Biochemical Principles: Molecular Forces, Acids & Bases, and Amino Acids

Biochemical Principles: A Review

Molecular Forces

Molecular forces are fundamental to the structure and function of biomolecules. They govern interactions between atoms, molecules, and macromolecules in biological systems.

• Electrostatic & Ionic Interactions: These occur between charged species, such as ions or polar molecules. Examples include salt bridges in proteins and interactions between metal ions and ligands.

• Hydrophobic Interactions: Nonpolar molecules or regions tend to aggregate in aqueous environments to minimize contact with water, driving protein folding and membrane formation.

• Hydrogen Bonds: Formed when a hydrogen atom covalently bonded to an electronegative atom (like O or N) interacts with another electronegative atom. Essential for the stability of DNA, proteins, and water structure.

Electrostatic Interactions

Electrostatic interactions include ion-ion, ion-dipole, and dipole-dipole forces. These are crucial for the stability and specificity of biomolecular complexes.

• Ion-Ion: Attraction between oppositely charged ions (e.g., Na+ and Cl-).

• Ion-Dipole: Interaction between an ion and a polar molecule (e.g., Zn2+ with a carbonyl group).

• Dipole-Dipole: Interaction between two polar molecules (e.g., water molecules).

Hydrophobic Interactions

Hydrophobic interactions arise from the tendency of nonpolar groups to avoid water, leading to aggregation and stabilization of macromolecular structures.

• Van der Waals Interactions: Weak attractions due to transient dipoles in molecules.

• π-Stacking Interactions: Occur between aromatic rings (π systems), such as those in nucleic acids and aromatic amino acids. These interactions stabilize DNA and protein structures.

• Example: Stacking of bases in DNA double helix.

Hydrogen Bonding

Hydrogen bonds are directional and contribute to the specificity and stability of biomolecular structures.

• Donor: Atom bonded to hydrogen (e.g., N-H or O-H).

• Acceptor: Electronegative atom with lone pair electrons (e.g., O or N).

• Example: Hydrogen bonding between water molecules or between base pairs in DNA.

Acids and Bases

Brønsted-Lowry Acids and Bases

The Brønsted-Lowry definition classifies acids as proton donors and bases as proton acceptors.

• Acid Example: $\mathrm{HBr} + \mathrm{H_2O} \rightleftharpoons \mathrm{Br^-} + \mathrm{H_3O^+}$

• Base Example: $\mathrm{NH_3} + \mathrm{H_2O} \rightleftharpoons \mathrm{NH_4^+} + \mathrm{OH^-}$

Weak Acids and pKa

Weak acids do not fully dissociate in water. Their strength is measured by the acid dissociation constant ($K_a$) and its logarithmic form, $pK_a$.

• Acid Dissociation Constant: $K_a = \frac{[A^-]}{[HA]}$

• Henderson-Hasselbalch Equation: $\mathrm{pH} = pK_a + \log \left( \frac{[A^-]}{[HA]} \right)$

• pKa: $pK_a = -\log K_a$

• Application: Buffer systems in biological fluids maintain pH using weak acids and their conjugate bases.

Lewis Acids and Bases

Lewis acids are electron pair acceptors, while Lewis bases are electron pair donors. This definition broadens the concept beyond protons.

• Lewis Acid: Accepts an electron pair.

• Lewis Base: Donates an electron pair.

• Example: $\mathrm{B} + \mathrm{A} \rightarrow \mathrm{B{-}A}$ (where B is the base and A is the acid)

Table: Measure of Acid Strength

The following table compares common acids, their pKa values, and their conjugate bases.

Acid	pKa	Conjugate Base
H2SO4	-9	HSO4-
HCl	-7	Cl-
CH3COOH	4.75	CH3COO-
HCOOH	-1.74	HCOO-
H2O	15.7	OH-
CH3CH2OH	16	CH3CH2O-
NH3	50	NH2-

Amino Acids

Side Chain pKa Values

Amino acids have characteristic pKa values for their α-carboxylic acid and α-amino groups, which influence their charge and behavior at physiological pH.

Amino Acid	Three-letter code	One-letter code	α-carboxylic acid	α-amino group
Alanine	Ala	A	2.35	9.87
Arginine	Arg	R	1.82	8.99
Asparagine	Asn	N	2.14	8.72
Aspartic Acid	Asp	D	1.99	9.90
Cysteine	Cys	C	1.92	10.70
Glutamic Acid	Glu	E	2.10	9.47
Glutamine	Gln	Q	2.17	9.13
Glycine	Gly	G	2.35	9.78
Histidine	His	H	1.80	9.33
Isoleucine	Ile	I	2.32	9.76
Leucine	Leu	L	2.33	9.74
Lysine	Lys	K	2.16	9.06
Methionine	Met	M	2.13	9.28
Phenylalanine	Phe	F	2.20	9.31
Proline	Pro	P	1.95	10.60
Pyrrolysine	Pyl	O	2.20	9.60
Serine	Ser	S	2.19	9.21
Selenocysteine	Sec	U	1.90	9.60
Threonine	Thr	T	2.09	9.10
Tryptophan	Trp	W	2.46	9.41
Tyrosine	Tyr	Y	2.20	9.72
Valine	Val	V	2.29	9.62

Side Chain Structures

Amino acids are classified based on the properties of their side chains, which determine their chemical behavior and role in proteins.

• Nonpolar Side Chains: Glycine, Alanine, Valine, Leucine, Isoleucine, Proline, Methionine, Phenylalanine, Tryptophan

• Polar Uncharged Side Chains: Serine, Threonine, Tyrosine, Cysteine, Selenocysteine, Asparagine, Glutamine, Pyrrolysine

• Positively Charged (Basic) Side Chains at Physiological pH: Histidine, Lysine, Arginine

• Negatively Charged (Acidic) Side Chains at Physiological pH: Aspartate, Glutamate

Example: The side chain of lysine contains an amino group, making it basic and positively charged at physiological pH.

Basic Principles of Redox Chemistry

Redox reactions involve the transfer of electrons between molecules, which is fundamental to energy production and metabolism in biochemistry.

• Oxidation: Loss of electrons by a molecule, atom, or ion.

• Reduction: Gain of electrons by a molecule, atom, or ion.

• Oxidizing Agent: Substance that oxidizes another by accepting electrons.

• Reducing Agent: Substance that reduces another by donating electrons.

• Example: NAD+ is reduced to NADH during glycolysis.

Thermodynamics and Reaction Energetics

Thermodynamics describes the energy changes in biochemical reactions and determines whether a reaction is spontaneous.

• Exergonic Reaction: Releases energy; equilibrium favors products ($\Delta G < 0$, $K > 1$).

• Endergonic Reaction: Requires energy input; equilibrium favors reactants ($\Delta G > 0$, $K < 1$).

• Gibbs Free Energy: $\Delta G$ is the change in free energy, indicating spontaneity.

• Rate vs. Thermodynamics: The spontaneity of a reaction ($\Delta G$) is not related to its rate (kinetics).

• Open Systems: The entropy of an open system can decrease as long as the total entropy of the universe increases.

Example: ATP hydrolysis is highly exergonic and drives many cellular processes.

Additional info: Some chemical structures and diagrams were inferred from context and standard biochemistry knowledge to provide complete explanations.