Chapter 3- The Energies of Life

1. Chapter 3- The Energies of Life Homework- 2, 3, 5, 9, 12, 15,

2. Bioenergetics • Cells do lots of things • They need energy to do these things • Bioenergetics- the quantitative analysis of how organisms gain and use energy • These processes follow the same fundamental rules of Thermodynamics

3. Basic Definitions • System- any part of the universe that we choose for study • Surroundings-everything outside of the defined system • Isolated system- unable to exchange energy and matter with surroundings • Closed System- able to exchange energy but not matter with surroundings

4. Open System- able to exchange both energy and matter with surroundings • Internal Energy, E- energy contained in any system • Internal energy is a function of state of a system • The thermodynamic state of a system is defined by the amounts of substances present and any two of three variables: P,V,T.

5. Closed and Open Systems, which we mostly deal with, exchange energy with surroundings, so we often talk about DE • For closed systems, this exchange can only occur through heat and/or work. • Work basically occurs when a force is exerted against a resistance to produce displacement.

6. Conventions for exchange of heat and work • Heat is denoted as q. A positive value of q indicates that heat is absorbed by the system from the surroundings. • Work is denoted as w. A positive value of w indicates that work is done by the system on the surroundings.

7. 1st law of Thermodynamics • Internal energy can only change by heat or work, therefore: DE = q - w • Basically a bookkeeping rule • Changes in E, as for any function of state, depend only on the initial and final states of a system and are independent of path.

8. Figure 3.1

9. Although heat and work exchanged with the surroundings depend on path, it is important to remember that DE does not, it only depends on the initial and final states.

10. Enthalpy • Virtually all biochemical reactions occur under conditions similar to the constant pressure example • Because heat does not equal DE in these conditions, we need a function of state, Enthalpy, H. H = E + PV • E, P, and V are functions of state, therefore H is as well

11. Change in Enthalpy DH = DE + PV • The value for DH is the same as the value of q calculated under constant pressure conditions • When the heat of reactions is measured at constant pressure, it is DH that is determined • In Biochemistry, most energy changes are expressed as DH

12. Recall that measures like DH and DE are path independent • That means that although processes in a biochemical system are different than those in a calorimeter, the values are the exact same!! • Lastly, due to conditions, the difference in DH and DE is very small and DH is commonly accepted at the energy change.

13. Entropy and 2nd Law • As important as enthalpy is, it doesn’t tell us the favored direction of a process • To answer this, we look at the reversibility of a process • Reversible processes are always almost at equilibrium • Irreversible processes are far from equilibrium and are driven toward equilibrium

14. Irreversible processes are often called spontaneous processes. • However, the term favorable process is preferred • Knowing whether a process is reversible, favorable or unfavorable is vital to bioenergetics

15. Entropy • The degree of randomness or disorder of a system is measured by a function of state called entropy (S) • Second Law of Thermodynamics- The entropy of an isolated system will tend to increase to a maximum value. • See table 3.1 for examples of low entropy vs high entropy

16. Free Energy: The 2nd law in open systems • All biological systems are open, not isolated as mentioned the in the 2nd law • Both energy and entropy changes will take place in many reactions, and both must be of importance in determining the direction of a favorable process • Gibbs Free Energy is a function of state that includes both energy and entropy

17. Free Energy DG = DH - T DS • So the 2nd Law can be restated: • The criterion for a favorable process in a nonisolated system, at constant pressure and temperature, is that DG be negative. • Positive DG values means a process is not favorable, BUT the reverse process is favorable

18. Processes with negative DG values are said to be exergonic • Processes with positive DG values are said to be endergonic • If DG is zero, the process is not favored to go either direction. • In fact, it is at equilibrium and reversible, meaning it can be pushed either way.

19. The Interplay of Enthalpy and Entropy • For all chemical and physical processes, it is the competition of enthalpy and entropy terms that determine the favorable direction. • Temperature will also play a critical role. • All three components must be considered.

20. H S Low T High T + + G + G – + - G + G + - + G - G – - - G - G +

21. 2 areas of confusion • The favorability of a process has nothing to do with the rate of the process Example: Diamond to Graphite • The entropy of an open system can decrease. Example: Living organisms

22. Not going over, but important • Free energy and useful work • Free energy and concentration • Free energy and chemical reactions

23. Sources of Free Energy • We just said that the processes that living organisms perform are endergonic processes. • So how does these occur? • In essence, we rely on the fact that every PROCESS must be thermodynamically favorable.

24. Reaction Coupling • We can achieve this by coupling reactions that are favorable with those that are not in a way that the net G will be negative.

25. Coupling of endergonic reactions or processes to exergonic reactions is one of the most important principles in biochemistry • Coupling of highly favored to unfavored processes is used not only to drive countless reactions, but also to transport materials across membranes, transmit nerve impulses, contract muscles, and carry out other physical changes.

26. High energy Phosphate Compounds • These compounds undergo reactions with large, negative G values. • This energy must be easily accessible • Most of these involve very simple hydrolysis reactions. • See figure 3.7 page 75 for examples.