2.1–2.5 Elements, Atoms, Bonds, Reactions, and Organic Compounds – OpenStax AP
2.1 Elements and Atoms: The Building Blocks of Matter
Matter and its definitions- Matter: the substance of the universe that occupies space and has mass.
Mass vs weight:
Mass: amount of matter; constant regardless of gravitational environment (Earth, Moon, zero gravity).
Weight: mass under the influence of gravity; e.g., a given mass weighs less on the Moon than on Earth due to weaker gravity.
Elements and compounds- Elements: 92 natural fundamental substances; cannot be created or broken down by ordinary chemical means.
Compounds: substances formed from two or more elements joined by chemical bonds.
Example: calcium (Ca) is an essential element; body uses calcium for bones; cannot be produced by the body, must come from environment.
Elements in the human body (most abundant first): O, C, H, N (with others like Ca, Na, Fe, etc.).
Elements in foods and air contribute to bodily elements; figure shows elemental abundance (O > C > H > N in typical body mass).
Elements, atoms, and subatomic particles- An element’s atom is the smallest unit retaining its properties.
Subatomic particles:
Protons (p): positive charge; located in nucleus.
Neutrons (n): neutral; located in nucleus.
Electrons (e): negative charge; orbit nucleus at high speed in electron shells.
Mass is primarily from protons and neutrons; electrons contribute negligibly to atomic mass (≈ 1/2000 of proton mass).
Neutral atoms have equal numbers of protons and electrons; overall charge is balanced.
Atomic number and mass number- Atomic number (Z): number of protons in the nucleus; defines the element (e.g., Carbon has Z = 6).
Mass number (A): total number of protons and neutrons in the nucleus; A = Z + neutrons.
In neutral atoms, the number of electrons ≈ number of protons (equal in most stable isotopes).
Isotopes: atoms of the same element with different numbers of neutrons (different A).
Isotopes and their significance- Isotopes can be stable or radioactive (radioisotopes).
Examples:
Carbon isotopes: C-12 (6p, 6n), C-13 (6p, 7n), C-14 (6p, 8n).
Hydrogen isotopes: Protium (1H): 1p, 0n; Deuterium (2H or D): 1p, 1n; Tritium (3H): 1p, 2n.
Radioactive isotopes decay over time; half-life is the time for half of a sample to decay (e.g., tritium ≈ 12 years).
Practical implications: medical imaging (PET), cancer therapy; excessive exposure is harmful (cell damage, cancer, birth defects) but controlled exposure provides diagnostic/therapeutic benefits.
Atomic identity and the periodic table- Protons (and electrons in neutral atoms) determine an element’s identity; the periodic table groups elements by atomic number and chemical properties.
In many elements, neutrons outnumber protons; mass number varies by isotope.
Valence electrons (electrons in the outermost shell) govern chemical reactivity.
The first column elements (e.g., H, Li, Na, K) typically have one valence electron; valence electrons participate in chemical reactions.
Electron shells and stability- Electron shells (energy levels) surround the nucleus.
Each shell has a maximum capacity; the first shell holds up to 2 electrons; each subsequent shell holds up to 8 electrons.
The number of shells depends on the atom’s total electrons; hydrogen and helium have only 1 and 2 electrons, respectively.
The octet rule: atoms are most stable when their valence shell contains eight electrons (except hydrogen and helium, which are stable with 2).
Nearest-neighbor interactions and medical relevance- In the body, atoms form bonds and interact (not by touching, but via electrical attractions).
The radiologic career connection (Interventional Radiology): use of radioisotopes for diagnosis and treatment; PET uses radioactive tracers (e.g., glucose) to visualize metabolic activity; radiologists leverage these technologies for non-surgical diagnostics and therapies.
Ethical/philosophical note: balance benefits of radioisotope use in medicine with potential risks of exposure; clinical decisions weigh risks/benefits for patient care.
Visual models of atoms- Planetary model (electrons in fixed orbits) vs. electron cloud model (electrons in probabilistic regions around the nucleus).
Quick terminologies to remember- Element, atom, proton, neutron, electron, atomic number Z, mass number A, isotope, ion, cation, anion, valence electrons, electron shells, octet rule.
2.2 Chemical Bonds
Atoms link to form molecules and compounds- Bond: a weak or strong electrical attraction that holds atoms together; atoms do not touch in a bond.
Molecule: two or more atoms held together by chemical bonds.
Compound: molecule composed of two or more different elements.
Ions and ionic bonds- Ions: atoms with a net electrical charge due to loss or gain of electrons.
Cation: positively charged ion (e.g., Na⁺, Mg²⁺).
Anion: negatively charged ion (e.g., Cl⁻, Se²⁻).
Ionic bonds: attraction between opposite charges (e.g., Na⁺ and Cl⁻); forms salts such as NaCl.
In aqueous solution, salts dissociate into ions (electrolytes); crucial for nerve impulses and muscle contraction.
Covalent bonds- Covalent bonds: atoms share electrons to fill valence shells; electrons are not permanently transferred.
Bond strengths: covalent bonds generally stronger than ionic bonds.
Bond types by sharing electrons:
Nonpolar covalent: equal sharing of electrons; no dipole moment (e.g., O₂, N₂, the C–H bonds in some hydrocarbons).
Polar covalent: unequal sharing; creates partial charges (e.g., H₂O).
Hydrogen bonding in covalently bound molecules
Weak interaction between a slightly positive hydrogen and a nearby electronegative atom (e.g., O, N).
Essential in water’s properties and the behavior of many biological molecules.
Water as a polar solvent and hydrogen bonding- Water’s polarity: δ⁺ on H, δ⁻ on O; facilitates dissolution of ionic and polar covalent compounds (hydrophilic substances).
Nonpolar substances (e.g., fats) do not dissolve; water can repel nonpolar molecules.
Hydrogen bonds contribute to water’s high cohesion, surface tension, and heat-handling capabilities.
Properties and implications of bonds in physiology- Ionic bonds form salts like NaCl; salts dissociate into ions in water (electrolytes) that conduct electricity; critical for nerve impulses and heart function.
Covalent bonds form molecules like water (H₂O) and carbon-based compounds; can be nonpolar or polar.
Hydrogen bonds link water molecules and stabilize many biological structures (protein folding, DNA double helix).
Quick terms to know- Ion, cation, anion, ionic bond, covalent bond, polar covalent bond, nonpolar covalent bond, hydrogen bond, electrolyte.
2.3 Chemical Reactions
What is a chemical reaction?- Bond breaking and bond formation; result is a new substance with different properties.
Energy changes accompany reactions; energy is neither created nor destroyed (conservation of energy/mass).
Forms of energy in chemistry- Kinetic energy: energy of motion.
Potential energy: energy of position or structure; chemical energy is a form of potential energy stored in bonds.
In metabolism, chemical energy is released from bonds to power cellular processes.
Exergonic vs. endergonic reactions- Exergonic: energy released; reactions proceed with a net release of energy (e.g., catabolic processes).
Endergonic: energy absorbed; reactions require energy input to proceed (often powered by other exergonic reactions).
How enzymes influence reactions- Enzymes: biological catalysts that speed up reactions by lowering activation energy (the energy needed to start a reaction).
Activation energy is the energy barrier to break initial bonds.
Enzymes work via the active site, often following an induced-fit model; substrate specificity means one enzyme typically acts on one type of substrate.
Monomers, polymers, and water in reactions- Dehydration synthesis (condensation): monomers covalently bond to form polymers; a molecule of water is released.
Hydrolysis: polymers are broken into monomers by adding water (H and OH across broken bonds).
Basic reaction examples and notations- Synthesis: A + B -> AB
Decomposition: AB -> A + B
Exchange: AB + CD -> AD + CB
Reversibility: many reactions are reversible, denoted with a double arrow: A + B <-> AB
Notable examples and equations- Ammonia synthesis (illustrative): N2 + 3H2 -> 2NH3
Glucose oxidation (energy release; simplified): C6H12O6 + 6O2 -> 6CO2 + 6H2O + ATP energy
Metabolic energy currency: ATP hydrolysis (activation energy and energy release):
ATP + H2O -> ADP + Pi + energy
Factors influencing the rate of chemical reactions (in humans)- Properties of reactants: surface area, gaseous state vs liquid/solid, molecule size, reactivity.
Temperature: higher temperature increases kinetic energy, increasing reaction rate.
Concentration and pressure: more particles and/or reduced volume (higher pressure) increases collision frequency.
Catalysts (enzymes): accelerate reactions by lowering activation energy; important in metabolism and physiology.
Examples: vinegar and baking soda reaction is rapid; enzymes catalyze many body reactions.
2.4 Inorganic Compounds Essential to Human Functioning
Water (H2O) in the body- Up to ~70% of body weight is water; present intracellularly and extracellularly.
Roles:
Lubricant and cushion in joints, digestive tract, and organs; protects brain and delicate tissues; cushions fetus.
Heat sink: absorbs heat from chemical reactions; helps regulate body temperature via sweating and blood flow.
Solvent: universal solvent; dissolves many ions and polar covalent compounds; water is polar, hydrophilic substances dissolve easily; nonpolar molecules are hydrophobic.
Water in mixtures
Solutions: homogeneous mixtures with solvent and solute; water is the solvent in most body fluids.
Hydration and solubility concepts: the ratio of solute to solvent influences solubility; molarity (M) defined as moles per liter:
M = n/V
Avogadro’s number: NA = 6.02 x 10^23 particles per mole
Solutions and solutes- Solute: substance dissolved in a solvent.
Hydrophilic vs hydrophobic:
Hydrophilic: water-loving (polar and ionic substances) dissolve in water.
Hydrophobic: water-fearing (nonpolar substances) do not dissolve well.
Salts and electrolytes- Salts: ionic compounds that dissociate into ions in water (e.g., NaCl -> Na⁺ + Cl⁻).
Electrolytes: ions in body fluids; conduct electrical currents; crucial for nerve impulses and muscle function.
Bile salts (fats emulsification) and calcium phosphate (teeth and bones) are examples of salts in physiology.
Acids, bases, and pH- Acids: Release hydrogen ions (H⁺) in solution; strong acids dissociate completely; weak acids dissociate incompletely.
Bases: Release hydroxyl ions (OH⁻) or bind H⁺; strong bases dissociate fully; weak bases dissociate partially.
pH: negative log of hydrogen ion concentration; pH = -log10 [H⁺]; pH scale ranges from 0 to 14 with 7 as neutral.
pH in the body: blood typically ~7.35–7.45; homeostasis via buffering systems, breathing, and renal excretion.
Buffers and homeostasis- Buffers: solutions of a weak acid and its conjugate base that resist pH changes.
In blood, buffers help maintain pH around 7.4; if pH drops, buffers bind excess H⁺; if pH rises, buffers release H⁺.
Dehydration synthesis and hydrolysis in inorganic context- Dehydration synthesis and hydrolysis also apply to organic chemistry; water is a byproduct in dehydration synthesis and is consumed in hydrolysis.
2.5 Organic Compounds Essential to Human Functioning
Overview of organic compounds- Characterized by carbon-based backbones with hydrogen, often with oxygen and other elements.
Four major classes in human structure/function: carbohydrates, lipids, proteins, nucleic acids.
Carbon’s covalent bonding flexibility enables formation of long chains, rings, and functional groups.
Carbon–hydrogen (hydrocarbon) motifs are common; carbon–oxygen/nitrogen functional groups modulate reactivity.
Functional groups important in physiology- Hydroxyl (–OH): polar; participates in dehydration synthesis and hydrolysis; part of all four major organic classes.
Carboxyl (–COOH): found in fatty acids and amino acids.
Amino (–NH₂): present in amino acids; contributes to buffering.
Methyl (–CH₃): common in organic molecules
Phosphate (–PO₄³⁻): in phospholipids and nucleotides; contributes to energy transfer (ATP).
Monomers and polymers- Macromolecules: large molecules built from repeating subunits.
Monomer: a single unit; Polymer: many monomer units.
Dehydration synthesis links monomers; hydrolysis breaks polymers with water.
Carbohydrates- General formula: (CH2O)n; carbohydrates are hydrated carbons.
Monosaccharides: key ones include glucose, fructose, galactose (hexoses; 6 carbons); ribose and deoxyribose (pentoses; 5 carbons).
Disaccharides: two monosaccharides linked by a glycosidic bond; examples: sucrose, lactose, maltose; digestion hydrolyzes them to monosaccharides before use.
Polysaccharides: starch (plants), glycogen (animals), cellulose (plant cell walls; indigestible by humans as fiber).
Functions: energy storage (glucose, glycogen), structural components (cellulose in plants), brain and RBC energy needs (glucose).
Lipids- General properties: hydrophobic; not soluble in water; include triglycerides, phospholipids, sterols, prostaglandins.
Triglycerides: glycerol backbone + three fatty acids; formed by dehydration synthesis; storage form of energy.
Formation: Glycerol + 3 Fatty Acids -> Triglyceride + 3H2O
Fatty acids: saturated (no double bonds; straight chains) vs. unsaturated (one or more double bonds; kinked chains).
Saturated fats: solid at room temperature; examples: butter, lard.
Unsaturated fats: oils; monounsaturated (one double bond) and polyunsaturated (two or more double bonds);
Omega-3 fatty acids: beneficial fatty acids in fish oils; influence prostaglandin production.
Phospholipids: glycerol + two fatty acids + phosphate group; hydrophobic tails and hydrophilic heads; form cell membranes.
Sterols: four-ring structures; cholesterol as a key sterol; roles in bile salts, hormones, and cell membranes.
Prostaglandins: derived from fatty acids; signaling molecules affecting inflammation and blood flow; NSAIDs reduce prostaglandin effects.
Proteins- Functions: structural (keratin, collagen), enzymes, antibodies, neurotransmitters, hormones, and buffering components.
Elements: predominantly C, H, O, N; many contain S as well.
Amino acids: building blocks; 20 standard amino acids; general structure includes amino group (–NH₂), carboxyl group (–COOH), hydrogen, and a variable side chain (R).
Polymers: proteins are polymers of amino acids linked by peptide bonds; dehydration synthesis forms peptide bonds.
Essential amino acids: nine cannot be synthesized by humans and must be obtained from the diet.
Protein structure levels:
Primary: sequence of amino acids.
Secondary: alpha-helix or beta-pleated sheet; stabilized by hydrogen bonds.
Tertiary: three-dimensional folding due to interactions among R-groups.
Quaternary: assembly of multiple polypeptide subunits (e.g., hemoglobin).
Denaturation: disruption of a protein’s shape due to heat, pH changes, or chemicals; loss of function.
Enzymes as proteins: catalysts that accelerate chemical reactions; specificity toward substrates; active sites; induced fit model.
Glycoproteins and proteoglycans: proteins bound to carbohydrates; contribute to membranes and extracellular matrix.
Nucleic acids and energy carriers- Nucleotides: monomers composed of a phosphate group(s), a pentose sugar (deoxyribose in DNA, ribose in RNA), and a nitrogenous base.
DNA (deoxyribonucleic acid): stores genetic information; bases A, C, G, T; double helix formed by base pairing and hydrogen bonds between backbones.
RNA (ribonucleic acid): ribose sugar; bases A, C, G, U; single-stranded; essential in translating genetic code to proteins.
Purines vs pyrimidines:
Purines: adenine (A) and guanine (G) — double-ring structures.
Pyrimidines: cytosine (C), thymine (T; in DNA), uracil (U; in RNA) — single-ring structures.
Backbones and bases: nucleotides linked by dehydration synthesis between the sugar of one nucleotide and the phosphate of another, creating a sugar-phosphate backbone with protruding bases; DNA bases pair via hydrogen bonds to form the double helix.
Genes and chromosomes: humans have ~22,000 genes within 46 chromosomes; genetic code guides protein synthesis; unique per individual (except identical twins).
ATP (adenosine triphosphate): primary energy currency of the cell; structure includes ribose, adenine, and three phosphate groups; high-energy phosphate bonds store energy.
ATP hydrolysis and phosphorylation:
Hydrolysis: ATP + H2O -> ADP + Pi + energy
Phosphorylation: transfer of a phosphate group from ATP to another molecule (e.g., glucose -> glucose phosphate) enabling energy storage or transfer.
Key takeaways for organic chemistry in physiology- Organic compounds rely on carbon’s tetravalence to form diverse macromolecules.
Functional groups dictate reactivity and properties of biomolecules.
Monomer–polymer relationships and energy dynamics (dehydration synthesis and hydrolysis) underpin metabolism.
Understanding the structure–function relationship is essential for predicting protein folding, enzyme activity, and genetic regulation.
Quick reference equations (LaTeX-friendly)
Atomic relationships
Z = number of protons
A = Z + neutrons
Avogadro’s number
NA = 6.02 x 10^23
Molarity
M = n/V where n= moles (mol), V= volume (L)
pH
pH = -log10 [H⁺]
Covalent bonding (general forms)
Synthesis: A + B -> AB
Decomposition: AB -> A + B
Reversibility: A + B <-> AB
Important biological reactions
Ammonia synthesis (illustrative): N2 + 3H2 -> 2NH3
Glucose oxidation (simplified): C6H12O6 + 6O2 -> 6CO2 + 6H2O + ATP energy
Dehydration synthesis (general): monomers formation of polymers formula:
ATP hydrolysis and phosphorylation
ATP + H2O -> ADP + Pi + energy
Triglyceride formation
Glycerol + 3 Fatty Acids -> Triglyceride + 3H2O
Carbohydrate generic formula
(CH2O)n
Glucose formula (example)
C6H12O6
DNA and RNA bases (illustrative)
DNA: bases A, C, G, T; RNA: A, C, G, U
Hydrogen bonding in water
Hydrogen bonds: H interacts with electronegative atoms (e.g., O in adjacent molecules); crucial for water’s properties and biomolecule structures.