1. Figure 6.1 The complexity of metabolism

2. METABOLISM Metabolism – totality of chemical reactions in cells
Catabolic pathways- breakdown molecules to produce energy
Anabolic pathways- build molecules by consuming energy

3. Figure The ATP cycle

4. Unit 2 – How do cells regulate chemical reactions?
Controlling reactions requires controlling thermodynamics and kinetics
Thermodynamics – can a reaction happen? How much energy will need to be added or will be released?
Kinetics – how fast will a reaction happen? (Determined by height of Ea); spontaneous rxns can be very slow

5. Figure 6.12 Energy profile of an exergonic reaction
Link
Free Energy (G) – which is favored at equilibrium – reactants or products? How much work can be obtained from a rxn?

6. Spontaneous Processes
Processes that are spontaneous in one direction are nonspontaneous in the reverse direction.
Spontaneous ≠ Fast

7. Key Concepts of Thermodynamics
Enthalpy, H - heat energy change in a reaction (at constant pressure); the result of making and breaking chemical bonds
Entropy, S - measure of disorder or # of possible arrangements of particles (microstates)
Gibb’s Free Energy, G – amount of energy available to do work; sign of ΔG determines whether a process will occur spontaneously

8. 3 Laws of Thermodynamics
1st Law (Conservation of Energy) – energy is neither created nor destroyed, only transformed.
2nd Law – The Entropy of the Universe always increases.
3rd Law – Defn of Zero Entropy – perfect crystalline solid at 0 K.

9. First Law of Thermodynamics “You can’t win”
Energy can be neither created nor destroyed; it can only be transformed from one form to another
Example: Transformation of potential energy -> kinetic energy
Chemical potential energy – energy stored in chemical bonds
Biological Example – Cellular respiration releases stored in energy in glucose, converts it into work and heat

10. Entropy Entropy(S)- measure of randomness or disorder
Mathematically entropy is determined by the # of arrangements or microstates possible

11. 3rd Law of Thermodynamics – Definition of Zero Entropy
Zero entropy = perfect crystalline solid at 0 K (absolute zero) – only 1 possible microstate
Microstate =
Atom =
Perfect crystal at 0 K = 1 microstate

12. Third Law of Thermodynamics
The entropy of a pure crystalline substance at absolute zero is 0.

13. ↑Microstates, ↑ Entropy
Atom =
0 K - 1 microstate
0.001 K – 3 microstates(entropy increasing)
0.002 K – 5 microstates (entropy increasing)

14. .001 K – 3 microstates (entropy ↑)
Atom =
Arrangement 1:
Arrangement 2:
Arrangement 3:

15. Entropy, letters and Shakespeare’s Julius Caesar
High Entropy
Low Entropy - Rich in Meaning and Symbolism
There is a tide in the affairs of men, Which, taken at the flood, leads on to fortune; Omitted, all the voyage of their life Is bound in shallows and in miseries.

16. Entropy and Molecules of Life
CONSIDER 3 AMINO ACIDS: Ala, Gly, Lys
Ala Gly Lys vs Ala – Gly - Lys
(separated) (bonded)
Bonded together = much more order, especially if a specific sequence is required
KEY CONCEPT: Biological molecules have extremely low entropy (highly ordered) because the components have to be bonded in a very specific sequence and have a very specific 3-D shape – conformation.

17. Entropy and the Second Law of Thermodynamics – “You can’t even break even”
The natural tendency in the universe is towards maximum disorder.
For any spontaneous process, the entropy of the universe increases
∆Suniv = ∆Ssystem + ∆Ssurroundings
2nd Law: ∆Suniv is always +

18. Does the existence of Life Violate the 2nd Law of Thermodynamics?
No. ∆Suniv = ∆Ssystem + ∆Ssurroundings
∆Suniv is always +.
∆Ssurroundings > ∆Ssystem
Heat flowing out of living organism increases entropy of its surroundings.
Metabolizing Food is the price we pay for overcoming entropy to produce order. Heat released ↑ entropy of surroundings.

19. Figure 6.4 Order as a characteristic of life

20. Does Entropy explain Death?
All living organisms that have ever existed on earth have eventually died.
There seems to be a limit as to the # of times cells can divide.
One theory – accumulated mutations (disorder!?) in DNA eventually causes too much chaos to sustain life

21. Gibbs Free Energy
If DG is negative, the forward reaction is spontaneous.
If DG is 0, the system is at equilibrium.
If G is positive, the reaction is spontaneous in the reverse direction.

22. Gibb’s Free Energy, G ∆G = Change in Free Energy for a reaction
Sign of ∆G indicates whether a process is spontaneous under a given set of conditions
Value of ∆G indicates the maximum amount energy of theoretically available to do work
Actual amount of work done is always less because some energy will be lost as heat (2nd Law of Thermodynamics)

23. Free Energy Changes Very key equation:
This equation shows how G changes with temperature.
(We assume S° & H° are independent of T.)

24. Free Energy and Temperature
There are two parts to the free energy equation:
H— the enthalpy term
TS — the entropy term (T in Kelvin)
The temperature dependence of free energy comes from the entropy term.

25. Spontaneous Processes
Processes that are spontaneous at one temperature may be nonspontaneous at other temperatures.
Above 0C it is spontaneous for ice to melt.
Below 0C the reverse process is spontaneous.

26. Sign of ∆G Predicts whether a process will occur spontaneously
If ∆G = -, rxn is spontaneous (can occur without outside intervention)
If ∆G = +, rxn is not spontaneous
If ∆G = 0, system is at equilibrium- no work can be obtained from reactions
Note: spontaneous does not necessarily mean fast; speed is the realm of kinetics

27. Figure 6.6 Energy changes in exergonic and endergonic reactions
Measure of maximum amount of work that can be obtained

28. Figure 6.5 The relationship of free energy to stability, work capacity, and spontaneous change