Metabolism: Overview

Definition and Scope

Metabolism refers to the sum of all chemical reactions occurring within a living cell or organism. These reactions are organized into biosynthetic (anabolic) and degradative (catabolic) pathways, which together maintain cellular function and homeostasis.

• Anabolic pathways build complex molecules from simpler ones, requiring energy input.

• Catabolic pathways break down complex molecules into simpler ones, releasing energy.

• Metabolic pathways can converge, diverge, or form cycles (e.g., the citric acid cycle).

Example: Glycolysis is a catabolic pathway that breaks down glucose to produce ATP.

Bioenergetics in Biochemistry

Energy Transformations and Thermodynamics

Bioenergetics is the quantitative study of energy transformations in living cells, focusing on the nature and function of chemical processes.

• Energy is required for biosynthetic reactions and released during degradative reactions.

• Understanding energy flow is essential for predicting reaction direction and feasibility.

First and Second Laws of Thermodynamics

• First Law: Conservation of energy – energy may change form or be transported, but cannot be created or destroyed.

• Second Law: Entropy – in all natural processes, the entropy of the universe increases.

Gibbs Free Energy and Reaction Spontaneity

The Gibbs free energy change () determines whether a reaction will occur spontaneously.

• = change in enthalpy (heat content)

• = change in entropy

• = absolute temperature (Kelvin)

• : Exergonic (spontaneous, energy released)

• : Endergonic (non-spontaneous, energy required)

Standard free energy change () is measured under standard conditions (1 M concentrations, pH 7, 25°C).

Relationship Between Free Energy and Equilibrium

• At equilibrium, .

• The relationship between standard free energy change and the equilibrium constant () is:

• = gas constant (8.314 J/mol·K)

• = temperature in Kelvin

• If , is negative (reaction proceeds forward).

• If , is positive (reaction proceeds in reverse).

Physical Constants in Thermodynamics

Chemical Logic in Biochemistry

Atomic Properties and Bonding

• Electronegativity: The tendency of an atom in a covalent bond to attract electrons. Increases across a period and decreases down a group in the periodic table.

• Bond cleavage can be homolytic (each atom gets one electron, forming radicals) or heterolytic (one atom gets both electrons, forming ions).

• Common intermediates: carbanions (negatively charged carbon), carbocations (positively charged carbon), and radicals.

Electron Flow in Reactions

• Curved arrows in reaction mechanisms show movement of electrons.

• Arrows start at the electron source (nucleophile) and point to the electron sink (electrophile).

• Full-headed arrows represent movement of electron pairs; half-headed arrows represent single electrons.

Nucleophiles and Electrophiles

• Nucleophile: Electron-rich species that donate an electron pair (e.g., OH-, NH3).

• Electrophile: Electron-deficient species that accept an electron pair (e.g., carbonyl carbon).

• Many biochemical reactions involve nucleophilic attack on electrophilic centers.

Chemical Properties of Carbonyl Groups

• Carbonyl group (C=O): Highly polarized, making the carbon electrophilic and susceptible to nucleophilic attack.

• Metals (e.g., Mg2+) and general acids can facilitate carbonyl group reactivity.

• Imines (C=N) can function similarly to carbonyls in facilitating electron withdrawal.

Categories of Biochemical Reactions

Major Reaction Types

• Carbon–carbon bond formation/breakage: Includes aldol condensation, Claisen ester condensation, and decarboxylation.

• Internal rearrangements, isomerizations, and eliminations: Involve changes within a molecule without altering its oxidation state; typically have small free energy changes.

• Group transfers: Transfer of functional groups from one molecule to another (e.g., phosphorylation).

• Oxidation-reduction (redox) reactions: Electron transfer between molecules, central to energy metabolism.

• Free-radical reactions: Involve unpaired electrons; less common in primary metabolism.

Additional info: These categories are not mutually exclusive; many metabolic reactions involve multiple types.

Examples of Biochemical Reactions

• Aldolase reaction: Formation of a C–C bond in glycolysis.

• Claisen condensation: Formation of C–C bonds in fatty acid synthesis.

• Decarboxylation: Removal of CO2 from a molecule (e.g., isocitrate dehydrogenase).

• Isomerization: Conversion of glucose-6-phosphate to fructose-6-phosphate.

• Elimination: Removal of water from malate to form fumarate.

Standard Free Energy Changes of Common Reactions

Additional info: Rearrangements, isomerizations, and eliminations tend to have small free energy changes because they do not alter the overall oxidation state of the molecule.

Calculating Free Energy Changes

Equations to Memorize

Where is the mass-action ratio (actual concentrations of reactants and products).

• If all concentrations are 1 M, and .

• Actual free energy change () determines reaction spontaneity in the cell.

Study Tips

• Understand the types of metabolic pathways and their regulation.

• Be able to calculate and interpret free energy changes and equilibrium constants.

• Familiarize yourself with the major categories of biochemical reactions and their chemical logic.

• Memorize key equations and physical constants.

• Practice assigned problems to reinforce understanding.