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Biochemistry and Atomic Structure Review

Isotopes

  • Isotopes exist because mass number can vary among atoms listed on the periodic table. The mass number indicates the most common type found in nature or man-made isotopes.
  • Altering mass numbers means different numbers of neutrons, creating isotopes.
  • Some isotopes are radioactive (radioisotopes), which can be dangerous in large amounts but are useful in small amounts for diagnostic imaging (radioactivity provides illumination for imaging).
  • In medical contexts, radioisotopes are used in diagnostic imaging in very tiny amounts (injectable or eliminable forms); they help illuminate structures for imaging purposes.

Solutions, Colloids, and Suspensions

  • Solution: particles are tiny, do not settle out or scatter light. Example: mineral water.
  • Colloid: particles scatter light and do not settle out; jelly-like in nature. Example: cytosol (inside cells) — considered a colloidal-type mixture.
  • Suspension: particles settle out. Example: blood in a hematocrit sample where red blood cells settle at the bottom, plasma on top, and a small buffy coat (white blood cells and platelets) in the middle.
  • In the body, many substances are in solution (water-based). Water is the primary solvent and the body’s universal solvent.
  • Solute: the substance dissolved in the solvent (e.g., hormones, minerals like calcium). In a fracture repair context, calcium may be sourced from intestines or kidneys and deposited where needed; processes can be reversible depending on physiology.
  • “Solvent” vs “solutes”: water is the solvent; solutes include the dissolved substances. Water is typically the body’s main solution.

Bonding, Electron Configuration, and Stability

  • Protons never change during ion formation; they remain constant (e.g., sodium has 11 protons).
  • Electrons can be lost or gained to form ions; this changes the atom’s charge.
  • Sodium example: neutral Na has 11 protons and 11 electrons. In the Na+ ion, sodium loses an electron (12? actually 11) to become Na+, resulting in 11 protons but only 10 electrons (outer shell not complete, but overall charge is +1).
    • Ion formation example: Na → Na+ + e−; electrons change but protons do not.
  • Chlorine example: neutral Cl has 17 protons and 17 electrons. Gaining one electron forms Cl− with 18 electrons, achieving a filled outer shell; this is a stable configuration.
  • Ionic bond: formed by transfer of electrons from one atom (e.g., Na) to another (e.g., Cl), creating oppositely charged ions (Na+ and Cl−) that attract each other.
  • Covalent bonds: electrons are shared between atoms. For example, carbon shares electrons with hydrogen to form covalent bonds.
  • Methane example: CH₄ forms when carbon shares electrons with four hydrogens.
  • Polar vs nonpolar covalent bonds:
    • Polar covalent bonds occur when there is unequal sharing of electrons (unequal pulling on electrons), creating partial charges.
    • Nonpolar covalent bonds occur when electrons are shared equally.
  • Water example (polar molecule): oxygen pulls more strongly on electrons than hydrogen, creating a polar covalent bond and an overall dipole.
  • Hydrogen bonds: not true bonds, but weak attractions between partially positive hydrogens and electronegative atoms (e.g., O, N). Essential for many biological structures (e.g., DNA base pairing).
  • Water’s polar nature enables it to dissolve salts and other ions (polar solvent properties).

Water: Properties and Biological Roles

  • Water is the basis of the universal solvent in the body; it supports many biochemical reactions and transport.
  • High heat capacity: it takes a lot of energy to change the temperature of water-rich tissues (plasma ~92% water). This stabilizes core body temperature.
  • High heat of vaporization: it takes a large amount of energy to convert liquid water to vapor, which helps cool the body (via sweating and evaporation).
  • Polar solvent properties: water easily dissociates salts and hydrates ions; example:
    • Dissociation of NaCl in water: ext{NaCl}
      ightarrow ext{Na}^+ + ext{Cl}^-
  • Water as cushioning and protection: cerebrospinal fluid (around brain and spinal cord); synovial fluid in joints; cushioning in bursae sacs (e.g., elbow, knee). Water-containing compartments cushion tissues and protect against mechanical stress.
  • Surfactant reduces surface tension in alveoli to prevent collapse during breathing; produced by type II alveolar cells; becomes critical around 28 weeks of gestation.
  • Surfactant therapy in preterm infants: if born before surfactant production is adequate, lungs may collapse; NICU therapies include surfactant administration (inhaled or injected) to promote alveolar expansion until the infant can breathe on its own.
  • Water’s role in digestion and metabolism: supports hydrolysis and dehydration synthesis; participates in many reactions; acts as a medium for transport and as a reactant/product in many processes.

Hydrolysis, Dehydration Synthesis, and Biochemical Reactions

  • Hydrolysis: adding water to split a molecule (break bonds). Example: hydrolysis reactions during digestion, where enzymes (e.g., lysosomes) may assist in breaking down molecules.
  • Dehydration synthesis: removing water to join molecules and build larger, more complex structures.
  • Example in digestion: dehydration synthesis in the formation of complex polymers and the formation of stool via water absorption in the large intestine (more compact waste).
  • Synthesis vs Decomposition:
    • Synthesis: build complex molecules from simpler ones (anabolic; energy may be required).
    • Decomposition: break down complex molecules into simpler components (catabolic; energy is released).
  • Exchange reactions (buffering context): a reaction where components trade partners to form new products; example: in physiology, carbon dioxide and water reaction (carbonic acid buffering system).
  • For example, CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻, which is central to CO₂ transport and buffering in blood.
  • In ANP2 (biochemistry focus), this exchange is a key buffering mechanism to manage carbon dioxide and acid–base balance.

Carbohydrates, Enzymes, and Energy: Catalysts and Digestion

  • Catalysts accelerate chemical reactions without being consumed; enzymes are the primary biological catalysts.
  • Enzymes reduce activation energy, allowing reactions to occur more readily and with less energy expenditure.
  • Example: Salivary amylase begins carbohydrate digestion in the mouth, breaking down starches into simpler sugars before they reach the small intestine. This reduces the workload on digestive organs downstream.

Acids, Bases, pH, and Buffering in the Body

  • Acids: substances that donate hydrogen ions (H⁺); in general, more H⁺ means more acidic.
  • Bases: substances that accept hydrogen ions or release hydroxide ions (OH⁻); in general, more OH⁻ means more basic/alkaline.
  • pH scale: ranges from 0 (strongly acidic) to 14 (strongly basic); 7 is neutral.
  • Blood pH: normally around 7.35–7.45 (slightly basic/alkaline).
  • Stomach pH: around 2 (highly acidic) to help digestion.
  • Acids and bases in the body must be tightly regulated; buffers help resist abrupt pH changes.
  • Buffering system example in blood: bicarbonate buffer system and hemoglobin act as buffers to maintain stable pH.
  • Carbon dioxide buffering and bicarbonate formation: the CO₂–H₂O equilibrium forming carbonic acid (H₂CO₃) and its dissociation to H⁺ and HCO₃⁻ helps buffer blood pH.
  • Key takeaway: the hydrogen ion concentration drives acidity; proton-donating substances lower pH; buffering systems minimize abrupt pH changes by shifting equilibrium (e.g., to the left or right) to tie up or release hydrogen ions as needed.
  • Practical acid–base context: excess stomach acidity can cause heartburn; common remedies include basic/antacid options (e.g., calcium carbonate-based Tums, milk, or medications like omeprazole) to raise local pH and protect the stomach lining; chronic acidity can lead to ulcers and hiatal hernias if not managed.

Practical Implications: Electrolytes, Nerves, and Muscles

  • Electrolytes are ions that conduct electrical impulses in the body; essential for muscle contraction and nerve signaling.
  • Key electrolytes discussed: potassium (K⁺), chloride (Cl⁻), magnesium (Mg²⁺).
  • Electrolyte balance is necessary for normal muscle function, including heart function. Imbalances can cause cramps or other issues.
  • Gatorade and similar drinks replenish electrolytes (K⁺, Cl⁻, Mg²⁺) and water to support hydration and impulse transmission.
  • Potassium is crucial for muscle relaxation and heart rhythm; sodium is crucial for depolarization and nerve impulse transmission.
  • A lack of ions can cause muscle fatigue or cramping (termed a “hit the wall” feeling).

Carbon Monoxide and Oxygen Transport (Blood Chemistry Context)

  • Carbon monoxide (CO) is a potent gas that binds to hemoglobin with higher affinity than oxygen, reducing oxygen delivery to tissues.
  • CO is odorless; detectors in homes (and in some clinical settings) help prevent exposure.
  • High CO exposure can lead to hypoxia; treatment may involve hyperbaric oxygen therapy to flush CO out of the system and restore oxygen delivery.

Respiratory Physiology: Surfactant, Alveoli, and Lung Mechanics

  • Alveoli are tiny air sacs in the lungs covered with water-based film; surface tension from water can cause alveolar collapse (atelectasis).
  • Surfactant reduces surface tension, enabling expansion of alveoli during breathing.
  • Surfactant production begins around 28 weeks of gestation in fetal development; insufficient surfactant in preterm infants risks respiratory distress syndrome.
  • Neonatal NICU interventions include surfactant administration to promote lung expansion until the infant can breathe independently.

Key Biochemical Concepts Highlighted in the Transcript

  • Synthesis: building more complex molecules from simpler ones.
  • Decomposition: breaking down complex molecules into simpler components.
  • Exchange reactions: rearranging components to form new products; crucial in buffering and transport.
  • Exergonic vs. Endergonic reactions:
    • Exergonic: release energy (ΔG < 0).
    • Endergonic: absorb energy (ΔG > 0).
  • Catabolic vs. Anabolic processes:
    • Catabolic: break down molecules to release energy.
    • Anabolic: build up molecules, often consuming energy.
  • Adenosine triphosphate (ATP) as cellular energy currency; anticipated in future discussions (A&P II).
  • Temperature, concentration, and particle size influence reaction rates:
    • Higher temperature increases rate (more collisions and faster kinetics).
    • Higher concentration increases rate due to more frequent collisions.
    • Smaller particles move faster, increasing reaction opportunities.
  • Catalysts (enzymes) speed reactions without being consumed, reducing energy expenditure.
  • Practical example of enzymes: salivary amylase speeds carbohydrate digestion in the mouth.

Quick Reference Equations and Key Values (LaTeX)

  • Ionic bond formation example: ext{NaCl}
    ightarrow ext{Na}^+ + ext{Cl}^-
  • Covalent bonding (methane): ext{CH}_4 ext{ (carbon and four hydrogens covalently bonded)}
  • Sodium ion example (electron loss): ext{Na}
    ightarrow ext{Na}^+ + e^-
  • Chlorine ion example (gain electron): ext{Cl} + e^-
    ightarrow ext{Cl}^-
  • Water dissociation and buffering (carbonic acid system): ext{CO}2 + ext{H}2 ext{O}
    ightleftharpoons ext{H}2 ext{CO}3
    ightleftrightharpoons ext{H}^+ + ext{HCO}_3^-
  • Methane, CH₄, as an example of covalent bonding and sharing electrons: ext{CH}_4
  • Acidity and pH references: ext{pH} ext{ scales 0–14; } pH < 7 ext{ acidic, } pH > 7 ext{ basic; } pH = 7 ext{ neutral}
  • Blood pH range: 7.35 ext{ to } 7.45
  • Stomach pH: ext{approximately } pH ext{ around } 2
  • Neutral sulfate environment and buffering: explicit mention of buffering systems; example of hemoglobin as buffer in blood (not a numeric equation here, but a functional role)
  • Alveolar surfactant concept: (no equation; functional role described)
  • General reaction rate dependencies (no fixed equation given beyond concepts): temperature, concentration, particle size, collision frequency

Final takeaways for Exam Preparation

  • Understand the differences between isotopes, isotopes’ stability, and the medical use of radioisotopes for imaging.
  • Be able to classify mixtures: solutions vs colloids vs suspensions, with examples from the body.
  • Explain ionic vs covalent bonds, and the role of polarity in determining molecular interactions (e.g., water’s polarity and hydrogen bonding).
  • Describe water’s essential properties (high heat capacity, high heat of vaporization, solvent abilities) and their biological significance (temperature regulation, alveolar mechanics, digestion, cushioning).
  • Distinguish hydrolysis vs dehydration synthesis and give body-relevant examples (digestion, stool formation).
  • Explain acids, bases, and pH, including the buffering role of bicarbonate and hemoglobin; be able to illustrate the CO₂/H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻ buffering reaction and discuss shifts to the left or right.
  • Recognize the importance of electrolytes for nerve and muscle function and the rationale for electrolyte-rich drinks like Gatorade.
  • Understand the concept of surfactant in the lungs and its clinical significance for preterm infants.
  • Be comfortable with the idea that enzymes lower activation energy and act as catalysts, with carbohydrate digestion as a concrete example.
  • Prepare to apply concepts to practical questions about acid-base balance, buffer shifts, and the roles of the major ions in physiological processes.