The Chemistry of Life
2.1 Atoms, Ions, and Molecules
Core idea: Matter in biology is organized around atoms, ions, isotopes, and bonds that determine structure and function of biomolecules.
Learning outcomes to guide study: identify element symbols, distinguish elements vs compounds, describe mineralsâ roles, understand radioactivity and ionizing radiation, differentiate ions, electrolytes, and free radicals, and define chemical bonds.
2.1a The Chemical Elements
A chemical element is the simplest form of matter with unique chemical properties.
Each element is identified by its atomic number Z (number of protons in the nucleus).
The periodic table orders elements by atomic number and uses 1â2 letter symbols.
There are 91 naturally occurring elements.
24 elements play roles in humans; 6 are the most abundant and account for about 98.5% of body weight: O, C, H, N, Ca, and P.
Trace elements are present in minute amounts but have vital roles.
Some trace elements are mineralsâ inorganic elements absorbed from soil by plants and transferred through the food chain to humans.
About 4% of body weight is minerals, primarily Ca and P; they contribute to body structure (bones, teeth) and functions (enzyme activity, nerve/muscle cells).
Important conceptual link: Elements combine to form molecules essential for life; biological function depends on the chemical properties of elements.
2.1b Atomic Structure
The term atom traces back to Greek âatomosâ meaning indivisible; Bohr proposed a planetary model in 1913 as a schematic, not a precise structure.
Nucleus: central core containing protons and neutrons.
Proton: charge +e; mass â 1 atomic mass unit (amu).
Neutron: charge 0; mass â 1 amu.
Atomic mass â total number of protons and neutrons: A â Z + N.
Electrons: orbit the nucleus in electron shells; charge âe; very small mass.
An atom is electrically neutral when the number of electrons equals the number of protons.
Valence electrons reside in the outermost shell and determine chemical bonding properties.
Conceptual takeaway: The arrangement of electrons, especially valence electrons, governs bonding and reactivity.
2.1c Isotopes and Radioactivity
Isotopes are variants of an element differing in neutrons (different atomic mass) but with the same number of protons.
Extra neutrons increase atomic weight; isotopes have similar chemical behavior due to identical valence electrons.
Atomic weight (relative atomic mass) reflects the mixture of isotopes.
Radioisotopes are unstable isotopes that decay, emitting radiation (radioactivity).
Ionizing radiation can eject electrons, destroy molecules, create free radicals, and cause genetic mutations and cancer. Examples: UV radiation, X-rays, alpha, beta, gamma rays.
Physical half-life: time for 50% of a radioisotope to decay to a stable state.
Biological half-life: time for 50% to be eliminated from the body.
Hydrogen isotopes example:
Protium ${}^1_1 ext{H}$: 1 proton, 0 neutrons, 1 electron
Deuterium ${}^2_1 ext{H}$: 1 proton, 1 neutron, 1 electron
Tritium ${}^3_1 ext{H}$: 1 proton, 2 neutrons, 1 electron
Standard radiation dose measures:
Sievert (Sv) as a dose unit; exposure â„5 Sv is usually fatal.
Background radiation average â ${2.4~ ext{mSv/year}}$; artificial sources â ${0.6~ ext{mSv/year}}$.
2.1d Ions, Electrolytes, and Free Radicals
Ion: a charged particle (atom or molecule) with unequal numbers of protons and electrons.
Ionization: transfer of electrons from one atom to another.
Anion: negative charge due to gain of electrons.
Cation: positive charge due to loss of electrons.
Oppositely charged ions attract each other.
Salts are ionic compounds that dissociate in water into ions and act as electrolytes.
Electrolytes: substances that ionize in water and conduct electricity; key roles include chemical reactivity, osmotic effects, and electrical excitability of nerve and muscle.
Electrolyte balance is critical in patient care (imbalance can lead to coma or cardiac arrest).
Free radicals: unstable, highly reactive species with unpaired electrons; produced by metabolism, radiation, and certain chemicals.
Free radicals can damage molecules, contribute to cancer and tissue damage; antioxidants neutralize free radicals (e.g., SOD converts superoxide to oxygen and hydrogen peroxide).
Dietary antioxidants include selenium, vitamins E and C, and carotenoids.
2.1e Molecules and Chemical Bonds
Molecule: two or more atoms bound together; can be a compound if it contains two or more different elements.
Molecular formula identifies constituent elements and their counts; structural formula shows atom locations.
Isomers: molecules with identical molecular formulas but different arrangements.
Molecular weight (MW) is the sum of atomic weights; example: glucose ${ ext{C}}6{ ext{H}}{12}{ ext{O}}_6$ has MW =
MW = 6(12) + 12(1) + 6(16) = 180 ext{ amu}.Chemical bonds hold atoms together within a molecule or link molecules:
Ionic bonds: attraction between cation and anion; relatively easily broken by water.
Covalent bonds: sharing of electron pairs; single bond (1 pair), double bond (2 pairs).
Polar covalent bonds: unequal sharing; e.g., hydrogenâoxygen bonds in water.
Nonpolar covalent bonds: equal sharing; e.g., bonds between carbon atoms.
Hydrogen bonds: weak attractions between a slightly positive hydrogen atom and a slightly negative atom (usually O or N); crucial in physiology; stabilize water and contribute to DNA/protein structure.
Van der Waals forces: very weak, brief attractions due to transient polarization; important for protein folding; ~1% strength of covalent bonds.
2.2 Water and Mixtures
Key learning outcomes for this section: define mixtures vs compounds; describe properties of water; distinguish three kinds of mixtures; define acids/bases and pH; understand concentration expressions.
2.2 Introduction
Body fluids are complex mixtures of chemicals.
Mixtures are substances blended physically but not chemically combined; each component retains its properties.
2.2a Water
Most bodily mixtures involve chemicals dissolved/suspended in water.
Water content: about 50â75% of body weight.
Waterâs properties derive from polar covalent bonds and a V-shaped geometry:
Solvency (universal solvent): dissolves more substances than any other solvent; basis for metabolic reactions.
Cohesion: water molecules stick to each other via hydrogen bonds.
Adhesion: water sticks to other surfaces, reducing friction in membranes.
Chemical reactivity: participates in hydrolysis, dehydration synthesis, etc.
Thermal stability: high heat capacity, resists temperature changes, stabilizing body temperature.
Hydrophilic substances are polar/charged; hydrophobic substances are nonpolar.
Hydration spheres: water molecules surround ions (e.g., Na+, Clâ) to dissolve salts; negative pole faces cation, positive pole faces anion.
Waterâs structure supports solubility, transport, and biochemical reactions.
Water also enables hydration layers around membranes to reduce friction.
2.2b Solutions, Colloids, and Suspensions
Solutions:
Solute particles < 1 nm; do not scatter light; pass through most membranes; do not separate on standing.
Colloids:
Particles ~1â100 nm; scatter light; do not separate easily; remain mixed.
In biology, often proteinâwater mixtures; can form gels.
Suspensions:
Particles >100 nm; cloudy or opaque; may separate on standing.
Emulsions are suspensions of one liquid in another (e.g., oil in water).
2.2c Acids, Bases, and pH
Acids donate protons (H+); bases accept protons.
pH scale measures acidity/alkalinity; pH < 7 acidic, pH > 7 basic, pH = 7 neutral.
Blood pH must be maintained within a narrow range for proper physiology; buffers resist pH changes.
pH formula:
ext{pH} = -\, ext{log}_{10}[H^+].A change of 1 unit in pH represents a tenfold change in hydrogen ion concentration:
ext{pH change} = 1
ightarrow [H^+] ext{ changes by a factor of } 10.Example: pH 4.0 is 10Ă more acidic than pH 5.0.
2.2d Other Measures of Concentration
Concentration expressions include:
Weight per volume: solute mass per volume of solution (e.g., 8.5 g NaCl per liter of solution).
Percent: weight/volume or volume/volume; e.g., 5% dextrose (5 g solute per 100 mL solution) or 70% ethanol (70 mL solute per 100 mL solution).
Molarity (M): moles of solute per liter of solution,
M = rac{n}{V}Millimolar (mM): common in physiology (1/1000 of a mole per liter).
Milliequivalents per liter (mEq/L): accounts for solute concentration and electrical charge; important for nerve/ muscle function and IV fluids.
2.3 Energy and Chemical Reactions
Core outcomes: define energy/work; understand chemical equations; list fundamental types of reactions; factors affecting reaction rates; define metabolism and its subdivisions; understand oxidation/reduction (redox).
2.3a Energy and Work
Energy: capacity to perform work (move a body or a molecule).
Types of energy:
Potential energy: stored energy due to position (e.g., water behind a dam).
Chemical energy: potential energy in chemical bonds.
Free energy: energy available to do work in a system.
Kinetic energy: energy of movement (motion, diffusion, etc.).
Heat: kinetic energy of molecular motion.
Electromagnetic energy: kinetic energy of photons.
Electrical energy: can be both potential and kinetic.
2.3b Classes of Chemical Reactions
A chemical reaction involves making or breaking covalent/ionic bonds.
Chemical equation format: Reactants â Products.
Major classes:
Decomposition: AB â A + B (large molecule breaks into smaller ones).
Synthesis: A + B â AB (two or more smaller molecules form a larger one).
Exchange: AB + CD â AC + BD (atoms or groups swap partners).
Reversible reactions can proceed in either direction; denoted by a double-headed arrow; follow the law of mass action; equilibrium when product/reactant ratio stabilizes.
Example in physiology: buffering of stomach acid with pancreatic bicarbonate leads to a reversible reaction under certain conditions.
2.3c Reaction Rates
Reactions occur when reactants collide with sufficient energy and proper orientation.
Factors increasing rate:
Higher reactant concentration.
Higher temperature.
Presence of a catalyst.
Enzymes are biological catalysts that:
Bind substrates at the active site to form an enzymeâsubstrate complex.
Lower activation energy, accelerating the reaction without being consumed.
Are highly specific (lock-and-key analogy).
Example: sucrose hydrolysis by sucrase yields glucose and fructose.
2.3d Metabolism, Oxidation, and Reduction
Metabolism: all chemical reactions in the body; comprises:
Catabolism: energy-releasing, exergonic decomposition that breaks covalent bonds and yields smaller molecules.
Anabolism: energy-storing, endergonic synthesis of larger molecules (e.g., proteins, fats).
Relationship: catabolism provides energy to drive anabolism; they are tightly linked.
Oxidation and reduction (redox):
Oxidation: loss of electrons; molecule is oxidized; oxidizing agent accepts electrons.
Reduction: gain of electrons; molecule is reduced; reducing agent donates electrons.
Redox reactions occur together; coupled changes ensure energy transfer in metabolism.
2.4 Organic Compounds
Organic chemistry studies compounds containing carbon; four major biomolecule categories:
Carbohydrates
Lipids
Proteins
Nucleic acids
2.4a Carbon Compounds and Functional Groups
Carbon is uniquely suited to form diverse backbones: four valence electrons allow four covalent bonds.
Carbon backbones form chains, branches, and rings; readily bonds with H, O, N, S, and other elements.
Functional groups are small clusters attached to carbon backbones that determine reactivity and properties:
Hydroxyl (-OH)
Methyl (-CH3)
Carboxyl (-COOH)
Amino (-NH2)
Phosphate (-OPO3^{2-})
Functional groups appear across carbohydrates, lipids, proteins, and nucleic acids; they largely dictate behavior in biological systems.
2.4b Monomers and Polymers
Macromolecules are large organic molecules; most are polymers.
Polymers are built from monomers (repetitive subunits).
Examples:
Starch is a polymer of glucose monomers (~3000 identical units).
DNA is a polymer of four different nucleotide monomers.
Polymerization: joining monomers via dehydration synthesis (condensation) with loss of water: monomer1âOH + monomer2âH â dimer + H2O.
Hydrolysis is the reverse: polymer is split into monomers with the addition of water; enzymes facilitate bond breakage.
Example reaction forms:
Dehydration synthesis: monomer1 + monomer2 â dimer + H2O
Hydrolysis: dimer + H2O â monomer1 + monomer2
2.4c Carbohydrates
Carbohydrates are hydrophilic organic molecules; general formula often written as $(CH2O)n$; glucose has $n=6$ and formula $C6H{12}O_6$.
Monosaccharides: simplest carbohydrates; main examples: glucose, galactose, fructose; all share $C6H{12}O_6$ and are isomers; ribose and deoxyribose are monosaccharides used in RNA and DNA respectively.
The three major monosaccharides: glucose, galactose, fructose (structural diagrams show ring forms).
Disaccharides: two monosaccharides covalently bonded; major ones: sucrose (glucose + fructose), lactose (glucose + galactose), maltose (glucose + glucose).
Oligosaccharides: short chains of 3+ monosaccharides.
Polysaccharides: long chains (>50 units) of monosaccharides; major examples:
Glycogen: energy storage in liver/muscle/brain; highly branched.
Starch: energy storage in plants; digestible by humans.
Cellulose: structural polymer in plants; dietary fiber for humans (indigestible).
Functions of carbohydrates:
Rapid energy source; all digested carbs convert to glucose and are oxidized to form ATP.
Can be conjugated to lipids and proteins to form glycolipids and glycoproteins; glycoproteins are a major component of mucus.
Proteoglycans: heavily carbohydrate-rich macromolecules that form gels to hold cells/tissues together, fill the umbilical cord and eye; contribute to joint lubrication and the rubbery texture of cartilage.
Moiety: each component of a conjugated macromolecule.
2.4d Lipids
Lipids are hydrophobic organic molecules with a high hydrogen-to-oxygen ratio; more calories per gram than carbohydrates.
Five primary lipid types in the body: fatty acids, triglycerides, phospholipids, eicosanoids, steroids.
Fatty acids: chains of 4â24 carbons with a carboxyl group on one end and a methyl group on the other; classed as saturated (no C=C bonds) or unsaturated (one or more C=C bonds); polyunsaturated fatty acids have multiple double bonds.
Triglycerides (neutral fats): glycerol + three fatty acids; formed by dehydration synthesis; energy storage; provide insulation and cushioning; oils are liquid at room temperature; fats are solid.
Trans fats and cardiovascular health: trans fats are triglycerides with trans fatty acids; they pack densely and resist breakdown, associated with higher cardiovascular risk.
Phospholipids: glycerol backbone with two fatty acids and a phosphate-containing head; amphipathic (hydrophobic tails, hydrophilic head); structural basis of cell membranes.
Eicosanoids: 20-carbon lipid signaling molecules derived from arachidonic acid; include prostaglandins; roles in inflammation, blood clotting, hormone action, labor, and vessel diameter.
Steroids: lipids with four ring structure; cholesterol is the parent steroid important for nervous system function and membranes; cholesterol balance: ~85% synthesized endogenously, ~15% from diet; other steroids include cortisol, progesterone, estrogens, testosterone, and bile acids.
Lipid health notes: HDL (high-density lipoprotein) = âgood cholesterol,â LDL (low-density lipoprotein) = âbad cholesterolâ; risk of cardiovascular disease correlates with lipid transport profiles.
2.4e Proteins
Proteins are polymers of amino acids; 20 standard amino acids have identical backbones but differ in the R group (side chain).
Amino acids: central carbon (alpha carbon) attached to amino group (-NH2), carboxyl group (-COOH), hydrogen, and a distinctive side chain (R).
Peptides: two or more amino acids linked by peptide bonds (formed via dehydration synthesis).
Peptide length nomenclature:
Dipeptide: 2 amino acids
Tripeptide: 3 amino acids
Oligopeptide: <10â15 amino acids
Polypeptide: >15 amino acids
Protein structure (conformation) is crucial for function. Denaturation: extreme conformational changes that disrupt function (e.g., cooking an egg).
Four levels of protein structure:
Primary: amino acid sequence; encoded by genes.
Secondary: hydrogen bonding leading to alpha helices or beta-pleated sheets.
Tertiary: three-dimensional folding driven by hydrophobic/hydrophilic interactions and van der Waals forces; disulfide bridges stabilize tertiary structure.
Quaternary: association of two or more polypeptide chains (e.g., hemoglobin with four subunits).
Globular vs fibrous proteins:
Globular: compact, functional proteins in membranes/fluids.
Fibrous: extended, structural proteins (e.g., keratin, collagen).
Conjugated proteins: proteins bound to non-amino acid moieties (prosthetic groups); example: hemoglobin contains heme prosthetic group.
Protein functions:
Structural: keratin, collagen.
Communication: signaling molecules (some are proteins) and receptors.
Membrane transport: channels and carriers.
Catalysis: enzymes (usually globular proteins).
Recognition and protection: antibodies, glycoproteins for immune recognition.
Movement: molecular motors.
Cell adhesion: proteins binding cells together.
2.4f Enzymes and Metabolism
Enzymes are biological catalysts; some enzymes are ribozymes (RNA-based) found in ribosomes.
Enzymes interact with one or more substrates at the active site to form an enzymeâsubstrate complex, speeding reactions by lowering activation energy.
Enzyme naming: many end with the suffix -ase, e.g., amylase (starch hydrolysis) and lactase (lactose hydrolysis).
Factors affecting enzyme activity:
Temperature and pH influence enzyme shape and function; optimum pH varies by enzyme (e.g., salivary amylase ~pH 7.0; pepsin ~pH 2.0); human enzymes usually have temperature optimum near 37°C.
Cofactors:
Nonprotein helpers required by some enzymes.
Inorganic cofactors: ions such as Fe, Cu, Zn, Mg, Ca.
Organic cofactors are called coenzymes (often vitamin-derived, e.g., niacin-derived NAD+/NADH).
Coenzyme role: act as electron shuttles between metabolic pathways (e.g., glycolysis and aerobic respiration).
Concept: enzyme activity can be regulated by cofactors and inhibitors, enabling control of metabolic flux.
2.4g ATP, Other Nucleotides, and Nucleic Acids
Nucleotides: basic units composed of a nitrogenous base, a sugar, and one or more phosphate groups.
ATP (adenosine triphosphate): a nucleotide with adenine, a ribose sugar, and three phosphates; the bodyâs primary energy-transfer molecule.
ATP function:
Stores energy from exergonic reactions and releases it for physiological work.
Energy is stored in high-energy phosphate bonds (especially the bonds linking the last two phosphates).
ATP hydrolysis:
ext{ATP}
ightarrow ext{ADP} + ext{Pi} + ext{energy} \ ext{ÎG}_{ ext{hydrolysis}} \ ext{â } -7.3 ext{ kcal/mol}.Phosphorylation: transfer of a free phosphate group to a molecule to activate it; typically catalyzed by kinases.
ATP synthesis primarily from glucose oxidation:
Glycolysis: glucose â 2 pyruvate; yields a small amount of ATP and NADH.
If oxygen is limited, pyruvate is reduced to lactate (anaerobic fermentation).
If oxygen is available, pyruvate enters mitochondria for aerobic respiration, producing CO2, H2O, and a larger yield of ATP.
Overall ATP yield per glucose in full aerobic respiration is commonly cited as about 30â32 ATP:
ext{ATP}_{ ext{total}} \,=\ 2 ext{ (glycolysis)} + 30\text{ (aerobic respiration)} = 32\text{ ATP}.Other nucleotides:
GTP (guanosine triphosphate) can donate a phosphate group in some reactions.
cAMP (cyclic adenosine monophosphate) is formed by removing two phosphates from ATP and acts as a second messenger in signal transduction.
Nucleic acids:
DNA (deoxyribonucleic acid): polymers of nucleotides; contain genes that encode protein synthesis.
RNA (ribonucleic acid): polymers of nucleotides; carries out genetic instructions to synthesize proteins; length ranges roughly from ~70 to 10,000 nucleotides.
End of Chapter Highlights
Core links across sections: atomic structure governs bonding; water properties enable life; macromolecule structure determines function; energy transfer (ATP) powers cellular processes; metabolism integrates catabolic and anabolic pathways; pH and electrolyte balance are essential for homeostasis; nucleotides provide energy transfer and genetic information.
Practical relevance: electrolyte imbalances affect nerve/muscle function; radiation safety and biological effects of isotopes; enzymes as drug targets; nutrition and lipid balance in health; carbohydrates as immediate energy; proteins as catalysts and structural components; nucleic acids as the blueprint of life.