Enzymes, Metabolism, and Energy Transformations
Enzymes and Metabolic Pathways
Enzymes Defined:
All enzymes are proteins.
However, not all proteins are enzymes.
Reaction Pathways - Stepped Processes:
Most biological reactions do not go from one reactant directly to a final product in a single step; they are broken down into multiple steps.
Benefits of Multi-Step Reactions:
Energy Harnessing: Allows for the gradual extraction of energy from molecules. Trying to extract all energy at once from a high-energy molecule (like in combustion) would be explosive and difficult to harness effectively.
Control Points: Provides multiple opportunities to regulate the process by turning specific enzymes on or off at different steps.
Catabolic vs. Anabolic Pathways
Catabolic Pathways:
Definition: Processes that break down complex molecules into simpler ones.
The root word "cata" means "falling apart" or "breaking apart" (e.g., catastrophe).
Function: Essentially digestive metabolism.
Examples:
Digestion of food in the stomach or intestines.
Cellular respiration: breaking down glucose into ATP (e.g., mitochondria catabolizing glucose).
Human body utilizes approximately its own body mass in ATP daily (not molecules but mass).
Energy Changes: Typically releases energy, often with a negative free energy change ($\Delta G$), making them spontaneous.
Anabolic Pathways:
Definition: Processes that consume energy to build complex molecules from simpler ones.
The root word "ana" means "building up" (e.g., anabolic steroids aim to build muscle).
Function: Constructive metabolism.
Examples:
Taking anabolic steroids (though often illegal/harmful in excess, steroids are natural in the body).
Dehydration synthesis: A process where smaller molecules are joined to form larger ones, releasing water.
Energy Changes: Requires energy input, meaning they typically have a positive free energy change ($\Delta G$) and are non-spontaneous.
Types of Energy
Kinetic Energy:
The energy of motion.
Thermal Energy (Heat): The kinetic energy associated with the random movement of molecules. Measuring average kinetic motion or thermal energy reflects this.
Example: A person jumping off a diving board converts potential energy (at height) into kinetic energy (as they fall and speed up).
The splash into water dissipates kinetic energy, producing sound/vibrational energy.
Light Energy:
A form of energy that can be harnessed to do work (e.g., plants converting sunlight into chemical energy via photosynthesis).
Thermodynamics
Definition: The study of energy transformations.
Systems and Surroundings:
Open Systems: Exchange both energy and matter with their surroundings.
Organisms are open systems: We take in air, food, etc., and release waste and energy.
Isolated Systems: Do not exchange energy or matter with their surroundings.
Organisms cannot be isolated systems; inability to exchange means inability to sustain life.
Second Law of Thermodynamics (Entropy):
States that every energy transfer or transformation increases the entropy (disorder or randomness) of the universe overall.
Even though organisms maintain an "island of low entropy" (highly organized structures), the metabolic reactions required to build and maintain this order contribute to an overall increase in the universe's entropy.
Spontaneous Processes and Entropy:
Processes that increase entropy (disorder) can occur without energy input (e.g., spraying perfume spontaneously diffuses throughout a room, even without air currents, though air currents speed it up).
This diffusion happens because molecules spread out, increasing randomness.
Organisms decrease their own internal entropy to maintain organization (e.g., highly structured cells, ordered biological molecules).
However, this local decrease in entropy within the organism happens at the expense of a greater increase in the entropy of the universe.
Free Energy Change ($\Delta G$)
Definition: The portion of a system's energy that can perform work when temperature and pressure are uniform throughout the system (standard biological conditions).
It essentially represents the amount of organization able to continuously do work, as part of a system tending towards entropy.
Calculation: The change in free energy ($\Delta G$) is given by the final state's free energy minus the initial state's free energy:
Historical Context: Prior to approximately 2019/2020, AP Biology curriculum required calculating $\Delta G$ (Gibbs free energy); now, the focus is more on understanding the concept.
Negative $\Delta G$ (Exergonic Reactions)
Characteristics:
Occurs when the final state of a reaction has lower free energy than the initial state.
Represented by a downhill energy profile (e.g., a skateboarder going down a hill).
Spontaneous: These reactions occur without a net input of energy.
Catabolic Reactions: Often have a negative $\Delta G$ because they break down high-energy, structured molecules into lower-energy components.
Energy Release: Releases free energy; the magnitude of $\Delta G$ represents the maximum amount of energy available to do work.
Examples: Most processes involving breakdown (like the wearing out of shoes, or the decomposition of objects) are spontaneous and involve a negative $\Delta G$.
Cannot spontaneously create order (e.g., shoes don't spontaneously become brand new).
Coupling: The energy released from an exergonic reaction can be used to drive an endergonic reaction (e.g., energy from the sun drives photosynthesis to build sugars).
Positive $\Delta G$ (Endergonic Reactions)
Characteristics:
Occurs when the final state of a reaction has higher free energy than the initial state.
These reactions absorb free energy from their surroundings.
Non-spontaneous: They will not occur by themselves without a continuous input of energy.
Anabolic Reactions: Typically have a positive $\Delta G$ because they build complex, higher-energy molecules from simpler ones.
Example: A car breaking down over time (negative $\Delta G$) but never spontaneously rebuilding itself (which would require a positive $\Delta G$).
Equilibrium
At equilibrium, the rates of the forward and reverse reactions are equal. There is no net change in free energy, meaning $\Delta G = 0$. This can be visualized as being on a flat path, with no energetic uphill or downhill. Consequently, a system at equilibrium can do no work.
Exergonic vs. Endergonic Reactions (Chemistry Application)
These terms are similar to the endothermic/exothermic concepts in chemistry (which deal with heat transfer).
Exergonic Reactions:
Energy Profile: Reactants have higher free energy than products.
Energy Change: Releases energy.
Spontaneity: Spontaneous.
$\Delta G$: Negative.
The energy released can be utilized by other processes; for example, coupling these reactions to drive endergonic (energy-requiring) reactions.
Endergonic Reactions:
Energy Profile: Products have higher free energy than reactants.
Energy Change: Absorbs energy from the surroundings.
Spontaneity: Non-spontaneous (requires energy input).
$\Delta G$: Positive.
Example: Building complex molecules like proteins or nucleic acids, which require energy to form higher-energy, more ordered structures.
Enzymes are proteins, and their function is highly dependent on their specific three-dimensional structure. When an enzyme is exposed to a pH outside its optimal range, the concentration of hydrogen ions () or hydroxide ions () in the solution can alter the charges on the amino acid side chains within the enzyme. This change in charge can disrupt the weak ionic and hydrogen bonds that maintain the enzyme's precise tertiary and quaternary structures. This disruption leads to a change in the enzyme's shape, particularly at its active site, a process known as denaturation. Once denatured, the enzyme's active site may no longer be able to bind effectively to its substrate, or it may lose its catalytic efficiency, thereby impairing or completely destroying its function. This loss of function is often irreversible if the pH change is extreme or prolonged.