2.1 Atoms, Ions, and Molecules — Vocabulary Flashcards
2.1 a The chemical elements
Definition: An element is the simplest form of matter with unique chemical properties. Water is not an element; it is a compound made of elements hydrogen and oxygen.
Breakdown hierarchy:
Elements are composed of atoms.
Atoms contain protons, neutrons, and electrons; protons determine the atomic number (Z).
Protons and neutrons are in the nucleus; electrons form electron clouds around the nucleus.
Atomic number (Z): the number of protons in the nucleus. Examples:
Carbon: Z = 6
Oxygen: Z = 8
Periodic table arrangement: Elements are ordered by atomic number (Z).
Symbol notation: Elements are represented by one- or two-letter symbols. Examples:
C for carbon, Mg for magnesium, Na for sodium, Cl for chlorine.
Some symbols derive from Latin names, e.g., K for potassium (Kalium), Fe for iron (Ferrum).
Natural abundance in the human body:
There are 24 elements that play physiological roles.
Major elements (≈98.5% of body weight): O, C, H, N, Ca, P.
Next 0.8%: S, K, Na, Cl, Mg, Fe.
Trace elements: the remaining ≈0.7% overall.
Trace elements: essential in small amounts but can be toxic in excess (e.g., lead, mercury).
Minerals vs. elements:
Minerals are inorganic elements obtained from soil via plants; they contribute to structure and function.
Minerals constitute about 4% of body weight. Roughly 75% of these minerals are Ca, P, Mg, and K (often abbreviated CAMP).
Roles of minerals:
Bone and teeth: crystals of Ca, phosphate, Mg, fluoride, sulfate.
Proteins: sulfur-containing amino acids; phosphorus in nucleic acids, ATP, membranes.
Enzymes: minerals enable enzyme function; iodine in thyroid hormone; Fe in hemoglobin; Mn, Zn, Cu as enzyme cofactors.
Electrolytes: mineral salts that dissociate into ions in body fluids and are essential for nerve and muscle function; they influence water distribution and osmotic balance.
Practical/clinical notes:
Electrolyte balance is critical in patient care; imbalances can cause cramps, bone brittleness, coma, or cardiac arrest.
Heavy metal poisoning (e.g., Pb, Hg) disrupts physiology.
2.1 b Atomic structure
Historical context:
Democritus (5th c. BCE) proposed atoms as indivisible particles.
Dalton (early 1800s) developed an experimental atomic theory.
Bohr (1913) proposed a planetary model; later quantum mechanics showed electrons occupy probabilistic regions rather than fixed orbits.
Nucleus:
Center of the atom contains protons (p+) and neutrons (n0).
Protons carry a positive charge; neutrons are neutral.
Each proton/neutron ~ 1 atomic mass unit (amu).
Atomic mass roughly equals the total number of protons and neutrons (A ≈ Z + N).
Electrons:
Electrons (e−) have a negative charge and very small mass; mass of electrons is negligible compared to protons/neutrons for most purposes (≈ 1/1836 of a proton).
Electrons occupy electron shells around the nucleus; shells have energy levels and hold a limited number of electrons.
The number of electrons equals the number of protons for electrical neutrality: Z electrons for a neutral atom.
Most physiologically relevant elements involve up to four electron shells (despite that some atoms can have up to seven).
Valence electrons:
Outer-shell electrons (valence electrons) determine chemical bonding properties of an atom.
Visual note: Bohr’s model is a useful visualization, but not a literal depiction of atomic structure in quantum terms.
2.1 c Isotopes and radioactivity
Isotopes: variants of a given element with the same number of protons (Z) but different numbers of neutrons (N); thus different atomic masses.
Hydrogen isotopes: protium 1H (1 p+), deuterium 2H (1 p+, 1 n), tritium 3H (1 p+, 2 n).
Carbon isotopes: 12C (6 p+, 6 n), 13C (7 n), 14C (8 n).
Chemical behavior: all isotopes of a given element generally behave the same chemically; differences are physical (e.g., mass, stability).
Example: Deuterium (2H) reacts with oxygen to form water just like ordinary hydrogen (1H).
Atomic weight vs. isotopic composition:
Atomic weight is the average mass of all isotopes weighted by their natural abundance (e.g., carbon’s atomic weight ≈ 12.011 amu due to heavier isotopes 13C, 14C).
Radioactivity:
Some isotopes are unstable and decay to more stable isotopes; this process emits high-energy radiation (ionizing radiation).
Common decay modes: alpha (α) particles, beta (β) particles, gamma (γ) rays.
An alpha particle is a helium nucleus (2 protons, 2 neutrons).
A beta particle is an electron.
Gamma rays are high-energy photons.
Half-life concepts:
Physical half-life: time required for 50% of a radioisotope to decay.
Example physical half-lives: 90Sr ≈ 28 years; 40K ≈ 1.3 × 10^9 years.
Biological half-life: time required for half of a radioisotope to disappear from the body (through decay plus excretion).
Example: Cs-137 has physical half-life ≈ 30 years but a biological half-life ≈ 17 days due to rapid excretion.
Health effects:
Ionizing radiation can eject electrons, create free radicals, and damage biomolecules (DNA), leading to mutagenesis and carcinogenesis.
High doses can be fatal; lower doses pose risks over time.
Measurements and units (brief):
Dose concepts are measured in sieverts (Sv) to reflect biological effect, accounting for radiation type and tissue sensitivity.
In occupational settings, limits are specified (e.g., around 50 mSv/year).
2.1 d Ions, electrolytes, and free radicals
Ions:
Charged atoms or molecules; cations are positively charged; anions are negatively charged.
Result from electron transfer during ionization or from chemical bonding in solution.
Example ion pair: Na⁺ (cation) and Cl⁻ (anion).
Some elements have multiple ionic states (e.g., iron: Fe²⁺ = ferrous, Fe³⁺ = ferric).
The charge on an ion is called its valence.
Electrolytes:
Substances that dissociate into ions in water; essential for chemical reactivity, osmotic balance, and electrical phenomena in nerves and muscles.
Common electrolytes include ions such as Na⁺, K⁺, Ca²⁺, Mg²⁺, Cl⁻, HCO₃⁻, and others.
Electrolytes enable many physiological processes, including nerve conduction and muscle contraction.
Free radicals:
Unstable, highly reactive chemical species with odd (unpaired) electrons, e.g., superoxide anion $ ext{O}_2^{ullet-}$.
Sources: normal metabolic reactions (mitochondrial oxidation), immune defense, and exposure to radiation or certain chemicals.
Dangers: react with fats, proteins, and DNA to form damaging radicals; implicated in aging and diseases such as cancer and heart disease.
Antioxidants and protection:
The body has natural antioxidant defenses, including the enzyme superoxide dismutase (SOD), which converts superoxide to oxygen and hydrogen peroxide.
Dietary antioxidants include selenium, vitamin E (α-tocopherol), vitamin C (ascorbic acid), and carotenoids (e.g., β-carotene).
2.1 e Molecules and chemical bonds
Molecules vs. compounds:
Molecule: two or more atoms bound together.
Compound: a molecule composed of at least two different elements.
Example: $ ext{O}2$ (molecule) vs. $ ext{CO}2$ (molecule and compound).
Isomers:
Molecules with the same molecular formula but different structures (e.g., ethanol vs. ethyl ether; both $ ext{C}2 ext{H}6 ext{O}$).
Molecular weight (MW): sum of atomic weights of all atoms in the molecule.
Glucose: $ ext{C}6 ext{H}{12} ext{O}_6$ → $MW
oughly = 6(12) + 12(1) + 6(16) = 180$ amu.
Bonds of greatest physiological interest:
Ionic bonds, covalent bonds, hydrogen bonds, and van der Waals forces.
Ionic bonds:
Attraction between cations and anions (e.g., Na⁺ and Cl⁻) forming salts like NaCl.
Ionic compounds tend to dissociate in water because ions are more attracted to water than to each other.
Question to consider: Are ionic bonds common in human body fluids? (Typically, salts exist as dissociated ions in body fluids rather than as intact ionic bonds in solution.)
Covalent bonds:
Single covalent bonds: sharing one pair of electrons (e.g., H–H).
Double covalent bonds: sharing two pairs of electrons (e.g., O=C=O in CO₂).
Nonpolar covalent bonds: electrons spent evenly around both nuclei.
Polar covalent bonds: electrons spend more time near one nucleus (e.g., O–H in water) creating partial charges.
Hydrogen bonds:
A weak attraction between a slightly positive hydrogen in one molecule and a slightly negative atom (usually O or N) in another or in the same molecule.
Essential for water's properties and for maintaining 3D structures of large biomolecules (proteins, DNA).
Represented by dotted lines in diagrams.
Van der Waals forces:
Weak, brief attractions between neutral atoms due to transient dipoles caused by electron density fluctuations.
Individually weak but can be significant collectively (e.g., protein folding, membrane interactions, gecko adhesion).
2.2 water and mixtures
2.2 a Water
Water structure and polarity:
Water is a polar molecule with a V-shaped geometry; bond angle ~ $105^ ext{o}$.
Oxygen bears a partial negative charge ($
abla^- ext{ delta minus}$) and hydrogens bear partial positive charges ($
abla^+ ext{ delta plus}$).
Key properties arising from polarity and hydrogen bonding:
Solvency: water dissolves many substances; water is the universal solvent.
Hydrophilic vs. hydrophobic:
Hydrophilic: substances that readily dissolve in water (e.g., NaCl).
Hydrophobic: substances that do not dissolve well in water (e.g., fats).
Cohesion: water molecules stick to each other via hydrogen bonds; explains surface tension and puddle formation.
Adhesion: water’s tendency to cling to surfaces, aiding lubrication and movement of certain tissues.
Chemical reactivity: water participates in reactions (hydrolysis and dehydration synthesis) and can ionize into $ ext{H}^+$ and $ ext{OH}^-$.
Thermal stability: high heat capacity; many reactions and temperature regulation rely on water’s ability to absorb heat without large temperature changes. When water vaporizes, it carries a lot of heat away, cooling the body (e.g., perspiration).
Hydration and solvation:
Hydration spheres form around ions in solution: e.g., Na⁺ ions attract water with the oxygen end; Cl⁻ ions attract water with the hydrogen ends.
Roles in physiology:
Transport medium for dissolved substances; solvent for metabolic reactions.
2.2 b Solutions, colloids, and suspensions
Mixtures vs. compounds:
Mixture: substances physically blended but not chemically combined; each retains properties.
Compound: elements chemically reacting to form a substance with new properties (e.g., NaCl).
Types of mixtures in the body:
Solution: particles < 1 nm; transparent; solute passes through most membranes; particles do not settle; examples include saline in plasma, glucose in blood.
Colloid: particles 1–100 nm; cloudy; do not settle; cannot pass through most membranes; examples include proteins in plasma.
Suspension: particles > 100 nm; cloudy or opaque; particles settle; examples include blood cells in plasma.
Emulsion: a suspension of one liquid in another (e.g., fats in breast milk, oil in vinegar). Blood can be described as a combination: a solution (NaCl), a colloid (protein), and a suspension (cells).
Table 2.4 (types of mixtures) provides representative examples for the body.
2.2 c Acids, bases, and pH
Definitions:
Arriving at acid/base concepts via proton transfer:
Acid: proton donor (releases H⁺ in water).
Base: proton acceptor (accepts H⁺).
Ammonia (NH₃) is a base because it accepts a proton to become NH₄⁺.
pH concept:
pH is the negative logarithm of the hydrogen ion concentration: ext{pH} = - ext{log}_{10}ig[H^+ig]
Pure water:
At equilibrium, $[H^+][OH^-] = 10^{-14}$ (at 25°C).
In pure water, each is $10^{-7} ext{ M}$, giving pH = 7 (neutral).
pH scale: 0–14; below 7 is acidic, above 7 is basic (alkaline).
The lower the pH, the higher the hydrogen ion concentration; the scale is logarithmic, so a change of 1 pH unit corresponds to a tenfold change in $[H^+]$.
Physiological relevance and buffers:
Blood normally maintained around pH 7.35–7.45; deviations disrupt physiology.
Buffers resist pH changes; discussed in detail in later chapters.
pH and drug action (clinical insight):
The pH of body fluids affects drug ionization and membrane permeability (e.g., aspirin in the stomach is uncharged and can cross membranes; in the bloodstream it ionizes and may get trapped).
Concept of ion trapping (pH partitioning) can be used to enhance excretion of poisons by manipulating urine pH.
2.2 d The other measures of concentration
Concentration units:
Weight per volume: grams or milligrams per liter (e.g., serum cholesterol 200 mg/dL).
Percentage: w/v (weight/volume) or v/v (volume/volume).
Molarity (M): number of moles of solute per liter of solution; reflects molecule number rather than mass.
Millimolar and milliequivalents:
Body fluids are often quantified in millimolar (mM) concentrations to reflect actual molar amounts.
Milliequivalents per liter (mEq/L) account for electrical charge of ions, important for nerve and muscle physiology.
Practical importance: units of concentration influence interpretation of electrolyte balance and IV fluid composition.
2.3 energy chemical reactions
2.3 a Energy and work
Definitions:
Energy: capacity to do work.
Work: movement of a physical object or, in physiology, movement of molecules/ions or mechanical processes (e.g., pumping blood, muscle contraction).
Energy forms:
Chemical energy: stored in chemical bonds; released during reactions to power physiological work.
Heat: kinetic energy of molecular motion; temperature measures the rate of motion.
Electromagnetic energy: moving quanta (photons), e.g., light.
Electrical energy: moving charges; potential when charges are separated (e.g., membrane potentials); kinetic when ions move to produce currents.
Free energy (G): the energy available to do useful work in a system.
In physiology, the primary free energy source is the chemical energy stored in bonds of organic molecules.
2.3 b Classes of chemical reactions
Chemical reactions involve bond formation or breaking; symbolized by chemical equations with reactants → products.
Examples:
Oxidation of ethanol by oxygen to form acetic acid and water (wine turning into vinegar).
Reaction types:
Decomposition: large molecules break into smaller pieces (e.g., starch → glucose).
Synthesis: smaller molecules combine into larger ones (e.g., amino acids into proteins).
Exchange: atoms or groups are exchanged between molecules (e.g., HCl neutralization in digestion).
Reversibility:
Many reactions are reversible and proceed in either direction depending on conditions.
Law of mass action governs direction: reactions proceed from higher to lower reactant/product concentrations.
Biochemical example: CO₂ + H₂O ⇌ H₂CO₃ ⇌ HCO₃⁻ + H⁺; equilibrium can shift with changes in CO₂, bicarbonate, or H⁺.
Equilibrium concept:
Under normal conditions, the carbonic acid/bicarbonate system maintains a physiological balance (e.g., CO₂ transport and blood pH).
2.3 c Reaction rates
Factors affecting reaction rates:
Concentration: higher reactant concentration → more frequent collisions → faster rate.
Temperature: higher temperature → faster molecular motion → more frequent/energetic collisions.
Catalysts: substances that increase reaction rate by providing a favorable orientation or lowering activation energy; are not consumed in the reaction.
Biological catalysts:
Enzymes are the most important biological catalysts; they speed up reactions under physiological conditions.
2.3 d Metabolism, oxidation, and reduction
Metabolism: all chemical reactions in the body.
Subdivisions: catabolism (energy-releasing) and anabolism (energy-storing).
Exergonic reactions release energy; endergonic reactions require energy input.
Oxidation-reduction (redox) reactions:
Oxidation: loss of electrons (molecule becomes more positive/less reduced); the oxidizing agent is the electron acceptor.
Reduction: gain of electrons (molecule becomes more reduced); the reducing agent is the electron donor.
In many cases electrons are transferred as hydrogen atoms; the presence of oxygen as electron acceptor is a common reason for using the term oxidation.
Redox couple notation:
A− → A is a reducing agent; B is oxidized; E− represents the transferred electrons.
Links to energy transfer:
Energy from exergonic oxidation reactions powers endergonic synthetic reactions in metabolism.
2.4 organic compounds
2.4 a Monomers and polymers
Organic chemistry basics:
Carbon is unique due to four valence electrons, allowing diverse bonding and the formation of long chains, branched structures, and rings.
Carbon forms covalent bonds with H, O, N, S; rich functional groups determine properties.
Polymers and monomers:
Macromolecules consist of repeating subunits called monomers.
Examples: starch (polysaccharide of glucose monomers) and proteins (polypeptides composed of amino acids).
Polymerization and hydrolysis:
Dehydration synthesis (condensation): monomers join, releasing water, forming covalent bonds and a polymer chain.
Hydrolysis: a water molecule breaks covalent bonds, yielding monomers; enzymes catalyze the process.
2.4 b Carbohydrates
General formula: carbohydrates have the form $CH2O$ repeating units; commonly written as $(CH2O)_n$.
Important monosaccharides: glucose, galactose, fructose; all C₆H₁₂O₆ isomers.
Other sugars:
Ribose (RNA) and deoxyribose (DNA).
Disaccharides: two monosaccharides linked, e.g., sucrose (glucose + fructose), lactose (glucose + galactose), maltose (glucose + glucose).
Polysaccharides: long chains of monosaccharides; 10–20 units can be oligosaccharides; 50+ units typically polysaccharides.
Glycogen: animal storage polysaccharide; liver, muscle, brain store glycogen; highly branched.
Starch: plant storage polysaccharide; digestible in humans.
Cellulose: structural polysaccharide in plants; not digestible by humans but provides dietary fiber.
Conjugated carbohydrates:
Glycolipids and glycoproteins: carbohydrates covalently bound to lipids or proteins; roles in cell membranes and mucus (protection, infection resistance).
Glycoconjugates:
Proteoglycans: protein–carbohydrate complexes important for gels that hold tissues together and lubricate joints; components in connective tissues and eye.
Functions of carbohydrates:
Major energy source; glucose oxidation generates ATP.
Some carbohydrates serve as cell surface markers and structural components.
2.4 c Lipids
General features:
Lipids are hydrophobic, with high hydrogen-to-oxygen ratio; mainly carbon, hydrogen, and oxygen.
Energy density is high (more calories per gram than carbohydrates).
Major lipid types in humans:
Fatty acids: carboxyl group at one end and methyl group at the other; saturated vs. unsaturated; polyunsaturated fats have multiple C=C bonds.
Triglycerides (neutral fats): glycerol backbone with three fatty acids; energy storage form; fats (solid at room temperature) vs oils (liquid at room temperature).
Phospholipids: glycerol + two fatty acids + a phosphate-containing head; amphipathic; major components of cell membranes; lecithin (phosphatidylcholine) as a common phospholipid.
Eicosanoids: 20-carbon signaling molecules derived from arachidonic acid; include prostaglandins involved in inflammation, clotting, hormone action, uterine contractions, and more.
Steroids: lipids with four fused rings; cholesterol is the main steroid, precursor to other steroids (cortisol, estrogens, testosterone) and bile acids; cholesterol is essential for membranes and nerve function; about 15% of cholesterol is from the diet; 85% is synthesized endogenously, mainly by the liver.
Cholesterol discussion:
LDL (low-density lipoprotein) is often labeled as “bad” cholesterol due to association with atherosclerosis; HDL (high-density lipoprotein) is considered “good” as it helps remove cholesterol from the bloodstream.
Trans fats: partially hydrogenated fats with trans double bonds; denser packing leads to higher risk of cardiovascular disease; FDA banned artificial trans fats in 2015 due to health concerns.
Practical notes:
Main roles of fats: energy storage, insulation, cushioning of organs; phospholipids form cell membranes; prostaglandins regulate various physiological processes.
2.4 d Proteins
Proteins overview:
Proteins are the most versatile biological macromolecules; they perform structural, catalytic, signaling, transport, and protective roles.
A protein is a polymer of amino acids.
Amino acids:
20 standard amino acids share the same backbone: an α-carbon attached to an amino group (–NH₂), a carboxyl group (–COOH), a hydrogen, and a variable R-group.
R-groups determine chemical properties (hydrophilic vs. hydrophobic) and function.
The amino acid sequence is indicated by the protein’s primary structure.
Peptides and polypeptides:
A peptide is a molecule with two or more amino acids joined by peptide bonds (dehydration synthesis).
Dipeptide (2 aa), tripeptide (3 aa); oligopeptides (<≈15–20 aa); polypeptides (>≈50 aa).
Protein structure levels:
Primary structure: amino acid sequence encoded by genes.
Secondary structure: regular folding patterns stabilized by hydrogen bonds, mainly alpha helices and beta sheets.
Tertiary structure: overall 3D folding driven by hydrophobic/hydrophilic interactions, hydrogen bonds, ionic interactions, and van der Waals forces.
Quaternary structure: assembly of multiple polypeptide chains (e.g., hemoglobin with four subunits).
Factors stabilizing structure:
Hydrogen bonds, ionic bonds, hydrophobic interactions, and van der Waals forces.
Disulfide bridges (covalent) between cysteine residues can stabilize tertiary and quaternary structures.
Protein function categories:
Structure: keratin, collagen.
Communication: hormones, receptors, signaling ligands.
Membrane transport: channels and carriers (e.g., Na⁺/K⁺-ATPase).
Catalysis: enzymes; most metabolic pathways are enzyme-catalyzed.
Recognition and protection: antibodies and other immune proteins.
Movement: motor proteins (e.g., myosin).
Adhesion: proteins that bind cells to each other.
Enzymes (2.4 f):
Enzymes are biological catalysts; most are proteins; some ribozymes (RNA enzymes) exist in some contexts.
Names: traditionally named after substrate (e.g., amylase acts on starch); modern naming uses suffix -ase and may include the substrate or class (e.g., carbonic anhydrase).
Active site: the region where substrates bind; enzyme–substrate complex forms via induced fit, where the enzyme changes shape to accommodate the substrate.
Enzyme specificity: each enzyme acts on a particular substrate; some enzymes catalyze two substrates at two active sites.
Denaturation: extreme conditions (temperature, pH) disrupt hydrogen bonds and other weak forces, altering enzyme shape and function.
Cofactors and coenzymes:
Cofactors: inorganic ions (Fe, Cu, Zn, Mg, Ca, etc.) that assist enzyme activity.
Coenzymes: organic molecules (often derived from vitamins) that transfer electrons or functional groups (e.g., NAD⁺, FAD).
Metabolic regulation and isoenzymes:
Isoenzymes (esoenzymes) are enzyme variants with the same function but different structures or regulatory properties; they help in diagnosing diseases (e.g., CK-MM for skeletal muscle, CK-MB for heart).
Enzymes in metabolism example:
Carbonic anhydrase hydrolyzes carbonic acid to CO₂ and water at very high rates (e.g., up to ~3.6×10^7 molecules per minute for a single enzyme).
Enzyme mechanisms and examples:
Sucrose breakdown by the enzyme sucrase (invertase) is an example of substrate binding to an active site and cleavage in the presence of water (hydrolysis).
2.4 g ATP, other nucleotides, and nucleic acids
Nucleotides:
Basic components: a nitrogenous base, a monosaccharide (ribose or deoxyribose), and one or more phosphate groups.
ATP (adenosine triphosphate) is a nucleotide with adenine base, ribose sugar, and three phosphates; a key energy-transfer molecule.
ATP in energy transfer:
ATP briefly stores energy from exergonic reactions and releases it for work (e.g., polymerization, muscle contraction, ion transport).
The high-energy bonds linking the phosphate groups are typically represented with a wavy line in structural formulas; hydrolysis releases energy.
Hydrolysis of the terminal phosphate bond yields ADP and inorganic phosphate (Pi):
Reaction: ext{ATP}
ightarrow ext{ADP} + ext{P}_i + ext{energy} \ ext{Energy release} \ ext{~7.3 kcal/mol}The energy release is used for cellular work; not all energy released is used as work—much is released as heat.
phosphorylation:
The free phosphate released can be transferred to enzymes or other molecules to activate them; this process is called phosphorylation and is carried out by kinases.
ATP turnover and energy balance:
The body consistently replenishes ATP; under moderate activity, a full day’s supply would weigh roughly 45 kg (about 99 lb).
The continuous need for ATP is why mitochondria are considered the cell’s powerhouses, where aerobic respiration occurs to regenerate ATP.
Other nucleotides and signaling:
GTP (guanosine triphosphate) can donate phosphate groups to ADP to regenerate ATP in some contexts.
cAMP (cyclic adenosine monophosphate) acts as a second messenger in many signaling pathways.
Nucleic acids:
DNA (deoxyribonucleic acid): polymers of nucleotides; carries genetic information; typically 100,000,000 to 1,000,000,000 nucleotides long.
RNA (ribonucleic acid): ribonucleotides; carries out coding and synthesis tasks; three major forms (mRNA, tRNA, rRNA) with lengths ranging from ~70 to ~10,000 nucleotides in various contexts.
Overview of nucleic acids in biology:
DNA stores genetic information and provides instructions for protein synthesis; transmission during cell division and heredity.
RNA translates genetic information into proteins and catalyzes synthesis in conjunction with ribosomes.
2.4 h Nucleic acids (continuation)
The genetic code and protein synthesis are detailed in Chapter 4, but the essential idea is:
DNA sequences encode amino acid sequences in proteins via transcription to mRNA and translation on ribosomes.
A, G, C, T (DNA) or U (RNA) nucleotides form the language of genetics.
2.4 f Enzymes in metabolism (summary)
Enzymes are biological catalysts that accelerate metabolic reactions under physiological conditions.
Cofactors and coenzymes expand the range of reactions that enzymes can catalyze and regulate their activity.
Enzyme activity is influenced by temperature, pH, and the presence of activators or inhibitors.
The concept of activation energy: even exergonic reactions require a minimum energy input to overcome the energy barrier; enzymes lower this barrier, allowing reactions to proceed rapidly at body temperature.
2.4 g Clinical insights and broader context
Trans fats and cardiovascular health:
Trans fats (partially hydrogenated oils) package fats in a way that resists enzymatic breakdown, remain in circulation longer, and tend to deposit in arteries.
Large studies linked higher trans fat intake to increased risk of coronary artery disease (CAD).
FDA banned artificial trans fats in processed foods due to health concerns.
Good vs. bad cholesterol:
LDL (low-density lipoprotein) is often labeled as 'bad' cholesterol; HDL (high-density lipoprotein) as 'good' cholesterol.
LDL contributes to atherosclerosis when elevated; HDL helps remove cholesterol from circulation.
Steroid use and health risks:
Anabolic/androgenic steroids (synthetic testosterone derivatives) can promote muscle growth but carry significant health risks: liver damage, cardiovascular disease, immune suppression, hormonal imbalances, and psychiatric effects.
Societal and ethical concerns accompany steroid use in sports.
Medical and physiological implications:
The body’s lipid profile, cholesterol balance, and steroid metabolism affect health; understanding lipids and proteins helps explain disease processes and treatments.
Quick conceptual recap (glossary-style)
Atom vs. element vs. isotope: atoms form elements; isotopes differ in neutron count but share chemical properties; radioisotopes are unstable and radioactive.
Ions and electrolytes: ions carry charges; electrolytes dissociate in water and enable electrical signaling and fluid balance.
Water’s role in life: solvent, thermal buffer, participant in reactions, and facilitator of transport.
Macromolecules: carbohydrates, lipids, proteins, nucleic acids; monomers → polymers via dehydration synthesis; water-driven hydrolysis = digestion.
Bonding types: ionic, covalent (polar and nonpolar), hydrogen bonds, van der Waals forces; all contribute to biomolecular structure and function.
Energy and metabolism: chemical energy stored in bonds; ATP as the primary energy currency; catabolism vs. anabolism; oxidation-reduction in energy transfer.
Enzymes and regulation: specificity and induced fit; cofactors and coenzymes; isoenzymes used in diagnostics.
pH and buffers: acid-base balance in the body; pH affects molecule ionization and consequently drug action and physiology.
Clinical relevance: trans fats, cholesterol transport, isotope usage, radiological safety, and enzyme-based diagnostics.
Note: The content above consolidates the major and supporting points from sections 2.1–2.4, including the examples, definitions, and implied clinical implications discussed in the transcript.