Week 4 - Sept 24

Energy and Enzymes

Introduction to Energy in Biology

All living organisms require energy to perform the essential processes of life. Energy transformations and the regulation of these processes are central to metabolism and cellular function.

• Energy is the capacity to do work.

• Two main types of energy: Potential energy (stored energy due to position or structure) and Kinetic energy (energy of motion).

Types of Energy

Potential Energy

Potential energy is stored in chemical bonds and is determined by the arrangement of electrons in molecules.

• Chemical potential energy is found in the bonds of molecules such as glucose.

• The potential energy in a covalent bond depends on the position of shared electrons.

• If shared electrons are far from the nuclei of both atoms, the bond has higher potential energy (e.g., C–H bonds).

• Equal sharing (nonpolar bonds) = weaker, higher energy bonds; Unequal sharing (polar bonds) = stronger, lower energy bonds.

Kinetic Energy

Kinetic energy is the energy of motion, such as the movement of molecules, light, or objects.

• Examples: Flowing water, sunlight, moving vehicles.

• Energy can be converted from one form to another (e.g., potential to kinetic).

Thermodynamics in Biological Systems

The First Law of Thermodynamics

The first law, also known as the law of conservation of energy, states that energy cannot be created or destroyed, only transformed or transferred.

• Energy is conserved in all processes.

• Energy can be transferred between objects or transformed into different forms (e.g., chemical to heat).

The Second Law of Thermodynamics

The second law states that the entropy (disorder) of a system and its surroundings will increase over time unless energy is input to maintain order.

• Energy spontaneously disperses from being localized to becoming more spread out if not hindered.

• Entropy (S) is a measure of disorder; systems tend toward higher entropy.

Examples of Increasing Entropy

• Increase in volume (e.g., phase change from solid to gas).

• Increase in number of molecules (e.g., breakdown of a macromolecule).

• Increase in molecular random motion (e.g., diffusion, heat transfer).

Enthalpy, Entropy, and Free Energy

Enthalpy (H)

Enthalpy is the total energy in a molecule, including both potential and kinetic energy.

• Changes in enthalpy are denoted as ΔH.

• In biological systems, changes in enthalpy often approximate changes in potential energy.

Entropy (S)

Entropy measures the amount of disorder or randomness in a system.

• Energy tends to disperse or spread out, increasing entropy.

Gibbs Free Energy (G)

Gibbs free energy is the portion of a system's energy available to do work. The change in free energy (ΔG) determines whether a reaction is spontaneous.

• The equation for Gibbs free energy is:

$\Delta G = \Delta H - T\Delta S$

• If ΔG < 0, the reaction is exergonic (spontaneous, releases energy).

• If ΔG > 0, the reaction is endergonic (non-spontaneous, requires energy input).

Calculating Free Energy Change

• $\Delta G = \text{Free energy in products} - \text{Free energy in reactants}$

Exergonic and Endergonic Reactions

Exergonic Reactions

Exergonic reactions have a negative free energy change and occur spontaneously.

• Release energy that can be used to do work.

• Products have lower enthalpy and/or higher entropy than reactants.

$\Delta G = \Delta H - T\Delta S < 0$

Endergonic Reactions

Endergonic reactions have a positive free energy change and require an input of energy to proceed.

• Products have higher enthalpy and/or lower entropy than reactants.

$\Delta G = \Delta H - T\Delta S > 0$

Example: Combustion of Glucose

The combustion of glucose is a spontaneous, exergonic reaction involving changes in enthalpy and entropy:

• $C_6H_{12}O_6 + 6O_2 \rightarrow 6CO_2 + 6H_2O$

• Releases energy as heat and increases disorder.

$\Delta G = \Delta H - T\Delta S$

Metabolism: Catabolism and Anabolism

Overview of Metabolism

Metabolism is the sum of all chemical reactions in a cell, divided into two main types:

• Catabolic reactions: Breakdown complex molecules into simpler ones, releasing energy (exergonic).

• Anabolic reactions: Use simple molecules to build more complex ones, requiring energy input (endergonic).

Catabolic Reactions	Anabolic Reactions
Breakdown of molecules (e.g., starch to glucose)	Synthesis of molecules (e.g., protein synthesis)
Release energy (exergonic)	Require energy (endergonic)

Energy Coupling in Cells

Cells power endergonic reactions by coupling them with exergonic reactions, often using energy released from catabolic pathways to drive anabolic pathways.

• Energy released by catabolic reactions is used to power anabolic reactions.

• Catabolism generates small, diffusible molecules (e.g., ATP, NADH) that act as energy shuttles.

Energy Transfer: Phosphate and Electron Carriers

Energy is transferred in cells via phosphate groups (e.g., ATP) or electrons (e.g., NADH, FADH2).

• Active carrier molecules (ATP, NADH, FADH2) store and transfer energy between reactions.

• Redox reactions (oxidation-reduction) are important for electron transfer.

ATP: The Energy Currency of the Cell

Structure and Function of ATP

ATP (adenosine triphosphate) stores a large amount of potential energy in its phosphate bonds.

• Composed of adenine, ribose, and three phosphate groups.

• Clustered negative charges on phosphate groups raise the potential energy.

ATP Hydrolysis

Hydrolysis of ATP releases energy that can be used to drive cellular processes.

• ATP + H2O → ADP + Pi + energy

• Standard free energy change: $\Delta G = -30.5$ kJ/mol (or $-7.3$ kcal/mol)

Coupling ATP Hydrolysis to Endergonic Reactions

ATP hydrolysis is often coupled to endergonic reactions to make them energetically favorable (exergonic overall).

• Example: Synthesis of glutamine from glutamic acid and ammonia is endergonic ($\Delta G = +3.4$ kcal/mol).

• Coupling with ATP hydrolysis ($\Delta G = -7.3$ kcal/mol) makes the overall reaction exergonic ($\Delta G = -3.9$ kcal/mol).

Mechanism: Phosphorylation

The key to coupling is the transfer of a phosphate group from ATP to a substrate (phosphorylation), increasing the substrate's free energy and reactivity.

• Phosphorylation often causes a change in the shape and activity of the substrate.

• Exergonic phosphorylation reactions are coupled to endergonic reactions in metabolism.

Summary Table: Key Concepts in Energy and Metabolism

Concept	Definition	Example
Potential Energy	Stored energy due to position or structure	Chemical bonds in glucose
Kinetic Energy	Energy of motion	Flowing water, moving molecules
Enthalpy (H)	Total energy in a molecule	Measured as heat content
Entropy (S)	Disorder/randomness in a system	Diffusion, phase changes
Gibbs Free Energy (G)	Energy available to do work	Cellular reactions
Exergonic Reaction	Releases energy, spontaneous	Cellular respiration
Endergonic Reaction	Requires energy input	Protein synthesis
ATP	Energy currency of the cell	Drives biosynthetic reactions

Key Equations

• Gibbs Free Energy: $\Delta G = \Delta H - T\Delta S$

• ATP Hydrolysis: $\text{ATP} + \text{H}_2\text{O} \rightarrow \text{ADP} + \text{P}_i + \text{energy}$

Conclusion

Understanding energy transformations, thermodynamic principles, and the role of ATP is essential for grasping how cells power life’s processes. These concepts form the foundation for studying metabolism, cellular respiration, and the regulation of biochemical pathways.