Molecule: two or more atoms bound together (can be the same element).
Compound: two or more elements bound together.
Water is a compound (H₂O) formed from two hydrogens and one oxygen; NaCl (sodium chloride) is another example.
Major vs. minor elements in the body
Four elements make up about 96% of body weight: CHON (carbon, hydrogen, oxygen, nitrogen).
The exact order in the list isn’t important; CHON are the major elements to remember.
Minor elements are still essential (e.g., sodium, potassium, iodine). For example, iodine is about 0.004% of body weight but is crucial for thyroid hormone production.
Iodine intake influences thyroid hormones, basal metabolic rate, weight, energy, and mood; both deficiency and excess have serious consequences.
Common test-ready points
You should know the names and chemical symbols of the four major elements CHON (carbon C, hydrogen H, oxygen O, nitrogen N).
You should know some minor elements by name and symbol (e.g., sodium Na, potassium K, iron Fe, iodine I).
No need to memorize every minor element, but be comfortable with symbols for commonly tested ones.
Atoms, nucleus, electrons, and isotopes
Atom basics
Atom: smallest unit of an element that preserves its chemical/physical properties.
Nucleus contains protons (positive charge) and neutrons (no charge).
Electrons (negative charge) orbit the nucleus.
Neutral atoms have equal numbers of protons and electrons.
Subatomic particles and charges
Protons: p⁺; Neutrons: n⁰; Electrons: e⁻.
Electric shells and the octet rule
Electrons occupy shells (orbitals).
The first shell holds 2 electrons; subsequent outer shells prefer 8 electrons (octet rule).
Atoms tend toward a full outer shell to be stable; this drives bond formation.
Isotopes
Isotopes differ in neutron number; most are radioactive (unstable).
Nuclear medicine often uses isotopes (e.g., technetium-99m, Tc-99m) to image or treat tissues.
Isotopes can be produced or altered using devices like cyclotrons.
Atomic number and mass number
Atomic number (top) = number of protons in the nucleus.
Mass number (bottom) = sum of protons and neutrons.
Example: Oxygen has Z = 8 (8 protons). If it has 8 neutrons, mass number ≈ 16.
Example in nuclear medicine
Technetium (Tc) with Z = 43; mass number 98 or 99 in practice; Tc-99m is metastable and used clinically.
Electron transfer and charge implications
When protons ≠ electrons, atoms become ions (positively charged cations or negatively charged anions).
Transferring electrons changes the atom’s chemistry and enables ionic bonding.
Chemical bonds: ionic, covalent, and hydrogen
Ionic bonds
Formed by transfer of electrons from one atom to another, creating oppositely charged ions that attract.
Example: sodium (Na) loses an electron to become Na⁺; chlorine (Cl) gains an electron to become Cl⁻; NaCl forms via ionic attraction.
Covalent bonds
Atoms share electrons rather than transfer them.
Example: H₂ (each hydrogen shares its single electron with the other to fill their outer shells).
Hydrogen bonds
A weaker attraction between polar molecules, not true full charges.
Water (H₂O) molecules form hydrogen bonds with each other; these are weak but numerous and important for structure and properties.
Hydrogen bonds occur between partial positive H and partial negative O in nearby molecules; they’re weaker than ionic or covalent bonds.
Summary of bond types
Ionic bonds: complete transfer of electrons;形成 ions; strong interactions.
Covalent bonds: equal/sharing of electrons; can be nonpolar (equal sharing) or polar (unequal sharing).
Hydrogen bonds: very weak, occur between polar molecules; season water’s properties and biological macromolecule structure.
Practical example: table salt
Na⁺ and Cl⁻ form an ionic lattice in solid NaCl, giving table salt its crystal structure.
Water: polarity, cohesion, and heat capacity
Water as a polar solvent
Water is the most abundant polar molecule in the body and a universal solvent for biochemical reactions.
Cohesion and surface tension
Water molecules stick to each other via hydrogen bonding, creating surface tension and cohesion.
Examples include organisms moving on water surfaces due to surface tension.
High specific heat
Water absorbs heat when bonds form and releases heat when bonds break.
This property helps stabilize body temperature during metabolic reactions.
Percentage of body weight
About 60–70% of total body weight is water.
Inorganic vs. organic molecules; electrolytes and pH
Inorganic molecules
Typically do not contain carbon-hydrogen (C–H) bonds.
Water (H₂O) and carbon dioxide (CO₂) are inorganic examples; CO₂ contains carbon but is still treated as inorganic because it doesn’t form C–H bonds.
CO₂ is an acid gas; it participates in acid-base balance in the body.
Electrolytes and ionization
Electrolytes dissolve in water to form ions (cations and anions).
Acids release H⁺ in solution; bases release OH⁻ (hydroxide) in solution.
pH is a measure of acidity/alkalinity based on hydrogen ion concentration.
Cation vs. anion (memory aid)
Cation: positively charged ion (e.g., Na⁺).
Anion: negatively charged ion (e.g., Cl⁻).
Mnemonic: cation contains a "T" (think of the plus sign like a cross); anion does not.
pH basics and the buffering concept
pH scale typically ranges 0–14 (some charts show slightly beyond this range).
Neutral pH = 7.0; below 7 is acidic; above 7 is basic (alkaline).
pH is defined by the negative logarithm of hydrogen ion concentration: extpH=−log10[H+]
Distilled water is neutral; blood is slightly basic/alkaline; body buffers help maintain homeostasis.
Everyday examples of acids and bases
Hydrochloric acid (HCl) in the stomach has pH ≈ 0–2; it helps digest proteins but requires buffers to protect stomach lining.
Lemon juice, citrus, tomatoes, and coffee are acidic (pH < 7).
Baking soda, ammonia, oven cleaners are basic (pH > 7).
Relationship to homeostasis
Blood pH is tightly regulated; significant deviations can disrupt enzyme activity and metabolism.
Organic macromolecules: carbohydrates, proteins, lipids, and nucleic acids
General rule for organics
Organic molecules always contain carbon; most also contain hydrogen and oxygen.
They are typically larger (macromolecules) than many inorganic molecules.
Carbohydrates
Building blocks: monosaccharides (e.g., glucose).
Can form disaccharides (two sugars, e.g., sucrose) or polysaccharides (complex carbohydrates; glycogen in animals).
Glucose is a universal energy source and readily enters cells to produce energy.
Complex carbs must be broken down to glucose to be used for energy.
Proteins
Building blocks: amino acids (20 total; 9 essential, 11 nonessential).
Essential amino acids must be obtained from the diet; nonessential can be synthesized by the body.
Peptide bonds link amino acids to form proteins.
Four levels of structure: primary (string of amino acids), secondary (alpha-helix or beta-pleated sheet), tertiary (folding of a single chain), quaternary (assembly of multiple folded chains).
Structural proteins provide support (e.g., collagen, keratin) and cell adhesion; functional proteins perform tasks (e.g., enzymes, hormones, transport channels, antibodies, troponin in muscle).
Proteins can denature under extreme conditions (pH changes, temperature, antibiotics, radiation), which disrupts structure and function.
Lipids
Also called fats or triglycerides; contain carbon and hydrogen in greater ratio than carbohydrates; hydrophobic and insoluble in water.
Structure: glycerol plus three fatty acids (triglyceride).
Saturated fats have all single C–C bonds; unsaturated fats have one or more C=C double bonds (fewer hydrogens).
Functions: long-term energy storage, protection against heat loss, insulation, and padding around organs.
Nucleic acids
DNA (deoxyribonucleic acid) and RNA (ribonucleic acid).
Building blocks: nucleotides; DNA uses deoxyribose sugar and bases A, T, C, G; RNA uses ribose sugar and bases A, U, C, G (uracil replaces thymine).
DNA is double-stranded and contains the information blueprint for protein synthesis; RNA is typically single-stranded and helps translate that blueprint to synthesize proteins.
Base pairing rules (across strands): A pairs with T in DNA; A pairs with U in RNA; C pairs with G in both.
The sugar-phosphate backbone and the arrangement of bases encode genetic information.
Protein structure in depth and the idea of denaturation
Primary to quaternary structure
Primary: linear sequence of amino acids joined by peptide bonds.
Secondary: local folding patterns such as alpha helices and beta-pleated sheets.
Tertiary: three-dimensional folding of a single polypeptide chain.
Quaternary: two or more tertiary structures combined into a functional protein.
Denaturation
Denaturation refers to disruption of the protein’s structure, not its amino acid sequence.
Causes include drastic pH changes, antibiotics, high/low temperatures, and radiation.
Denaturation can disable function (e.g., in radiation therapy where tumor proteins are denatured).
Proteins in biology
Antibodies: proteins that recognize and help destroy foreign substances.
Troponin: involved in muscle contraction regulation.
Membrane transport proteins: form channels and tunnels to regulate movement of substances in and out of cells.
Structural proteins: collagen, keratin; provide support and form tissues.
Functional proteins: hormones, enzymes, transporters, etc., that drive biological processes and signaling.
Nucleic acids in the cell: DNA vs RNA specifics
DNA vs RNA structural differences
DNA: double-stranded, forms a helical structure.
RNA: single-stranded; can form complex shapes to assist in protein synthesis.
Nucleotide bases
DNA bases: A, T, C, G (Thymine replaces with Uracil in RNA).
RNA bases: A, U, C, G (Uracil pairs with Adenine).
Roles in protein synthesis
DNA: blueprint with directions for all proteins; located in the nucleus.
RNA: carries the genetic message to the cellular machinery that builds proteins; can fit through cellular doors (nuclear pores) to reach ribosomes.
Practical note on lab/clinical context
The base-pairing rules and the difference between DNA and RNA underpin transcription and translation processes (to be covered in protein synthesis).
Key conceptual links and real-world relevance
Homeostasis and chemical reactions
Reversible reactions (synthesis, decomposition, exchange) balance to maintain homeostasis in the body.
Energy considerations: synthesis requires energy; decomposition releases energy; exchange can require or release energy depending on context.
Why water matters in biology
Water’s polarity drives solvent properties, reaction chemistry, temperature regulation, and transport in organisms.
Water’s cohesive and thermal properties help maintain body temperature and protect organs.
Practical health implications
Iodine deficiency leads to thyroid problems, hypothyroidism (weight gain, fatigue) or excess thyroid activity (anxiety, hair loss).
Proper balance of acids and bases is critical; buffers keep physiological pH near neutral to support enzyme function.
Understanding carbohydrates, fats, and proteins informs nutrition and metabolic health (e.g., glucose as key energy source; fats as energy storage; proteins as enzymes and structural components).
Medical and clinical examples mentioned
Nuclear medicine uses isotopes (e.g., Tc-99m) for imaging.
Radiation therapy denatures tumor proteins to reduce tumor mass over time.
PET scans use glucose uptake to identify metabolically active tissues (tumors).
Ethical and practical considerations in medical contexts
The need for proper education and caution when handling isotopes, radiation therapy, and patient care.
The importance of patient information and safety (e.g., knowing diabetic status for emergency care, carrying medical IDs).
Quick reference recap (equations and quick rules)
Bonding and electrons: ionic vs covalent vs hydrogen bonds; electron transfer vs sharing.