1. Introduction to Biochemical Principles
Section 1
Introduction to Biochemical Principles

2. Chapter 4
Energy

3. Section 4.1: Thermodynamics
Energy is the basic constituent of the universe
Energy is the capacity to do work
In living organisms, work is powered with the energy provided by ATP (adenosine triphosphate)
Thermodynamics is the study of energy transformations that accompany physical and chemical changes in matter
Bioenergetics is the branch of TD that deals with living organisms

4. Section 4.1: Thermodynamics
Bioenergetics is especially important in understanding biochemical reactions
Reactions are affected by three factors:
Enthalpy (H)—total heat content
Entropy (S)—state of disorder
Free Energy(G)—energy available to do chemical work

5. Section 4.1: Thermodynamics
Three laws of thermodynamics:
First Law of Thermodynamics—Energy cannot be created nor destroyed, but can be transformed
Second Law of Thermodynamics—Disorder always increases
Third Law of Thermodynamics—As the temperature of a perfect crystalline solid approaches absolute zero, disorder approaches zero

6. Section 4.1: Thermodynamics
First two laws are powerful biochemical tools
Thermodynamic transformations occur in a universe composed of a system and its surroundings
Energy exchange between a system and its surroundings can happen in two ways: heat (q) or work (w)
Work is the displacement or movement of an object by force
Figure 4.2 A Thermodynamic Universe

7. Section 4.1: Thermodynamics
First Law of Thermodynamics
Expresses the relationship between internal energy (E) in a closed system and heat (q) and work (w)
Total energy of a closed system (e.g., our universe) is constant
DE = q + w
Unlike a human body, which is an open system
Enthalpy (H) is related to internal energy by the equation: H = E + PV
PV work is usually negligible in biochemical systems, DH is often equal to DE (DH = DE)

8. Section 4.1: Thermodynamics
First Law of Thermodynamics Continued
If DH is negative (DH <0) the reaction gives off heat: exothermic
If is DH positive (DH >0) the reaction takes in heat from its surroundings: endothermic
In isothermic reactions (DH =0) no heat is exchanged
Reaction enthalpy can also be calculated:
DHreaction = SDHproducts  SDHreactants
Standard enthalpy of formation per mole (25°C, 1 atm) is symbolized by DHf°

9. Section 4.1: Thermodynamics
First Law of Thermodynamics Continued
DHreaction = SDHproducts  SDHreactants
Calculate the standard enthalpy of combustion for the following reaction:
C2H5OH (l) + (7/2) O2 (g) ---> 2 CO2 (g) + 3 H2O (l)
ΔH°f CO2 = kJ/mol
ΔH°f H2O = kJ/mol
ΔH°f C2H5OH = kJ/mol
ΔH°f O2 = 0 kJ/mol
Answer? kJ/mol

10. Section 4.1: Thermodynamics
Figure 4.3 A Living Cell as a Thermodynamic System
Second Law of Thermodynamics
Physical or chemical changes resulting in a release of energy are spontaneous
Nonspontaneous reactions require constant energy input

11. Section 4.1: Thermodynamics
As a result of spontaneous processes, matter and energy become more disorganized
Gasoline combustion
The degree of disorder is measured by the state function entropy (S)
Figure 4.4 Gasoline Combustion

12. Section 4.1: Thermodynamics
Second Law of Thermodynamics Continued
Entropy change for the universe is positive for every spontaneous process
DSuniv = DSsys + DSsurr
Living systems do not increase internal disorder; they increase the entropy of their surroundings
For example, food consumed by animals to provide energy and structural materials needed are converted to disordered waste products (i.e., CO2, H2O and heat)
Organisms with a DSuniv = 0 or equilibrium are dead

13. Free energy is the most definitive way to predict spontaneity
Section 4.2: Free Energy
Free energy is the most definitive way to predict spontaneity
Gibbs free energy change or DG
Negative DG indicates spontaneous and exergonic
Positive DG indicates nonspontaneous and endergonic
When DG is zero, it indicates a process at equilibrium
Figure 4.5 The Gibbs Free Energy Equation

14. Standard Free Energy Changes
Section 4.2: Free Energy
Standard Free Energy Changes
Standard free energy, DG°, is defined for reactions at 25°C,1 atm, and 1.0 M concentration of solutes
Standard free energy change is related to the reactions equilibrium constant, Keq, at equilibrium (ΔG = 0)
DG° = -RT ln Keq
Allows calculation of DG° if Keq is known
Because most biochemical reactions take place at or near pH 7.0 ([H+] = 1.0  10-7 M), this exception can be made in the 1.0 M solute rule in bioenergetics
The free energy change, relative to pH, is expressed as DG°′, as follows:
Δ𝐺°𝑓=Δ𝐺°+𝑅𝑇ln 𝐻 +

15. Section 4.2: Free Energy Δ𝐺°=−𝑅𝑇 ln 𝐾 𝑒𝑞
For the reaction HC2H3O2 ↔ C2H3O2 + H+, calculate the standard free energy change at equilibrium (ΔG°) and free energy change (ΔG°f). Assume that T = 25°C, P = 1 atm, R = J/mol•K, and Keq = 1.80 ×10-5. Is this reaction spontaneous?
Δ𝐺°=−𝑅𝑇 ln 𝐾 𝑒𝑞
Δ𝐺°=− 𝐽 𝑚𝑜𝑙•𝐾 298𝐾 ln 1.80 ×10−5 =27.1 kJ/mol
Not spontaneous.
Δ𝐺°𝑓=Δ𝐺°+𝑅𝑇ln 𝐻 +
=27.1× 𝐽 𝑚𝑜𝑙 𝐽 𝑚𝑜𝑙•𝐾 298𝐾 ln 10 −7 =−12.9 𝑘𝐽/𝑚𝑜𝑙
Rxn is spontaneous when pH drops below physiological pH.

16. Section 4.2: Free Energy Coupled Reactions
Figure 4.6 A Coupled Reaction
Coupled Reactions
Many reactions have a positive DG°′
Free energy values are additive in a reaction sequence
If a net DG°′ is sufficiently negative, forming the product(s) is an exergonic process

17. The Hydrophobic Effect Revisited
Section 4.2: Free Energy
The Hydrophobic Effect Revisited
TD principles help us better understand spontaneous aggregation of nonpolar substances
NP substances disrupt water H-bond interactions (which are energetically favorable)
Entropy of water decreases when molecules “cage” the NP molecules
Aggregation of NP molecules decreases the surface area of their contact with water, increasing water’s entropy (i.e., free energy is negative, process is spontaneous)
Spontaneous exclusion of water contributes to membrane formation and protein folding

18. Section 4.3: The Role of ATP
Figure 4.7 Hydrolysis of ATP
Adenosine triphosphate is a nucleotide that plays an extraordinarily important role in living cells
Hydrolysis of ATP  ADP + Pi (orthophosphate) or AMP + PPi (pyrophosphate) provides free energy through transfer of phosphoryl group

19. Section 4.3: The Role of ATP
ATP produced from ADP + Pi.
Drives reactions of several types:
1. Biosynthesis of biomolecules
2. Active transport across membranes
3. Mechanical work such as muscle contraction
Figure 4.8 The Role of ATP

20. Section 4.3: The Role of ATP
Structure of ATP is ideally suited for its role as universal energy currency
Its two terminal phosphoryl groups are linked by phosphoanhydride bonds
Specific enzymes facilitate ATP hydrolysis
Figure 4.9 Structure of ATP

21. Section 4.3: The Role of ATP
Figure 4.10 Transfer of Phosphoryl Groups
The tendency of ATP to undergo hydrolysis is an example of its phosphoryl group transfer potential
ATP acts as energy currency, because it can carry phosphoryl groups from high-energy compounds (e.g., PEP) to low-energy compounds (e.g., glucose, as part of glycolysis).
Process is described in detail in Chapter 8 of the textbook.

22. Section 4.3: The Role of ATP

23. Section 4.3: The Role of ATP
Figure 4.11 Contributing Structure of the Resonance Hybrid of Phosphate
Several factors need to be considered to understand why ATP is so exergonic:
1. At physiological pH, ATP has multiple negative charges
2. Because of resonance stabilization, the products of ATP hydrolysis are more stable than resonance-restricted ATP
Resonance is when a molecule has two or more alternative structures that differ only in the position of their electrons
3. Hydrolysis products of ATP are more easily solvated
4. Increase in disorder with more molecules