The science of matter — its composition, structure, properties, and transformations
The study of carbon-based compounds. Over 10 million known organic compounds. Includes hydrocarbons, polymers, pharmaceuticals, and the molecules of life.
Compounds not based on carbon. Metals, minerals, coordination compounds, and catalysts. Includes ionic salts, acids, bases, and transition metal chemistry.
Applies physics and math to chemical systems. Thermodynamics, kinetics, quantum chemistry, spectroscopy, and electrochemistry.
Identifies and quantifies matter. Spectroscopy (NMR, MS, IR), chromatography (HPLC, GC), titrations. Underpins quality control and forensics.
Chemistry of living systems. Proteins, DNA, carbohydrates, lipids, enzymes, metabolism. The bridge between chemistry and biology.
Radioactive decay, nuclear reactions, isotopes. Alpha, beta, gamma radiation. Fission and fusion. Radiometric dating and nuclear medicine.
| Law / Constant | Value / Meaning |
|---|---|
| Avogadro's Number (N_A) | 6.022 × 10²³ particles per mole |
| Gas Constant (R) | 8.314 J·mol⁻¹·K⁻¹ |
| Faraday's Constant (F) | 96,485 C/mol (charge per mole of electrons) |
| Ideal Gas Law | PV = nRT |
| Law of Conservation of Mass | Mass is neither created nor destroyed in reactions |
| Law of Definite Proportions | A compound always has the same ratio of elements by mass |
| Law of Multiple Proportions | Two elements can form multiple compounds with whole-number ratios |
| Gibbs Free Energy | ΔG < 0: spontaneous; ΔG > 0: non-spontaneous; ΔG = 0: equilibrium |
118 confirmed elements, organized by increasing atomic number. Rows (periods) indicate electron shells; columns (groups) share chemical properties.
Alkali Metals — Group 1, highly reactive, soft, low melting points (Li, Na, K…)
Alkaline Earth — Group 2, reactive metals (Be, Mg, Ca…)
Transition Metals — d-block, many oxidation states, catalytic
Halogens — Group 17, highly electronegative (F, Cl, Br, I)
Noble Gases — Group 18, stable, full outer shells (He, Ne, Ar…)
| Rank | Element | In Earth's crust |
|---|---|---|
| 1 | Oxygen (O) | 46.6% |
| 2 | Silicon (Si) | 27.7% |
| 3 | Aluminum (Al) | 8.1% |
| 4 | Iron (Fe) | 5.0% |
| 5 | Calcium (Ca) | 3.6% |
| — | In human body: O, C, H, N dominate |
| Element | Fact |
|---|---|
| Hydrogen (H) | Lightest, most abundant in universe (75% of all matter) |
| Carbon (C) | Basis of all life; forms graphite, diamond, graphene, fullerenes |
| Gold (Au) | Inert, highly conductive; won't corrode under normal conditions |
| Mercury (Hg) | Only liquid metal at room temperature |
| Fluorine (F) | Most electronegative element; reacts with almost everything |
| Uranium (U) | Heaviest naturally occurring element; radioactive |
Atomic Radius: Decreases left→right (more protons pull electrons in); increases top→bottom (more shells).
Electronegativity: Increases left→right, decreases top→bottom. Fluorine is highest (3.98 Pauling).
Ionization Energy: Energy to remove an electron. Increases left→right, decreases top→bottom.
Metallic Character: Increases left and down. Most metallic: Fr, Cs, Rb.
| Era | Elements Known | How Discovered |
|---|---|---|
| Ancient | C, S, Fe, Cu, Ag, Au, Hg, Sn, Pb | Naturally occurring, isolated by smelting |
| 1700s | H, O, N, Cl, Mn, Mo, W, and ~20 more | Combustion experiments, electrolysis |
| 1800s | Na, K, Ca, Ba, Li, Al, He, F, noble gases, rare earths | Electrolysis (Davy), spectroscopy |
| 1900s | Radioactive elements, transuranic (Pu, Am, Cm…) | Radioactive decay, nuclear reactors, particle accelerators |
| 2000s | Fl (114), Mc (115), Ts (117), Og (118) | Heavy-ion collisions (Dubna, GSI, RIKEN) |
Nucleus: Contains protons (positive) and neutrons (neutral). ~99.97% of atom's mass. Diameter ~1 fm (10⁻¹⁵ m).
Electrons: Negative, in orbitals around nucleus. Mass = 1/1836 of proton. Determine chemical behavior.
Electrons fill orbitals in order of increasing energy (Aufbau principle): 1s, 2s, 2p, 3s, 3p, 4s, 3d…
Each orbital holds 2 electrons (Pauli). Electrons fill orbitals singly before pairing (Hund's rule).
| Orbital | Shape | Max e⁻ |
|---|---|---|
| s | Sphere | 2 |
| p | Dumbbell (3 orientations) | 6 |
| d | Cloverleaf (5 orientations) | 10 |
| f | Complex (7 orientations) | 14 |
Isotopes have the same number of protons but different neutrons. ¹H, ²H (deuterium), ³H (tritium).
Radioactive decay types:
α decay: emits He nucleus (⁴He) — penetration: paper
β⁻ decay: neutron → proton + electron + antineutrino — penetration: aluminum sheet
γ radiation: high-energy photon — penetration: lead/concrete
Electrons don't orbit like planets — they exist as probability distributions described by wave functions ψ. The probability of finding an electron at a given point is |ψ|².
The four quantum numbers (n, l, m_l, m_s) fully describe each electron's state. No two electrons in an atom can share all four (Pauli Exclusion).
When electrons absorb or emit photons, they jump between energy levels. Each element produces a unique spectrum — its "fingerprint."
Emission spectra: excited atoms release specific wavelengths. Used to identify elements in distant stars, in lab analysis (flame tests), and in neon signs.
Formed when one atom transfers electrons to another. Metal + nonmetal. Produces oppositely charged ions held together by electrostatic attraction.
Examples: NaCl (table salt), MgO, CaCO₃
Properties: high melting points, conduct electricity when dissolved or molten, brittle crystalline solids.
Atoms share electron pairs. Nonmetal + nonmetal. Can be single (1 pair), double (2 pairs), or triple (3 pairs) bonds.
Polar: unequal sharing. Nonpolar: equal sharing (H₂, N₂, O₂).
Metal atoms release valence electrons into a "sea" shared by all atoms. This "electron sea" model explains metals' properties.
✓ Electrical conductivity (mobile electrons)
✓ Thermal conductivity
✓ Malleability and ductility (layers slide)
✓ Metallic luster (electrons reflect light)
Hydrogen bonding: N, O, or F bonded to H creates strong dipole. Explains water's high boiling point and DNA's double helix.
Dipole-dipole: Polar molecules attract. CO, HCl.
London dispersion: Temporary dipoles in all molecules. Only force in noble gases. Increases with molecular mass.
Strength: ionic > H-bond > dipole-dipole > dispersion
Valence Shell Electron Pair Repulsion: electron pairs arrange to minimize repulsion, determining molecular geometry.
| Pairs | Shape | Example |
|---|---|---|
| 2 | Linear (180°) | CO₂, BeCl₂ |
| 3 | Trigonal planar (120°) | BF₃, SO₃ |
| 4 | Tetrahedral (109.5°) | CH₄, NH₄⁺ |
| 4+1 lone | Trigonal pyramidal | NH₃, PCl₃ |
| 4+2 lone | Bent (104.5°) | H₂O |
Electronegativity (EN) measures an atom's pull on shared electrons (Pauling scale: 0.7–4.0).
ΔEN < 0.5: nonpolar covalent | 0.5–1.7: polar covalent | > 1.7: ionic
Water (EN_O=3.44, EN_H=2.20, ΔEN=1.24): polar covalent. This polarity makes water an extraordinary solvent — the "universal solvent."
| Type | Pattern | Example |
|---|---|---|
| Synthesis | A + B → AB | 2H₂ + O₂ → 2H₂O |
| Decomposition | AB → A + B | 2H₂O₂ → 2H₂O + O₂ |
| Single replacement | A + BC → AC + B | Zn + 2HCl → ZnCl₂ + H₂ |
| Double replacement | AB + CD → AD + CB | NaCl + AgNO₃ → NaNO₃ + AgCl |
| Combustion | CₓHᵧ + O₂ → CO₂ + H₂O | CH₄ + 2O₂ → CO₂ + 2H₂O |
| Acid-base | HA + BOH → BA + H₂O | HCl + NaOH → NaCl + H₂O |
| Redox | Electron transfer | 2Fe + 3Cl₂ → 2FeCl₃ |
Exothermic: ΔH < 0 — releases heat (combustion, respiration, neutralization)
Endothermic: ΔH > 0 — absorbs heat (photosynthesis, dissolving ammonium nitrate)
Spontaneous when ΔG < 0. A reaction can be spontaneous even if endothermic if entropy gain is large enough.
Activation energy Eₐ is the energy barrier reactions must overcome. Catalysts lower Eₐ without being consumed. Temperature increase speeds most reactions (~doubles per 10°C).
Le Chatelier's Principle: if a system at equilibrium is disturbed, it shifts to counteract the disturbance.
K > 1: products favored | K < 1: reactants favored
Haber process: N₂ + 3H₂ ⇌ 2NH₃. High pressure, moderate temp, iron catalyst optimize yield.
Arrhenius: Acid produces H⁺; base produces OH⁻ in water.
Brønsted-Lowry: Acid is a proton donor; base is a proton acceptor.
Lewis: Acid is an electron-pair acceptor; base is an electron-pair donor (broadest definition).
Strong acids: HCl, H₂SO₄, HNO₃, HBr, HI, HClO₄. Fully dissociate in water.
Galvanic cell: Spontaneous redox reaction generates electricity (batteries).
Electrolytic cell: Uses electricity to drive non-spontaneous reactions (electrolysis of water, electroplating).
Carbon has 4 valence electrons, can form 4 bonds, bonds stably with itself and H/N/O/S, forms chains, rings, and branched structures. This versatility creates millions of distinct compounds.
Carbon allotropes: graphite (layered hexagons), diamond (tetrahedral lattice), graphene (single atom sheet), fullerenes (C₆₀ "Buckyballs"), carbon nanotubes.
| Group | Formula | Class |
|---|---|---|
| Hydroxyl | −OH | Alcohol |
| Carbonyl | −C=O | Aldehyde/Ketone |
| Carboxyl | −COOH | Carboxylic acid |
| Amino | −NH₂ | Amine |
| Ester | −COO− | Ester |
| Phosphate | −OPO₃²⁻ | DNA backbone |
| Thiol | −SH | Thiol (cysteine) |
Alkanes (CₙH₂ₙ₊₂): All single bonds. Methane (CH₄), ethane, propane, butane. Saturated; used as fuels.
Alkenes (CₙH₂ₙ): One C=C double bond. Ethylene (C₂H₄), propene. Reactive; monomers for polyethylene.
Alkynes (CₙH₂ₙ₋₂): One C≡C triple bond. Acetylene (ethyne) used in welding.
Aromatics: Benzene ring (C₆H₆) — delocalized π electrons. Stable; basis for pharmaceuticals and dyes.
Giant molecules formed by linking many monomers. Backbone can be carbon-carbon or include O, N, Si.
| Polymer | Monomer | Use |
|---|---|---|
| Polyethylene | Ethylene | Plastic bags, bottles |
| Nylon | Diamines + diacids | Textiles, gears |
| DNA | Nucleotides | Genetic code |
| Protein | Amino acids | Enzymes, structure |
| Cellulose | Glucose | Plant cell walls |
Carbohydrates: (CH₂O)ₙ — energy storage and structure. Glucose, sucrose, starch, cellulose.
Lipids: Triglycerides, phospholipids, steroids. Hydrophobic; cell membranes, energy reserves.
Proteins: Amino acid polymers. Enzymes, antibodies, hormones, structural (collagen, keratin).
Nucleic Acids: DNA (deoxyribose) and RNA (ribose) — store and transmit genetic information.
Substitution (SN1/SN2): One group replaces another. SN2 is concerted (backside attack); SN1 forms a carbocation intermediate.
Addition: Atoms add across a double/triple bond. Hydrogenation: C=C + H₂ → C−C.
Elimination: Removes atoms to form a double bond. E2 and E1 pathways.
Condensation: Two molecules join with loss of water. Forms esters, peptides, glycosidic bonds.