Chemistry Regents Review Summary

Overview

• Full review of chemistry topics for Regents exam covering 12 units.
• Emphasis on practice problems (~150+), use of reference table, and attempting problems before viewing solutions.
• Notes summarize key concepts, formulas, and typical problem approaches from the transcript.

States Of Matter (Particle Diagrams)

• Gas: molecules far apart, random motion; draw separate pairs with arrows for movement.
• Liquid: molecules close together but not fixed; often drawn within a container.
• Solid vs gas: solids ordered and fixed, gases disordered and random.

Mixtures And Separation

• Filtration: separates sand by particle size (sand insoluble/large particles).
• Homogeneous vs Heterogeneous:
  • Homogeneous: uniform distribution (e.g., salt dissolved in water).
  • Heterogeneous: components not uniformly distributed (e.g., sand in water).
• Evidence of varying proportion: different masses of NH4Cl and sand in mixtures A and B.

Solutions, Solubility, Saturation

• Solubility curves: use table G; values usually per 100 g H2O.
• Saturation classification: compare actual solute per 100 g H2O to curve (under = unsaturated; on = saturated; above = supersaturated).
• Solubility increases/decreases with temperature depending on solute (check curve).
• Example: K3 solubility at 50°C ≈ 84 g per 100 g H2O (interpolation).

Phase Changes, Heat Flow, Entropy

• Sublimation: solid → gas directly (dry ice CO2(-78°C) in air 21°C inflates balloon).
• Heat flow direction: from higher temperature to lower (air → dry ice).
• Entropy: measure of disorder; solid CO2 (low entropy) → gas CO2 (higher entropy).

Heat, Specific Heat, Thermal Energy

• ΔQ = m × c × ΔT
• Example setup: 800 g water, c = 4.18 J/g·°C, ΔT = 10°C → numerical setup m·c·ΔT.
• Direction of heat transfer: surroundings → water when water temperature increases.
• Average kinetic energy ∝ temperature (higher T → greater average kinetic energy).

Atomic/Periodic Concepts

• Isotopes: same protons, different neutrons (H-3 has 1 proton, 2 neutrons; H-1 has 0 neutrons).
• Atomic number = protons; mass number = protons + neutrons.
• Noble-gas electron configurations: positive ions often adopt noble gas configuration (e.g., K+ → Ar).
• Group trends:
  • Valence electrons same within group (Group 14 → 4 valence electrons).
  • Atomic radius increases down group due to added electron shells.
  • Group 18 elements rarely form compounds (full valence shells).

Models Of Atom & Historical Experiments

• Rutherford gold foil: atom mostly empty space; nucleus is small and positive; protons identified as positive nuclear particles.
• Thomson model (plum pudding): first to include electrons.
• Bohr and later models: electrons in energy levels; emission/absorption produce line spectra.

Spectra And Light Emission

• Emission: electrons absorb energy → move to higher levels; return to lower levels → emit photons (discrete wavelengths).
• Spectral lines independent of mass of sample; depend on electronic transitions.
• Presence/absence: spectral line matching indicates element present; missing lines indicate element absent.

Bonding, Lewis Structures, Polarity

• Ionic bonding: metal + nonmetal (electron transfer); e.g., K → Br forms ionic compound.
• Covalent bonding: nonmetal + nonmetal (electron sharing); e.g., Br2 single covalent bond, Cl2.
• Lewis diagrams: count total valence electrons, form bonds, place remaining electrons as lone pairs.
• Polarity:
  • Molecule polar if unequal charge distribution (e.g., HCl is polar).
  • Nonpolar if symmetrical charge distribution (e.g., CH4, symmetric COH structures).
• Bond polarity determined by electronegativity difference (greater difference → more polar bond).

Intermolecular Forces And Physical Properties

• Stronger intermolecular forces → higher boiling and melting points.
• Example evidence: NH3 has higher boiling/melting values than CF4 → NH3 has stronger forces (hydrogen bonding/dipole interactions).
• Solubility: "like dissolves like" (polar solutes dissolve in polar solvents).

Gas Laws

• Ideal gas suggestions: high temperature, low pressure make real gas behave ideally.
• Combined gas law: P1V1/T1 = P2V2/T2 (use for conversions to STP).
• Molar volume at STP: 1 mole gas ≈ 22.4 L.
• Density = mass/volume (use formula mass and 22.4 L for gas at STP).
• Examples:
  • Density of HBr at STP: mass(80.9 g) / 22.4 L ≈ 3.61 g/L.
  • Use combined gas law for volume changes to STP (setup P1V1/T1 = P2V2/T2).

Stoichiometry And Reaction Types

• Moles ↔ grams: use molar mass (grand formula mass).
• Stoichiometric ratios from balanced equations used in multi-step conversions (mass → moles → mole ratio → product mass).
• Percent composition: (mass of element in formula / molar mass of compound) × 100.
• Empirical formula: simplify subscripts by greatest common divisor.
• Reaction types:
  • Decomposition (one compound → simpler substances).
  • Addition (unsaturated hydrocarbons plus halogen → addition product).
  • Neutralization: acid + base → water + salt (write salt from remaining ions).

Example stoichiometry:

• C6H12O6 → 2 C2H5OH + 2 CO2: 270 g glucose → calculate moles glucose → mole ratio → mass ethanol (final example: 138 g ethanol).

Concentration, Molarity, Titration

• Molarity M = moles solute / liters solution.
• Titration relation: MaVa = MbVb for monoprotic acid-base titrations (use volumes in liters).
• Sig figs for concentration: round to least significant figures from measurements involved.
• pH and hydronium concentration: each pH unit = factor of 10 change in [H3O+].
• Indicators: color depends on pH (reference Table M for indicator ranges).

Colligative Properties And Freezing Point Depression

• Freezing point depression: higher solute concentration → lower freezing point.
• Compare solutions by molality/molarity: greater concentration → larger depression.

Kinetics And Catalysts

• Reaction rate increased by: higher temperature (more frequent, energetic collisions), increased concentration (more collisions), higher surface area (more collision opportunities), catalysts.
• Catalyst effect: lowers activation energy; does not change ΔH of reaction; increases reaction rate.

Equilibrium And Le Chatelier’s Principle

• Equilibrium: forward and reverse rates equal in closed system; concentrations constant.
• Le Chatelier: system shifts to counteract stress (change in concentration, pressure, temperature).
  • Increase pressure → equilibrium shifts toward side with fewer gas moles.
  • Decrease in reactant concentration → equilibrium shifts to produce more of that reactant side accordingly.
• Closed container required to maintain equilibrium (prevent loss of matter).

Electrochemistry And Cells

• Voltaic (galvanic) cell: spontaneous chemical reaction produces electrical energy (no external power source).
  • Anode: oxidation (mass decreases as metal atoms become ions).
  • Cathode: reduction (mass increases as ions gain electrons).
  • Electrons flow through wire from anode → cathode.
  • Salt bridge allows ion flow to maintain charge balance.
• Electrolysis: requires external power source (battery) to drive non-spontaneous reactions.
• Half-reactions: balance mass and charge; electrons appear on appropriate side (oxidation loses electrons; reduction gains electrons).

Organic Chemistry: Hydrocarbons And Functional Groups

• Alkane general formula: CnH2n+2 (single bonds).
• Alkenes: contain C=C double bonds (unsaturated).
• Alkynes: contain C≡C triple bonds (unsaturated).
• Naming and structure:
  • Prefix indicates number of carbons (meth-, eth-, prop-, but-, pent-).
  • Suffix indicates functional group (‑ane, ‑ene, ‑yne, ‑anol, etc.).
  • Branched naming: numbers indicate positions of substituents (e.g., 2,2,4‑trimethylpentane).
• Functional groups:
  • Alcohol (‑OH): ethanol → alcohol class.
  • Ester: contains C=O adjacent to O single bond (R‑COO‑R).

Nuclear Chemistry And Radioactivity

• Alpha decay: emits He nucleus (mass -4, protons -2).
• Beta decay: neutron → proton + beta particle (electron), increases proton number by 1.
• Nuclear notation: top = mass number, bottom = atomic number, symbol = element.
• Half-life calculations:
  • Fraction remaining = (1/2)^(number of half-lives).
  • Number of half-lives = elapsed time / half-life.
• Nuclear fission releases far more energy per mole than chemical combustion.
• Penetration: gamma rays most penetrating; beta particles less penetrating and deflect in magnetic fields.
• Medical uses: radioisotopes in cancer treatment risk damaging healthy tissue.

Key Terms And Definitions

• Sublimation: solid → gas.
• Entropy: measure of disorder.
• Molarity (M): moles solute per liter solution.
• ΔQ (thermal energy change): m·c·ΔT.
• Half-life: time for half of radioactive sample to decay.
• Isomer: same molecular formula, different structural formula.
• Catalyst: lowers activation energy (increases reaction rate).
• Le Chatelier’s Principle: equilibrium shifts to oppose change.

Action Items / Exam Preparation Tips

• Always carry and use the reference table (tables G, H, M, N, P, Q, R, etc.).
• Practice particle diagrams for gas, liquid, and solid phases.
• Practice Lewis structures, molecular polarity, and counting valence electrons.
• Memorize common molar masses and molar volume at STP (22.4 L).
• Practice combined gas law problems and stoichiometry multi-step conversions.
• Review half-life problems and nuclear notation balancing exercises.
• Practice titration calculations and sig-fig rules for final reported values.
• Work through representative practice problems for each unit before exam day.