ANP101 Wk 2

Elements and Matter

• Chemistry: the science that examines matter’s composition & properties; foundation for understanding normal/abnormal body function.
• Matter: anything that has mass & occupies space; all substances in the universe.
• Elements:
  • Fundamental types of matter; each identified by a name, symbol, & unique atomic number.
  • Represented in the periodic table; e.g. Carbon (C, $6$), Sodium (Na, $11$).
  • The body is composed chiefly of O, C, H, N (≈$96\%$ of body weight) plus Ca, P, K, S, Cl, Mg, I, Fe, etc.
• Atoms: smallest units of elements that retain chemical properties; cannot be broken down by ordinary chemical/physical means.

Atomic Structure

• Nucleus:
  • Protons: positively charged.
  • Neutrons: no charge.
• Electrons: negatively charged; orbit nucleus in energy levels; number of electrons = number of protons in a neutral atom.
• Atomic number = number of protons (and electrons) in the atom; no two elements share the same atomic number.
• Atomic weight ≈ protons + neutrons (electrons contribute negligible mass).

Energy Levels & Valence

• Energy (electron) levels: regions where electrons orbit.
  • 1st level: max $2$ e⁻.
  • 2nd level: max $8$ e⁻.
  • Stability when outermost (valence) level is full.
• Valence: number of additional electrons an atom needs to fill its outer shell (or the number it can donate/share).
  • Example: Carbon has $4$ valence electrons → valence of $4$ (can form four bonds).
  • Hydrogen has one electron; needs one more to fill its only shell (valence $1$).

Chemical Bonds

• Purpose: enable atoms to fill outer shells by transferring or sharing electrons.
• Ionic Bonds:
  • Formed by electron transfer.
  • Donor becomes positive cation (e.g. $Na^+$), acceptor becomes negative anion (e.g. $Cl^-$).
  • Attraction between opposite charges = ionic bond.
  • Compounds that dissociate into ions in water are electrolytes.
• Covalent Bonds:
  • Formed by electron sharing; most common in the body.
  • Non-polar: equal sharing (e.g. $H<em>2$, $O</em>2$).
  • Polar: unequal sharing (e.g. $H_2O$ – O attracts electrons more strongly).
• Molecules vs. Compounds:
  • Molecule: two + atoms joined by covalent bonds (O$<em>2$, N$</em>2$, H$_2O$).
  • Compound: two + DIFFERENT atoms joined by ionic or covalent bonds (NaCl, CO$<em>2$, H$</em>2O$).
• Chemical equations:
  • Irreversible: $A + B \rightarrow C$ (single arrow).
  • Reversible: $D + E \rightleftharpoons F$ (double arrow).

Electrolytes

• Definition: substances that dissociate in solution into ions able to conduct electricity; the ions themselves are also termed electrolytes.
• Clinical importance:
  • Electrical activity of heart (ECG) & brain (EEG) depends on ion flow.
• Homeostasis maintains ion levels within narrow limits.
• Key body ions & roles:
  • $Ca^{2+}$ – muscle contraction, blood clotting.
  • $HCO_3^-$ – major buffer; pH regulation.
  • $Na^+$, $K^+$, $Cl^-$ etc.

Mixtures & Water

• Mixture: combination of substances physically blended, not chemically bonded.
  1. Solutions – homogeneous; solute dissolves in solvent (salt ⁄ sugar in water).
  2. Suspensions – heterogeneous; particles settle unless mixed (RBCs in plasma, milk of magnesia).
  3. Colloids – heterogeneous; particles remain evenly dispersed due to small size/opposing charges (cytosol, plasma).
• Water: most abundant body compound; universal solvent, stable liquid at body temp, participant in many reactions, critical for temperature regulation & transport.
• Dehydration threatens health; water deficiency alters blood pressure, heart rate, ion concentrations.

Acids, Bases, Salts & pH

• Acid: releases $H^+$ (e.g. $HCl \rightarrow H^+ + Cl^-$).
• Base: releases $OH^-$ and/or accepts $H^+$ (e.g. $NaOH \rightarrow Na^+ + OH^-$).
• Salt: product of acid–base reaction (e.g. $HCl + NaOH \rightarrow NaCl + H_2O$).
• pH Scale:
  • $0$ (strong acid) → $14$ (strong base); each unit = $10$-fold $[H^+]$ change.
  • Neutral $=7$; normal body fluids $7.35–7.45$.
  • Acidosis: $pH<7.35$; Alkalosis: $pH>7.45$.
• Buffers: mixtures of weak acid/base pairs that resist sharp pH changes (e.g. bicarbonate buffer $H<em>2CO</em>3 / HCO_3^-$).

Isotopes & Radioactivity

• Isotopes: atoms of same element with identical proton number but different neutron number → different atomic weight (e.g. Carbon-12,-13,-14).
• Radioisotopes: unstable; emit radiation (α, β, γ) as they decay.
• Medical uses:
  • Cancer therapy: radiation destroys tumor cells.
  • Diagnostic tracers: e.g. radioactive Iodine-131 for thyroid imaging; PET scans.

Organic Compounds

• Built on carbon; main categories:
  1. Carbohydrates – energy & structural molecules.
  2. Lipids – energy storage, membranes, hormones.
  3. Proteins – structure, enzymes, transport, signaling.

Carbohydrates

• Monosaccharides: single sugars (glucose – key fuel & building block).
• Disaccharides: two monosaccharides (sucrose = glucose+fructose; lactose = glucose+galactose).
• Polysaccharides: long chains (glycogen – animal storage; starch – plant storage).

Lipids

• Triglycerides: glycerol + 3 fatty acids (3 carbon atoms in glycerol); insulate, protect, energy reserve.
• Phospholipids: glycerol, 2 fatty acids, phosphate; major membrane component → amphipathic (hydrophilic head, hydrophobic tails).
• Steroids: four carbon rings; include cholesterol (membrane stability), cortisol, sex hormones.

Proteins

• Chains of amino acids (contain amino group –NH$_2$ with nitrogen, carboxyl group –COOH, side R-group).
• Peptide bonds link amino acids; folding yields 3-D shape & function.
• Denaturation: loss of shape/function via heat, pH extremes, chemicals.

Enzymes

• Biological catalysts (proteins) ending in “-ase” (e.g. lipase, sucrase).
• Highly specific: “lock-and-key” fit to substrate.
• Lower activation energy; increase reaction rate; unchanged after reaction.

Nucleotides & Nucleic Acids

• Each nucleotide = nitrogenous base + sugar (ribose/deoxyribose) + phosphate.
• Polymers form DNA (stores genetic code) & RNA (protein synthesis).
• ATP (adenosine triphosphate; tri = three phosphates) – nucleotide that stores usable chemical energy.

Metabolism & Energy

• Metabolism: sum of all chemical reactions in the body.
  • Catabolism: breakdown reactions; release energy; important for ATP synthesis (e.g. digestion).
  • Anabolism: building reactions; require energy/ATP (e.g. protein synthesis).
• Energy forms:
  • Kinetic: energy of movement (electric current, radiant light).
  • Potential: stored energy (chemical bonds, gravitational position); ATP contains potential chemical energy & releases kinetic form when phosphate bond breaks.
• Energy conversion underlies physiological processes (e.g. cellular respiration converts chemical to ATP; muscles convert chemical to mechanical kinetic).

Case Study – Body Fluid Regulation

• Patient Margaret: dehydration led to ↓fluid volume → effects:
  • Elevated hematocrit (concentrated blood), hypernatremia (high Na$^+$), hypotension (low BP), compensatory tachycardia (↑HR).
• Treatment: IV isotonic saline (solvent = water; solute = NaCl) plus dextrose (glucose) – a monosaccharide carbohydrate.
• Blood $pH = 7.28$ indicated acidosis (<$7.35$).
• IV inserted in antebrachium (forearm) region.

Word Anatomy – Key Parts

• co- “together” → covalent (shared electrons).
• aqu/e, hydr/o “water”; dehydration (loss of water), aqueous solution.
• hetero- “different” / homo- “same” → heterogeneous vs homogeneous mixtures.
• phil “love” / phob “fear” → hydrophilic vs hydrophobic.
• ‑ase (enzyme), de- (remove; denature), di- (two), mono- (one), poly- (many), tri- (three), glyc/o (glucose), sacchar/o (sugar).
• ana- (up/build), cata- (down/break), kine-/kinet- (movement).

Review & Self-Test Connections

• Which element forms the largest % body weight? → Oxygen.
• Atom most likely to react: one with incomplete outer shell (e.g. $6$ valence electrons vs fully filled $8$).
• Ionic vs covalent, polar vs non-polar, electrolytes in solution, pH & buffers, radioisotope uses, metabolic pathways – all integrate chemistry with physiology & pathology.