Biochemistry Foundations Notes

Biochemistry Foundations (Transcript Notes)

Major molecular classes and elemental context

  • Four major macromolecules in the body: proteins, carbohydrates, nucleic acids, and lipids. These are built from amino acids and have a carbon backbone. The instructor emphasizes these as the major macromolecules we’ll discuss moving forward.
  • Major elements in the human body vs others:
    • The question is posed: What are the four major elements? The transcript acknowledges there are others beyond the basics.
    • Examples mentioned as other significant elements in physiology: chloride (Cl^-), potassium (K^+), calcium (Ca^2+), phosphorus (P).
    • Calcium and phosphate are specifically noted as the reason bones and teeth are hard, but they are described as not being the major elements for structural parameter purposes.
  • Organic vs inorganic distinction (introduced later): organic compounds all have a carbon backbone; inorganic compounds include salts, water, acids, bases, etc.
  • Water as the first inorganic compound discussed: the most abundant in the body (roughly 70% of cell volume is water). Its properties set up many physiological processes.

Atomic structure, protons, neutrons, and electrons

  • Hydrogen as a primary example:
    • Atomic number (Z) = 1 ⇒ 1 proton and 1 electron (neutral atom).
    • Neutrons = 0 in the most common hydrogen isotope (protium); atomic mass is about 1 (rounded).
    • If you change the number of protons, you change the element (e.g., 1 proton becomes helium with Z = 2).
    • Changing the number of electrons produces ions; hydrogen losing an electron forms a hydrogen ion (H^+).
    • A hydrogen ion is essentially a proton; hydrogen ions and protons are used interchangeably in this context.
  • Neutrons and isotopes:
    • Changing the number of neutrons changes atomic mass and creates isotopes (e.g., carbon-12 vs carbon-13).
  • Ionization concepts:
    • An ion is a charged particle formed by loss or gain of electrons.
    • Hydrogen ion (H^+) is a positive ion after electron loss.
    • The elemental identity remains defined by protons; electrons determine charge.
  • Electron arrangement in atoms:
    • Nucleus contains protons and neutrons; electrons orbit in shells/orbitals.
    • Shell capacities (as described in transcript):
    • 1st shell: 2 electrons
    • 2nd shell: 10 electrons, but stability is achieved at 8
    • 3rd shell: 18 electrons, but stability is achieved at 8
    • Stability is often described in terms of achieving a full outer shell (the octet rule as a guiding principle).
    • Reactivity is driven by outer-shell (valence) electron count.

Electron shells, octet rule, and reactivity

  • Octet concept:
    • Elements strive to achieve a full outer shell (stability). The transcript emphasizes octet stability (8 electrons in the valence shell as a common target).
  • What drives chemical reactivity:
    • The number of electrons in the outer shell (valence electrons) determines how atoms bond with others to become stable.

Types of chemical bonds and representative examples

  • Covalent bonds (share electrons):
    • Example: Methane, CH₄
    • Carbon (Z = 6; e = 6) forms four covalent bonds by sharing electrons with four hydrogens (each H has 1 proton and 1 electron).
    • This sharing results in carbon achieving a full outer shell (octet) and hydrogen achieving its 2-electron shell.
    • Methane is covalently bonded; bonds are strong.
    • Double bonds (two shared electron pairs) can occur, as in O₂ (oxygen gas) where a double bond forms between two oxygens, giving each atom a full outer shell.
  • Ionic bonds (transfer of electrons):
    • Example: Sodium chloride, NaCl (table salt)
    • Sodium (Na, Z = 11; e = 11) tends to lose one electron to achieve a stable configuration (Na⁺).
      • Distribution described: 2 electrons in the 1st shell, 8 in the 2nd shell, and 1 in the 3rd shell.
      • Losing one electron yields Na⁺ with a full 2 + 8 = 10 electrons in the first two shells and a stable third shell configuration for the ion.
    • Chlorine (Cl, Z = 17; e = 17) tends to gain one electron to achieve a stable configuration (Cl⁻).
      • Distribution described: 2 in the 1st shell, 8 in the 2nd shell, 7 in the 3rd shell; needs 1 more electron to complete the 3rd shell to 8 (outer shell becomes full at 8).
    • Ionic bond is the electrostatic attraction between Na⁺ and Cl⁻ resulting from electron transfer.
  • Hydrogen bonds (special, weaker type of interaction):
    • Involves hydrogen and an electronegative atom (commonly oxygen, nitrogen, or fluorine).
    • Not a bond in the same sense as covalent or ionic bonds, but a strong dipole-dipole interaction that is crucial in biology.
    • Water (H₂O) forms hydrogen bonds and explains water’s liquid state at physiological temperatures.
    • Hydrogen bonding also underpins interactions between drugs and enzymes (e.g., atorvastatin example: hydrogen bonds between the drug and enzyme’s oxygens).

Inert gases and chemical inertness

  • Noble gases (e.g., helium, neon) are chemically inert due to full outer electron shells:
    • Helium (Z = 2): first shell filled (2 electrons).
    • Neon (Z = 10): two in the first shell, eight in the second shell, which is considered a filled outer shell here.
    • Because their outer shells are complete, they have little tendency to gain, lose, or share electrons.
  • Conceptual takeaway: full outer electron shells confer chemical inertness in these elements.

Water: properties, structure, and physiological roles

  • Water is inorganic and polar; it dissolves many substances (universal solvent concept).
  • Key properties discussed:
    • High heat capacity: water can absorb or release heat with minimal change in body temperature.
    • High heat of vaporization: requires substantial heat to evaporate, enabling cooling via sweating.
    • Polar solvent: dissolves salts and other ionic compounds, forming hydrated shells around ions (e.g., Na⁺, Cl⁻) and around proteins.
    • Reactive and cushioning roles: participates in chemical reactions (e.g., hydrolysis, dehydration synthesis) and cushions joints and organs (synovial fluid, cerebrospinal fluid, pericardial fluid).
  • Hydration shells and electrolyte dissociation:
    • Salts dissociate into ions in water (electrolytes): NaCl → Na⁺ + Cl⁻.
    • Hydrated layers form around ions and molecules in blood and urine, allowing transport in the circulatory system.
  • Practical physiology examples:
    • Blood plasma is largely water with dissolved electrolytes and proteins (e.g., albumin).
    • Synovial fluid and cerebrospinal fluid rely on water content for cushioning and shock absorption.
    • Urine contains water plus dissolved nitrogenous wastes and other solutes.

Metabolism, acids, bases, and pH

  • Metabolic reactions:
    • Catabolic reactions use hydrolysis (water-breaking) to break molecules down.
    • Anabolic reactions use dehydration synthesis (removing water) to build molecules.
    • Metabolism is the sum of all these processes required for life.
  • Acids, bases, and pH basics:
    • Acids donate hydrogen ions (protons) and increase hydrogen ion concentration, lowering pH.
    • Bases accept hydrogen ions (sponge up protons), decreasing hydrogen ion concentration, raising pH.
    • pH is a measure of hydrogen ion concentration:
    • extpH=− log⁡10[H+]ext{pH} = -\,\log_{10} [H^+]
    • [H+]=10−extpH[H^+] = 10^{- ext{pH}}
    • The pH scale ranges from 0 to 14; 7 is neutral; values below 7 are acidic; above 7 are basic (alkaline).
    • In body chemistry:
    • Stomach acid: HCl contributes protons to create an acidic environment that aids digestion.
    • Carbonic acid-bicarbonate buffer system: H₂CO₃ ⇌ H⁺ + HCO₃⁻ helps maintain blood pH.
    • Bases mentioned: sodium hydroxide NaOH (dissociates into Na⁺ and OH⁻) and ammonia NH₃; OH⁻ can accept H⁺, forming water.
  • pH as a logarithmic scale:
    • Small numerical changes in pH reflect large changes in hydrogen ion concentration: e.g., a shift from pH 3 to pH 4 is a 10-fold change in [H⁺].
    • The typical blood pH is about 7.4; even a shift to 7.8 or 6.9 is physiologically meaningful due to the logarithmic nature.
  • Acids and buffers in physiology:
    • H₂CO₃ (carbonic acid) and HCO₃⁻ (bicarbonate) buffer changes in blood pH.
    • The body maintains homeostasis via buffering systems to prevent dangerous pH shifts.
  • Important ions and electrolytes:
    • Sodium (Na⁺), potassium (K⁺), chloride (Cl⁻), calcium (Ca²⁺) are major electrolytes that dissociate in water and are essential for resting membrane potential, muscle contraction, and various signaling processes.
    • Electrolyte balance is critical for cardiac function and neural activity; imbalance can be life-threatening (e.g., seizures, cardiac arrest).
  • Salts and electrolytes in physiology:
    • Common salts discussed include calcium phosphate (bone enamel), potassium chloride (KCl), and others contributing to bone mineralization and physiological processes.
    • Rehydration after sweating often involves electrolyte-containing drinks (e.g., Gatorade/Propel) to replenish Na⁺, K⁺, and Cl⁻.

Special fluids and anatomical cushioning

  • Synovial fluid in joints provides lubrication and shock absorption in knees, elbows, wrists, hips.
  • Cerebrospinal fluid bathes the brain and spinal cord to cushion against trauma.
  • Pericardial fluid surrounds the heart, providing cushioning and lubrication.
  • These fluids are water-rich and rely on water’s properties for protective buffering and mechanical support.

Quick takeaways for exam preparation

  • Remember the bond types and key examples:
    • Covalent: share electrons (e.g., CH₄, O₂ double bond, N₂ triple bond).
    • Ionic: transfer of electrons to form ions (NaCl → Na⁺ + Cl⁻).
    • Hydrogen bonds: hydrogen involvement and electronegativity differences (water’s structure and drug–enzyme interactions).
  • Noble gases are chemically inert due to full outer shells (He, Ne).
  • Water’s properties underpin much of physiology: high heat capacity and vaporization, polar solvent, ability to form hydration shells, and cushioning roles.
  • Acids and bases regulate pH; the body uses buffers (carbonic acid/bicarbonate) to maintain homeostasis.
  • Metabolic processes rely on hydrolysis and dehydration synthesis (utilizing water).
  • Electrolytes are essential for membrane potentials, muscle function, and neural signaling; proper hydration and electrolyte balance are crucial for health.

Key equations and symbols to memorize

  • pH and hydrogen ion concentration:
    • extpH=−log⁡10[H+]ext{pH} = -\log_{10} [H^+]
    • [H+]=10−extpH[H^+] = 10^{- ext{pH}}
  • Hydrolysis and dehydration synthesis (conceptual, no specific equation provided here):
    • Hydrolysis: bonds are broken with water input
    • Dehydration synthesis: bonds formed with water removal
  • Carbonic acid–bicarbonate buffer relationship (conceptual):
    • H<em>2CO</em>3⇌H++HCO3−\mathrm{H<em>2CO</em>3 \rightleftharpoons H^+ + HCO_3^-}
  • Ionic dissociation example (salt in water):
    • NaCl→Na++Cl−\text{NaCl} \rightarrow \text{Na}^+ + \text{Cl}^-
  • Representative electron shell distributions (as described):
    • Sodium (Na, Z = 11): 2 in first shell, 8 in second shell, 1 in third shell; tendency to lose 1 electron to become Na⁺.
    • Chlorine (Cl, Z = 17): 2 in first shell, 8 in second shell, 7 in third shell; tendency to gain 1 electron to become Cl⁻.