and out (A Cell!)

4. Internal energy is a state function • The thermodynamic state is defined by prescribing the amounts of all substances present, and two of these variables: temperature (T), Pressure (P), and Volume (V) of the system. • The internal energy (E) of the system reflects all of the kinetic energy of motion, vibration, and rotation and all of the energy contained within chemical bonds and non-covalent interactions

5. How do cells make and use chemical energy? • Bioenergetics must follow the laws of thermodynamics • First Law: the total amount of energy in the universe remains constant; energy may change form or location, but cannot be created or destroyed. • Second Law: Entropy is always increasing

6. First Law of Thermodynamics DE = q – w q = heat; positive q indicates heat is absorbed by the system, negative q indicates heat given off by system w = work; positive w means the system is doing work, negative w means work is being done on the system

7. Oxidation of palmitic acid

8. A “bomb” calorimeter allows reactions to be carried out at constant volume • Because the reaction in (a) is carried out at constant V, no work is done on the surroundings • Therefore, DE = q • In this case, DE = -9941.4 kJ/mole • The negative sign indicates the reaction releases energy stored in chemical bonds and transfers heat to the surroundings

9. Reactions at constant pressure • In reaction (b), the reaction proceeds at 1 atm pressure • The system is free to expand or contract, the final state has contracted because the amount of gas has changed from 23 moles to 16 • The decrease in volume means that work has been done on the system by the surroundings

10. PV work appears as extra heat released • When volume is changed against a constant pressure, w = PDV • Assumptions: constant T, gases are ideal, which allows us to use PV = nRT • w = DnRT = -17.3 kJ/mol • SO, under constant pressure q = DE + w = DE + DnRT = -9941.4 kJ/mol – 17.3 kJ/mol = -9958.7 kJ/mol – In (b) the surroundings can do work on the system, this (PV) work looks like extra heat

11. Most biochemical reactions occur under constant pressure, not constant volume • Because q does not equal DE, we need to account for PV work done • We define a new quantity, enthalpy (H) • H = E + PV DH = DE + PDV • When the heat of a reaction is measured at constant pressure, DH is determined

12. DE and DH measurements are useful for biochemists • Although oxidation of palmitic acid occurs very differently in the human body than in a calorimeter, the values of DE and DH are the same regardless of the pathway • Average human expends ~6000 kJ or roughly 1500 kcal for bodily function, with exercise that figure easily doubles

13. DE, DH, is there a big distinction? • For most chemical reactions the difference between these two quantities is negligible • Typically, PDV is a tiny quantity • For instance, it’s about 0.2% difference for palmitic acid oxidation DH is generally considered a direct measure of the energy change in a process and is the heat evolved in a reaction at constant P

14. Entropy and the second law The minimal value of entropy is a perfect crystal at absolute zero

15. Diffusion is an entropy driven process

16. Increase in entropy can lead to -DG

17. Thermodynamic quantities DH = enthalpy, the heat content of the system exothermic = negative, endothermic = positive; Units: Joules/mole DS = entropy, randomization of energy and matter; positive sign indicates increased entropy; Units: joules/mole(K) DG = Gibbs Free energy, amount of energy that is available to do work at constant T and P; Units: Joules/mole Note 1 calorie = 4.184 Joule

18. Gibbs-Helmholtz equation DG = DH – TDS Positive DG is endergonic, requires energy for reaction to occur, this is unfavorable Negative DG is exergonic, releases energy, this is a favorable process; spontaneous but not necessarily rapid A decrease in energy (-DH) and/or increase in entropy (+DS) make favorable processes DG =0 indicates the system is at equilibrium

19. Thermodynamics of melting ice Ice is a crystal lattice held together by H-bonds, bonds must be broken to form water Energy for breakage of H-bonds is almost entirely the DH for this reaction and this term is positive Entropy favors water over ice But recall DG is also temperature dependent

20. Entropy and Enthalpy contributions to melting ice

21. Biochemical reactions can have different contributions

22. Why is DG called “free” energy? DG represents the portion of an energy change DH that is available or free to do useful work. TDS is amount of energy that is unavailable to do work DG = DH – TDS

23. A DG Warning! • You will see many different DG’s DG – Gibbs Free Energy DG’o or DGo – Standard State Free Energy energy per mole in standard state (1M) DGo – Standard state Free Energy of Activation enzyme catalysis