Thermodynamics and energetics of life Flashcards
(27 cards)
Principles of Bioenergetics
Metabolism
Catabolism
Exergenic: Reactions involved in breaking down foodstuffs. Energy is released.
Anabolism
Endergenic: Biosynthetic reactions. Energy is required.
Bioenergetics
Quantitative study of the energy transductions that occur in the living cell.
Biological transformations follow the laws of thermodynamics
1st Law of Thermo: Principle of conservation of energy. In any physical or chemical change the total amount of energy in the universe is constant.
2nd Law of Thermo: The total entropy of the universe is always increasing. The universe tends towards more and more disorder.
Living organisms are more highly organized than surrounding materials; organisms maintain and produce order.
Do they violate 2nd law?
Reacting System + surroundings = universe (organism)
Living cells are open systems: exchange both materials and energy. Not at equilibrium with their surroundings.
The order produced within cells as they grow is compensated by the disorder they create in the environment.
By combining the two laws we get:
DG = DH – TDS (eq. 1)
Gibbs Free Energy (Delta G)
Amount of energy capable of doing work during a reaction at T and P are constant
Delta G values relative to reactions
DG < 0 - System releases free energy (exergenic).
- Spontaneous.
- System changes as to posses less free energy.
DG = 0 - System at equilibrium, the rates of the forward and reverse reactions are exactly equal and no further net change occurs in the system
DG > 0 - System gains (endergenic) free energy.
- Non-spontaneous.
Enthalpy (H)
Heat content of the reactive system.
Delta H values relative to reactions
DH < 0 - chemical reaction releases heat: exothermic
- heat content of Product’s < that content of Reactant’s.
DH > 0 - chemical reaction gains heat: endothermic.
DeltaH = DeltaE + PDeltaV (eq. 2)
since in biochemical reactions DeltaV is so small and P is constant then eq. 2 converts to DeltaH = DeltaE
and eq. 1 changes to
DeltaG = DeltaE - TDeltaS (eq. 3)
Entropy (S)
DeltaS is entropy change. Randomness. When the Products are less complex (more disordered) than the reactants: Reaction gains entropy and DeltaS > 0.
If the randomness of a system increases then DeltaS = (+) and if energy is released DeltaH = (-) (ideal for a rxn to take place, positive DeltaS and negative DeltaH). Makes DeltaG negative
Standard free energy
is directly related to the equilibrium constant.
Consider: aA + bB <—> cC+ dD
Keq = [C]c [D]d /
[A]a [B]b
DG = DGo’ + RT ln Keq
When a reaction is not at equilibrium, the tendency to move towards equilibrium represents a driving force the magnitude of which can be expressed as Free energy change for the reaction, DG
Standard Conditions
T = 298 K (25 oC) initial []’s of P’s and R’s = 1 M
Standard free energy change, DGo
The force driving the system towards equilibrium
By this definition, [H+1] or pH = 0
Since most biochem. reactions occur where [H+] = 10-7
or pH = 7, we denote a new way to express DG
DGo’: biochemical standard state
For a rxn at equillibrium
There is no net change because the free energy of the forward reaction exactly balances that of the reverse reaction. Consequently, DG = 0. System is at equilibrium and no net change can take place.
DGo’ = - RT ln K’eq
DGO’ and Keq are constants for each reaction
R = gas constant = 1.987 cal/mol K or 8.31 J/mol K
The standard free energy change (DGo’) of a reaction is simply an alternative mathematical way of expressing its equilibrium constant. It is constant for every reaction.
DGo’ is a constant for every reaction.
When K’eq > 1 DGo’ = (-): P’s contain less free energy than R’s under standard conditions reaction is shifted to the right.
Reaction proceeds spontaneously to form P’s. [P’s]»_space;>[R’s]
When K’eq < 1 DGo’ = (+): P’s contain more free energy than R’s under standard conditions reaction is shifted to the left.
Reaction is not spontaneous. [R’s]»_space;>[P’s]
In regards to a chemical rxn
All chemical reactions tend to go in the direction that results in a decrease in the free energy of the system. The system is more stable when the free energy content is lowest.
The DGo’ of any reaction proceeding spontaneously toward its equilibrium is always (-), becomes less (-) as the reaction proceeds, and is zero at equilibrium, indicating that no more work can be done by the reaction.
Actual Free-Energy Changes Depend on Reactant and Product Concentrations
Each chemical reaction has a characteristic standard free-energy change, which may be (+), (-), or zero, depending on the equilibrium constant of the reaction.
It tells us in which direction and how far a given reaction must go to reach equilibrium when the initial concentration of each component is 1.0 M, the pH is 7.0, the T is 25 C, and the P is 1 atm. (standard conditions).
Thus DG0’ is a constant for a given reaction
The criterion for spontaneity for a reaction is DG not DGo’
But the actual free-energy change, DG, is a function of reactant and product []’s and of the T prevailing during the reaction, which will not necessarily match the standard conditions as defined above.
DG and DGo’ for any reaction A + B <—> C + D are related by the equation
DG = DGo’ + RT ln [C][D] /
[A] [B]
in which the terms in red (after ln) are those actually prevailing in the system under observation.
DG varies with []’s of R and P and with T and pH.
The criterion for spontaneity of a reaction is the value of DG, not DGo’.
A reaction with a positive DGo’ can go in the forward direction if DG is negative.
This is possible if the term RT ln ([products]/[reactants]) in equation is negative, [P]«[R], and has a larger absolute value than DGo’.
For example, the immediate removal of the products of a reaction can keep the ratio [products]/[reactants] well below 1, such that the term RT ln([products]/[reactants]) has a large, negative value.
Thermodynamics in Cells
Thermodynamics in Cells
Driving energy: Minimum energy/Maximum entropy.
Need negative DG to go forward with a reaction
If DGo’ is negative but small, the [R’s] may affect direction (For/Rev).
If the negative DGo’ is large, the reaction is irreversible since the []’s in a cell cannot be altered to a large degree (10-3 to 10-4 M).
Why is it important?
Thermodynamics determines which direction the pathway goes.
Regulate flow of metabolites along pathways.
Need irreversible reactions to prevent “futile” cycles.
STAGES IN THE EXTRACTION OF ENERGY (CATABOLISM) FROM FOODSTUFFS
Stage 1: Large biopolymers are broken down into smaller units (nonomers). No useful energy generated.
Stage 2: Monomers are degraded to a few simple units (pyruvate, acetyl-CoA or Krebb cycle intermediates) that play a central role in metabolism.
Stage 3: Energy producing stage. Oxidation of acetyl-CoA and production of ATP through oxpho. Production of 3 principal end products, NH3, CO2 and H2O.
OXIDATION AND REDUCTION
Oxidation: loss of e’s Addition of oxygen. Loss of hydrogen.
Reduction: gain of e’s. Addition of hydrogen. Loss of oxygen.
Oxidation is the source of metabolic energy.
O2 is the electron acceptor, reduced to H2O
Glc + 6O2 —> 6CO2 + H2O
DGo = - 2870 kJ/mol
Glc is oxidized to CO2 and H2O. The more reduced a compound is, the higher the amount of energy it contains.
In Biological Systems
Oxidation of fuel molecules yields e’s that are transferred to O2 in aerobic organisms.
Electrons are not transferred directly to O2, instead there are a Series of coupled redox reactions.
Carriers such as NAD + and FAD carry electrons which eventually get transferred to oxygen. Called e- carriers.
ATP: UNIVERSAL ENERGY CURRENCY
ATP (adenosine triphosphate). Carrier of free energy
Releases large amounts of Gibbs free energy upon hydrolysis.
Active form complex with Mg+2 or Mn+2
Triphosphate form contains 2 phosphoanhydride bonds.
Has a group transfer potential which is intermediate among biologically important phosphorylated molecules. Function as an efficient carrier of phosphorylated groups.
