Biological Structure and Aqueous Environment Flashcards

Introduction to Chemistry of Human Cells

• This lecture introduces the chemistry of human cells, focusing on the molecular basis of cellular function. It examines water's properties as a solvent and acid-base changes in aqueous environments.
• The lecture aims to provide a foundation for understanding cellular and organismal structure, function, and biological specificity.

Relationships of Function to Structure

• Examples illustrating the relationship between structure and function include:
  • Enzyme specificity
  • Genetic information
  • Membrane properties

Cellular Environment

• Water: Makes up 70\% of a cell.
• Most abundant intracellular cation: Potassium (K^+).
• Intracellular anions: Chloride (Cl^-) and phosphate (PO_4^{3-}).
• Cells in vivo reside in an aqueous environment.
• Cell culture attempts to mimic the physiological environment by providing essential extracellular ions like sodium (Na^+$), calcium (Ca^{2+}$), chloride (Cl^−), and bicarbonate (HCO_3^−).

Organic Molecules

• Life on Earth is carbon-based; all organic molecules contain carbon.
• Sugars (e.g., glucose) and fatty acids (e.g., nervistate, palmitate) contain carbon, hydrogen, and oxygen.
• Amino acids contain nitrogen.
• Nucleotides contain sugars, nitrogenous bases, and phosphate.

Covalent Bonds

• Small organic molecules are held together by covalent bonds.
• Covalent bonds involve sharing electrons and are generally strong, broken only during specific chemical reactions.

Macromolecules

• Macromolecules are large molecules assembled from smaller organic building blocks (subunits).
• The primary macromolecules in cells are proteins, nucleic acids, polysaccharides, and lipids.
  • Proteins: Covalently bound chains of amino acids.
  • Polysaccharides: Assembled from sugar units.
  • Nucleic acids: Chains of nucleotides.
• The sequence of subunits determines the physical and functional characteristics of macromolecules. For example, the amino acid sequence dictates a protein's properties.
• Glycoproteins are an example of molecules containing more than one type of component (carbohydrates attached to a protein).

Formation and Breakdown of Macromolecules

• Macromolecules are formed by adding subunits to one end with the removal of water (condensation reaction).
• The reverse reaction, hydrolysis, breaks down the polymer by adding water.
• Digestive enzymes catalyze the hydrolysis of macromolecules.

Molecular Shapes and Models

• DNA is typically found as a double helix.
• Protein molecules vary greatly in shape and size.
• Models for visualizing molecules:
  • Space-filling models (atoms as colored spheres)
  • Ball-and-stick models (depicting bonds between atoms)
• Color-coding:
  • Hydrogen: White
  • Oxygen: Red
  • Carbon: Gray
  • Phosphate: Yellow
  • Nitrogen: Blue
  • Sulfate: Green

Non-Covalent Bonds and Interactions

• Non-covalent bonds/forces affect interactions between molecules and maintain macromolecules in specific 3D structures.
• Types of non-covalent interactions:
  • Hydrogen bonds
  • Electrostatic (ionic) interactions
  • Hydrophobic interactions
  • Van der Waals interactions

Hydrogen Bonds

• Hydrogen bonds mediate the binding of two DNA strands to form a double helix.
• Electrostatic interactions occur between charged atoms or molecules; opposite charges attract, while similar charges repel.
• Sodium chloride is a classic example of an inorganic molecule held together by an ionic bond. In aqueous solution, ionic bonds are weakened as ions are separated and surrounded by water molecules.
• Hydrogen bonds form between water molecules due to the polar nature of the O-H covalent bonds. Oxygen has a stronger attraction for electrons, resulting in partial charges (δ+ and δ-).
• Water is a liquid at body temperature due to the strength of hydrogen bonds between water molecules.
• Electropositive hydrogen atoms are partially shared by two electronegative atoms (oxygen).
• Molecules with polar bonds (e.g., alcohols) can form hydrogen bonds with water and dissolve readily.
• Water molecules surround ions, with negative oxygen attracted to cations and positive hydrogen attracted to anions.
• Hydrogen bonds between nucleotide bases enable DNA pairing, and they stabilize protein secondary structures.
• Water can break hydrogen bonds within a macromolecule and replace them with new bonds to water molecules, changing the molecule's conformation.

Hydrophobic Interactions

• Carbon-hydrogen bonds are nonpolar and do not associate with water.
• Nonpolar molecules or regions are forced out of the hydrogen-bonded water network and are called hydrophobic.
• Fats (triglycerides) are nonpolar hydrophobic molecules that separate from water, like oil and vinegar.

Amphipathic Molecules

• Molecules with both hydrophilic and hydrophobic portions aggregate in aqueous solutions.
• They form structures where nonpolar portions are internal and polar portions interact with the aqueous environment (e.g., cell membrane).
• Cell membranes are composed of phospholipids in a bilayer arrangement, with nonpolar hydrocarbon chains forming the membrane's interior and polar head groups facing outwards.

Protein Folding

• Biological structures fold to minimize unfavorable interactions with water.
• In functional proteins, nonpolar side chains are folded into the molecule's core, and polar groups are on the surface to form hydrogen bonds with water or other polar groups.
• Disrupting the native folding denatures the protein, causing loss of function.

Van der Waals Interactions

• Van der Waals interactions result from transient, flickering polarization of electrons around nonpolar atoms, inducing polarization in nearby atoms.
• Individually weak, the aggregate effect is significant when surfaces are in close contact.
• These forces attract atoms but repel them when they get too close.

Ionization of Water, Acids, and Bases

• Water can ionize as positively charged hydrogen atoms move from one molecule to another, forming hydronium (H_3O^+) and hydroxyl (OH^−) ions.
• This is a spontaneous and reversible process.
• The process is in equilibrium and quantitatively expressed as:[H^+][OH^-]=10^{-14}
• In pure water, the concentrations of hydronium and hydroxyl ions are equal (10^{-7} M).
• The pH scale uses the negative logarithm of the hydrogen ion concentration.
• Pure water has a pH of 7.
• The number is the exponent when a concentration is expressed as a power of 10.

pH Scale

• Adding an acidic substance increases hydrogen ion concentration, resulting in a pH less than 7.
• Basic (alkaline) solutions have a higher hydroxyl ion concentration and a pH greater than 7.
• The higher the pH, the lower the hydrogen ion concentration.

pH Values of Different Solutions

• Gastric fluid: pH 1 (acidic).
• Vinegar: pH 3.
• Pure water: pH 7 (neutral).
• Seawater: pH 8 (slightly alkaline).
• Household ammonia: pH 11.

pH of Common Biological Fluids

• Cytoplasm/cytosol: Neutral.
• Lysosomes: Acidic (low pH for macromolecule breakdown).
• Stomach acid: Acidic (for digestion).
• Tissue culture media: Slightly basic (pH 7.2-7.4).
• Phenol red is used as a pH indicator in culture media (red at pH 7.4, yellow with acidity, purple when more basic).

Strong vs. Weak Acids and Bases

• Strong acids and bases (e.g., hydrochloric acid, sodium hydroxide, sulfuric acid) are completely dissociated in aqueous solutions.
• Biologically important acids/bases are weak and only partially ionized.

Buffers

• Weak acids (e.g., acetic acid) are in equilibrium with their conjugate base (acetate ion).
• Weak acids and bases can act as buffers by binding excess hydrogen or hydroxyl ions to maintain a stable pH.
• Each weak acid has a characteristic dissociation constant (Ka), expressed as: Ka = \frac{[H^+][conjugate \ base]}{[undissociated \ acid]}
• pKa is the negative logarithm of Ka.

Henderson-Hasselbalch Equation

• The Henderson-Hasselbalch equation relates pH and pKa: pH = pKa + log \frac{[conjugate \ base]}{[acid]}
• When pH = pK_a, the concentrations of the acid and its conjugate base are equal, providing the greatest buffering capacity.

Physiological Buffers

• Carbon dioxide-bicarbonate system buffers blood pH.
• Proteins in blood also provide buffering capacity.
• Some culture media contain organic buffers (e.g., HEPES) to enhance buffering capacity and reduce dependence on carbon dioxide.
• Inorganic phosphate regulates cytosolic pH.
• Phosphate buffers can't be used in cell culture media with millimolar calcium concentrations because they form insoluble calcium phosphate complexes.

Attachments and Contact Information

• Review Chapters 4 and 5 from Marks' Basic Medical Biochemistry book.
• Contact the instructor with any questions. Thank you.