Introduction to Biochemistry: Cells, Chemical Bonds, Functional Groups, and Thermodynamics

Introduction to Biochemistry

Definition and Scope

Biochemistry is the study of the chemical processes and substances that occur within living organisms. It explains life in terms of the atomic structure of biological molecules and the chemical changes resulting from electron rearrangement.

• Bios means life in Greek.

• Al kimya (Arabic): Refers to transmutation, highlighting chemistry's role in explaining change at the atomic level.

• All chemical change is a consequence of electron rearrangement.

The Cell

Animal Cell vs Plant Cell

Cells are the fundamental units of life, with animal and plant cells sharing many structures but also having key differences.

• Animal Cell: Contains nucleus, cytosol, cytoskeleton, mitochondrion, lysosome, peroxisome, Golgi apparatus, vesicles, endoplasmic reticulum, and plasma membrane.

• Plant Cell: Contains all animal cell structures plus chloroplasts, a cell wall, and a large central vacuole.

The Periodic Table of the Elements

Elemental Composition of Life

Life depends on a select group of elements, classified by abundance and biological necessity.

• Most Abundant Elements: C, H, O, N, P, S

• Essential Ions: Na, K, Ca, Mg, Cl

• Common Trace Elements: Fe, Zn, Cu, Mn, I, Se

• Less Common Trace Elements: Mo, Co, F, etc.

Diversity of Life

Biomolecular Composition

There are approximately $10^{10}$ different species in the biosphere, ranging from unicellular to complex multicellular organisms. All living cells use the same types of biomolecules and share common metabolic features, suggesting a common ancestor.

Biomolecule	% in E. coli
Protein	15
Nucleic Acid	7
Carbohydrate	3
Lipid	2
Water	70
Other	3

Key Point: Evolution leads to change over time, but the chemical basis of life is conserved.

Metabolism

Overview

Metabolism is the sum of all catabolic and anabolic reactions in the cell.

• Catabolism: Degradation of nutrient molecules, releasing energy stored in chemical bonds.

• Anabolism: Assembly of large molecules from smaller ones, requiring energy input.

E. coli contains about 1000 metabolites interconverted by ~2000 enzymes.

Types of Biochemical Reactions

1. Group Transfer

2. Internal Rearrangement

3. Cleavage

4. Condensation

5. Oxidation-Reduction

Chemical Bonds

Electronegativity

Electronegativity measures an atom's ability to attract electrons in a chemical bond.

• H: 2.1, C: 2.5, N: 3.0, O: 3.5, P: 2.1, S: 2.5

Saturated Hydrocarbons

• Contain only C–C and C–H covalent bonds (e.g., CH4).

• Highly non-polar, with equal electron sharing.

Functional Groups

Key Functional Groups in Biochemistry

• C=O (carbonyl)

• O–C=O (carboxyl)

• H–O (alcohol/hydroxyl)

• H–N (amino)

• H–S (thiol)

• P–O (phosphate)

Nucleophiles and Electrophiles

• Nucleophiles: Electron-rich species that donate electrons (e.g., alkoxide, carboxylate, carbanion, hydroxide, thiolate, amine, imidazole).

• Electrophiles: Electron-deficient species that accept electrons (e.g., aldehyde, ketone, carboxylic acid, ester, protonated imine, phosphate, proton).

Functional Group Reactivity

Nucleophile and Electrophile Reactivity

Most biochemical reactions occur when nucleophiles and electrophiles interact, often depicted with arrow-pushing mechanisms.

• Group Transfer: Nucleophilic addition and substitution reactions.

• Condensation Reactions: Formation of a new bond with the loss of a small molecule (often water).

Thermodynamics in Biochemistry

First Law of Thermodynamics

The total energy of the universe is constant. Energy is the capacity to do work or release heat, measured in joules.

Forms of Energy

• Kinetic Energy: Due to motion.

• Potential Energy: Due to position (e.g., gravity).

• Enthalpy (H): Heat energy at constant pressure.

• Heat: Transfer of energy from high to low temperature.

Second Law of Thermodynamics

• The entropy (S) of the universe increases.

• Entropy is a measure of disorder, in J/K.

Gibbs Free Energy

Gibbs Free Energy (G) is the energy available to do work at constant temperature and pressure.

• $\Delta G = \Delta H - T\Delta S$

• If $\Delta G < 0$, the reaction is spontaneous (exergonic).

• If $\Delta G > 0$, the reaction is non-spontaneous (endergonic).

Example Equations

• For a reaction: $A + B \rightarrow C + D$

• $\Delta H = H_C + H_D - (H_A + H_B)$

• $\Delta S = S_C + S_D - (S_A + S_B)$

• $\Delta G = G_C + G_D - (G_A + G_B)$

Equilibrium and Free Energy

• At equilibrium, the reaction quotient equals the equilibrium constant ($K_{eq}$):

$K_{eq} = \frac{[C]^c [D]^d}{[A]^a [B]^b}$

• The relationship between free energy change and equilibrium:

$\Delta G = -RT \ln(K_{eq})$

• Where $R = 8.31$ J/mol·K and $T$ is temperature in Kelvin.

Entropy and Cells

• Cells decrease their internal entropy by increasing the entropy of their surroundings, in accordance with the second law of thermodynamics:

$\Delta S_{univ} = \Delta S_{cell} + \Delta S_{surr} > 0$

• Energy within glucose and other nutrients is stored as kinetic and potential energy of electrons in chemical bonds.

Summary Table: Thermodynamic Parameters

$\Delta H$	$\Delta S$	$\Delta G$	Spontaneity
-	+	-	Spontaneous
+	-	+	Non-spontaneous
+	+	?	Depends on $T$
-	-	?	Depends on $T$

Example: The oxidation of glucose to carbon dioxide and water is exergonic and spontaneous under physiological conditions.

Additional info: These foundational concepts are essential for understanding the chemical logic of metabolism, enzyme catalysis, and the molecular basis of life.