MTEL Chemistry cheat sheet

Atomic structure: Z = protons = electrons in a neutral atom; mass number A = protons + neutrons; isotopes differ only in neutron count, so they share chemical behavior but differ in mass.

Electron configuration follows Aufbau, Pauli, and Hund rules; valence electrons (outermost s and p) govern periodic trends and bonding; transition metals lose 4s electrons before 3d when ionized.

Periodic trends: atomic radius decreases left-to-right across a period and increases down a group; ionization energy and electronegativity increase across a period and decrease down a group.

Mole relationships: n = m/M; N = n × 6.022×10^23; molarity M = mol solute / L solution; dilution M1V1 = M2V2; at STP 1 mol gas occupies 22.4 L.

Stoichiometry: balance the equation, convert grams to moles, apply the mole ratio, convert back; identify the limiting reactant by comparing available mole ratios, not masses.

Gas laws: PV = nRT with R = 0.0821 L·atm/(mol·K); combined gas law P1V1/T1 = P2V2/T2; Dalton's law Ptotal = sum of partial pressures; temperature must be in kelvin.

Thermochemistry: q = mcΔT; ΔH negative = exothermic, positive = endothermic; Hess's law sums step enthalpies; bond breaking absorbs energy, bond formation releases energy.

Solutions and colligative properties: molality m = mol solute / kg solvent; boiling point elevation ΔTb = i·Kb·m; freezing point depression ΔTf = i·Kf·m; van't Hoff factor i counts dissolved particles.

Kinetics: rate depends on concentration, temperature, surface area, and catalysts; rate law is determined experimentally, not from coefficients; catalysts lower activation energy without changing ΔH.

Equilibrium: K expresses products over reactants raised to coefficients; Q compared with K predicts direction (Q<K forward, Q>K reverse); only temperature changes K; Le Chatelier predicts shifts from stress.

Acid-base: pH = -log[H3O+]; pOH = -log[OH-]; pH + pOH = 14 at 25 °C; Kw = 1.0×10^-14; Ka·Kb = Kw for a conjugate pair; strong acids and bases ionize completely.

Buffers resist pH change via a weak acid/conjugate base pair; Henderson-Hasselbalch pH = pKa + log([base]/[acid]); buffer capacity is greatest near pH = pKa.

Redox: oxidation is loss of electrons (anode), reduction is gain (electrons, cathode); assign oxidation numbers; balance by half-reactions conserving charge and atoms.

Electrochemistry: E°cell = E°cathode - E°anode; positive E°cell means spontaneous galvanic cell; ΔG° = -nFE°cell; electrolysis uses external voltage to drive nonspontaneous reactions.

Nuclear chemistry: alpha decay loses He-4 (A-4, Z-2); beta-minus decay converts a neutron to a proton (Z+1, A unchanged); half-life t1/2 relates to decay constant by t1/2 = ln2/k.

Bonding and intermolecular forces: ionic bonds transfer electrons, covalent bonds share them; polarity from electronegativity difference; London dispersion < dipole-dipole < hydrogen bonding in strength.

VSEPR geometry from electron domains: 2 = linear, 3 = trigonal planar, 4 = tetrahedral (109.5°), 3 bonds + 1 lone pair = trigonal pyramidal, 2 bonds + 2 lone pairs = bent.

Hybridization: steric number 2 = sp, 3 = sp2, 4 = sp3; sp2 centers give trigonal planar geometry and are found in alkenes and aromatic rings.

Organic functional groups: alcohols (-OH), aldehydes (terminal C=O with H), ketones (internal C=O), carboxylic acids (-COOH), esters (-COO-), amines (-NH2), amides (-CONH-).

Organic reaction patterns: addition across double bonds, substitution at saturated carbons, elimination forming alkenes, oxidation of alcohols to aldehydes/ketones/carboxylic acids, esterification of acid plus alcohol.

Isomerism: structural isomers differ in connectivity; stereoisomers include enantiomers (mirror images, chiral centers) and diastereomers including cis/trans geometric isomers.

Laboratory safety: add acid to water, never water to acid; use fume hood for volatiles; know SDS hazard categories; eyewash and safety shower locations; never pipette by mouth.

Separation techniques: filtration removes solids, distillation separates by boiling point, chromatography separates by differential affinity, extraction uses solubility differences, recrystallization purifies by solubility.

Data and measurement: significant figures follow addition (decimal places) and multiplication (least sig figs) rules; percent error = |measured - accepted| / accepted × 100; calibrate instruments before use.

Spectroscopy: IR identifies functional groups (C=O near 1700 cm^-1, O-H broad 3200-3600 cm^-1); NMR reveals hydrogen environments and splitting; mass spectrometry gives molecular mass and fragmentation.

Beer-Lambert law A = εbc: absorbance is proportional to concentration and path length; used for colorimetric concentration determination and calibration curves.

Solubility rules: most nitrate, sodium, potassium, and ammonium salts are soluble; most chlorides soluble except Ag, Pb, Hg; most carbonates, phosphates, and hydroxides are insoluble.

Precipitation and net ionic equations: cancel spectator ions; a precipitate forms when ion product exceeds Ksp; Ksp expresses dissolution equilibrium for a sparingly soluble salt.

Collision theory and reaction energy diagrams: effective collisions need proper orientation and energy ≥ activation energy; transition state is the highest-energy point; intermediates are local minima.

Periodic group properties: alkali metals react vigorously with water, halogens are diatomic oxidizing agents decreasing in reactivity down the group, noble gases are inert due to full valence shells.

Empirical and molecular formulas: empirical from percent composition by dividing moles by the smallest value; molecular formula = empirical formula × (molar mass / empirical mass).

Phase changes and heating curves: melting and boiling occur at constant temperature with energy going to breaking intermolecular forces; heat of fusion and vaporization apply at plateaus.

Scientific inquiry and experimental design: identify independent, dependent, and controlled variables; use a control group; replicate trials; distinguish hypothesis, theory, and law.

Error analysis: random error affects precision, systematic error affects accuracy; report mean with standard deviation; distinguish percent error from percent difference.

Graphing conventions: independent variable on x-axis, dependent on y-axis; slope of a linear plot gives a rate or constant; R² near 1 indicates good linear fit.

Models and explanations: use particle-level diagrams to explain macroscopic observations; distinguish observations from inferences; recognize that models are revised with new evidence.

Safety and ethics in chemistry: dispose of chemicals per regulations, never down the drain; report spills; use personal protective equipment; handle unknowns with appropriate caution.

Chemistry in society: connect content to real applications such as water treatment, batteries, pharmaceuticals, and materials; evaluate benefits and risks using evidence.

Mathematical reasoning: dimensional analysis cancels units; check that final units match the quantity requested; estimate orders of magnitude to catch calculation errors.

Reading and interpreting data: distinguish correlation from causation; identify outliers with justification; recognize that equilibrium and rate data require different interpretations.

Matter classification: pure substances (elements, compounds) versus mixtures (homogeneous, heterogeneous); physical changes alter state or form, chemical changes form new substances.

Atomic models and evidence: Dalton's atomic theory, Thomson's cathode-ray discovery of electrons, Rutherford's gold-foil nucleus, Bohr quantized orbits, quantum mechanical model.

Periodic table organization: groups share valence electron counts; periods correspond to principal energy levels; metals, metalloids, and nonmetals are separated by a staircase boundary.

Chemical equations: conserve mass and charge; state symbols (s, l, g, aq) indicate phase; combination, decomposition, single replacement, double replacement, and combustion are common types.

Thermodynamics and spontaneity: ΔG = ΔH - TΔS; spontaneous when ΔG < 0; entropy increases with disorder and phase changes toward gas; second law governs direction.

Solutions and concentration units: percent by mass, mole fraction, molarity, and molality differ; molarity changes with temperature due to volume expansion, molality does not.

Gas behavior and kinetic molecular theory: gas particles move randomly, collisions are elastic, average kinetic energy is proportional to kelvin temperature; real gases deviate at high pressure and low temperature.

Laboratory measurement tools: graduated cylinders for volume, analytical balances for mass, burettes for titration delivery, volumetric flasks for precise dilution; read meniscus at eye level.

Titration analysis: at equivalence moles of acid equal moles of base adjusted by stoichiometry; indicator choice depends on pH at equivalence; strong-strong equivalence is pH 7.

Electron configuration notation: use noble gas shorthand; exceptions include Cr ([Ar]4s1 3d5) and Cu ([Ar]4s1 3d10) due to stability of half-filled and filled subshells.

Molecular polarity: a molecule is polar only if bond dipoles do not cancel; CO2 is nonpolar despite polar bonds, while H2O is polar due to bent geometry.

Reaction yield: percent yield = actual/theoretical × 100; theoretical yield comes from limiting reactant; losses from side reactions, transfer, and incomplete reaction reduce yield.

Acid strength and structure: stronger acids stabilize the conjugate base; oxyacids strengthen with more electronegative central atoms and more oxygen atoms; binary acids strengthen down a group.

Radioactive decay and dating: carbon-14 dating applies to once-living material within roughly 50,000 years; isotope half-lives are constant regardless of conditions.

Safety data and hazard communication: GHS pictograms signal flammability, corrosivity, toxicity, and reactivity; SDS sections describe handling, storage, and first aid.

Interpreting experimental results: distinguish precision (repeatability) from accuracy (closeness to true value); significant figures communicate measurement uncertainty.

Redox and oxidation numbers: free elements are 0; monatomic ions equal charge; oxygen usually -2, hydrogen usually +1; sum equals overall charge for polyatomic species.

Electrochemical cell notation: anode oxidation on left, cathode reduction on right; salt bridge maintains charge balance; electrons flow anode to cathode in the external circuit.

Reaction rates and factors: increasing temperature raises collision frequency and energy; catalysts provide alternate pathway; concentration and surface area affect frequency of collisions.

Equilibrium calculations: ICE tables track initial, change, and equilibrium concentrations; small K permits approximation; check validity of approximation against Ka or Kb.

Organic nomenclature basics: IUPAC names use longest carbon chain, number to give lowest locants, and suffix for principal functional group; alkanes end in -ane.

Bond energy and reaction enthalpy: ΔH ≈ sum of bonds broken minus sum of bonds formed; strong bonds in products release more energy and favor exothermic reactions.

Chemical equilibrium and solubility: common ion effect lowers solubility; pH affects solubility of salts with basic or acidic anions; temperature shifts Ksp.

Scientific literacy in chemistry: evaluate sources, distinguish peer-reviewed evidence from anecdote, and interpret units and scales correctly when reading chemical data.

Acid-base definitions: Arrhenius produces H+ or OH- in water; Brønsted-Lowry transfers protons; Lewis accepts or donates electron pairs; definitions broaden across contexts.

pH calculations: strong acid pH = -log(concentration); weak acid requires Ka and equilibrium; salt of weak acid with strong base gives basic solution.

Molecular geometry and polarity prediction: draw Lewis structure, count electron domains, apply VSEPR, then assess dipole cancellation to determine net polarity.

Thermochemical equations: reversing a reaction changes sign of ΔH; multiplying coefficients multiplies ΔH; adding reactions adds ΔH under Hess's law.

Gas stoichiometry: convert gas volume to moles using PV = nRT or molar volume at STP, then apply mole ratios; account for limiting reactant and percent yield.

Periodic table and electron shielding: effective nuclear charge increases across a period, pulling electrons closer; shielding by inner electrons increases down a group.

Chemical formulas and naming: ionic compounds name cation then anion; transition metals require Roman numerals; molecular compounds use prefixes such as mono, di, tri.

Laboratory calculations: percent composition, molar mass, and empirical formula derive from measured masses; round only at the final step to avoid compounding error.

Enthalpy of solution: lattice energy released and hydration energy absorbed determine whether dissolution is exothermic or endothermic; entropy also drives dissolution.

Reaction spontaneity and temperature: when ΔH and ΔS have the same sign, temperature determines spontaneity; high temperature favors entropy-driven processes.

Nuclear equations: conserve mass number and atomic number; positron emission decreases Z by 1 with A unchanged; gamma emission changes neither A nor Z.

Organic reaction mechanisms: nucleophiles donate electron pairs, electrophiles accept; substitution versus elimination depends on substrate, base strength, and solvent.

Spectroscopy interpretation strategy: IR for functional groups, NMR for hydrogen environments and coupling, MS for molar mass; combine evidence rather than relying on one technique.

Laboratory safety for heating and mixing: use boiling chips, point test tubes away from people, never heat sealed containers, add reagents slowly with stirring.

Data interpretation in kinetics: zero-order plots concentration versus time linearly; first-order plots ln[concentration] versus time; second-order plots 1/[concentration] versus time.

Chemical bonding and properties: ionic compounds have high melting points and conduct when molten; molecular covalent compounds have lower melting points and often do not conduct.

Mole concept applications: molar mass converts between mass and moles; Avogadro's number converts moles to particles; molar volume applies only to gases at specified conditions.

Solution preparation: dissolve solute in less than final volume, then dilute to the mark; volumetric glassware is calibrated to contain at a stated temperature.

Titration calculations: moles = molarity × volume in liters; use stoichiometric ratio from balanced equation; account for dilution and aliquot factors.

Electron transitions and light: absorbed or emitted photon energy equals energy difference between levels; E = hν = hc/λ; higher energy corresponds to shorter wavelength.

Periodic table and ion formation: metals lose electrons to form cations with noble-gas configurations; nonmetals gain electrons to form anions; ion charge relates to group number.

Chemical equilibrium and catalysts: catalysts speed approach to equilibrium but do not change equilibrium position or K; only temperature alters the value of K.

Chemical equations and conservation: balance atoms and charge; for redox in acidic or basic solution add H+, OH-, and H2O as needed to balance.

Scientific practices: pose questions, plan investigations, analyze data, construct explanations, and communicate results using evidence and appropriate chemical models.

Colligative properties and electrolytes: van't Hoff factor i equals number of ions per formula unit for strong electrolytes; weak electrolytes have i between 1 and the maximum.

Acid-base titration curves: strong-strong has sharp equivalence near pH 7; weak acid-strong base has buffer region and equivalence above pH 7; half-equivalence pH equals pKa.

Redox spontaneity: compare standard reduction potentials; the species with higher E° is reduced; calculate E°cell and relate sign to ΔG and spontaneity.

Organic isomer counting: consider chain branching, position of functional groups, and stereocenters; enantiomers share physical properties except optical rotation.

Gas collection and measurement: correct for water vapor pressure when collecting gas over water; dry gas pressure equals total minus vapor pressure.

Error and uncertainty in measurement: report uncertainty with units; use propagation rules for addition and multiplication; distinguish absolute from relative uncertainty.

Periodic trends and reactivity: metallic character increases down and left; nonmetallic character increases up and right; reactivity patterns follow these trends.

Chemical safety and waste: segregate incompatible wastes, label containers, neutralize acids and bases before disposal per local rules, and document incidents.

Molar mass determination: sum atomic masses from formula; use for stoichiometry; molar mass of a gas can be found from density using M = dRT/P.

Empirical formula from combustion data: convert CO2 and H2O masses to moles of C and H; oxygen by difference; divide by smallest mole value.

Reaction types and predictions: use solubility rules and activity series to predict products; activity series ranks metals by tendency to oxidize.

Electron configuration and periodic position: group number relates to valence electrons for main-group elements; period number equals highest principal quantum level.

Thermodynamic calculations: use standard enthalpies of formation ΔH°rxn = sum products minus sum reactants; elements in standard states have ΔH°f = 0.

Laboratory technique: use proper significant figures, tare balances, read burette to two decimal places, and record observations immediately.

Chemical equilibrium applications: Haber and other industrial processes adjust temperature and pressure to optimize yield based on Le Chatelier's principle.

Scientific explanation quality: claims must be supported by data and chemical principles; distinguish hypothesis from conclusion and cite evidence explicitly.

Atomic mass and isotopes: average atomic mass is the weighted average of isotope masses by natural abundance; use percent abundances as weights.

Quantum numbers: n principal, l angular momentum (0 to n-1), ml magnetic, ms spin; Pauli exclusion limits two electrons per orbital with opposite spins.

Periodic table and bonding: ionic bonding occurs between metals and nonmetals; covalent between nonmetals; metallic bonding involves delocalized electrons.

Reaction stoichiometry with solutions: combine molarity and volume to get moles; use mole ratios; convert back to mass or volume as requested.

Acid-base and equilibrium: strong acid fully dissociates; weak acid partially; use Ka to find [H+] and pH; check assumptions and approximations.

Nuclear stability: neutron-to-proton ratio predicts stability; too many neutrons favor beta decay; too many protons favor positron emission or electron capture.

Organic chemistry and isomers: identify functional groups, degree of unsaturation, and stereocenters; use IUPAC naming rules to communicate structure.

Experimental design and safety: control variables, use appropriate glassware, wear goggles and gloves, and know emergency procedures for spills and exposure.

Solutions and solubility: like dissolves like; temperature affects solid solubility; gas solubility decreases with temperature and increases with pressure (Henry's law).

Chemical kinetics and mechanisms: rate-determining step controls overall rate; intermediates appear in mechanism but not overall equation; catalysts appear in mechanism.

Electrochemistry applications: batteries, corrosion, and electrolysis rely on redox; sacrificial anodes protect metals by being oxidized preferentially.

Periodic table and atomic size: cations are smaller than parent atoms; anions are larger; isoelectronic species shrink as nuclear charge increases.

Measurement and significant figures: exact numbers have infinite sig figs; leading zeros are not significant; trailing zeros after a decimal are significant.

Laboratory calculations and yield: theoretical yield from limiting reactant; actual yield measured; percent yield assesses efficiency and purity.

Chemical bonding and molecular shape: multiple bonds count as one electron domain; lone pairs repel more than bonding pairs, reducing bond angles.

Scientific literacy: interpret chemical formulas, units, and data tables accurately; evaluate claims using evidence and identify sources of bias.

Thermochemistry and calorimetry: q = mcΔT for solutions; qreaction = -qcalorimeter; use bomb calorimetry for constant-volume heat; ΔH relates to q at constant pressure.

Equilibrium and acid-base: buffer pH depends on ratio of conjugate base to acid; adding strong acid consumes base; capacity limited by buffer component amounts.

Redox balancing in basic solution: balance as in acid, then add OH- to neutralize H+ and combine to form water; check atoms and charge.

Organic reactions and functional groups: alcohols oxidize to aldehydes or ketones; carboxylic acids react with alcohols to form esters; amines are basic.

Nuclear chemistry applications: fission splits heavy nuclei releasing energy; fusion combines light nuclei; both conserve mass-energy; E = mc² quantifies mass defect.

Laboratory safety and chemical handling: know hazard pictograms, use fume hood for volatile or toxic substances, and never return unused chemicals to stock bottles.

Data analysis and graphing: calculate slope and intercept; interpret units; use linear regression for calibration; identify systematic versus random deviations.

Chemical education and communication: use models, diagrams, and equations to explain phenomena; connect macroscopic observations to particle-level reasoning.

Periodic properties and ionization: successive ionization energies increase; large jumps indicate noble-gas core removal; first ionization energy trends predict reactivity.

Molecular orbital and bonding concepts: sigma bonds form head-on overlap; pi bonds form side-by-side overlap; double bonds contain one sigma and one pi bond.

Stoichiometry and limiting reactants: compare mole ratios to identify limiting reactant; excess reactant remains; product yield depends on limiting reactant.

Acid-base pH of salts: salt of strong acid and strong base is neutral; strong acid and weak base is acidic; weak acid and strong base is basic.

Electrochemistry and cell potential: standard potentials are reduction potentials; reverse sign for oxidation; E°cell positive indicates spontaneous galvanic reaction.

Laboratory measurement and precision: use appropriate glassware; volumetric flasks for dilution, graduated cylinders for approximate volumes, burettes for titration.

Chemical safety and emergency response: know location of eyewash, shower, fire extinguisher, and spill kit; report all incidents; follow SDS instructions.

Scientific reasoning: distinguish observation from inference; support conclusions with data; recognize that models are approximations revised by evidence.

Atomic structure and ions: protons determine element identity; neutrons determine isotope; electrons determine charge and chemical behavior; ions form by gaining or losing electrons.

Periodic table and electron configuration: s, p, d, f blocks correspond to subshells; valence electrons determine group chemistry; noble gases have full valence shells.

Chemical reactions and equations: balance atoms and charge; identify reaction type; predict products using solubility rules, activity series, and acid-base neutralization.

Mole and stoichiometry: use molar mass, Avogadro's number, molar volume, and mole ratios; identify limiting reactant; calculate theoretical and percent yield.

Solutions and concentration: molarity, molality, percent composition, dilution; colligative properties depend on particle count; electrolytes increase i.

Acids, bases, and equilibrium: pH, pOH, Ka, Kb, Kw; buffers resist pH change; Le Chatelier predicts shifts; K changes only with temperature.

Redox and electrochemistry: oxidation numbers; half-reactions; E°cell; spontaneity via ΔG = -nFE°cell; electrolysis drives nonspontaneous reactions.

Laboratory safety and data: use PPE, add acid to water, dispose of waste properly; record measurements with correct significant figures and units.

Organic chemistry and isomers: structural isomers differ in connectivity; geometric isomers differ in arrangement around double bonds; enantiomers are mirror images.

Spectroscopy and structure: IR identifies functional groups; NMR identifies hydrogen environments; MS gives molecular mass; combine evidence for structure determination.

Thermochemistry and spontaneity: ΔG = ΔH - TΔS; negative ΔG spontaneous; entropy increases with gas formation and temperature; enthalpy from bond energies or formation.

Kinetics and equilibrium: rate laws experimental; catalysts lower activation energy; equilibrium constant expression from balanced equation; Q versus K predicts direction.

Nuclear chemistry: alpha, beta, gamma decay; half-life; mass defect and binding energy; fission and fusion release energy; dating applications.

Laboratory techniques: filtration, distillation, chromatography, extraction, recrystallization; choose based on physical property differences.

Chemical safety and hazard communication: GHS pictograms; SDS sections; PPE; fume hood; waste segregation; emergency procedures.

Scientific practices: ask questions, design investigations, analyze data, construct explanations, and communicate evidence-based conclusions.

Periodic trends and bonding: electronegativity difference predicts ionic, polar covalent, or nonpolar covalent bonding; bond polarity affects molecular properties.

Molecular geometry and VSEPR: count electron domains; lone pairs reduce bond angles; geometry determines polarity and intermolecular forces.

Stoichiometry and gas laws: combine PV = nRT with mole ratios; account for temperature in kelvin; use limiting reactant and percent yield.

Acid-base and buffers: Henderson-Hasselbalch; buffer capacity; titration curves; equivalence point pH depends on acid-base strength.

Redox and electrochemistry: assign oxidation numbers; balance half-reactions; calculate E°cell; relate to ΔG and spontaneity; applications in batteries and corrosion.

Laboratory measurement and error: accuracy versus precision; significant figures; percent error; calibration; proper glassware selection.

Chemical safety and waste: add acid to water; use fume hood; know SDS; segregate waste; report spills and exposures.

Scientific reasoning and data: distinguish observation from inference; use evidence; identify variables; interpret graphs and tables correctly.

Atomic models and evidence: Dalton, Thomson, Rutherford, Bohr, quantum mechanical; each model supported by experimental evidence such as cathode rays and gold foil.

Periodic table organization: groups and periods; metals, nonmetals, metalloids; valence electrons determine chemical properties; periodic trends predict reactivity.

Chemical bonding: ionic, covalent, metallic; bond polarity; intermolecular forces; molecular geometry and polarity from VSEPR.

Stoichiometry and mole concept: molar mass; Avogadro's number; molar volume; mole ratios; limiting reactant; theoretical and percent yield.

Solutions and concentration: molarity, molality, dilution; colligative properties; solubility rules; precipitation and net ionic equations.

Acids, bases, and equilibrium: pH, pOH, Ka, Kb, Kw; buffers; titration; Le Chatelier's principle; equilibrium constant expression.

Redox and electrochemistry: oxidation numbers; half-reactions; E°cell; spontaneity; electrolysis; applications.

Laboratory safety and data: PPE; add acid to water; waste disposal; significant figures; accuracy and precision; graphing conventions.

Organic chemistry: functional groups; IUPAC naming; isomers; reaction types such as addition, substitution, elimination, oxidation, esterification.

Spectroscopy: IR for functional groups; NMR for hydrogen environments; MS for molecular mass; combine techniques for structure determination.

Thermochemistry: q = mcΔT; ΔH; Hess's law; standard enthalpies of formation; ΔG = ΔH - TΔS; spontaneity.

Kinetics and equilibrium: rate laws; activation energy; catalysts; equilibrium constant; Q versus K; Le Chatelier's principle.

Nuclear chemistry: alpha, beta, gamma decay; half-life; mass defect; fission and fusion; dating; E = mc².

Laboratory techniques: filtration, distillation, chromatography, extraction, recrystallization; choose based on physical properties.

Chemical safety: GHS pictograms; SDS; PPE; fume hood; waste segregation; emergency procedures.

Scientific practices: design investigations; analyze data; construct explanations; communicate evidence-based conclusions.

Atomic structure: protons, neutrons, electrons; isotopes; ions; electron configuration; quantum numbers; periodic trends.

Periodic table: groups, periods, blocks; valence electrons; metals, nonmetals, metalloids; periodic trends in radius, ionization energy, electronegativity.

Chemical bonding: ionic, covalent, metallic; bond polarity; VSEPR geometry; hybridization; intermolecular forces.

Stoichiometry: molar mass; Avogadro's number; molar volume; mole ratios; limiting reactant; theoretical and percent yield.

Solutions: molarity, molality, dilution; colligative properties; solubility rules; precipitation; net ionic equations.

Acids and bases: pH, pOH, Ka, Kb, Kw; buffers; titration; Henderson-Hasselbalch; Le Chatelier's principle.

Redox and electrochemistry: oxidation numbers; half-reactions; E°cell; spontaneity; electrolysis; applications.

Laboratory safety and data: PPE; add acid to water; waste disposal; significant figures; accuracy and precision; graphing.

Organic chemistry: functional groups; IUPAC naming; isomers; reaction types; reaction mechanisms; stereochemistry.

Spectroscopy: IR; NMR; MS; combine evidence for structure determination; Beer-Lambert law for concentration.

Thermochemistry: q = mcΔT; ΔH; Hess's law; standard enthalpies of formation; ΔG = ΔH - TΔS; spontaneity.

Kinetics and equilibrium: rate laws; activation energy; catalysts; equilibrium constant; Q versus K; Le Chatelier's principle.

Nuclear chemistry: alpha, beta, gamma decay; half-life; mass defect; fission and fusion; dating; E = mc².

Laboratory techniques: filtration, distillation, chromatography, extraction, recrystallization; choose based on physical properties.

Chemical safety: GHS pictograms; SDS; PPE; fume hood; waste segregation; emergency procedures.

Scientific practices: design investigations; analyze data; construct explanations; communicate evidence-based conclusions.

Atomic structure and periodic trends: Z, A, isotopes, ions, electron configuration, valence electrons, periodic trends.

Chemical bonding and geometry: ionic, covalent, metallic; polarity; VSEPR; hybridization; intermolecular forces.

Stoichiometry and mole concept: molar mass; Avogadro's number; molar volume; mole ratios; limiting reactant; yield.

Solutions and concentration: molarity, molality, dilution; colligative properties; solubility rules; precipitation.

Acids, bases, and equilibrium: pH, pOH, Ka, Kb, Kw; buffers; titration; Le Chatelier's principle.

Redox and electrochemistry: oxidation numbers; half-reactions; E°cell; spontaneity; electrolysis.

Laboratory safety and data: PPE; add acid to water; waste disposal; significant figures; accuracy and precision.

Scientific practices: design investigations; analyze data; construct explanations; communicate conclusions.

Organic chemistry: functional groups; IUPAC naming; isomers; reaction types; mechanisms; stereochemistry.

Spectroscopy: IR; NMR; MS; Beer-Lambert law; combine evidence for structure determination.

Thermochemistry: q = mcΔT; ΔH; Hess's law; standard enthalpies of formation; ΔG = ΔH - TΔS; spontaneity.

Kinetics and equilibrium: rate laws; activation energy; catalysts; equilibrium constant; Q versus K; Le Chatelier's principle.

Nuclear chemistry: alpha, beta, gamma decay; half-life; mass defect; fission and fusion; dating; E = mc².

Laboratory techniques: filtration, distillation, chromatography, extraction, recrystallization.

Chemical safety: GHS pictograms; SDS; PPE; fume hood; waste segregation; emergency procedures.

Scientific practices: design investigations; analyze data; construct explanations; communicate conclusions.

Atomic structure: protons, neutrons, electrons; isotopes; ions; electron configuration; quantum numbers; periodic trends.

Periodic table: groups, periods, blocks; valence electrons; metals, nonmetals, metalloids; periodic trends.

Chemical bonding: ionic, covalent, metallic; bond polarity; VSEPR geometry; hybridization; intermolecular forces.

Stoichiometry: molar mass; Avogadro's number; molar volume; mole ratios; limiting reactant; yield.

Solutions: molarity, molality, dilution; colligative properties; solubility rules; precipitation.

Acids and bases: pH, pOH, Ka, Kb, Kw; buffers; titration; Henderson-Hasselbalch; Le Chatelier's principle.

Redox and electrochemistry: oxidation numbers; half-reactions; E°cell; spontaneity; electrolysis.

Laboratory safety and data: PPE; add acid to water; waste disposal; significant figures; accuracy and precision.

Organic chemistry: functional groups; IUPAC naming; isomers; reaction types; mechanisms; stereochemistry.

Spectroscopy: IR; NMR; MS; Beer-Lambert law; combine evidence for structure determination.

Thermochemistry: q = mcΔT; ΔH; Hess's law; standard enthalpies of formation; ΔG = ΔH - TΔS; spontaneity.

Kinetics and equilibrium: rate laws; activation energy; catalysts; equilibrium constant; Q versus K; Le Chatelier's principle.

Nuclear chemistry: alpha, beta, gamma decay; half-life; mass defect; fission and fusion; dating; E = mc².

Laboratory techniques: filtration, distillation, chromatography, extraction, recrystallization.

Chemical safety: GHS pictograms; SDS; PPE; fume hood; waste segregation; emergency procedures.

Scientific practices: design investigations; analyze data; construct explanations; communicate conclusions.

Atomic structure: protons, neutrons, electrons; isotopes; ions; electron configuration; quantum numbers; periodic trends.

Periodic table: groups, periods, blocks; valence electrons; metals, nonmetals, metalloids; periodic trends.

Chemical bonding: ionic, covalent, metallic; bond polarity; VSEPR geometry; hybridization; intermolecular forces.

Stoichiometry: molar mass; Avogadro's number; molar volume; mole ratios; limiting reactant; yield.

Solutions: molarity, molality, dilution; colligative properties; solubility rules; precipitation.

Acids and bases: pH, pOH, Ka, Kb, Kw; buffers; titration; Henderson-Hasselbalch; Le Chatelier's principle.

Redox and electrochemistry: oxidation numbers; half-reactions; E°cell; spontaneity; electrolysis.

Laboratory safety and data: PPE; add acid to water; waste disposal; significant figures; accuracy and precision.

Organic chemistry: functional groups; IUPAC naming; isomers; reaction types; mechanisms; stereochemistry.

Spectroscopy: IR; NMR; MS; Beer-Lambert law; combine evidence for structure determination.

Thermochemistry: q = mcΔT; ΔH; Hess's law; standard enthalpies of formation; ΔG = ΔH - TΔS; spontaneity.

Kinetics and equilibrium: rate laws; activation energy; catalysts; equilibrium constant; Q versus K; Le Chatelier's principle.

Nuclear chemistry: alpha, beta, gamma decay; half-life; mass defect; fission and fusion; dating; E = mc².

Laboratory techniques: filtration, distillation, chromatography, extraction, recrystallization.

Chemical safety: GHS pictograms; SDS; PPE; fume hood; waste segregation; emergency procedures.

Scientific practices: design investigations; analyze data; construct explanations; communicate conclusions.

Atomic structure: protons, neutrons, electrons; isotopes; ions; electron configuration; quantum numbers; periodic trends.

Periodic table: groups, periods, blocks; valence electrons; metals, nonmetals, metalloids; periodic trends.

Chemical bonding: ionic, covalent, metallic; bond polarity; VSEPR geometry; hybridization; intermolecular forces.

Stoichiometry: molar mass; Avogadro's number; molar volume; mole ratios; limiting reactant; yield.

Solutions: molarity, molality, dilution; colligative properties; solubility rules; precipitation.

Acids and bases: pH, pOH, Ka, Kb, Kw; buffers; titration; Henderson-Hasselbalch; Le Chatelier's principle.

Redox and electrochemistry: oxidation numbers; half-reactions; E°cell; spontaneity; electrolysis.

Laboratory safety and data: PPE; add acid to water; waste disposal; significant figures; accuracy and precision.

Organic chemistry: functional groups; IUPAC naming; isomers; reaction types; mechanisms; stereochemistry.

Spectroscopy: IR; NMR; MS; Beer-Lambert law; combine evidence for structure determination.

Thermochemistry: q = mcΔT; ΔH; Hess's law; standard enthalpies of formation; ΔG = ΔH - TΔS; spontaneity.

Kinetics and equilibrium: rate laws; activation energy; catalysts; equilibrium constant; Q versus K; Le Chatelier's principle.

Nuclear chemistry: alpha, beta, gamma decay; half-life; mass defect; fission and fusion; dating; E = mc².

Laboratory techniques: filtration, distillation, chromatography, extraction, recrystallization.

Chemical safety: GHS pictograms; SDS; PPE; fume hood; waste segregation; emergency procedures.

Scientific practices: design investigations; analyze data; construct explanations; communicate conclusions.

Atomic structure: protons, neutrons, electrons; isotopes; ions; electron configuration; quantum numbers; periodic trends.

Periodic table: groups, periods, blocks; valence electrons; metals, nonmetals, metalloids; periodic trends.

Chemical bonding: ionic, covalent, metallic; bond polarity; VSEPR geometry; hybridization; intermolecular forces.

Stoichiometry: molar mass; Avogadro's number; molar volume; mole ratios; limiting reactant; yield.

Solutions: molarity, molality, dilution; colligative properties; solubility rules; precipitation.

Acids and bases: pH, pOH, Ka, Kb, Kw; buffers; titration; Henderson-Hasselbalch; Le Chatelier's principle.

Redox and electrochemistry: oxidation numbers; half-reactions; E°cell; spontaneity; electrolysis.

Laboratory safety and data: PPE; add acid to water; waste disposal; significant figures; accuracy and precision.

Organic chemistry: functional groups; IUPAC naming; isomers; reaction types; mechanisms; stereochemistry.

Spectroscopy: IR; NMR; MS; Beer-Lambert law; combine evidence for structure determination.

Thermochemistry: q = mcΔT; ΔH; Hess's law; standard enthalpies of formation; ΔG = ΔH - TΔS; spontaneity.

Kinetics and equilibrium: rate laws; activation energy; catalysts; equilibrium constant; Q versus K; Le Chatelier's principle.

Nuclear chemistry: alpha, beta, gamma decay; half-life; mass defect; fission and fusion; dating; E = mc².

Laboratory techniques: filtration, distillation, chromatography, extraction, recrystallization.

Chemical safety: GHS pictograms; SDS; PPE; fume hood; waste segregation; emergency procedures.

Scientific practices: design investigations; analyze data; construct explanations; communicate conclusions.

Atomic structure: protons, neutrons, electrons; isotopes; ions; electron configuration; quantum numbers; periodic trends.

Periodic table: groups, periods, blocks; valence electrons; metals, nonmetals, metalloids; periodic trends.

Chemical bonding: ionic, covalent, metallic; bond polarity; VSEPR geometry; hybridization; intermolecular forces.

Stoichiometry: molar mass; Avogadro's number; molar volume; mole ratios; limiting reactant; yield.

Solutions: molarity, molality, dilution; colligative properties; solubility rules; precipitation.

Acids and bases: pH, pOH, Ka, Kb, Kw; buffers; titration; Henderson-Hasselbalch; Le Chatelier's principle.

Redox and electrochemistry: oxidation numbers; half-reactions; E°cell; spontaneity; electrolysis.

Laboratory safety and data: PPE; add acid to water; waste disposal; significant figures; accuracy and precision.

Organic chemistry: functional groups; IUPAC naming; isomers; reaction types; mechanisms; stereochemistry.

Spectroscopy: IR; NMR; MS; Beer-Lambert law; combine evidence for structure determination.

Thermochemistry: q = mcΔT; ΔH; Hess's law; standard enthalpies of formation; ΔG = ΔH - TΔS; spontaneity.

Kinetics and equilibrium: rate laws; activation energy; catalysts; equilibrium constant; Q versus K; Le Chatelier's principle.

Nuclear chemistry: alpha, beta, gamma decay; half-life; mass defect; fission and fusion; dating; E = mc².

Laboratory techniques: filtration, distillation, chromatography, extraction, recrystallization.

Chemical safety: GHS pictograms; SDS; PPE; fume hood; waste segregation; emergency procedures.

Scientific practices: design investigations; analyze data; construct explanations; communicate conclusions.

Atomic structure: protons, neutrons, electrons; isotopes; ions; electron configuration; quantum numbers; periodic trends.

Periodic table: groups, periods, blocks; valence electrons; metals, nonmetals, metalloids; periodic trends.

Chemical bonding: ionic, covalent, metallic; bond polarity; VSEPR geometry; hybridization; intermolecular forces.

Stoichiometry: molar mass; Avogadro's number; molar volume; mole ratios; limiting reactant; yield.

Solutions: molarity, molality, dilution; colligative properties; solubility rules; precipitation.

Acids and bases: pH, pOH, Ka, Kb, Kw; buffers; titration; Henderson-Hasselbalch; Le Chatelier's principle.

Redox and electrochemistry: oxidation numbers; half-reactions; E°cell; spontaneity; electrolysis.

Laboratory safety and data: PPE; add acid to water; waste disposal; significant figures; accuracy and precision.

Organic chemistry: functional groups; IUPAC naming; isomers; reaction types; mechanisms; stereochemistry.

Spectroscopy: IR; NMR; MS; Beer-Lambert law; combine evidence for structure determination.

Thermochemistry: q = mcΔT; ΔH; Hess's law; standard enthalpies of formation; ΔG = ΔH - TΔS; spontaneity.

Kinetics and equilibrium: rate laws; activation energy; catalysts; equilibrium constant; Q versus K; Le Chatelier's principle.

Nuclear chemistry: alpha, beta, gamma decay; half-life; mass defect; fission and fusion; dating; E = mc².

Laboratory techniques: filtration, distillation, chromatography, extraction, recrystallization.

Chemical safety: GHS pictograms; SDS; PPE; fume hood; waste segregation; emergency procedures.

Scientific practices: design investigations; analyze data; construct explanations; communicate conclusions.

Atomic structure: protons, neutrons, electrons; isotopes; ions; electron configuration; quantum numbers; periodic trends.

Periodic table: groups, periods, blocks; valence electrons; metals, nonmetals, metalloids; periodic trends.

Chemical bonding: ionic, covalent, metallic; bond polarity; VSEPR geometry; hybridization; intermolecular forces.

Stoichiometry: molar mass; Avogadro's number; molar volume; mole ratios; limiting reactant; yield.

Solutions: molarity, molality, dilution; colligative properties; solubility rules; precipitation.

Acids and bases: pH, pOH, Ka, Kb, Kw; buffers; titration; Henderson-Hasselbalch; Le Chatelier's principle.

Redox and electrochemistry: oxidation numbers; half-reactions; E°cell; spontaneity; electrolysis.

Laboratory safety and data: PPE; add acid to water; waste disposal; significant figures; accuracy and precision.

Organic chemistry: functional groups; IUPAC naming; isomers; reaction types; mechanisms; stereochemistry.

Spectroscopy: IR; NMR; MS; Beer-Lambert law; combine evidence for structure determination.

Thermochemistry: q = mcΔT; ΔH; Hess's law; standard enthalpies of formation; ΔG = ΔH - TΔS; spontaneity.

Kinetics and equilibrium: rate laws; activation energy; catalysts; equilibrium constant; Q versus K; Le Chatelier's principle.

Nuclear chemistry: alpha, beta, gamma decay; half-life; mass defect; fission and fusion; dating; E = mc².

Laboratory techniques: filtration, distillation, chromatography, extraction, recrystallization.

Chemical safety: GHS pictograms; SDS; PPE; fume hood; waste segregation; emergency procedures.

Scientific practices: design investigations; analyze data; construct explanations; communicate conclusions.