Levels of Organization and Chemical Principles in the Human Body

Levels of Organization in the Human Body

• Atom
  • Smallest unit of matter that defines an element; composed of subatomic particles (protons, neutrons, electrons).
• Molecule
  • Two or more atoms bonded together (e.g.,
    $<br /> \mathrm{H_2}<br />$
    ).
• Macromolecule
  • Large molecules formed by polymerization (e.g., nucleic acids, proteins).
• Organelle
  • Specialized structures within a cell (e.g., Golgi apparatus).
• Cellular level
  • Cells are the basic units of life.
• Tissue level
  • Tissues are groups of cells with a common function. Examples: epithelial tissue, connective tissue.
• Organ level
  • Organs are structures composed of two or more tissue types working together (e.g., stomach, liver, small intestine, gallbladder, large intestine).
• Organ System level
  • Groups of organs that cooperate to perform a common function.
• Organismal level
  • The human organism as a whole.
• Examples shown in transcript include: small intestine, liver, stomach, gallbladder, large intestine; tissues like epithelial and connective; organ systems.

Chapter 2: The Chemical Levels of Organization

• Chemistry: science of structure and interactions of matter.
  • All living things consist of matter.
• Matter: anything that takes up space and can be measured.
• Mass: measure of how much matter is contained within matter; relates to inertia.
• Weight: force of gravity acting on matter/mass; heavier object experiences greater gravitational force.

Matter is Organized into Elements and Atoms

• Elements
  • Unique substances that cannot be broken down by ordinary means.
  • Have specific physical properties and chemical properties.
• Atoms
  • Smallest unit of an element.
• Molecules
  • Two or more atoms bonded together (e.g., $\mathrm{H_2}$).
• Compounds
  • Two or more different atoms joined together.
• Molecular formula
  • Shows the number of atoms of each element in a molecule (e.g., $\mathrm{H<em>2O}$ vs $\mathrm{H</em>2}$).

Chemical Elements in the Body

• 112 elements exist; 92 occur naturally.
• About 26 elements are present in the human body.
• Four elements form ~96% of the body's mass: $\mathrm{C, H, O, N}$.
• Trace elements are present in tiny amounts but are required for normal functioning.

Subatomic Particles and Atomic Structure

• Subatomic particles
  • Protons: $^+$
  • Neutrons: neutral
  • Electrons: $^-$
• Nucleus: center of the atom; electrically neutral overall because the charges balance.
• Orbitals: regions around the nucleus where electrons reside.

Atomic Number and Mass Number

• Atomic number (Z): number of protons in the nucleus; defines the element.
• Mass number (A): average mass of all naturally occurring isotopes; equals protons + neutrons.
  • Formula: $A = Z + N$ where $N$ is the number of neutrons.

Radioactivity, Isotopes, and Radioisotopes

• Isotopes: atoms with a different number of neutrons but the same number of protons and electrons; same chemical properties, may have different masses.
• Radioisotopes: radioactive isotopes that are unstable and decay over time.
• Half-life: time required for a quantity to reduce to half its initial amount.
• Example in clinical imaging:
  • Radioactive iodine is used in medical procedures to trace metabolic reactions.
  • Thyroid uptake imaging shows darker areas where more radioactive iodine is taken up, indicating higher metabolism; can help locate a hyperactive nodule.
• Nuclear medicine thyroid scan: distinguishes abnormally decreased vs. increased uptake and detects gland size abnormalities.
• Isotope decay examples can be described with the half-life formula:
  $N(t) = N<em>0 \left( \tfrac{1}{2} \right)^{\tfrac{t}{t</em>{1/2}}}$

Electron Shells, Valence, and the Octet Rule

• Electron shells (energy levels) surround the nucleus; the outermost shell is the valence shell.
• Valence shell is critical for chemical bonding and reactions.
• Octet Rule: the valence shell tends to have 8 electrons for stability.
• What this implies:
  • Atoms strive to fill their valence shells through sharing or transfer of electrons during bonding.

Inert Elements and Noble Gases

• Inert elements: chemically stable; do not react because their valence shells are full.
• Noble gases: elements in Group 18 of the Periodic Table; located on the right side; helium is an exception in some contexts due to its full shell but unique properties.

Chemical Reactivity and Bonding

• Atoms with incomplete valence shells are unstable and seek to fill them via chemical reactions.
• Chemical bonds include:
  • Intramolecular bonds (WITHIN a molecule): ionic or covalent (polar or nonpolar); often the strongest bonds.
  • Intermolecular bonds (BETWEEN molecules).

Intramolecular Bonds: Ionic Bonds

• Formed by ionization: creation of anions (-) and cations (+).
• Anions and cations attract each other via electrostatic forces.
• Ions: atoms with a different number of electrons than protons, resulting in a net charge.
• Ionization is a type of chemical reaction (donation or gain of electrons).
• Example: Sodium chloride: $\mathrm{Na^+ \; Cl^-}$ forming an ionic bond.
• Electrolytes: compounds that form ions in solution and conduct electricity.
• When valence shells are full, atoms become stable.

Common Anions in the Body (Table 2.1 highlights)

• Chloride ion: $\mathrm{Cl^-}$
  • Physiologic significance: component of stomach acid (HCl) and extracellular fluid balance; involved in chloride shifts in erythrocytes.
• Bicarbonate ion: $\mathrm{HCO_3^-}$
  • Physiologic significance: intracellular buffer; important for maintaining pH; participates in buffering of blood.
• Phosphate ion: $\mathrm{PO_4^{3-}}$
  • Physiologic significance: component of bone and teeth as calcium phosphate; component of phospholipids, nucleotides (including ATP) and nucleic acids (DNA/RNA); intracellular buffer.
• Calcium phosphate: $\mathrm{Ca<em>3(PO</em>4)_2}$
  • Note: contributes to bone and teeth hardness; part of intracellular signaling and structural components.
• Other common anions are present in bodily fluids and contribute to homeostasis.

Intramolecular Bonds: Covalent Bonds

• Atoms share electrons in covalent bonds.
• Examples:
  • Methane: $\mathrm{CH_4}$
  • Ethanol: $\mathrm{C<em>2H</em>5OH}$
  • Carbon dioxide: $\mathrm{CO_2}$
• Covalent bonds can be single, double, or triple, depending on the number of electron pairs shared.
• Covalent bonds form the carbon skeleton of organic molecules.

Covalent Bonds: Polar vs Nonpolar

• Covalent bonds can be polar (unequal sharing) or nonpolar (equal sharing).
• Electronegativity: tendency of an atom to attract electrons.
• Polar covalent bonds: electrons shared unequally; results in partial charges (δ+ and δ−).
• Nonpolar covalent bonds: electrons shared equally; no partial charges.
• Examples:
  • Polar covalent: water molecule involves unequal sharing; polarity leads to dipole moments.
  • Nonpolar covalent: diatomic oxygen, nitrogen gas, etc. (in our context, example structures illustrate equal sharing).

Comparison of Intramolecular Bond Types

• Ionic bonds: complete transfer of electrons; formation of ions; strong interactions through electrostatic forces.
• Polar covalent bonds: unequal sharing; partial charges on atoms.
• Nonpolar covalent bonds: equal sharing; no charges.
• Representations include structural diagrams showing charge distribution and electron transfer.

Amphipathic Molecules

• Amphipathic: molecules with both hydrophilic (polar) and hydrophobic (nonpolar) parts.
• Example conceptually: phospholipids have polar head groups and nonpolar tails.

Types of Intermolecular Attractions

• Hydrogen bonds: weak attractions between a partially positive H and a electronegative atom (often O or N) in a different polar molecule; e.g., in water; strength ~5–10% that of a covalent bond; typically drawn as dotted lines.
• Van der Waals forces: momentary unequal distribution of electrons causing transient dipoles between adjacent nonpolar molecules; about 1% strength of covalent bonds.
• Hydrophobic interactions: nonpolar molecules tend to associate with each other in polar environments (e.g., oil & water) to minimize contact with water.

Water: The Universal Solvent

• Inorganic vs Organic distinction:
  • Water is a polar solvent; dissolves many substances (polar molecules, ions).
  • Hydration spheres form around dissolved ions and polar molecules.
• Water’s biological roles:
  • Transportation (blood, lymph), lubrication (serous fluid, mucus), cushioning (cerebrospinal fluid), excretion (urine, sweat).
• Special properties of water:
  • Cohesion: water molecules stick to each other; adhesion: water molecules stick to other surfaces; surface tension.
  • High specific heat: large energy needed to raise water’s temperature; helps organisms resist temperature changes.
  • High heat of vaporization: significant heat required to vaporize water; evaporative cooling.

Premature Infants and Surfactant (Clinical View)

• Water tends to stick together, causing tiny lung sacs (alveoli) to collapse (atelectasis).
• Surfactant: slippery substance produced by lung cells; reduces surface tension; helps alveoli stay open during breathing.
• Premature infants may lack sufficient surfactant; breathing is difficult or impossible without treatment.
  • Treated with artificial surfactant via breathing tube.

Interactions with Water and Biological Molecules

• Polar molecules and ions dissolve in water; nonelectrolytes (e.g., glucose) dissolve via hydration spheres.
• Electrolytes dissolve and dissociate into ions; example: $\mathrm{NaCl}\rightarrow \mathrm{Na^+} + \mathrm{Cl^-}$.
• Solvent vs solute:
  • Solvent: the dissolving medium (water in biological systems).
  • Solute: the substance being dissolved.
• Water as universal solvent supports biochemical processes.
• Hydrophobic or nonpolar substances do not dissolve in water (e.g., triglycerides, cholesterol; oil and water separation).
• Amphipathic molecules partially dissolve in water (e.g., phospholipids in membranes).
• Types of mixtures:
  • Suspension: large molecules in water that do not stay mixed.
  • Colloid: proteins remain mixed in water (e.g., gelatin, plasma proteins).
  • Solution: very small molecules dissolved in water (e.g., salts, sugars).
  • Emulsion: forced suspension of nonpolar substances in water (e.g., salad dressing).

Acids, Bases, and Salts in Water

• Acids: proton donors.
• Bases: proton acceptors.
• Salts: dissociate into ions in water; neither ion is a proton or hydroxide ion alone.
• Acid + Base = Salt + Water (neutralization).

pH and Buffer Systems

• pH scale: 0 to 14; acidic < neutral < alkaline.
• Lower pH indicates higher [H⁺].
• Typical body fluid ranges:
  • Gastric juice: $1.2 \le pH \le 3.0$
  • Saliva: $6.35 \le pH \le 6.85$
  • Bile: $7.6 \le pH \le 8.6$
  • Blood: $7.35 \le pH \le 7.45$
• Buffer systems (e.g., carbonic acid-bicarbonate buffer): help resist pH changes in body fluids.
• Carbonic acid-bicarbonate system:
  • Balance between $\mathrm{H<em>2CO</em>3}$ and $\mathrm{HCO_3^-}$ maintains blood pH.

Biological Macromolecules (Organic Molecules)

• Four classes of biomolecules:
  • Carbohydrates: energy source; include glycogen, DNA/RNA components; typically 3%–4% of body weight.
  • Lipids: diverse class including triglycerides, phospholipids, steroids, eicosanoids; mostly carbon, hydrogen, and oxygen; hydrophobic; energy storage and membrane structure.
  • Nucleic Acids: DNA and RNA; ATP, NAD, FAD are important energy-related nucleotides.
  • Proteins: built from 20 amino acids; structural and functional roles.

Carbon and Functional Groups

• Carbon has the ability to form bonds with other carbon atoms, creating long carbon chains and rings.
• Functional groups: specific atom groups attached to carbon skeletons that confer distinct chemical properties and reactivity.
• Functional groups enable a diversity of compounds with varying pharmacological, structural, and metabolic roles.

Carbohydrates

• General formula per monomer: $\mathrm{CH_2O}$ (1 C : 2 H : 1 O).
• Primary function: energy for ATP formation.
• Proportion in body weight: relatively small (about 2–3% of body weight) but essential for energy storage and genetic components.
• Three sizes:
  • Monosaccharides (simple sugars): easiest to absorb; 3 can be absorbed without further digestion depending on context.
  • Disaccharides
  • Polysaccharides
• Dehydration synthesis forms disaccharides/polysaccharides by linking monosaccharides (release of water).
• In animals: glycogen stored in liver and skeletal muscle.
• In plants: starch and cellulose.
• Polysaccharides are large storage molecules for sugar.

Lipids

• Lipids include four primary classes:
  • Triglycerides
  • Phospholipids
  • Steroids
  • Eicosanoids
• Triglycerides
  • Backbone: a single $\mathrm{glycerol}$ molecule + 3 fatty acids.
  • Very concentrated source of energy.
  • Saturated fats: single bonds (packed densely) and tend to be solid at room temperature.
  • Unsaturated fats: one or more C=C double bonds; looser packing; healthier profile.
• Clinical View: Fats and fatty acids
  • Most animal fats are saturated; most vegetable fats are unsaturated.
  • Trans fats can be produced by partial hydrogenation; associated with higher risk of heart disease.
• Phospholipids
  • Have polar (hydrophilic) heads and nonpolar (hydrophobic) tails; amphipathic.
  • Critical components of cell membranes forming a phospholipid bilayer.
• Steroids
  • Carbon rings; include cholesterol; bile salts; sex hormones; some vitamins.
  • Cholesterol is found in animal cell membranes.
• Eicosanoids
  • Derived from arachidonic acid; 20-carbon fatty acids; functions in inflammation and signaling.
  • Classes include prostaglandins, prostacyclins, thromboxanes, leukotrienes.

Nucleotides and Nucleic Acids

• Nucleotides: monomers of nucleic acids; composed of
  • A nitrogenous base (A, C, G, T, U in RNA)
  • A five-carbon sugar (pentose): deoxyribose in DNA, ribose in RNA
  • A phosphate group
• Five nitrogen bases: Adenine (A), Thymine (T), Guanine (G), Cytosine (C), Uracil (U).
• Nucleic Acids
  • DNA: genetic material; long chains of nucleotides; bases A, T, G, C; deoxyribose sugar; double-stranded helix.
  • RNA: one strand; ribose sugar; Uracil replaces Thymine; three major types: Messenger RNA (mRNA), Ribosomal RNA (rRNA), Transfer RNA (tRNA).
• Other important nucleotides: ATP, NAD, FAD (energy-related roles).

Flow of Genetic Information

• DNA is used to synthesize RNA (transcription).
• RNA is used to synthesize proteins (translation).
• Proteins determine physical characteristics and control most bodily functions.
• Therefore, DNA serves as the genetic code specifying protein structure and function.

DNA and RNA: Structures in Brief

• DNA
  • Double-stranded helix; bases: ${A, T, G, C}$; sugar: $\text{deoxyribose}$.
• RNA
  • Single-stranded; bases: ${A, U, G, C}$; sugar: $\text{ribose}$; backbone formed by phosphodiester bonds.
• Visual representations show 5' and 3' ends and general nucleotide structure.

ATP: Immediate Energy Currency

• Adenosine Triphosphate (ATP): a nucleotide that provides immediately usable cellular energy.
• Structure includes adenine base, ribose sugar, and three phosphate groups.
• Emphasized as a key energy molecule explored more in later lectures.

Proteins

• composition:
  • Elements: $\mathrm{C, H, O, N}$
  • Built from 20 standard amino acids.
  • Forms include dipeptides (two amino acids) and polypeptides (10–2000 amino acids).
• Amino Acid Structure
  • Central carbon (alpha carbon) bonded to:
  • Amino group $\mathrm{-NH_2}$
  • Carboxyl group $\mathrm{-COOH}$
  • Side chain (R group) that differs among amino acids.
• Protein Structure Levels
  • Primary structure: linear sequence of amino acids.
  • Secondary structure: alpha helix or beta-pleated sheet.
  • Tertiary structure: three-dimensional shape of a single polypeptide; two major types: globular and fibrous.
  • Quaternary structure: assembly of multiple polypeptide chains; e.g., hemoglobin with four chains.
• Bonds that stabilize higher-order structures
  • Hydrogen bonds, disulfide bridges, ionic bonds, hydrophobic/hydrophilic interactions.
• Protein Denaturation
  • A change in a protein’s three-dimensional shape that destroys function.
  • Causes include heat, changes in pH, radiation, heavy metals, alcohol.
• Practical example: pH changes can disrupt electrostatic interactions and other bonds, potentially lethal in blood if not buffered.

Key Equations and Formulas (LaTeX-Format)

• Atomic Mass and Isotopes:
  • Mass number: $A = Z + N$
• Half-life decay:
  • $N(t) = N<em>0 \left( \tfrac{1}{2} \right)^{\tfrac{t}{t</em>{1/2}}}$
• pH definition:
  • $pH = -\log [H^+]$
• Carbohydrate empirical formula:
  • $\mathrm{CH_2O}$ (per unit)
• Ionic dissociation example (NaCl):
  • $\mathrm{NaCl \rightarrow Na^+ + Cl^-}$
• Covalent bond types (illustrative):
  • Single, double, triple bonds indicate one, two, or three shared electron pairs respectively.
• Hydration concept (solvent-solute):
  • Hydration sphere forms around dissolved ions and polar molecules.

Quick Connections and Relevance

• Foundational principles connect chemistry to biology:
  • Elemental composition of the body (C, H, O, N) underpins biomolecule structure.
  • Bonding types (ionic, covalent, hydrogen) dictate molecule stability, interaction, and function in metabolism.
  • Water’s properties enable dissolution, transport, temperature regulation, and hydration, which are essential for nearly all physiological processes.
  • Macromolecules (carbohydrates, lipids, nucleic acids, proteins) provide energy, structure, signaling, genetic information, and catalysis.
  • pH and buffering maintain homeostasis and prevent extreme shifts that could disrupt biological processes.

Connections to Real-World Applications

• Medical imaging using radioisotopes to assess organ function (e.g., thyroid imaging).
• Understanding nutrition and metabolism through carbohydrate and lipid chemistry.
• Pharmacology and biochemistry rely on knowledge of functional groups and molecular interactions.
• Clinical issues such as surfactant therapy in premature infants connect biochemistry to patient care.

Summary of Key Concepts

• Levels of organization progress from atoms to the whole organism, with increasing complexity and specialization.
• Elements, atoms, isotopes, and radioactivity underpin how matter behaves in biological systems.
• Bonding types (ionic, covalent, polar, nonpolar) determine molecular structure and interactions.
• Water is the universal solvent with unique properties that support life.
• Biological macromolecules (carbs, lipids, nucleic acids, proteins) are built from monomers and functional groups, and their structures dictate function.
• DNA and RNA govern genetic information flow; proteins execute most cellular functions.
• pH, buffers, and homeostasis are critical for maintaining stable internal environments.
• Clinical connections illustrate how biochemical principles apply to diagnostics and therapy.