1. Thermodynamics in Biology
Lecture 24
Thermodynamics in Biology
From a University of Wisconsin lecture course

2. A Simple Thought Experiment
Just think of the biochemistry that occurs when one E. coli cells becomes one billion cells in 48 hours from simple material contained in a solution.

3. Driving Forces for Natural Processes
Enthalpy
Tendency toward lowest energy state
Form stablest bonds
Entropy
Tendency to maximize randomness
Energy and order.

4. Enthalpy and Bond Strength
Enthalpy = ∆H = heat change at constant pressure
Units
cal/mole or joule/mole
1 cal = 4.18 joule
Sign
∆H is negative for a reaction that liberates heat
Heat is a form of energy.

5. Entropy and Randomness
This goes against the grain; must supply lots of energy to take 153 free amino acids and convert them into the ordered sequence that is in myoglobin.

6. Entropy and Randomness
Entropy = S = measure of randomness
cal/deg·mole
T∆S = change of randomness
For increased randomness, sign is “+”
Order versus randomness. Just think of the energy input that is required to convert the random biochemical facts in your mind into a metabolic chart - don’t let entropy overcome you.

7. “System” Definition
If your brain is a closed, you can not gain knowledge. Consider how your attitude influences what you can do.

8. “System” Definition
When isolated, you can exchange energy.

9. “System” Definition
When open, both mass and energy are exchanged in a system.

10. Cells and Organisms: Open Systems
Material exchange with surroundings
Fuels and nutrients in (glucose)
By-products out (CO2)
Energy exchange
Heat release (fermentation)
Light release (fireflies)
Light absorption (plants)
Life occurs in an open system.

11. 1st Law of Thermodynamics
Energy is conserved, but transduction is allowed
Transduction
Energy is conserved; can change forms.

12. 2nd Law of Thermodynamics
In all spontaneous processes, total entropy of the universe increases
Randomness increase. A fact of life, just look on how messy your room gets.

13. 2nd Law of Thermodynamics
∆Ssystem + ∆Ssurroundings = ∆Suniverse > 0
A cell (system) can decrease in entropy only if a greater increase in entropy occurs in surroundings
C6H12O6 + 6O2  6CO2 + 6H2O
complex simple
Balanced equation.

14. Entropy: A More Rigorous Definition
From statistical mechanics:
S = k lnW
k = Boltzmann constant = 1.3810–23 J/K
W = number of ways to arrange the system
S = 0 at absolute zero (-273ºC)
Definition.

15. Gibbs Free Energy Unifies 1st and 2nd laws ∆G ∆G = ∆H – T∆S
Useful work available in a process
∆G = ∆H – T∆S
∆H from 1st law
Kind and number of bonds
T∆S from 2nd law
Order of the system
Gibbs free energy connects the first and second laws of thermodynamics.

16. ∆G Driving force on a reaction
Work available  distance from equilibrium
∆G = ∆H – T∆S
State functions
Particular reaction T P
Concentration (activity) of reactants and products
Free energy is the driving force for reactions.

17. Equilibrium ∆G = ∆H – T∆S = 0 So ∆H = T∆S
∆H is measurement of enthalpy
T∆S is measurement of entropy
Enthalpy and entropy are exactly balanced at equilibrium
Equilibrium balanced.

18. Effects of ∆H and ∆S on ∆G
Variations
Voet, Voet, and Pratt. Fundamentals of Biochemistry

19. Standard State and ∆Gº Arbitrary definition, like sea level
[Reactants] and [Products]
1 M or 1 atmos (activity)
T = 25ºC = 298K
P = 1 atmosphere
Standard free energy change = ∆Gº
Standard state, but not physiological.

20. Biochemical Conventions: ∆Gº
Most reactions at pH 7 in H2O
Simplify ∆Gº and Keq by defining [H+] = 10–7 M
[H2O] = unity
Biochemists use ∆Gº and Keq
Biochemical conventions (assumptions).

21. Relationship of ∆G to ∆Gº
∆G is real and ∆Gº is standard
For A in solution
GA = GA + RT ln[A]
For reaction aA + bB  cC + dD
∆G = ∆Gº + RT ln
Constant Variable
(from table) º
[C]c [D]d }
Real & standard.
[A]a [B]b

22. Relationship Between ∆Gº and Keq
[C]c [D]d
∆G = ∆Gº + RT ln
At equilibrium, ∆G = 0, so
∆Gº = –RT ln
∆Gº = –RT ln Keq
[A]a [B]b
[C]c [D]d
Relationship between equilibrium constant.
[A]a [B]b

23. Relationship Between Keq and ∆Gº
Equilibrium position and energy

24. Will Reaction Occur Spontaneously?
When:
∆G is negative, forward reaction tends to occur
∆G is positive, back reaction tends to occur
∆G is zero, system is at equilibrium
∆G = ∆Gº + RT ln
Spontaneous
[C]c [D]d
[A]a [B]b

25. A Caution About ∆Gº
Even when a reaction has a large, negative ∆Gº, it may not occur at a measurable rate
Thermodynamics
Where is the equilibrium point?
Kinetics
How fast is equilibrium approached?
Enzymes change rate of reactions, but do not change Keq
Caveats

26. ∆Gº is Additive (State Function)
Reaction
A  B
B  C
Sum: A  C
Also: B  A
Free energy change
∆G1º
∆G2º
∆G1º + ∆G2º
– ∆G1º
Additive

27. Coupling Reactions Glucose + HPO42–  Glucose-6-P ATP  ADP + HPO42–
ATP + Glucose  ADP + Glucose-6-P
∆Gº
kcal/mol kJ/mol
– –30.5
– –16.7
A biochemical key - how to get unfavorable to occur couple favorable with unfavorable.

28. Resonance Forms of Pi – – – –
Stabilization by resonance gives higher energy. –

29. Phosphate Esters and Anhydrides
Anhydride - high energy bonds; esters ordinary energy.

30. Hydrolysis of Glucose-6-Phosphate
Glucose 6-phosphatase
∆Gº = –3.3 kcal/mol
= –13.8 kJ/mol