Principles of Biological Structure Notes

Principles of Biological Structure

Key Principles

• Provide the foundation for understanding the structure and function of cells and organisms.
• Provide the foundation for understanding biological specificity.
• Examples of the relationship between function and structure:
  • Enzyme specificity
  • Genetic information
  • Membrane properties

Small Molecules in Cells - Part 1

• Water and small inorganic ions constitute the majority of small molecules in cells.
• Major intracellular ions:
  • Potassium (K^+
  • Chloride (Cl^-
• The primary extracellular cation (positively charged ion) found in biological fluids and culture media is sodium (Na^+

Small Molecules in Cells - Part 2

• Cells encompass a variety of small organic molecules.
• These molecules consist mainly of carbon (C), hydrogen (H), oxygen (O), and often nitrogen (N) atoms, interconnected via covalent bonds.

Macromolecules in Cells

• Macromolecules are constructed by covalently linking a series of smaller molecules.
• Major macromolecules in cells:
  • Proteins
  • Nucleic Acids (DNA, RNA)
  • Polysaccharides
  • Lipids
• Many molecules incorporate more than one of the above components:
  • Example: Glycoproteins, which are proteins with carbohydrate (sugar) chains attached.

Hydrolysis of Macromolecules

• Macromolecules are formed through the addition of subunits at one end, with a water molecule being removed during each addition (condensation reaction).
• The reverse process, which involves the breakdown of the polymer, occurs through hydrolysis. This process requires the addition of water.

Macromolecule Shapes

• DNA takes the form of a double helix.
• Proteins exhibit significant diversity in their shapes and sizes.

Non-Covalent Interactions

• Non-covalent forces influence interactions between molecules and maintain macromolecules in their characteristic shapes.
• Types of non-covalent interactions:
  • Electrostatic (Ionic) Interactions
  • Hydrogen Bonds
  • Hydrophobic Interactions
  • Van der Waals Interactions

Electrostatic (Ionic) Interactions

• Occur between charged molecules (ions):
  • Small ions: Na^+, Cl^-
  • Charged (ionic) groups on large molecules like proteins and nucleic acids
• Opposite charges attract, and like charges repel.
• Example:
  • Negatively charged group of a substrate interacting with a positively charged group of an enzyme.

Hydrogen Bonds

• Water molecules are polar, with oxygen having a slight negative charge due to its higher electronegativity and hydrogen atoms having a partial positive charge.
• Partially positively charged hydrogen atoms are attracted to negatively charged unshared electron pairs of oxygen.
• Hydrogen bonds between water molecules contribute to the physical properties of water, such as high melting and boiling points.

Water as a Solvent

• Water is effective at dissolving charged molecules, such as salt.
• This is due to electrostatic interactions with ions.

Types of Hydrogen Bonds

• Hydrogen bonds occur between many types of polar groups.
• Hydrogen bonding occurs between groups that contain oxygen or nitrogen with unshared electron pairs and partially positively charged hydrogen:
  • -OH
  • -NH
• Hydrogen bonds stabilize many biological molecules:
  • DNA
  • Proteins

Water's Role in Breaking Hydrogen Bonds

• Hydrogen bonds stabilize the folded structure of macromolecules, such as proteins.
• Water can disrupt and replace these hydrogen bonds.

Hydrophobic Interactions

• Nonpolar molecules are hydrophobic.
• Triacylglycerols (fats) are nonpolar molecules with long hydrocarbon chains.
• Hydrophobic molecules attract each other and dissolve poorly in water.

Membranes

• Membranes have a hydrophobic interior and a hydrophilic exterior.
• Phospholipids constitute the basic structure of cellular membranes, possessing polar (or sometimes charged) head groups and nonpolar tails.

Biological Structures and Water Interactions

• Biological structures fold to maximize favorable and minimize unfavorable interactions with water.
• Proteins can exist in unfolded (denatured) or folded (native) states.

Van der Waals Interactions

• Short-range interactions between molecules.
• Involve positive attraction between molecules.
• Include repulsion of molecules when they get too close.

Ionization of Water

• Positively charged hydrogen atoms move from one water molecule to another, creating hydronium (H_3O^+) and hydroxyl (OH^-) ions.
• The process can be simplified as: H_2O
  ightharpoonup H^+ + OH^-
• Equilibrium constant: K{eq} = { [H^+][OH^-] \over [H2O] } = 1.8 \times 10^{-16} M = {10^{-14} \over 55.5 M}
• When [H^+] increases, [OH^-] decreases.
• [H^+] \times [OH^-] = 10^{-14}

The pH Scale

• Pure water: [H^+] = [OH^-] = 10^{-7} M
• The pH scale reflects logarithmic changes in [H^+].
• Pure water is neutral and has a pH of 7.
• Solutions with higher [H^+] (and lower [OH^-]) are acidic.

pH of Biological Fluids

• The cytosol is neutral.
• Lysosomes are acidic.
• Most culture media are slightly basic, with a pH of 7.2 – 7.4.
• Phenol red is used as a pH indicator in culture media:
  • Red at pH 7.4
  • Turns yellow with increased acidity (↓pH)
  • Turns purple with increased pH

Acids and Bases

• Acid (HA) dissociates into H^+ + Conjugate base (A^-).
• Weak acids and bases are only partially dissociated in aqueous solution.
• Acetic acid (vinegar) is a weak acid: CH3COOH + H2O \rightleftharpoons H3O^+ + CH3COO^-
  • Vinegar has a pH of ~3.
• Ammonia (NH_3) is a weak base;
  • The conjugate acid is the ammonium ion (NH4^+): NH4^+ \rightleftharpoons H^+ + NH_3
  • Household ammonia has a pH of ~11.
• Many biologically important molecules are weak acids or bases.

Buffers

• Weak acids and bases are buffers.
• There is an equilibrium between an acid and its conjugate base, allowing them to either donate or accept protons (H^+).
• Weak acids and bases can act as buffers by binding excess H^+ or OH^- ions, maintaining a relatively constant pH or [H^+].

Henderson-Hasselbalch Equation

• Each weak acid has a characteristic dissociation constant or Ka.
  • The pKa is the negative log of Ka.
• A compound is most effective as a buffer when the pH of the solution is within one unit of its pka.
• K_a = { [H^+][A^-] \over [HA] }\n* Where HA is the undissociated acid, and A- is the conjugated base.
• Henderson-Hasselbalch equation: pH = pK_a + log{ [A^-] \over [HA] }
• pKa, analogous to pH, is the negative log of the dissociation constant Ka.
• When the concentration of HA and A- are equal, pH = pKa.

Physiological Buffers

• The carbon dioxide-bicarbonate acid-base couple is a major regulator of blood pH. Cell culture media often use CO2 as a buffer:
  • CO2 + H2O \rightleftharpoons H2CO3 (carbonic acid)
  • H2CO3 \rightleftharpoons H^+ + HCO_3^- (bicarbonate)
• Inorganic phosphate is a major regulator of cytosolic pH:
  • H2PO4^- \rightleftharpoons H^+ + HPO_4^{-2}