Chemistry
Master atomic structure, chemical bonding, reactions, stoichiometry, and thermodynamics. This guide covers all content areas of the CLEP Chemistry exam, equivalent to a full-year introductory college chemistry course.
Atomic Structure & Periodicity
~20%Atomic Theory and Subatomic Particles
Atoms consist of protons (positive charge, in nucleus), neutrons (neutral, in nucleus), and electrons (negative charge, surrounding nucleus). Atomic number (Z) = number of protons; defines the element. Mass number (A) = protons + neutrons. Isotopes are atoms of the same element with different numbers of neutrons; they have the same chemical properties but different masses. Atomic mass on the periodic table is a weighted average of isotope masses.
Electron Configuration
Electrons occupy orbitals in energy levels. Orbital types: s (2 e⁻), p (6 e⁻), d (10 e⁻), f (14 e⁻). Fill according to the Aufbau principle (lowest energy first), Hund's rule (maximize unpaired electrons within a subshell), and Pauli Exclusion Principle (no two electrons have the same four quantum numbers; each orbital holds max 2 electrons with opposite spins). Order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p… The four quantum numbers: n (principal), ℓ (angular momentum, 0 to n−1), mℓ (magnetic, −ℓ to +ℓ), ms (spin, ±½).
Periodic Trends
Atomic radius decreases across a period (more protons pulling electrons in) and increases down a group (more electron shells). Ionization energy: the energy to remove an electron — increases across a period, decreases down a group. Electron affinity: energy released when an electron is added — generally increases across a period. Electronegativity (Pauling scale): tendency to attract electrons in a bond — increases up and to the right (highest: F = 4.0). Metallic character decreases left to right, increases down groups.
The Periodic Table
Groups (columns) have similar chemical properties due to identical valence electron configurations. Periods (rows) correspond to principal quantum number. Key regions: alkali metals (Group 1, very reactive), alkaline earth metals (Group 2), transition metals (Groups 3–12), halogens (Group 17), noble gases (Group 18, full valence shell, largely inert). Metalloids (B, Si, Ge, As, Sb, Te) have intermediate properties.
Electromagnetic Radiation and Spectra
Electrons can absorb energy and jump to higher energy levels (excited states) then emit photons when returning to ground state. Energy of photon: E = hν = hc/λ, where h = Planck's constant = 6.626 × 10⁻³⁴ J·s, ν is frequency, c is speed of light, λ is wavelength. Atomic emission spectra are unique fingerprints of each element. The Bohr model describes the hydrogen atom's quantized energy levels; the quantum mechanical model extends this to multi-electron atoms using wave functions (orbitals).
Chemical Bonding & Molecular Structure
~20%Types of Chemical Bonds
Ionic bonds form between metals and nonmetals: the metal loses electrons (becomes cation) and the nonmetal gains them (becomes anion); electrostatic attraction holds the ions together. Covalent bonds form between nonmetals sharing electrons; polar covalent if electronegativity difference is ~0.4–1.7, nonpolar if near 0. Metallic bonds involve a "sea of electrons" delocalized among metal cations, explaining conductivity, malleability, and luster.
Lewis Structures
Lewis structures show valence electrons as dots and bonds as lines. Steps: (1) count total valence electrons, (2) connect atoms with single bonds, (3) complete octets on outer atoms using lone pairs, (4) place remaining electrons on central atom, (5) form multiple bonds if central atom lacks an octet. Exceptions to octet rule: expanded octets for period 3+ elements (PCl₅, SF₆), incomplete octets for boron compounds (BF₃), and radicals with odd electron counts.
VSEPR Theory
Valence Shell Electron Pair Repulsion (VSEPR): electron groups around the central atom orient to maximize distance. Electron group geometries and molecular geometries: 2 groups = linear (180°); 3 = trigonal planar (120°) or bent if one lone pair; 4 = tetrahedral (109.5°), trigonal pyramidal (3 bonds + 1 lone pair, ~107°), or bent (2 bonds + 2 lone pairs, ~104.5°); 5 = trigonal bipyramidal; 6 = octahedral. Lone pairs compress bond angles more than bonding pairs.
Polarity and Intermolecular Forces
A molecule is polar if it has polar bonds AND an asymmetric geometry (dipole moments don't cancel). Intermolecular forces (weakest to strongest): London dispersion forces (instantaneous dipole, present in all molecules, strength increases with size/polarizability); dipole-dipole interactions (polar molecules); hydrogen bonding (N-H, O-H, or F-H···N/O/F — strongest IMF for small molecules). These forces determine boiling point, surface tension, viscosity, and solubility.
Hybridization and Molecular Orbital Theory
Hybridization explains the geometry of bonded atoms. sp³ (tetrahedral), sp² (trigonal planar, includes one unhybridized p → π bond possible), sp (linear, two unhybridized p orbitals). In MO theory, atomic orbitals combine to form bonding and antibonding molecular orbitals. Bond order = (bonding − antibonding electrons)/2. Benzene and other aromatic compounds have delocalized π electrons via resonance.
Chemical Reactions & Types
~15%Types of Chemical Reactions
Synthesis (combination): A + B → AB. Decomposition: AB → A + B. Single displacement: A + BC → AC + B (A displaces B if A is more reactive). Double displacement (metathesis): AB + CD → AD + CB — drives to completion by forming precipitate, water, or gas. Combustion: fuel + O₂ → CO₂ + H₂O. Acid-base neutralization: acid + base → salt + water. Redox: electrons transfer between species.
Balancing Chemical Equations
Conservation of mass requires equal numbers of each atom on both sides. Balance by inspection or by the half-reaction method for redox. Net ionic equations show only the species that change; spectator ions are excluded. A balanced equation must have equal atoms AND equal charges for ionic equations. Combustion of hydrocarbons: CₓHᵧ + O₂ → CO₂ + H₂O — balance C first, then H, then O.
Acid-Base Chemistry
Arrhenius: acid produces H⁺, base produces OH⁻ in water. Brønsted-Lowry: acid is a proton donor, base is a proton acceptor. Lewis: acid accepts electron pair, base donates electron pair. Strong acids (HCl, HBr, HI, HNO₃, H₂SO₄, HClO₄) and strong bases (Group 1 hydroxides, Ca(OH)₂, Ba(OH)₂) fully dissociate. Weak acids/bases partially dissociate — governed by Ka and Kb. pH = −log[H⁺]; pOH = −log[OH⁻]; pH + pOH = 14 at 25°C. Buffers resist pH changes (weak acid + its conjugate base).
Oxidation-Reduction (Redox) Reactions
In redox reactions, oxidation numbers change. Oxidation: increase in oxidation number (loss of electrons — LEO). Reduction: decrease in oxidation number (gain of electrons — GER). LEO the lion says GER. Rules for assigning oxidation numbers: pure element = 0; monatomic ion = charge; O = −2 (except in peroxides); H = +1 (except metal hydrides); sum = overall charge. Oxidizing agent is reduced; reducing agent is oxidized.
Precipitation and Solubility Rules
Key solubility rules: All Group 1 salts and NH₄⁺ salts are soluble. All nitrates (NO₃⁻), acetates, perchlorates are soluble. Chlorides, bromides, iodides are soluble except with Ag⁺, Pb²⁺, Hg₂²⁺. Sulfates are soluble except with Ba²⁺, Sr²⁺, Pb²⁺, Ca²⁺. Carbonates, phosphates, sulfides, hydroxides are generally insoluble except with Group 1 and NH₄⁺. The solubility product Ksp = [ions] product at equilibrium with undissolved solid.
Stoichiometry, Solutions & Gases
~20%The Mole Concept
One mole = 6.022 × 10²³ particles (Avogadro's number). Molar mass (g/mol) numerically equals atomic/molecular mass in amu. Key conversions: moles = mass / molar mass; particles = moles × 6.022 × 10²³; for gases at STP (0°C, 1 atm): 1 mole = 22.4 L. Empirical formula gives the simplest whole-number ratio of atoms; molecular formula is a whole-number multiple of the empirical formula.
Stoichiometric Calculations
Mole ratio from balanced equation. Limiting reagent: the reactant that is completely consumed first — determines theoretical yield. Percent yield = (actual yield / theoretical yield) × 100%. Strategy for stoichiometry: convert grams → moles (divide by molar mass) → use mole ratio → convert to desired units. For limiting reagent: calculate moles of product from each reactant; the smaller amount indicates the limiting reagent.
Solution Chemistry
Concentration = molarity (M) = moles of solute / liters of solution. Dilution: M₁V₁ = M₂V₂. Electrolytes dissociate in water (produce ions); strong electrolytes fully dissociate. Colligative properties depend on the number of particles: vapor pressure lowering (Raoult's Law: P_solution = χ_solvent · P°_solvent), boiling point elevation (ΔTb = Kb · m · i), freezing point depression (ΔTf = −Kf · m · i), and osmotic pressure (π = MRT). Van't Hoff factor i accounts for dissociation.
Gas Laws
Boyle's Law: P₁V₁ = P₂V₂ (constant T, n). Charles's Law: V₁/T₁ = V₂/T₂ (constant P, n). Gay-Lussac's Law: P₁/T₁ = P₂/T₂ (constant V, n). Combined Gas Law: P₁V₁/T₁ = P₂V₂/T₂. Ideal Gas Law: PV = nRT, where R = 0.08206 L·atm/(mol·K). Always use Kelvin for gas law calculations. Dalton's Law of Partial Pressures: P_total = Σ P_i. Graham's Law of effusion: rate₁/rate₂ = √(M₂/M₁) — lighter gases move faster.
Thermochemistry Basics
Enthalpy (H) is the heat content of a system. ΔH < 0 = exothermic (releases heat); ΔH > 0 = endothermic (absorbs heat). Hess's Law: ΔH for a reaction is the sum of ΔH values for any series of steps that add up to the overall reaction. Standard enthalpy of formation ΔH°f is the enthalpy change for forming 1 mole of compound from elements in their standard states. ΔH°rxn = Σ ΔH°f(products) − Σ ΔH°f(reactants). Heat capacity: q = mcΔT (m = mass, c = specific heat, ΔT = temperature change).
Kinetics, Equilibrium & Thermodynamics
~15%Reaction Kinetics
Reaction rate = change in concentration / time. Rate law: rate = k[A]ᵐ[B]ⁿ, where m and n are determined experimentally (not from coefficients). The overall reaction order is m + n. First-order reactions: [A] = [A]₀e^(−kt); half-life t₁/₂ = 0.693/k (constant). Second-order: 1/[A] = 1/[A]₀ + kt; half-life = 1/(k[A]₀). The Arrhenius equation: k = Ae^(−Ea/RT), where Ea is activation energy. Catalysts lower Ea without being consumed — they speed up the reaction without changing equilibrium.
Chemical Equilibrium
At equilibrium, forward and reverse reaction rates are equal; concentrations are constant (not equal). The equilibrium constant Keq = [products]^coefficients / [reactants]^coefficients (pure solids/liquids omitted). Kc uses concentration; Kp uses partial pressures (Kp = Kc(RT)^Δn). K > 1: products favored; K < 1: reactants favored; K = 1: roughly equal. Reaction quotient Q: if Q < K, reaction proceeds forward; Q > K, proceeds reverse; Q = K, at equilibrium.
Le Chatelier's Principle
If a system at equilibrium is disturbed, it shifts to partially counteract the disturbance. Adding a reactant → shifts right. Removing a product → shifts right. Increasing pressure (decreasing volume) → shifts toward the side with fewer moles of gas. Increasing temperature for an endothermic reaction → shifts right (increasing K). Adding a catalyst → equilibrium is reached faster but position doesn't change. Inert gas at constant volume: no shift.
Thermodynamics: Entropy and Free Energy
Entropy (S) is a measure of disorder/randomness. Processes increase entropy: dissolution, gas expansion, more particles formed, higher temperature. The Second Law of Thermodynamics: the entropy of the universe always increases for spontaneous processes. Gibbs Free Energy: ΔG = ΔH − TΔS. Spontaneous if ΔG < 0; nonspontaneous if ΔG > 0; at equilibrium if ΔG = 0. ΔG° = −RT ln K. At the temperature where ΔG = 0: T = ΔH/ΔS — this is the crossover temperature for spontaneity.
Acid-Base Equilibria
Ka is the acid dissociation constant: HA ⇌ H⁺ + A⁻; Ka = [H⁺][A⁻]/[HA]. pKa = −log Ka; weaker acid = larger pKa. Kb for bases; pKa + pKb = 14 for conjugate acid-base pairs. Henderson-Hasselbalch equation for buffers: pH = pKa + log([A⁻]/[HA]). Buffer is most effective when pH ≈ pKa. Titration: at the equivalence point, moles of acid = moles of base; half-equivalence point pH = pKa for weak acid titrations.
Electrochemistry, Nuclear & Organic Overview
~10%Electrochemical Cells
Galvanic (voltaic) cells convert chemical energy to electrical energy (spontaneous redox). Electrolytic cells use electrical energy to drive nonspontaneous redox. Cell potential: E°cell = E°cathode − E°anode (use standard reduction potentials). Positive E°cell → spontaneous → ΔG < 0. ΔG° = −nFE°cell, where n = moles of electrons, F = Faraday's constant = 96,485 C/mol. Nernst equation: E = E° − (RT/nF)ln Q, at 25°C: E = E° − (0.0592/n)log Q.
Activity Series and Electrolysis
The activity series ranks metals by their tendency to be oxidized (lose electrons). Metals higher on the list displace those lower in single displacement reactions. In electrolytic cells (non-spontaneous), an external power source drives the reaction. Faraday's laws of electrolysis: moles of substance deposited = (current × time) / (n × F). Practical application: electroplating, refining metals, producing Al from Al₂O₃.
Nuclear Chemistry
Nuclear reactions involve changes to the nucleus. Types of radioactive decay: alpha (α) — 4He nucleus emitted, Z decreases by 2; beta (β⁻) — neutron → proton + electron emitted, Z increases by 1; positron emission (β⁺) — proton → neutron + positron, Z decreases by 1; gamma (γ) — high-energy photon, no change in Z or A; electron capture — inner electron absorbed, Z decreases by 1. Half-life: t₁/₂ = 0.693/λ. Nuclear binding energy: mass defect × c² (Einstein's E = mc²). Fission: heavy nuclei split; fusion: light nuclei combine — both release energy.
Organic Chemistry Overview
Carbon forms 4 bonds; the backbone of organic molecules. Alkanes (CₙH₂ₙ₊₂, single bonds), alkenes (C=C double bond), alkynes (C≡C triple bond), aromatics (benzene ring). Functional groups determine reactivity: hydroxyl (−OH, alcohols), carbonyl (C=O), carboxyl (−COOH, acids), amine (−NH₂), ester, ether, halide. IUPAC naming: find the longest carbon chain (parent), number to give substituents lowest numbers, add prefixes for substituents. Isomers: structural (same formula, different connectivity) vs. stereoisomers (same connectivity, different spatial arrangement).
Laboratory Techniques and Measurements
Significant figures: the number of meaningful digits in a measurement. In addition/subtraction, round to the least number of decimal places; in multiplication/division, round to the fewest significant figures. Scientific notation: a × 10ⁿ. Percent error = |experimental − theoretical| / theoretical × 100%. Common lab techniques: titration (determine concentration), chromatography (separate mixtures), spectroscopy (identify substances by light absorption). SI base units: meter (length), kilogram (mass), kelvin (temperature), mole (amount), second (time).
Key Figures in the History of Chemistry
| Chemist / Scientist | Era / Nationality | Contribution to Chemistry |
|---|---|---|
| Antoine Lavoisier | 1743–1794, French | "Father of modern chemistry"; Law of Conservation of Mass; named oxygen and hydrogen; first modern list of elements |
| John Dalton | 1766–1844, English | Atomic theory: all matter consists of atoms; atoms of the same element are identical; introduced atomic weights |
| Amedeo Avogadro | 1776–1856, Italian | Avogadro's hypothesis: equal volumes of gas at same T and P contain equal numbers of molecules; Avogadro's number |
| Dmitri Mendeleev | 1834–1907, Russian | Created the periodic table organized by atomic mass; predicted undiscovered elements with remarkable accuracy |
| Lothar Meyer | 1830–1895, German | Independently developed periodic table; showed periodic variation in atomic volume |
| J.J. Thomson | 1856–1940, English | Discovered the electron via cathode ray experiments; proposed the plum pudding model of the atom |
| Ernest Rutherford | 1871–1937, New Zealander | Gold foil experiment proving the nuclear model; discovered the nucleus and proton; named alpha and beta radiation |
| Niels Bohr | 1885–1962, Danish | Bohr model of the hydrogen atom with quantized electron orbits; explained atomic emission spectra |
| Erwin Schrödinger | 1887–1961, Austrian | Wave equation for electrons; quantum mechanical model replacing Bohr model with probability-based orbitals |
| Werner Heisenberg | 1901–1976, German | Uncertainty principle: cannot simultaneously know exact position and momentum of an electron |
| Linus Pauling | 1901–1994, American | Nature of the chemical bond; electronegativity scale; resonance theory; alpha-helix protein structure |
| Gilbert Lewis | 1875–1946, American | Lewis dot structures; covalent bond theory; Lewis acid-base definition; cubical atom model |
| Svante Arrhenius | 1859–1927, Swedish | Arrhenius acid-base theory; Arrhenius equation for reaction rates; electrolyte theory |
| Johannes Brønsted | 1879–1947, Danish | Co-developed Brønsted-Lowry acid-base theory (acids as proton donors, bases as proton acceptors) |
| Thomas Lowry | 1874–1936, English | Co-developed Brønsted-Lowry acid-base theory independently and simultaneously with Brønsted |
| Robert Boyle | 1627–1691, Irish | Boyle's Law relating gas pressure and volume; early champion of the experimental method in chemistry |
| Marie Curie | 1867–1934, Polish-French | Discovered polonium and radium; coined "radioactivity"; first woman to win a Nobel Prize (won two) |
| Henry Moseley | 1887–1915, English | Used X-ray spectra to determine atomic number; reorganized periodic table by atomic number instead of mass |
| Jacobus van't Hoff | 1852–1911, Dutch | First Nobel Prize in Chemistry; stereochemistry of carbon; van't Hoff factor for colligative properties |
| Friedrich Wöhler | 1800–1882, German | Synthesized urea from inorganic materials (1828), disproving vitalism; bridged organic and inorganic chemistry |
Key Terms & Definitions
Video Resources
Practice Questions (150)
A) Cl neutral atom
B) Cl⁻ ion
C) Ar neutral atom
D) S²⁻ ion
A) Ca
B) Zn
C) Cu
D) Ni
A) Na
B) Mg
C) Al
D) Si
A) Linear
B) Trigonal planar
C) Bent
D) Tetrahedral
A) 1, 1, 1
B) 2, 3, 1
C) 4, 3, 2
D) 2, 1, 1
A) 0.5 mol
B) 1.0 mol
C) 2.0 mol
D) 44 mol
A) 3 L
B) 6 L
C) 12 L
D) 1 L
A) 4
B) −4
C) 10
D) 0.0001
A) Synthesis
B) Decomposition
C) Double displacement
D) Single displacement
A) H₂
B) O₂
C) H₂O
D) Neither; they're equal
A) London dispersion only
B) Dipole-dipole only
C) Hydrogen bonding
D) Ionic bonding
A) Rate doubles
B) Rate quadruples
C) Rate is halved
D) Rate stays the same
A) Shift equilibrium left
B) Shift equilibrium right
C) Have no effect
D) Increase K
A) H₂O(g) → H₂O(l)
B) NaCl(s) → Na⁺(aq) + Cl⁻(aq)
C) 4NH₃(g) → 2N₂(g) + 6H₂(g)
D) Both B and C
A) +3
B) +4
C) +6
D) +7
A) The reaction is nonspontaneous
B) Cu is oxidized
C) Zn is oxidized at the anode
D) ΔG° = +nFE°
A) sp³
B) sp²
C) sp
D) p
A) 0.1 M
B) 0.2 M
C) 0.4 M
D) 8.0 M
A) Alpha decay
B) Beta (β⁻) decay
C) Gamma emission
D) Electron capture
A) Nonspontaneous and endothermic
B) Spontaneous
C) At equilibrium
D) Endothermic
A) They increase activation energy
B) They change the equilibrium constant
C) They are consumed in the reaction
D) They lower activation energy and are regenerated
A) CHO
B) CH₂O
C) C₂H₄O₂
D) C₃H₆O₃
A) 24.6 L
B) 44.8 L
C) 49.3 L
D) 22.4 L
A) Calculate cell potential
B) Calculate activation energy
C) Find pH of a buffer solution
D) Determine molecular geometry
A) AgCl is a strong acid
B) Equilibrium between solid AgCl and its dissolved ions
C) The rate of AgCl dissolving
D) AgCl has no solubility
A) A covalent bond within water molecules
B) An intermolecular force between H and a highly electronegative atom (N, O, or F)
C) An ionic interaction between metal and nonmetal
D) Electron sharing between two hydrogen atoms
A) 120°
B) 109.5°
C) ~107°
D) 180°
A) +130 kJ
B) −30 kJ
C) +30 kJ
D) −130 kJ
A) 1/2
B) 1/4
C) 1/6
D) 1/8
A) NH₃
B) OH⁻
C) BF₃
D) H₂O
A) 80%
B) 75%
C) 90%
D) 125%
A) Oxidation occurs
B) Electrons are produced
C) Reduction occurs
D) The salt bridge connects
A) CCl₄
B) CO₂
C) BF₃
D) HCl
A) Reaction shifts forward (to products)
B) Reaction shifts backward (to reactants)
C) Reaction is at equilibrium
D) Q cannot be greater than K
A) Vapor pressure lowering
B) Boiling point elevation
C) Freezing point depression
D) Osmotic pressure
A) H₃CO₃⁺
B) HCO₃⁻
C) CO₃²⁻
D) H₂CO₄
A) PV = nRT for all gases
B) Rate of effusion is inversely proportional to √M
C) Total pressure = sum of partial pressures of individual gases
D) Volume is inversely proportional to pressure
A) Increases down a group, decreases across a period
B) Decreases down a group, increases across a period
C) Increases in both directions
D) Decreases in both directions
A) Concentration to time
B) Rate constant to temperature and activation energy
C) pH to concentration
D) Cell potential to concentration
A) The second bond in a C=C double bond
B) The bond in a C≡C triple bond formed by end-to-end overlap
C) A bond formed by p orbital side-to-side overlap
D) A bond only found in aromatic rings
A) Increases vapor pressure
B) Decreases vapor pressure
C) Has no effect
D) Effect depends on temperature only
A) 25 mL
B) 50 mL
C) 100 mL
D) 200 mL
A) Glucose (C₆H₁₂O₆)
B) NaCl
C) CaCl₂
D) Acetic acid (HC₂H₃O₂)
A) Find equilibrium constant from cell potential
B) Calculate cell potential under non-standard conditions
C) Balance redox half-reactions
D) Determine pH of a buffer
A) The 2nd ionization energy is always less than the 1st
B) There is a large jump in IE when removing an electron from a stable noble gas configuration
C) Ionization energy decreases as you move right across a period
D) Metals have higher ionization energies than nonmetals
A) C has two single bonds to O, each O has 3 lone pairs
B) O=C=O with 2 lone pairs on each O and no lone pairs on C
C) C has one double bond and one single bond to O
D) C has two single bonds and two lone pairs
A) Evaporation
B) Sublimation
C) Condensation
D) Deposition
A) 0.0347 min⁻¹
B) 0.0500 min⁻¹
C) 14.4 min⁻¹
D) 20 min⁻¹
A) Constitutional (structural) isomers
B) Stereoisomers
C) Conformational isomers
D) Tautomers
A) ΔG° = RT/ln K
B) ΔG° = −nFE°
C) ΔG° = −RT ln K
D) ΔG° = K/RT
A) ²³⁴Th + ⁴He
B) ²³⁴Pa + ⁴He
C) ²³⁴U + ⁴He
D) ²³⁶U + ²n
A) 11,460 years
B) 17,190 years
C) 5,730 years
D) 22,920 years
A) An alpha particle and a neutrino
B) An electron and an antineutrino
C) A positron and a neutrino
D) Two gamma photons
A) Fission releases more per gram
B) Fusion releases more per gram
C) They release exactly equal energy per gram
D) Neither releases energy; both absorb energy
A) 'a' corrects for temperature; 'b' corrects for pressure
B) 'a' corrects for intermolecular attractions; 'b' corrects for finite molecular volume
C) 'a' corrects for finite molecular volume; 'b' corrects for intermolecular attractions
D) Both correct for pressure deviations only
A) H₂ effuses 4 times faster than O₂
B) H₂ effuses 16 times faster than O₂
C) O₂ effuses faster because it's heavier
D) They effuse at the same rate
A) The highest temperature at which a liquid can exist
B) The conditions where solid, liquid, and gas coexist in equilibrium
C) The temperature at which a gas cannot be liquefied
D) The normal boiling point
A) A slanted segment where q = mcΔT
B) A flat segment at 100°C where q = m·ΔHvap
C) A flat segment at 0°C where q = m·ΔHfus
D) Any segment where temperature increases
A) CH₄ < NH₃ < HF < H₂O
B) HF < NH₃ < CH₄ < H₂O
C) CH₄ < HF < NH₃ < H₂O
D) CH₄ < H₂O < NH₃ < HF
A) Ice molecules move faster than liquid water molecules
B) Hydrogen bonding in ice creates an open hexagonal lattice structure less dense than liquid
C) Ice has stronger London dispersion forces
D) Ice contains fewer water molecules per unit volume due to lower temperature
A) −0.93°C
B) −1.86°C
C) −3.72°C
D) −0.46°C
A) 2.45 atm
B) 0.245 atm
C) 24.5 atm
D) 0.82 atm
A) P_solvent = χ_solute × P°_solvent
B) P_solvent = χ_solvent × P°_solvent
C) P_solvent = P°_solvent + χ_solvent
D) P_solvent = P°_solvent / χ_solute
A) Zero order
B) First order
C) Second order
D) Third order
A) 0.500 M
B) 0.250 M
C) 0.693 M
D) 0.100 M
A) High temperature and low pressure
B) Sufficient energy (≥ Eₐ) and proper molecular orientation
C) High concentration and slow molecular speed
D) Catalyst present and high concentration
A) The energy of the reactants
B) The energy of the products
C) The energy maximum along the reaction pathway
D) The heat of reaction
A) High temperature and low pressure
B) Low temperature and high pressure
C) High temperature and high pressure
D) Low temperature and low pressure
A) Kp = Kc·(RT)²
B) Kp = Kc·(RT)⁻²
C) Kp = Kc
D) Kp = Kc·(RT)⁴
A) Increase AgCl solubility
B) Decrease AgCl solubility
C) Have no effect on solubility
D) Increase Ksp
A) pH = 4.44
B) pH = 5.04
C) pH = 4.74
D) pH = 5.74
A) Neutral (pH = 7)
B) Acidic (pH < 7)
C) Basic (pH > 7)
D) The same pH as the original weak acid
A) Basic, because NH₃ is a weak base
B) Acidic, because NH₄⁺ is the conjugate acid of a weak base
C) Neutral, because Cl⁻ and NH₄⁺ cancel
D) Basic, because NaCl analogy
A) MnO₄⁻ + 4H⁺ + 3e⁻ → Mn²⁺ + 2H₂O (reduction)
B) MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O (reduction)
C) MnO₄⁻ → Mn²⁺ + 4O²⁻ (oxidation)
D) MnO₄⁻ + 5e⁻ → Mn²⁺ (reduction)
A) +147 kJ/mol
B) −147 kJ/mol
C) +73.5 kJ/mol
D) −73.5 kJ/mol
A) +0.28 V
B) +0.40 V
C) +0.34 V
D) −0.34 V
A) 0.635 g
B) 1.27 g
C) 2.54 g
D) 3.18 g
A) Primary alcohol → ketone → carboxylic acid
B) Primary alcohol → aldehyde → carboxylic acid
C) Primary alcohol → ether → ester
D) Primary alcohol → alkene → alkane
A) An alcohol and a ketone
B) A carboxylic acid and an alcohol, producing an ester and water
C) An aldehyde and an amine
D) A fat and NaOH
A) Polyethylene (from CH₂=CH₂)
B) Polystyrene (from styrene monomer)
C) Nylon (from diamine + diacid)
D) Natural rubber (from isoprene)
A) Aldehyde (−CHO)
B) Ketone (C=O between two carbons)
C) Ether (−O−)
D) Carboxylic acid (−COOH)
A) Volume
B) Pressure
C) Absolute temperature
D) Molar mass
A) High temperature and low pressure
B) Low temperature and high pressure
C) Low temperature and low pressure
D) High temperature and high pressure
A) Mass of products minus mass of reactants in chemical reactions
B) The energy required to remove one proton from a nucleus
C) The energy equivalent of the mass lost when a nucleus forms from its constituent nucleons (E = Δmc²)
D) The heat released during radioactive decay
A) Increases by 1
B) Decreases by 1
C) Stays the same
D) Decreases by 2
A) Boiling point elevation
B) Freezing point depression
C) Color of a solution
D) Osmotic pressure
A) Rate doubles
B) Rate halves
C) Rate stays the same
D) Rate quadruples
A) HCl (hydrochloric acid)
B) NaOH (sodium hydroxide)
C) H₂SO₄ (sulfuric acid)
D) H₂O (water only)
A) Release water as a byproduct
B) Form only from two different monomers
C) Form by opening double bonds with no small-molecule byproduct
D) Have lower molecular weights
A) rate = k[A][B]
B) rate = k[B]²
C) rate = k[B]
D) rate = k[A]²[B]
A) 2 times
B) 10 times
C) 100 times
D) 1,000 times
A) The nucleus emits an electron
B) An inner-shell electron is captured by the nucleus, converting a proton to a neutron
C) Two electrons annihilate, releasing gamma rays
D) A neutron captures an electron and becomes a proton
A) Dipole-dipole forces
B) Hydrogen bonding
C) London (dispersion) forces
D) Ion-dipole forces
A) 2.77 min
B) 4.00 min
C) 0.693 min
D) 1.38 min
A) −OH (hydroxyl)
B) −COOH (carboxyl)
C) −NH₂ (amino)
D) −CHO (aldehyde)
A) Proceed forward (toward products)
B) Proceed in reverse (toward reactants)
C) Be at equilibrium
D) Stop completely
A) Catalysts increase the value of Kc
B) Catalysts are consumed in the reaction
C) Catalysts lower the activation energy but do not shift equilibrium
D) Catalysts only affect the reverse reaction
A) pH = 2
B) pH = 1
C) pH = 10
D) pH = 12
A) CH₄
B) CCl₄
C) H₂S
D) HF
A) Both require an external power source
B) Galvanic cells produce electrical energy from spontaneous reactions; electrolytic cells use electrical energy to drive non-spontaneous reactions
C) Electrolytic cells are spontaneous; galvanic cells are non-spontaneous
D) Both have the same electrode designations (anode positive, cathode negative)
A) Mass and charge
B) Position and momentum
C) Spin and charge
D) Energy and charge
A) The 3d sublevel fills before 4s in period 4 elements
B) A half-filled 3d sublevel (d⁵) provides extra stability, making the actual configuration lower in energy
C) The 4s orbital can hold only one electron when 3d is partially filled
D) Chromium violates the Pauli exclusion principle
A) S has a larger atomic radius than P
B) The paired electron in S's 3p subshell experiences greater electron–electron repulsion, making it easier to remove
C) P has a higher nuclear charge than S
D) S's 3p electrons are shielded less effectively than P's
A) Tetrahedral
B) Trigonal bipyramidal
C) Seesaw (disphenoidal)
D) Square planar
A) Octahedral; square planar
B) Trigonal bipyramidal; seesaw
C) Octahedral; octahedral
D) Tetrahedral; square planar
A) sp³; 4 hybrid orbitals
B) sp³d; 5 hybrid orbitals
C) sp³d²; 6 hybrid orbitals
D) sp²d; 5 hybrid orbitals
A) Bond order 3; diamagnetic
B) Bond order 2; diamagnetic
C) Bond order 2; paramagnetic
D) Bond order 1; paramagnetic
A) All single bonds; minimizes electron sharing
B) Double bonds to oxygen; formal charges closer to zero are preferred, and electronegative atoms should have negative formal charges
C) Single bonds are always preferred to avoid expanded octets
D) Both structures are equally valid and contribute equally to the resonance hybrid
A) +67.6 kJ
B) −293.6 kJ
C) +293.6 kJ
D) +180.6 kJ
A) +248 kJ
B) −183 kJ
C) −248 kJ
D) +183 kJ
A) H₂O(l) → H₂O(s)
B) CaCO₃(s) → CaO(s) + CO₂(g)
C) N₂(g) + 3H₂(g) → 2NH₃(g)
D) NaCl(s) → NaCl(aq) at low concentration
A) Above 480 K
B) Below 480 K
C) At all temperatures
D) At no temperature, because ΔH is positive
A) +146.7 kJ/mol
B) −146.7 kJ/mol
C) −73.3 kJ/mol
D) +73.3 kJ/mol
A) 8,000 J/mol
B) 66,520 J/mol
C) 96,120 J/mol
D) 961,200 J/mol
A) High temperature and high pressure only
B) Sufficient activation energy AND proper molecular orientation
C) Catalyst presence AND aqueous solution
D) High concentration AND large molecular size
A) The energy of reactants increases
B) The energy of products decreases
C) The activation energy (Ea) decreases, but ΔH remains the same
D) Both Ea and ΔH decrease
A) Increasing temperature
B) Decreasing pressure
C) Removing NH₃ as it forms
D) Adding an inert gas at constant volume
A) Kp = Kc always
B) Kp = Kc(RT)^Δn, so Kp > Kc at this temperature
C) Kp = Kc/RT, so Kp < Kc
D) Kp and Kc are unrelated quantities
A) 4.74
B) 4.92
C) 4.56
D) 5.10
A) Solubility increases due to increased ionic strength
B) Solubility decreases because the common ion Cl⁻ shifts equilibrium left (common ion effect)
C) Solubility is unchanged because Ksp is constant
D) Solubility increases because Ag⁺ forms complexes with Cl⁻
A) E°cell = +1.10 V; Cu is oxidized
B) E°cell = +1.10 V; Zn is oxidized
C) E°cell = −0.42 V; Zn is oxidized
D) E°cell = −1.10 V; Cu is oxidized
A) E increases above E°
B) E decreases below E°
C) E remains equal to E°
D) E becomes negative regardless of E°
A) 0.99 g
B) 1.78 g
C) 3.56 g
D) 7.12 g
A) ²³⁸U → ²³⁸Np + ⁰e (beta)
B) ²³⁸U → ²³⁴Th + ⁴He (alpha)
C) ²³⁸U → ²³⁴Pa + ⁴He
D) ²³⁸U → ²³⁸U + γ (gamma)
A) 1/2
B) 1/8
C) 1/16
D) 1/32
A) It escapes as neutrinos
B) It is converted to energy according to E = mc²
C) It becomes binding energy stored in the daughter nuclei
D) It is transferred to the neutrons as kinetic energy
A) Principal quantum number (n); values 1, 2, 3
B) Angular momentum quantum number (ℓ); values 0, 1, 2
C) Magnetic quantum number (mℓ); values −3 to +3
D) Spin quantum number (ms); values +1/2, −1/2
A) Electrons are being added to inner shells, pushing outer electrons farther out
B) Increasing nuclear charge pulls valence electrons closer, with shielding roughly constant across the period
C) Valence electrons repel each other more strongly as more electrons are added
D) The number of electron shells decreases from Na to Cl
A) Rate doubles
B) Rate increases by a factor of 4
C) Rate increases by a factor of 8
D) Rate is unchanged
A) Fission combines light nuclei; fusion splits heavy nuclei
B) Fusion requires extremely high temperatures; fission does not
C) Both fission and fusion release energy by decreasing binding energy per nucleon
D) Only fission is used in current commercial nuclear power plants
A) The temperature and pressure at which solid and liquid are in equilibrium
B) The temperature and pressure at which all three phases (solid, liquid, gas) coexist in equilibrium
C) The maximum temperature at which liquid can exist
D) The point at which a substance transitions from gas to plasma
A) Surface tension; ΔTb = Kb/m
B) Boiling point elevation; ΔTb = Kb × m × i
C) Vapor pressure; ΔP = P°solvent × Xsolute only for gases
D) Osmotic pressure; Π = MRT only for macromolecules
A) Fe₂O₃ + 3CO → 2Fe + 3CO₂; CO is reduced (gains oxygen)
B) Fe₂O₃ + 3CO → 2Fe + 3CO₂; CO is oxidized (loses electrons/gains oxygen)
C) Fe₂O₃ + 3CO → 2Fe + 3CO₂; Fe³⁺ is oxidized
D) This is not a redox reaction because no metal dissolves in acid
A) Gas molecules have negligible volume
B) There are no intermolecular forces between gas molecules
C) The average kinetic energy of gas molecules is directly proportional to absolute temperature
D) Gas molecules move in straight lines between collisions
A) HCl and NaOH
B) H₂SO₄ and SO₄²⁻
C) H₃O⁺ and OH⁻
D) NH₃ and NH₄⁺
A) k = 0.050 min⁻¹
B) k = 0.0347 min⁻¹
C) k = 0.693 min⁻¹
D) k = 20 min⁻¹
A) London dispersion forces
B) Dipole-dipole interactions
C) Hydrogen bonding
D) Ion-dipole forces
A) 50.00 mL
B) 25.00 mL
C) 12.50 mL
D) 5.00 mL
A) C–C (single) is shorter and stronger than C=C (double)
B) C≡C (triple) is shorter and stronger than C=C (double), which is shorter and stronger than C–C (single)
C) Bond length increases as bond order increases
D) Bond strength is independent of bond length
A) 3 times faster
B) 9 times faster
C) 6 times faster
D) Equal rates
A) Gas molecules slow down at high pressure
B) Intermolecular attractions become significant, and molecular volume is no longer negligible
C) High pressure causes gas molecules to decompose
D) The gas constant R changes at high pressures
A) Enantiomers
B) Constitutional (structural) isomers
C) Cis-trans (geometric) isomers
D) Conformational isomers
A) –OH attached to carbonyl carbon = alcohol
B) –COOH (carboxyl group) = carboxylic acid
C) –NH₂ attached to carbonyl = ester
D) –CHO (aldehyde group) = ketone
A) Ethanol and acetic acid
B) Sodium acetate and ethanol
C) Sodium ethoxide and acetaldehyde
D) Ethyl acetate is not hydrolyzable by base
A) Strong electrolytes are ionic compounds; weak electrolytes are covalent
B) Strong electrolytes completely dissociate in water; weak electrolytes only partially dissociate
C) Strong electrolytes conduct electricity better because they are larger molecules
D) Weak electrolytes do not dissolve in water at all
A) Second-order; rate depends on collision of two H₂O₂ molecules
B) First-order; rate depends on [H₂O₂] linearly, and the reaction likely proceeds by a mechanism with a unimolecular rate-determining step
C) Zeroth-order; rate is independent of concentration
D) The rate law must match the stoichiometry (second-order since coefficient is 2)
A) 1.0 m glucose (C₆H₁₂O₆) in water
B) 1.0 m NaCl in water
C) 1.0 m CaCl₂ in water
D) 1.0 m sucrose in water
A) sp
B) sp²
C) sp³
D) sp³d
A) NaOH (forms colored precipitate)
B) AgNO₃ (forms white AgCl precipitate)
C) BaCl₂ (forms white BaSO₄)
D) H₂SO₄ (liberates gas)
A) The entropy of the universe decreases in all spontaneous processes
B) Entropy of a system always increases during a spontaneous process
C) The total entropy of the universe (system + surroundings) increases for every spontaneous process
D) A process can be spontaneous only if it is exothermic
A) Principal quantum number (n)
B) Angular momentum quantum number (ℓ)
C) Magnetic quantum number (mℓ)
D) Spin quantum number (ms)
A) Tetrahedral
B) See-saw (disphenoidal)
C) Trigonal pyramidal
D) Square planar
A) 6.0 L
B) 12.0 L
C) 3.0 L
D) 1.5 L
A) HCl and NaCl
B) H₂O and OH⁻
C) H₂SO₄ and NaOH
D) NH₄⁺ and N₂H₄
A) Cathode; positive (+)
B) Anode; negative (−)
C) Anode; positive (+)
D) Cathode; negative (−)
A) 278 g/mol
B) 310 g/mol
C) 182 g/mol
D) 246 g/mol
A) Kc = [CH₄][H₂O] / ([CO][H₂]³)
B) Kc = [CO][H₂]³ / ([CH₄][H₂O])
C) Kc = [CH₄][H₂O] / [CO][H₂]
D) Kc = [CO][H₂]³ / [CH₄]
A) +2
B) +4
C) +7
D) +6
A) 0.100 M
B) 0.200 M
C) 1.00 M
D) 0.050 M
A) Is always exothermic if the reaction goes to completion
B) Can be calculated by adding the enthalpy changes of a series of steps that sum to the overall reaction, regardless of the pathway taken
C) Depends on the physical state of only the products
D) Equals the activation energy of the reaction divided by temperature
A) Fission joins small nuclei; fusion splits large nuclei
B) Fission splits heavy nuclei (like U-235) into smaller fragments, releasing energy; fusion combines light nuclei (like H isotopes) to form heavier nuclei, also releasing energy
C) Both fission and fusion absorb energy and require extreme activation energy input with no net energy release
D) Fission occurs naturally in stars; fusion is used in nuclear power plants
A) London dispersion forces, which are stronger in H₂O due to its smaller size
B) Hydrogen bonding in H₂O — the highly electronegative oxygen and small hydrogen atom size allow unusually strong H···O interactions between molecules
C) Ion-dipole forces that form when water ionizes in solution
D) Covalent network bonding that links water molecules into a macromolecular structure
A) A strong acid and its conjugate base in equal molar quantities
B) A weak acid and its conjugate base (or a weak base and its conjugate acid) in significant amounts, which can neutralize added strong acid or base
C) Only distilled water, which has equal concentrations of H⁺ and OH⁻
D) A strong acid and a strong base that neutralize each other to maintain constant pH
A) ΔG = ΔH + TΔS
B) ΔG = ΔH − TΔS
C) ΔG = ΔS − TΔH
D) ΔG = TΔS / ΔH
A) 2
B) 12
C) 7
D) 4
A) Zinc (Zn)
B) Potassium (K)
C) Copper (Cu)
D) Chromium (Cr)
A) O₂ is limiting; 4 moles H₂O produced
B) H₂ is limiting; 4 moles H₂O produced
C) O₂ is limiting; 6 moles H₂O produced
D) H₂ is limiting; 2 moles H₂O produced
A) 1/2
B) 1/4
C) 1/8
D) 1/6
A) CH₄ (methane)
B) HCl (hydrogen chloride)
C) O₃ (ozone)
D) H₂O (water)
A) Left, because increasing pressure favors more moles of gas
B) Right, because increasing pressure favors the side with fewer moles of gas (2 mol vs. 4 mol)
C) Neither direction, because pressure does not affect equilibrium of gas-phase reactions
D) Right only at low temperatures, and left at high temperatures
A) Hydrogen bonding
B) Dipole-dipole forces
C) London dispersion forces (induced dipole-induced dipole)
D) Ion-dipole forces
A) 1
B) 3
C) 5
D) 7
A) The quantum mechanical behavior of electrons at low temperatures
B) Intermolecular attractions (a/V² correction to P) and the finite volume occupied by gas molecules themselves (b correction to V)
C) The effect of gravity on heavier gas molecules at high altitudes
D) The conversion between mass and moles for non-ideal conditions
A) Colligative properties depend on the chemical identity of the solute particles
B) Colligative properties (boiling point elevation, freezing point depression, osmotic pressure, vapor pressure lowering) depend only on the number of solute particles, not their chemical identity
C) Only ionic solutes affect colligative properties; molecular solutes have no effect
D) Colligative properties are only applicable to gaseous solutions
A) A neutron converts to a proton, emitting a beta particle (electron) and antineutrino
B) The nucleus emits a helium-4 nucleus (²₂He), reducing atomic number by 2 and mass number by 4
C) The nucleus absorbs energy from gamma rays and ejects a proton
D) Two protons annihilate with two electrons, releasing energy as gamma radiation
A) Ksp = [Ag⁺][CrO₄²⁻] = s²
B) Ksp = [Ag⁺]²[CrO₄²⁻] = (2s)²(s) = 4s³
C) Ksp = [Ag⁺][CrO₄²⁻]² = (s)(s²) = s³
D) Ksp = [Ag⁺]²[CrO₄²⁻]² = 4s⁴
A) 90°
B) 109.5°
C) 120°
D) 180°
A) Primary alcohols have one –OH group; secondary have two; tertiary have three
B) Primary alcohols have the –OH on a carbon bonded to one other carbon; secondary on a carbon bonded to two; tertiary on a carbon bonded to three other carbons
C) Primary alcohols are the simplest (methanol), secondary are 6–carbon chain alcohols, and tertiary are over 12 carbons
D) The classification depends on the number of hydrogen atoms directly bonded to the oxygen
A) Cell potential increases as reactants are consumed and products accumulate
B) Cell potential decreases as the reaction proceeds — the concentrations move toward equilibrium (Q increases), reducing Ecell until Ecell = 0 at equilibrium
C) Cell potential remains constant regardless of concentration changes during discharge
D) Cell potential first increases then decreases, reaching a maximum at 50% discharge
A) Combination (synthesis) reaction
B) Decomposition reaction
C) Combustion reaction
D) Single displacement reaction
A) The total pressure equals the pressure of the most abundant gas
B) The total pressure of the mixture equals the sum of the partial pressures of each individual gas (Ptotal = P₁ + P₂ + P₃ + ...)
C) Each gas in a mixture has the same partial pressure regardless of its mole fraction
D) Heavier gases exert greater pressure than lighter gases at the same temperature and volume
A) HCl + NaOH → NaCl + H₂O (proton transfer)
B) BF₃ + NH₃ → F₃B–NH₃ (electron pair donation)
C) Na + Cl₂ → NaCl (electron transfer/redox)
D) CO₂ + H₂O → H₂CO₃ (gas absorption)
A) pH = 0
B) pH = pKa
C) pH = 14
D) pH = 2pKa
A) 11.1%
B) 88.9%
C) 50.0%
D) 33.3%
A) Only the concentration of the substrate (first order overall)
B) Both the concentration of the nucleophile and the substrate (second order overall); the reaction proceeds in one step with simultaneous bond formation and breaking via a backside attack
C) Temperature only, with no dependence on concentration
D) The stability of the carbocation intermediate formed after the leaving group departs
A) High-intensity light of any frequency will eject electrons from a metal surface
B) Light must exceed a threshold frequency (not intensity) to eject electrons, because light energy is quantized in photons (E = hν); Einstein's explanation demonstrated the particle nature of light
C) The photoelectric effect proved that electrons orbit the nucleus in discrete shells with fixed energies
D) It demonstrated that electrons have wavelike properties by showing diffraction patterns
A) 0.280 M
B) 0.140 M
C) 0.350 M
D) 0.070 M
A) Ether and water
B) Ester and water (Fischer esterification)
C) Amide and water
D) Aldehyde and water
A) −1
B) 0
C) +1
D) +4
A) 1.00 L/mol
B) 22.4 L/mol
C) 8.314 L/mol
D) 6.022 L/mol
A) Diamond has sp² hybridized carbon; graphite has sp³ hybridized carbon
B) Diamond has a 3D covalent network with sp³ carbons (4 single bonds) — hardest natural substance; graphite has sp² carbons in 2D hexagonal layers with delocalized π electrons — soft and electrically conducting
C) Diamond is ionic; graphite is metallic in bonding
D) They have the same crystal structure but different atomic masses
A) Doubles (×2)
B) Quadruples (×4)
C) Remains the same (×1)
D) Halves (×0.5)
A) Components separate based on their density — heavier components settle faster in the paper matrix
B) Components separate based on their differential affinities for the stationary phase (paper/water) versus the mobile phase (solvent) — more polar substances travel less far; less polar travel farther with organic solvents
C) Components are separated by electric charge — positively charged molecules migrate toward the negative electrode embedded in the paper
D) Separation depends entirely on molecular size — smaller molecules migrate faster through the paper pores
A) Rate ∝ molar mass; O₂ effuses 4× faster than H₂
B) Rate ∝ 1/√(molar mass); H₂ effuses 4× faster than O₂
C) Both effuse at the same rate because they are at the same temperature
D) Rate ∝ √(molar mass); H₂ effuses 4× slower than O₂
A) Electrons are added to higher energy subshells that are larger and farther from the nucleus
B) As atomic number increases across a period, the nuclear charge (protons) increases while electrons are added to the same principal energy level (same shell), pulling the electron cloud closer to the increasingly positive nucleus
C) Electrons repel each other more strongly across a period, compressing all orbitals
D) The principal quantum number increases across a period, placing electrons farther from the nucleus
A) 7.00 (neutral)
B) The pKa of the weak acid
C) 14 minus the pKb of the conjugate base
D) The pH at the equivalence point
A) H₂ gas at cathode; Cl₂ gas at anode
B) Na metal at cathode; Cl₂ gas at anode
C) Na metal at cathode; O₂ gas at anode
D) Cl₂ gas at cathode; Na metal at anode
A) CH₂O
B) C₂H₄O₂
C) C₆H₁₂O₆
D) C₃H₆O₃
A) ΔH°f is always exothermic; ΔH°rxn can be endothermic or exothermic
B) ΔH°f is the enthalpy change for forming 1 mole of a compound from its elements in standard states; ΔH°rxn = Σ ΔH°f(products) − Σ ΔH°f(reactants), and includes stoichiometric coefficients
C) ΔH°f applies only to ionic compounds; ΔH°rxn applies to molecular reactions
D) They are identical values expressed in different units
A) E°cell = RT·ln K (no dependence on number of electrons transferred)
B) E°cell = (RT/nF)·ln K = (0.0592/n)·log K at 25°C, where n = moles of electrons transferred; a positive E°cell means K > 1 (products favored)
C) K = E°cell/nF at all temperatures
D) E°cell and K are unrelated; one is thermodynamic and the other is kinetic