FTCE Chemistry 6-12 cheat sheet

Significant figures: multiplication/division keeps the least number of sig figs in any factor; addition/subtraction keeps the least number of decimal places.

Accuracy = closeness to true value; precision = reproducibility among replicates. Random error affects precision; systematic error shifts accuracy.

Dimensional analysis: multiply by conversion factors equal to 1; track units so they cancel to the desired unit.

Percent error = |experimental − accepted| / accepted × 100%. Percent yield = actual / theoretical × 100%.

Density = mass/volume; a common derived measurement used to identify substances and check purity.

A controlled experiment varies one independent variable, holds others constant, and measures a dependent variable; controls establish baseline.

Linear calibration: y = mx + b; slope gives sensitivity, intercept shows systematic offset; R² near 1 indicates good linearity.

Safety: read SDS before use, add acid to water (never water to acid), use fume hood for volatiles, know eyewash and shower locations.

Scientific method cycle: question → hypothesis → experiment → data analysis → conclusion → revision; conclusions must be falsifiable and repeatable.

Classification: pure substances (elements, compounds) vs mixtures (homogeneous, heterogeneous); mixtures separate by physical means.

Physical change alters state or form without changing composition; chemical change produces new substances (gas, precipitate, color, heat, light).

States: solid (fixed shape/volume), liquid (fixed volume, variable shape), gas (variable both); plasma is ionized gas.

Phase changes: melting, freezing, vaporization, condensation, sublimation, deposition; each is accompanied by energy absorption or release.

Dalton: atoms indivisible; Thomson: electron (plum pudding); Rutherford: dense positive nucleus; Bohr: quantized energy levels; quantum model: orbitals.

Atomic number Z = protons; mass number A = protons + neutrons; isotopes differ in neutrons; ions differ in electrons.

Periodic trends: atomic radius decreases left→right, increases top→bottom; ionization energy and electronegativity increase left→right and up a group.

Groups: alkali metals (1), alkaline earth (2), halogens (17), noble gases (18), transition metals (3–12).

Mole concept: 1 mol = 6.022×10²³ particles = molar mass in grams; molarity M = mol solute / L solution.

First law of thermodynamics: ΔE = q + w; energy is conserved, only transformed. Second law: entropy of the universe increases for spontaneous processes.

Enthalpy: ΔH = q_p at constant pressure; exothermic ΔH < 0, endothermic ΔH > 0.

Hess's law: ΔH for a reaction is the sum of ΔH for any sequence of steps; reverse reaction changes sign of ΔH.

Calorimetry: q = mcΔT; for a bomb calorimeter, q = CΔT (C = heat capacity).

Bond energy: ΔH_rxn = Σ(bonds broken) − Σ(bonds formed).

Gibbs free energy: ΔG = ΔH − TΔS; spontaneous when ΔG < 0, nonspontaneous when ΔG > 0, equilibrium when ΔG = 0.

Kinetics: rate depends on concentration, temperature, catalyst, surface area; Arrhenius k = A·e^(−Ea/RT).

Catalysts lower activation energy without changing ΔH or equilibrium position; they speed both forward and reverse rates.

Le Chatelier: adding reactant shifts right; removing product shifts right; increasing pressure shifts toward fewer gas moles; increasing temperature shifts endothermic direction.

Equilibrium constant K: for aA + bB ⇌ cC + dD, K = [C]^c[D]^d / [A]^a[B]^b; K > 1 favors products.

Ionic bonds: metal + nonmetal, electron transfer, high melting point, conduct when molten or dissolved. Covalent bonds: nonmetal + nonmetal, electron sharing.

Electronegativity difference: >1.7 mostly ionic, 0.4–1.7 polar covalent, <0.4 nonpolar covalent (guideline).

Lewis structures: count valence electrons, minimize formal charges, satisfy octets (H = 2); resonance shows delocalized electrons.

VSEPR: electron pairs repel; lone pairs occupy more space than bonding pairs, reducing bond angles.

Common geometries: linear (2), trigonal planar (3), tetrahedral (4), trigonal bipyramidal (5), octahedral (6); lone pairs give bent, trigonal pyramidal, seesaw.

Hybridization: sp (linear, 180°), sp² (trigonal planar, 120°), sp³ (tetrahedral, 109.5°).

Intermolecular forces strength: ion–dipole > hydrogen bond > dipole–dipole > London dispersion; stronger IMF → higher boiling point.

Hydrogen bonding requires H bonded to N, O, or F; explains water's high boiling point and surface tension.

Polarity: a molecule is polar if bond dipoles do not cancel by symmetry; CO₂ nonpolar, H₂O polar.

Metallic bonding: sea of delocalized electrons; explains conductivity, malleability, and luster.

Reaction types: synthesis (A + B → AB), decomposition, single replacement, double replacement, combustion (hydrocarbon + O₂ → CO₂ + H₂O).

Balanced equations conserve atoms and charge; coefficients give mole ratios used in stoichiometry.

Stoichiometry path: grams A → mol A → mol B (via coefficients) → grams B; use molar mass and mole ratio.

Limiting reactant: the one that produces the least product; excess reactant remains after reaction.

Percent yield = (actual mass / theoretical mass) × 100%; theoretical mass comes from the limiting reactant.

Molarity and dilution: M₁V₁ = M₂V₂; mol = M × V; titration uses n_acid × V_acid × M_acid = n_base × V_base × M_base.

Solubility rules: most nitrates, acetates, alkali salts soluble; AgCl, PbCl₂, BaSO₄, CaCO₃ insoluble (common exceptions).

Net ionic equations: cancel spectator ions; insoluble product forms a precipitate.

Acid–base: Arrhenius (H⁺/OH⁻), Brønsted–Lowry (proton donor/acceptor); pH = −log[H⁺]; pH + pOH = 14 at 25 °C.

Redox: oxidation loses electrons (LEO), reduction gains electrons (GER); balance by half-reactions and equalize electron transfer.

Significant figures: multiplication/division keeps the least number of sig figs in any factor; addition/subtraction keeps the least number of decimal places.

Accuracy = closeness to true value; precision = reproducibility among replicates. Random error affects precision; systematic error shifts accuracy.

Dimensional analysis: multiply by conversion factors equal to 1; track units so they cancel to the desired unit.

Percent error = |experimental − accepted| / accepted × 100%. Percent yield = actual / theoretical × 100%.

Density = mass/volume; a common derived measurement used to identify substances and check purity.

A controlled experiment varies one independent variable, holds others constant, and measures a dependent variable; controls establish baseline.

Linear calibration: y = mx + b; slope gives sensitivity, intercept shows systematic offset; R² near 1 indicates good linearity.

Safety: read SDS before use, add acid to water (never water to acid), use fume hood for volatiles, know eyewash and shower locations.

Scientific method cycle: question → hypothesis → experiment → data analysis → conclusion → revision; conclusions must be falsifiable and repeatable.

Classification: pure substances (elements, compounds) vs mixtures (homogeneous, heterogeneous); mixtures separate by physical means.

Physical change alters state or form without changing composition; chemical change produces new substances (gas, precipitate, color, heat, light).

States: solid (fixed shape/volume), liquid (fixed volume, variable shape), gas (variable both); plasma is ionized gas.

Phase changes: melting, freezing, vaporization, condensation, sublimation, deposition; each is accompanied by energy absorption or release.

Dalton: atoms indivisible; Thomson: electron (plum pudding); Rutherford: dense positive nucleus; Bohr: quantized energy levels; quantum model: orbitals.

Atomic number Z = protons; mass number A = protons + neutrons; isotopes differ in neutrons; ions differ in electrons.

Periodic trends: atomic radius decreases left→right, increases top→bottom; ionization energy and electronegativity increase left→right and up a group.

Groups: alkali metals (1), alkaline earth (2), halogens (17), noble gases (18), transition metals (3–12).

Mole concept: 1 mol = 6.022×10²³ particles = molar mass in grams; molarity M = mol solute / L solution.

First law of thermodynamics: ΔE = q + w; energy is conserved, only transformed. Second law: entropy of the universe increases for spontaneous processes.

Enthalpy: ΔH = q_p at constant pressure; exothermic ΔH < 0, endothermic ΔH > 0.

Hess's law: ΔH for a reaction is the sum of ΔH for any sequence of steps; reverse reaction changes sign of ΔH.

Calorimetry: q = mcΔT; for a bomb calorimeter, q = CΔT (C = heat capacity).

Bond energy: ΔH_rxn = Σ(bonds broken) − Σ(bonds formed).

Gibbs free energy: ΔG = ΔH − TΔS; spontaneous when ΔG < 0, nonspontaneous when ΔG > 0, equilibrium when ΔG = 0.

Kinetics: rate depends on concentration, temperature, catalyst, surface area; Arrhenius k = A·e^(−Ea/RT).

Catalysts lower activation energy without changing ΔH or equilibrium position; they speed both forward and reverse rates.

Le Chatelier: adding reactant shifts right; removing product shifts right; increasing pressure shifts toward fewer gas moles; increasing temperature shifts endothermic direction.

Equilibrium constant K: for aA + bB ⇌ cC + dD, K = [C]^c[D]^d / [A]^a[B]^b; K > 1 favors products.

Ionic bonds: metal + nonmetal, electron transfer, high melting point, conduct when molten or dissolved. Covalent bonds: nonmetal + nonmetal, electron sharing.

Electronegativity difference: >1.7 mostly ionic, 0.4–1.7 polar covalent, <0.4 nonpolar covalent (guideline).

Lewis structures: count valence electrons, minimize formal charges, satisfy octets (H = 2); resonance shows delocalized electrons.

VSEPR: electron pairs repel; lone pairs occupy more space than bonding pairs, reducing bond angles.

Common geometries: linear (2), trigonal planar (3), tetrahedral (4), trigonal bipyramidal (5), octahedral (6); lone pairs give bent, trigonal pyramidal, seesaw.

Hybridization: sp (linear, 180°), sp² (trigonal planar, 120°), sp³ (tetrahedral, 109.5°).

Intermolecular forces strength: ion–dipole > hydrogen bond > dipole–dipole > London dispersion; stronger IMF → higher boiling point.

Hydrogen bonding requires H bonded to N, O, or F; explains water's high boiling point and surface tension.

Polarity: a molecule is polar if bond dipoles do not cancel by symmetry; CO₂ nonpolar, H₂O polar.

Metallic bonding: sea of delocalized electrons; explains conductivity, malleability, and luster.

Reaction types: synthesis (A + B → AB), decomposition, single replacement, double replacement, combustion (hydrocarbon + O₂ → CO₂ + H₂O).

Balanced equations conserve atoms and charge; coefficients give mole ratios used in stoichiometry.

Stoichiometry path: grams A → mol A → mol B (via coefficients) → grams B; use molar mass and mole ratio.

Limiting reactant: the one that produces the least product; excess reactant remains after reaction.

Percent yield = (actual mass / theoretical mass) × 100%; theoretical mass comes from the limiting reactant.

Molarity and dilution: M₁V₁ = M₂V₂; mol = M × V; titration uses n_acid × V_acid × M_acid = n_base × V_base × M_base.

Solubility rules: most nitrates, acetates, alkali salts soluble; AgCl, PbCl₂, BaSO₄, CaCO₃ insoluble (common exceptions).

Net ionic equations: cancel spectator ions; insoluble product forms a precipitate.

Acid–base: Arrhenius (H⁺/OH⁻), Brønsted–Lowry (proton donor/acceptor); pH = −log[H⁺]; pH + pOH = 14 at 25 °C.

Redox: oxidation loses electrons (LEO), reduction gains electrons (GER); balance by half-reactions and equalize electron transfer.