An Introduction to Energy and Metabolism in Biological Systems

An Introduction to Energy in Biological Systems

Key Concepts and Definitions

• Energy: The capacity to cause change, perform work, or move matter against opposing forces. Energy exists in various forms and is fundamental to all biological processes.

• Kinetic Energy: The energy of motion. Examples include a moving car, an antelope running, or the heat and light produced by a fire.

• Potential Energy: Stored energy due to position or structure. Examples include water held behind a dam (gravitational potential energy) and the chemical energy stored in the bonds of an apple (chemical potential energy).

• Heat: A form of kinetic energy associated with the random movement of atoms or molecules.

• Chemical Reaction: A process that transforms one set of chemical substances (reactants) into another (products).

• Reactant: A starting substance in a chemical reaction.

• Product: A substance formed as a result of a chemical reaction.

Types of Work Performed by Cells

• Mechanical work (e.g., muscle contraction, movement of cilia and flagella)

• Transport work (e.g., pumping substances across membranes against concentration gradients)

• Chemical work (e.g., synthesis of polymers from monomers)

Thermodynamics in Biological Systems

The Laws of Thermodynamics

• First Law of Thermodynamics: Energy cannot be created or destroyed, only transformed or transferred. The total energy of the universe remains constant.

• Second Law of Thermodynamics: Every energy transfer or transformation increases the entropy (disorder) of the universe. Some energy is always lost as heat, making processes less efficient. This is sometimes called the "you always lose rule."

Example: When cells convert chemical energy from food into kinetic energy for movement, some energy is lost as heat, increasing entropy.

Spontaneous Processes

• A spontaneous process occurs without the input of additional energy once it has started. These processes are typically exergonic (energy-releasing).

Free Energy and Chemical Reactions

Free Energy (G)

• Free Energy (G): The portion of a system's energy that is available to do work at constant temperature and pressure.

• Symbol: G

• The change in free energy during a reaction is denoted as ΔG.

Exergonic vs. Endergonic Reactions

• Exergonic Reaction: Releases free energy; ΔG is negative. These reactions occur spontaneously.

• Endergonic Reaction: Requires an input of energy; ΔG is positive. These reactions are non-spontaneous.

Examples:

• Cellular Respiration: Exergonic; ΔG is negative.

• Photosynthesis: Endergonic; driven by light energy from the sun.

Determining Spontaneity of Reactions

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

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

Equation for Free Energy Change:

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

• Where ΔH is the change in enthalpy (total energy), T is the temperature in Kelvin, and ΔS is the change in entropy.

Equilibrium and Metabolism

• At equilibrium, the rates of the forward and reverse reactions are equal, and there is no net change in free energy.

• Cells maintain a state of metabolic disequilibrium to keep life processes running; equilibrium would mean death for a cell.

Activation Energy and Catalysis

Activation Energy (EA)

• Activation Energy (EA): The minimum amount of energy required to start a chemical reaction. It is needed to reach the transition state, where bonds are unstable and ready to be broken or formed.

Catalysts and Enzymes

• Catalyst: A substance that increases the rate of a chemical reaction without being consumed in the process.

• Enzymes: Biological catalysts that lower the activation energy barrier, allowing reactions to proceed more rapidly and at lower temperatures.

Why Spontaneous Reactions May Not Occur Instantly

• Even if a reaction is spontaneous (ΔG < 0), it may not occur quickly because it must overcome the activation energy barrier.

• Enzymes help overcome this barrier, enabling essential biochemical reactions to occur at rates compatible with life.

Summary Table: Types of Energy and Examples

Type of Energy	Description	Example
Kinetic Energy	Energy of motion	Running antelope, heat from fire
Potential Energy	Stored energy due to position or structure	Water behind a dam, chemical bonds in apple

Key Terms Glossary

• Energy: Capacity to do work or cause change.

• Kinetic Energy: Energy of motion.

• Potential Energy: Stored energy.

• First Law of Thermodynamics: Energy conservation law.

• Second Law of Thermodynamics: Entropy increases in energy transfers.

• Heat: Random movement of particles.

• Reactant: Starting material in a reaction.

• Product: Resulting material from a reaction.

• Free Energy (G): Energy available to do work.

• Chemical Reaction: Transformation of substances.

• Equilibrium: State where forward and reverse reactions are balanced.

• Endergonic Reaction: Energy-requiring reaction (ΔG > 0).

• Exergonic Reaction: Energy-releasing reaction (ΔG < 0).

• Spontaneous Reaction: Occurs without added energy once started.

• Activation Energy (EA): Minimum energy to start a reaction.

• Transition State: Unstable state during a reaction when bonds are being broken/formed.

• Catalyst: Substance that speeds up a reaction without being consumed.

Additional info: For further understanding, students are encouraged to review textbook sections on energy transformations, exergonic and endergonic reactions, and enzyme function, as well as animations referenced in the study guide.