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A Few Things You Might Want To Know For The 2013 AP CHEMISTRY EXAM —Florian König 2012-2013 AP Chemistry Hastings High School 2 "For a man to attain to an eminent degree in learning costs him time, watching, hunger, nakedness, dizziness in the head, weakness in the stomach, and other inconveniences." – Miguel de Cervantes “It’s not having been in the dark house, it’s having left it, that counts.” – Theodore Roosevelt 3 TABLE OF CONTENTS Basic Concepts Atomic Theory and the Periodic Table Electronic Structure of the Atom Quantum Mechanics Ionic Bonding Covalent Bonding VSEPR Model Hybridization MO Theory Reaction Types and Stoichiometry Gases Liquids Solids Phase Diagrams Solutions and Solubility Rules Acids and Bases in Water Redox Reactions Kinetics Equilibrium Aqueous Equilibria I (Acids and Bases) Aqueous Equilibria II (Solutions) Thermodynamics Electrochemistry Organic Chemistry Nuclear Chemistry Before the Test During the Test 4 5 6 7 10 11 13 15 16 18 19 21 23 25 26 29 30 31 33 34 40 41 45 48 50 52 52 Appendix A: Appendix B: Appendix C: Appendix D: Appendix E: Appendix F: 53 100 102 104 105 109 Equations in the AP Examinations Solubility Rules For Ionic Compounds in Water Nomenclature of Inorganic Compounds Nomenclature of Organic Compounds Miscellaneous Helpful Websites 4 BASIC CONCEPTS 1. Physical changes do not affect the nature of a compound; chemical changes do. Physical and chemical properties go along with according changes. 2. Mixtures can be heterogeneous or homogeneous (= solutions). They consist of substances that can be separated by physical changes (distillation, crystallization, chromatography). Substances can be either elements or compounds. Compounds can be separated into elements by chemical changes (redox reactions). 3. Law of Definite Proportions (Law of Constant Composition) The composition of a compound is constant. The relative masses of elements in a compound form simple, whole-number ratios. 4. Law of Multiple Proportions If two elements combine with each other in more than one ratio, the relative masses of the varying element form simple, whole-number ratios. 5. Intensive and Extensive Properties Intensive properties do not depend on the amount of matter (solubility, density, color, reactivity, conductivity, etc.); extensive properties do (mass, volume, etc.). 6. Significant Figures Significant figures = all certain digits + 1 uncertain digit Determination of significant figures: Atlantic-Pacific Rule Calculation Significant figures are the … Addition/subtraction least number of decimal places Multiplication/division least number of significant figures pH only decimal places are significant Notes: 5 ATOMIC THEORY AND THE PERIODIC TABLE 7. Atomic Theory Contributions by LAVOISIER and DALTON ( 3. Law of Definite Proportions) 8. THOMSON: “cathode rays” Proved existence of electrons Determination of charge/mass ratio of e Leads to “plum pudding model” of atoms 9. MILLIKAN: “oil-drop experiment” Determination of e- charge (see Appendix E) Determination of e- mass (see Appendix E) 10. RUTHERFORD: “gold foil experiment” Atoms consist of small, heavy, positive nuclei. Electrons are negative, and circle the nucleus. The nucleus-electron distance far exceeds the nucleus size. (“Atoms are mostly empty space.”) 11. Nuclides are nuclei with Specific proton and neutron number (same #p + = isotopes) 12. The Periodic Table (see: DÖBEREINER, LOWLAND, MENDELEEV, MOSELEY) Periodic arrangement by atomic number and properties Group numbers and names, representative elements Distinction between metals, semimetals and nonmetals Trends in metallic/nonmetallic character, ionization energy, electron affinity, electronegativity, activity series of metals Notes: 6 ELECTRONIC STRUCTURE OF THE ATOM 13. Electronic Structure of the Atom Speed of light c = (frequency ) x (wavelength ) Visible spectrum: 400 nm (blue) 700 nm (red) Energy is found in “quanta” with E = h (h = PLANCK’s constant). Line spectrum: distinct wavelengths, i.e., electron transitions between distinct energy levels Continuous spectrum: no distinct wavelengths 14. BOHR’s Model of the Hydrogen Atom Absorption or release of energy = jump of e- between discrete energy levels All e- in lowest possible state: “ground state” Flaw: electrons in ground-state atoms do not lose energy as predicted by classical mechanics 15. Matter-Wave Dualism (deBROGLIE) Very small particles (photons, electrons) can exhibit wave behavior h = —— mv 16. m = mass (kg); v = velocity (m/s) “Allowed orbit” = orbit length is a multiple of the electron wavelength HEISENBERG’s Uncertainty Principle It is impossible to know position and momentum of an electron exactly. Uncertainty of momentum (p = mv): p Uncertainty of position: q h p q —— 2 Notes: 7 QUANTUM MECHANICS 17. Quantum-mechanical Description of the Atom: SCHRÖDINGER: Electrons are described by wave-functions (exact only in one-electron systems, e.g., H, He+). 2 = probability of finding e- = deBROGLIE “orbital “ (replaces BOHR’s “orbits”) Quantum numbers: principal quantum number ( energy): n = 1, 2, 3, 4, … azimuthal quantum number ( orbital shape): l = 0, 1, …, (n-1) corresponds to s, p, d, f Orbital type s p d Orbital shape spherical “dumbbells” “double dumbbells” magnetic quantum number ( orbital orientation): ml = - l, …, 0, … , + l spin quantum number ( orientation of e- in magnetic field) ms = + ½ , - ½ 18. PAULI Exclusion Principle No two electrons in an atom have the same four quantum numbers. 19. The Aufbau Principle Arrange orbitals in energetic sequence Start filling sublevels with lowest energy first 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p 8 20. Electron arrangement: all electrons are in lowest possible state ground state one or more electrons occupy higher levels excited state s- and p-electrons of the outermost principal energy level (“valence shell”) are called “valence electrons.” (determine reactivity) HUND’s Rule Orbitals with the same energy (= same sublevel) are called “degenerate.” They are filled with one electron each before pairing up occurs. 9 Notes: 10 IONIC BONDING 21. “Octet Rule” Atoms tend to attain noble-gas configuration, i.e., a filled valence shell with 8 (or 2) valence electrons. 22. Transition Metals s-electrons ionize first; half-filled or filled sublevels (especially dsublevels) are relatively stable. 23. The Ionic Bond Octet Rule is followed by gaining or losing valence electrons. Metal (low EN) + nonmetal (high EN) Determining factor for stability of compound: ionization energy + electron affinity + lattice energy Cations are smaller than metal atoms, anions bigger than nonmetal atoms. Ionic compounds: hard, brittle, high m.p./ b.p., conductors when liquid or dissolved in polar solvents (= electrolytes) Notes: 11 COVALENT BONDING 24. The Covalent Bond Nonmetal + nonmetal Octet Rule is followed by sharing electrons (1-3 per atom). 25. Multiple Bonds Up to six electrons (= triple bond) are shared between two atoms. Multiple bonds are shorter and stronger than single bonds. 26. Lewis Structures Add up available valence e-, add/subtract extra charges. Derive central atom (often C, N, P, S). Draw single bonds first. Put excess valence e- at central atom. Electron deficiencies? (if yes: form multiple bonds) Structures with low formal charges are favored. Negative charges on electronegative atoms are favored. Formal charges: bonds are broken homolytically (= evenly) (compare # of valence e- to “normal” # in a neutral atom). 27. Resonance Degenerate “resonance formulas” add to stability. 28. Exceptions to Octet Rule Odd number of valence e- (e.g., NO, NO2) Less than 8 valence e- (Per. 2 elements: e.g., BeH2, BF3) More than 8 valence e- require available d-orbitals (Per. 3+ elements: e.g., ICl4-, SF6) 29. Bond and Molecule Polarity and Electronegativity EN < 0.4 nonpolar 0.4 < EN < 1.7 polar EN > 1.7 ionic Polar molecules only if polar bonds and unsymmetrical geometry Bond and molecule polarities are not discrete properties, i.e., they change gradually. The EN differences serve as rough guidelines only for the determination of bind polarity. 12 30. Oxidation Numbers Bonds are formally broken heterolytically, i.e., both e - are counted towards the more electronegative atom Elements: 0 Monoatomic ions: oxidation number = charge (e.g., Na +, N3-) Polyatomic compounds: sum of oxidation numbers = total ionic charge Notes: 13 VSEPR MODEL 31. Maximum Repulsion ( angle) between electron pairs VSEPR = Valence Shell Electron Pair Repulsion Repulsion: lone pair > bonding pair > unpaired single electron Distinction between molecular geometry and electron-pair geometry “Electron-pair geometry” is also called “electron-domain (ED) geometry.” Multiple bonds are treated as one ED, repel stronger (higher edensity). 32. Common Electron-Domain and Molecule Geometries # of ED 2 3 4 5 6 Example CO2 BeH2 BCl3 NO2CH4 NH3 H2O PF5 SF4 ClF3 XeF2 SF6 BrF5 XeF4 ED geometry linear linear trigonal planar trigonal planar tetrahedral tetrahedral tetrahedral trigonal bipyramidal trigonal bipyramidal trigonal bipyramidal trigonal bipyramidal octahedral octahedral octahedral Molecule Geometry linear linear trigonal planar bent tetrahedral pyramidal bent trigonal bipyramidal seesaw T-shaped linear octahedral square pyramidal square planar The ideal bond angles and electron-domain geometries are: 180o (linear) o 120 (trigonal planar) o 109.5 (tetrahedral) 90o/120o (trigonal bipyramidal) 90o (octahedral) Actual bond angles can vary due to the repulsive effect of 1) lone pair < single bond and 2) single bond < double bond < triple bond 14 Notes: 15 HYBRIDIZATION 33. Symmetrical Overlap -bond (between nuclei) 34. Unsymmetrical Overlap -bond (above nucleus-nucleus axis) 35. Orbitals s-orbitals: only -bonds p-orbitals: - and -bonds possible 36. Hybridization At least 2 orbitals from degenerate hybrid orbitals Number of hybridizing orbitals = Number of hybrid orbitals 37. Common Hybridizations Hybridization sp sp2 sp3 sp2d sp3d sp3d2 Notes: Example BeH2, HgCl2, C2H2 C2H4, BF3 CH4, NH3, H2O PdBr42PF5, SF4 SF6, SbCl6-, ICl4- Electron-domain geometry linear trigonal planar tetrahedral square planar trigonal bipyramidal octahedral 16 MO THEORY 38. “O2 problem” Valence-bond and hybridization approach provide no explanation for the paramagnetism (= unpaired e-) of O2 39. Atomic Orbitals (AO) and Molecular Orbitals (MO) AOs combine to form bonding (lower energy) and antibonding (higher energy) MOs see textbook for diagram Bonding: , ; antibonding: * , * (subscript = constituting AOs) e.g.: 1s — *1s — 2s — *2s — 2p — 2p (2) — 2p* (2) — 2p* 40. Bond Order (B.O.) B.O. = ½ [(# of bonding e-) – (# of antibonding e-)] B.O. > 0 stable bond B.O. 0 no stable bond 17 41. Rules # of AO = # of MO AOs combine effectively if they have good overlap and similar energies. MOs belong to the whole molecule. HUND’s Rule and the PAULI Exclusion Principle apply to MOs. 42. Delocalization p-orbitals that don’t participate in hybridization form a -MO that stretches across the molecule. e- are evenly distributed (example: benzene C6H6, nitrate NO3-). Stabilization of molecule leads to lower reactivity. Notes: 18 REACTION TYPES AND STOICHIOMETRY 43. Reaction Types: see Appendix A 44. Gram Atomic Mass (gam) / gram molecular mass (gmm) / gram formula mass (gfm) = mass of 1 mole of substance 45. Conversions (# of moles) x NA = # of particles (atoms, molecules, ions) (# of moles) x gfm = mass (# of moles) x 22.4 l mol-1 = volume (for ideal gases at STP) 46. Empirical Formula from Quantitative Analysis Assume 100g-sample. Change % into masses. Change masses into moles. Write formula. Divide by smallest subscript. Multiply by integer, if necessary, to remove simple fractions such as 0.5 or 0.33. 47. Limiting Reactant Check stoichiometry of reaction for “perfect” ratio. Check reactant that is “used up” first. 48. Concentrations moles solute (n) Molarity M = ———————— liters solution (V) Dilution: n1 = M1V1 = M2V2 = n2 Notes: 19 GASES 49. Properties take shape of container can be easily compressed 50. Pressure Manometer (standard pressure: 760 mm Hg = 760 Torr = 1 atm = 101.3 kPa) 51. Gas Laws BOYLE’s Law: pV = constant (T = const.) V CHARLES’ Law : —— = constant (p = const.) T AVOGADRO’s Hypothesis: Equal volumes of gas at equal temperature and equal pressure contain equal numbers of molecules. AVOGADRO’s Law: The volume of a gas is proportional to the quantity. (V n) 52. The Ideal Gas Equation pV = nRT (p = pressure, V = volume, n = number of moles, R = Universal Gas Constant, T = absolute temperature [Kelvin]) 53. DALTON’s Law of Partial Pressures ptotal = p1 + p2 + p3 + … + pn (total pressure = sum of partial pressures) 54. Molecular Weight and Gas Densities 55. Molecular mass MM = dRT —— p Gas Pressure (of gas collected over water) external pressure = ptotal = pgas + pwater d = density (gram/liter) 20 56. The Kinetic-Molecular Theory of Ideal Gases Gases consist of large numbers of molecules that are in constant, random, straight-lined motion. The volume of gas molecules is negligible compared to the total volume. There are only negligible intermolecular forces between molecules. Collisions are elastic, i.e., no thermal energy is lost. Average kinetic energy temperature 57. GRAHAM’s Law of Diffusion and Effusion At equal temperatures heavier molecules move slower. 58. m2 ——— (v1, v2, m1, m2 = molecular speeds & masses) m1 v1 ——— = v2 Effusion: rate of movement through small hole Diffusion: rate of movement through space Nonideal Gases Deviations at high pressure and/or low temperature an2 (p + ——) (V- nb) = RT V2 a = corrective term for intermoleclar attractions (reduces pressure) b = corrective term for finite molecular volume (increases overall volume) Manifestations of nonideal behavior: condensation at high pressure and/or low temperatures condensate takes up space Notes: 21 LIQUIDS 59. Solids Regular arrangement (crystalline lattice) Particles close together Little vibration, no free motion 60. Liquids Partial disorder (clusters and particles) Still relatively close together 61. Gases Maximum disorder Free motion Big intermolecular distance 62. Equilibria Between phases: solid liquid gas 63. Phase Changes For each phase change there is a molar enthalpy, e.g., Hfusion, Hvapor., etc.) No temperature change, i.e., enthalpy change = change in potential energy 64. Properties of Liquids Strong intermolecular forces high enthalpy of vaporization low volatility low vapor pressure high boiling point 65. Critical Temperature Tc / Critical Pressure pc There is a maximum temperature at which a gas can be liquefied. T > Tc no condensation possible T = Tc condensation at pc T < Tc gas will condensate at p < pc 66. Viscosity How easily molecules move around each other Depends on attractive forces and structure (long hydrocarbons!) Decreases with increasing temperature 67. Surface Tension Minimizes surface, makes liquids “bead up” 22 68. Cohesive Forces Between molecules of a sample Causes surface tension and viscosity Affects vapor pressure, boiling point, Hvap. 69. Adhesive forces Between different substances, e.g., liquid and container Intermolecular forces: ion-dipole (esp. in aqueous solutions) dipole-dipole (between polar molecules) London dispersion forces (temporary, weak dipoles) hydrogen bonding (between molecules with H+), require electronegative atoms, usually N, O, F Notes: 23 SOLIDS 70. Crystalline Solids Highly regular arrangement Crystalline planes Show cleavage 71. Amorphous Solids Only short-range order, long-range disorder Example: glass, soot 72. Atomic Solids Atoms London dispersion forces Soft, low m.p., poor conductors Example: noble gases 73. Molecular Solids Molecules London dispersion forces, dipole-dipole, hydrogen bond Soft, low to moderate m.p., poor conductor Example: CH4, carbohydrates, CO2, H2O 74. Ionic Solids Ions (mono- or polyatomic) Electrostatic attraction (ionic bond) Hard and brittle, high m.p., poor conductors as solids Electrolyte strength depends on solubility. Example: NaCl, FeSO4, Ag2Cr2O7 75. Network Solids Atoms Network of covalent bonds Hard, high m.p., poor conductor Example: diamond C, quartz SiO2, β-boron nitride (BN)∞ 76. Metallic Solids Metallic bond (ions and delocalized valence e-) Soft to very hard, low to high m.p., good conductors Example: Cu, Fe, Na, Ag, Au, alloys 24 Notes: 25 PHASE DIAGRAMS 77. See Diagrams in Textbook, Notes and Review Books 78. State of matter is a function of temperature AND pressure 79. Parts of Phase Diagram Three curves connecting four points A, B, C, D solid liquid gas AB: vapor pressure curve AD: melting curve positive slope = solid is denser than liquid negative slope = solid is less dense than liquid (e.g., H2O) AC: vapor pressure curve of solid (describes sublimation) A: triple point (only point where all three phases coexist) B: critical point (above B liquid and gas are indistinguishable) Notes: 26 SOLUTIONS AND SOLUBILITY RULES 80. Solutions = Homogeneous Mixtures solute (smaller part, often undergoes phase change) solvent (bigger part) 81. Concentrations parts per million (ppm) = mass solute —————— x 106 mass solution parts per billion (ppb) = mass solute —————— x 109 mass solution = mass solute —————— x 100% mass solution 82. Weight percent Volume percent = volume solute ——————— x 100% volume solution Mole fraction X = moles solute ——————— total # of moles Molarity M = moles solute ——————— (titrations) liters of solution Molality M = moles solute ——————— (colligative properties) kg of solvent Solution Process Solute-solute interaction (e.g., ionic bond) H1 Solvent-solvent interaction (e.g., hydrogen bond) H2 Solute-solvent interaction (e.g., ion-dipole forces) H3 Hsolv. = H1 + H2 + H3 Hsolv. > 0 (exothermic, e.g., NaOH) Hsolv. < 0 (endothermic, e.g., NH4NO3) 27 83. Saturated Solution Dynamic equilibrium of solution with undissolved solvent 84. Factors Affecting Solubility Molecular structure (“likes dissolves like”) Pressure (gases only; Henry’s Law: cg = kpg) Temperature 85. Electrolytes Electrolytes are substances that conduct electricity in aqueous solutions or when molten. Electrolytes in water dissociate into hydrated ions. Strong electrolytes: strong acids/ bases (e.g., HNO3, H2SO4, NaOH, KOH) soluble salts (e.g., NaNO3, FeBr3) osmosis (osmotic pressure = MRT) Weak electrolytes: weak acids/ bases (e.g., CH3COOH, HNO2, HF, Ca(OH)2, Ba(OH)2) poorly soluble salts (e.g., CaCO3, SrSO4, Hg2Cl2) Nonelectrolytes: organic compounds (e.g., carbohydrates, proteins, alcohols) 86. Vapor Pressure of Solutions Volatile solvents (e.g., alcohol) increase vapor pressure, reduce b.p. Nonvolatile solvents (e.g., ionic compounds) decrease vapor pressure, increases b.p. 87. RAOULT’s Law pA = XApAo pA = partial vapor pressure of A above the solution XA = mole fraction of solvent pAo = vapor pressure of pure solvent 88. Colligative Properties Depend only on number of solvent molecules Concentration is given in molality m (except for osmosis) Examples: boiling point elevation (Tb = kb i m) freezing point depression(Tf = kf i m) osmosis (osmotic pressure = MRT) i: dissociation factor (“van’t HOFF factor”) kb: boiling-point elevation constant (solvent-specific) kf: freezing-point depression constant (solvent-specific) 28 89. Related Systems Solution homogeneous transparent particle size 10-9 m particles do not settle no Tyndall effect (scattering of light) no filtration possible Suspension heterogeneous particle size > 10-7 m particles settle can be filtrated Tyndall effect Colloid particle size = 10-9 10-7 m no filtration possible no settling Tyndall effect Emulsion: colloid of liquid in liquid, e.g. oil in water, milk 90. Solubility Rules: see Appendix B 91. Factors Driving a Chemical Reaction Formation of a precipitate (e.g., Ba2+ + SO42- BaSO4) Neutralization (formation of a nonelectrolyte (e.g., H+ + OH- H2O) Formation of a weak electrolyte (e.g., Na2S + H+ H2S + Na+) Redox reactions (e- transfer) (e.g., CrO42- Cr3+) Rupture or formation of covalent bonds (e.g., CH3-CHO CH3-COOH) Notes: 29 ACIDS AND BASES IN WATER 92. Acid Dissociation forms H3O+ and bigger adducts such as H5O2+ 93. Oxides (Anhydrides) in Water (oxidation state remains constant) Nonmetal oxides in water yield acids, e.g., P2O5 + H2O H3PO4 Metal oxides in water yield bases, e.g., CaO + H2O Ca(OH)2 94. BRØNSTED-LOWRY Theory Acids donate H+. Bases accept H+. 95. Conjugate Acids and Bases Strong acids (big Ka) yield weak conjugate bases (small Kb) and vice versa Ka (acid) x Kb (conjugate base) = Kw = 1.0 x 10-14 Notes: 30 REDOX REACTIONS 96. Balancing Redox Reactions Write reactants/products, omit spectator ions if present. Write oxidation number, determine # of transferred electrons. Write separate oxidation/reduction reactions. Balance charges with H+ (acidic solutions) or OH- (basic solutions). Balance mass for each half-reaction with water. Multiply half reactions equal # of e- for reduction and oxidation Add reduction and oxidation equations. Simplify and add spectator ions, if necessary. 97. Redox Reactants Common oxidizing and reducing agents: see Appendix A Notes: 31 KINETICS 98. Reaction Rate Reaction rate = (change of measurable property) / time Property: mostly concentration ([…]), also pH, color, temperature Rate decreases with time: rate (t = 0) > rate (t) Rate has to be referred to one reactant or product. Rates of appearance or disappearance of substances relate like stoichiometric coefficients. e.g.: 4 NH3 + 5 O2 4 NO + 6 H2O rate (NH3) = - 2/3 rate (H2O) 99. Collision Theory Proper atoms collide Proper orientation Sufficient energy ( EA) 100. Effect of Concentration reaction order rate law integrated rate law half-life [A]t = kt + [A]0 [A]0 t1/2 = ——— 2k rate = k [A] ln [A]t = kt + ln [A]0 ln 2 t1/2 = —— k rate = k[A]2 1 1 —— = —— + kt [A]t [A]0 1 t1/2 = ——— k[A]0 0 rate = k 1 2 32 101. Effect of Temperature 102. 103. EA Arrhenius equation: k = A x exp ( ——) RT EA Plot lnT vs. 1/T yields straight line with slope = —— R Reaction Mechanisms Series of small, elementary steps Intermediates = formed and consumed during reaction Intermediates do not show up in overall equation. ONLY the rate laws of elementary reactions can be deduced from stoichiometry. Multi-step reaction: stoichiometry of elementary reactions adds up rate law of slowest step = rate law of overall reaction Catalyst Reduces EA for forward and reverse reaction Decreases time needed to reach equilibrium; does not affect position of equilibrium Takes part in reaction, but is reformed at the end of the reaction Can be homo- or heterogeneous Examples: enzymes, NH3 synthesis (Fe) Notes: 33 EQUILIBRIUM 104. Dynamic Equilibrium Rate (forward reaction) = rate (reverse reaction) Examples: phase, solution, chemical equilibrium At equilibrium: all concentrations remain constant 105. Mass-Action Expression For aA + bB —> cC + dD Q = [C]c [D]d [A]a [B]b If concentrations are equilibrium concentrations: Q = K eq Only concentrations included: solutions and gases 106. Equilibrium Constants Kc and Kp Relationship between Kc (concentrations) and Kp (partial pressures) Kp = Kc x (RT)n n = (moles product) – (moles reactant) 107. Direction of Reaction Q > Keq reaction proceeds towards reactants Q = Keq reaction is at equilibrium Q < Keq reaction proceeds towards products 108. LeChatêlier’s Principle A system at equilibrium will counteract – if possible – any stress, i.e., change in temperature, pressure or concentrations. Notes: Increased pressure favors side with fewer gas molecules Increased temperature favors endothermic reaction Increased reactant concentration stronger formation of products Add inert gas at const. V partial pressures remain constant (= no change of equilibrium) Add inert gas at const. p V increases (= favors side with more gas molecules) Add catalyst no change Keq will only be affected by a temperature change 34 AQUEOUS EQUILIBRIA I (ACIDS AND BASES) (That’s the chapter of assumptions and negligences.) 109. Ion Product of Water Autodissociation of water: H2O H+ + OH Kw (25 oC) = [H+][OH-] = 1.0 x 10-14 pH = - log[H+] pOH = -log[OH-] pH + pOH =14 110. Indicators Weak acids HInd H+ + Ind- (pH-dependent equilibrium) Color of HInd and Ind- are different Indicator methyl orange methyl red litmus bromothymol blue phenolphthalein Color change red yellow red yellow red blue yellow blue colorless pink at pH 3.1 – 4.4 4.4 – 6.2 5.0 – 8.0 6.0 – 7.6 8.2 – 10.0 111. Assumptions Weak acids (less than 5% dissociation): [HA] = constant Polyprotic acids, e.g., H3PO4, H2SO4: If Ka(1) > 1000 x Ka (2), the pH is only determined by the first deprotonation step. 112. Acid Strength of Polyprotic Acids pKa(1) < pKa(2) < pKa (3) (acid strength decreases) 113. Hydrolysis Salt + water acid + base Strong acid + strong bases Weak acid + strong base Strong acid + weak base Weak acid + weak base neutral basic acidic (pH depends on relative strengths) 35 114. Prediction of Strength Binary acids/bases HX: metal hydrides (bases): base strength increases with period #, e.g., LiH < NaH < KH nonmetal hydrides (acids): acid strength increases with period #, e.g., HF < HCl < HBr < HI (Reason: increasing EN of element X, increasing bond length X—H) Oxyacids HnXOm same structure, different EN of central atom X: higher EN stronger acid (H2SeO3 < H2SO3) same central atom X, different number m of oxygen atoms: increasing number m stronger acid HClO < HClO2 < HClO3 < HClO4 HNO2 < HNO3 H2SO3 < H2SO4 H3PO3 < H3PO4 115. Lewis Theory of Acids and Bases Lewis acids: electron-pair acceptors (esp. small, highly charged metal ions, e.g., Al3+, Fe3+, Cr3+, also BCl3 and analogous compounds) Lewis bases: electron-pair donors (e.g., NH3, H2O, CN-, SCN- etc.) 116. Common-Ion Effect Addition of an already present ion to an equilibrium shifts it according to LeChatêlier’s Principle. Change of solubility 117. Buffer Solutions Aqueous solutions of weak acid and its conjugate base Resists pH change upon addition of little H+/OH Henderson-Hasselbalch Equation: [X-] pH = pKA + log ——— [HX] or (for base B and conjugated acid BH+) [BH+] pOH = pKB + log ——— [B] 36 118. Biological function: keep pH in blood and cells constant Examples: HCO3-/CO32-, H2PO4-/HPO42- Titration Curves Plot of pH as a function of added titrant Strong acid + strong base (e.g., HCl + NaOH): (indicator range around pH = 7) Titration curve strong acid/strong base (50.00 ml 0.100 M HCl with 0.100 M NaOH) 14.00 12.00 245.000 10.00 195.000 8.00 pH 145.000 6.00 95.000 4.00 45.000 2.00 0.00 -5.000 0.00 10.00 20.00 30.00 40.00 50.00 60.00 70.00 80.00 90.00 100.0 0 0.100 M NaOH (ml) 37 Weak acid + strong base or weak base/strong acid: Theindicator range has to be around pH = pKB / pKA of conjugate base/acid Titration curve weak acid/strong base (50.00 ml 0.100 M HC2H3O2 with 0.100 M NaOH) 14.00 45.000 pH 12.00 10.00 35.000 8.00 25.000 6.00 15.000 4.00 5.000 2.00 0.00 -5.000 0.00 10.00 20.00 30.00 40.00 50.00 60.00 70.00 80.00 90.00 100.0 0 0.100 M NaOH (ml) Titration curve weak base/strong acid (50.00 ml 0.100 M NH3 with 0.100 M HCl) 14.00 98.000 12.00 10.00 78.000 pH 8.00 58.000 6.00 38.000 4.00 18.000 2.00 0.00 0.00 -2.000 10.00 20.00 30.00 40.00 50.00 60.00 70.00 80.00 90.00 100.0 0 0.100 M HCl (ml) 38 Titration and pH at equivalence point: “Equivalence point” means equal moles of acid and base have been added. It does not necessarily mean a pH of 7. Acid strong strong weak weak Base strong weak strong weak pH at equivalence point 7 <7 >7 7 (depends on relative strength) Deviations from pH = 7 at equivalence point are due to hydrolysis of salt Polyprotic acid, e.g. H3PO4 (loss of all H+ different equiv. points) H3PO4 H2PO4- HPO42- PO43pH (1), pH (2), and pH (3) correspond to the eq. points #1, 2, 3. Titration curve multiprotic acid/strong base (50.00 ml 0.100 M H3PO4 with 0.100 M NaOH) 14.00 4.000 12.00 3.500 3.000 10.00 2.500 pH 8.00 2.000 6.00 1.500 4.00 1.000 2.00 0.00 0.00 0.500 25.00 50.00 75.00 100.00 125.00 0.100 M NaOH (ml) 150.00 175.00 0.000 200.00 39 At the halfway equivalence points: [H3PO4] = [H2PO4-] pH = pKA (H3PO4) [H2PO4-] = [HPO42-] pH = pKA (H2PO4-) [HPO42-] = [PO43-] pH = pKA (HPO42-) Note that the equivalence point “jumps” become smaller. This can be explained by the decreasing acid strength H3PO4 > H2PO4- > HPO42-. The third equivalence point becomes undetectable due to the interference by the autodissociation of water. Notes: 40 AQUEOUS EQUILIBRIA II (SOLUTIONS) (We’re still assuming and neglecting.) 119. Equilibria Acid-base: homogeneous equilibrium Solution-precipitate: heterogeneous equilibrium 120. Dissociation of Ionic Compounds and Solubility Product Ionic compound AaBb a A+ + b B Ksp = [A+]a [B-]b Calculation of (molecular) solubility based on relative concentrations of A+ vs. B If a common ion is added, the contribution via dissociation is negligible. Ion product Q = (product of actual ion concentrations) Solubility product Ksp = (product of equilibrium ion concentrations) Q > Ksp precipitation occurs until Q = Ksp Q = Ksp saturated solution Q < Ksp unsaturated solution, extra solute dissolves 121. Solubility and pH Solubility of metal sulfides (CuS, ZnS, FeS, MnS) depends on pH. Notes: 7.4 x 10-22 In a saturated H2S solution ( 0.1 M): [S2-] = —————— [H+]2 41 THERMODYNAMICS 122. Energy Ability to do work and/or Transfer heat 123. Forms of Energy Kinetic energy Ek = ½ mv2 Potential energy Ep = mgh (gravitational potential) q1q2 Potential Energy Ep = k —— (electrostatic potential) d2 (q1, q2 = charges; d = distance between charges) 124. 125. Unit Unit: 1 J(oule) = 1 kg m2 s-2 1 J = 4.184 cal The First Law of Thermodynamics (Law of Conservation of Energy) The sum of all kinetic and potential energies is constant. E = Efinal - Einitial E < 0 system loses energy E > 0 system gains energy Energy change E = q + w (q = heat; w = work) 126. State Functions The value of a state function depends only on the present state, not on the initial conditions. Examples: internal energy, volume, pressure, temperature, density, mass, refractive index, amount of matter 127. Path Functions Depend on how the present state was reached Examples: work, heat 128. p-V work and enthalpy Volume work against constant outside pressure w = – pV constant volume qV = E (heat change = change of internal energy) constant pressure qP = H (heat change = change of enthalpy) reaction enthalpy H = Hproducts – Hreactants (extensive quantity!) 42 129. Hess’ Law If a reaction is carried out in a series of steps, the individual enthalpies add up to the total enthalpy. For n steps: Htotal = H1 + H2 + … + Hn Hess’ Law allows the indirect determination of reaction enthalpies 130. Heat of Formation Standard heat of formation is the enthalpy change during the formation of one mole of substance from its elements in their standard state. Hf0 (elements) = 0 (by definition) Reaction enthalpy = (sum of heats of formation of products) – (sum of heats of formation of reactants) H0 = n Hf0 (products) – m Hf0 (reactants) (n, m: stoichiometric coefficients of products and reactants) 131. Calorimetry Transferred heat = heat capacity x temperature change q = C x T (C = heat capacity) Heat capacity = (mass) x (specific heat) C=mxs (m = mass; s = specific heat) 132. The Second Law of Thermodynamics The entropy of the universe increases constantly. An entropy decrease is always accompanied by a bigger entropy increase somewhere else. Heat never flows spontaneously from a cold object to a warmer one. 133. Entropy S Synonyms: randomness, chaos, disorder, state of higher probability, degree of dispersion of heat Unit: J(oule) / K(elvin) S = Sfinal – Sinitial Suniverse = Ssystem + Ssurroundings > 0 Ssolid < Sliquid < Ssolution < Sgas Entropy increases with temperature, phase changes and number of particles. 134. Driving Forces of Reactions Enthalpy towards minimum Entropy towards maximum 43 135. The Third Law of Thermodynamics Absolute zero cannot be reached in a finite number of steps. The entropy S of a pure, crystalline substance at T = 0 K is 0. 136. Standard Entropies S0 Are tabulated and allow calculation of Standard Entropy of Reaction S0 Standard Entropy of Reaction S0 = n S0 (p) – m S0 (r) n,m: stoichiometric coefficients of products and reactants) 137. The Free-Energy Function (Gibbs Function) G = H – TS H = enthalpy T = absolute temperature S = entropy 138. G = H – TS G < 0 spontaneous reaction G > 0 nonspontaneous reaction At low temperatures: –TS < H Enthalpy (H) determines spontaneity. At high temperatures: -TS > H Entropy (S) determines spontaneity. Gf0 = n Gf0 (products) - m Gf0 (reactants) (n,m: stoichiometric coefficients of products and reactants) Spontaneity and Free Energy “Some things happen spontaneously, some things don't.” — P.W. Atkins, Physical Chemistry 139. Sign of G determines direction in which system moves G < 0 (forward) reaction is spontaneous G = 0 reaction has reached equilibrium G > 0 reaction is not spontaneous (reverse reaction is!) Definition of Standard State Standard state pure solid pure liquid gas at p = 1 atm c=1M at 25 0C, Gf0 = 0 State of matter solid liquid gas solutions elements 44 140. Free Energy, Temperature and Concentrations G = Go + RT lnQ = Go + 2.303 RT logQ Equilibrium: Q = Keq; G = 0 (reaction stops), so Go = -RTlnKeq Sign of G0 allows semiquantitative statement about Keq Go is negative Keq > 1 o G is positive Keq < 1 Go is zero Keq = 1 Notes: 45 ELECTROCHEMISTRY 141. Electrochemical Cell Converts chemical energy into electrical energy Two half-cells connected with wire (= electron flow) and salt bridge/porous cup (= ion flow to ensure charge balance) Reduction occurs at the cathode, oxidation at the anode. (anode oxidation, reduction cathode) 142. Cell EMF (electromotive force) 1 J(oule) Unit: 1 V(olt) = —————— 1 C(oulomb) Standard EMF E0 at standard cond.: p = 1 atm, c = 1 M, T = 298 K Listed: standard reduction potentials E0red Connect two half-reactions: bigger E0red will be the reduction smaller E0red will be the oxidation (reverse sign!!) multiply (if necessary) to eliminate e do not multiply the reduction potentials add equations and (unmultiplied) potentials E0cell = E0red + E0ox Reference half-cell: S(tandard) H(ydrogen) E(lectrode) Hydrogen bubbling over a Pt-electrode in acidic solution p (H2) = 1 atm; [H+] = 1 M pH = 0 143. Reducing and Oxidizing Agents Oxidizing agents are reduced. Reducing agents are oxidized. Strong oxidizing agents have big E0red (F2, Au3+, Pb4+, Cl2). Strong reducing agents have small E0red (Li, K, Na, Ca). 144. EMF and Spontaneity E0 > 0 spontaneous (G < 0) 0 E < 0 not spontaneous (G > 0) 46 145. EMF and Free Energy G = nFE (n = number of electrons transferred; F = Faraday constant; E = cell potential) 0 G = n F E0 146. EMF and Equilibrium Constant/Mass-Action Expression 0.0591 – ———— log Q n NERNST Equation: If electrochemical cell has used up one or more reactants: G = 0 E= E = 0 and E0 E= E0 0.0591 – ———— log Keq n 147. Commercial Voltaic Cells Lead battery: Pb + PbO2 + 4 H+ + 2 SO42- 2 PbSO4 + 2 H2O Dry cell: Zn + NH4Cl + MnO2 ZnCl2 + Mn2O3 + 2 NH3 + H2O 148. Electrolysis Reversal of electrochemical cell Converts electrical energy into chemical energy Same process (an ox, red cat) Electrolysis of aqueous solutions: often water is reduced 2 H2O + 2 e- H2 + 2 OH- 149. Electrodes Active electrodes = electrode material participates in electrolysis Application: electroplating, electroraffination Examples: Cu, Zn, Co, Ni Inert electrodes = electrode material does not participate Examples: C (graphite), Pt 150. Quantitative Analysis Amount of charge of 1 mol e- = 1 F = 96,500 C 1 C(oulomb) = 1 A(mpère) x 1 s(econd) 47 151. Electrical Work wmax = n F E Work done by system = voltaic cell Work done to system = electrolytical cell Electrical work: 1 W(att) = 1 J/s 1 kWh = 3.6 x 106 J 152. Corrosion Common form of oxidation of metals Often slowed down by formation of inert oxides, e.g. Al, Cr, Mg Sped up by salts (electrolytes!) pH dependent, e.g., Fe does not corrode if pH > 9 Corrosion protection coating with metal that is inert to H+ (Sn) coating with metal that corrodes first (Zn) “sacrificial anode” (Mg) Notes: 48 ORGANIC CHEMISTRY 153. Central Feature Carbon-carbon bond Second common element: hydrogen 154. Nomenclature See Appendix D 155. Homologous Series Compounds that differ by number of methylene units –CH2– Examples: alkanes 156. Isomerism Same molecular formula Different structure Structural isomerism CH3 CH3–C–CH3 CH3 CH3CH2CH2CH2CH3 vs. n-pentane 2,2-dimethylpropane Geometrical isomerism CH3CH2CH2 \ CH2CH2CH3 CH3CH2CH2 H / \ / C==C vs. C==C / \ / \ H H H CH2CH2CH3 cis Substituents on trans same side of double bond: different side of double bond: cis trans 49 157. Aromatic Hydrocarbons Mostly 6-membered rings with delocalized -electrons Stabilization due to resonance Examples: benzene C6H6 look up formula notice sextet of electrons toluene C6H5CH3 C6H5 = phenyl 158. Reactions of Hydrocarbons See Appendix A 159. Derivatives Functional groups (often polar) affect properties, e.g. solubility, b.p. (R = alkyl group or other organic fragment) ROH alcohols CH3CH3 + <O> CH3CH2OH 1 2 R OR ethers R1OH + R2OH R1OR2 + H2O O // RC \ H aldehydes R—CHO (R = H formaldehyde CH3 acetaldehyde) O R1CR2 ketones R1COR2 (R1 = R2 = CH3 acetone (R1 = R2 = C2H5 diethylketone) RCOOH carboxylic acids R1COOR2 ester R1COOH + R2OH R1COOR2 + H2O RNH2 amines (organic ammonia derivatives) Notes: (R = H formic CH3 acetic C2H5 propanoic) 50 NUCLEAR CHEMISTRY 160. Radioactivity Spontaneous decay of nuclei (radioisotopes) 161. Decay modes Alpha () decay = loss of 42He (charge: +2) 222 Rn 218 Po + 4 He 86 84 2 Beta () decay = loss of electrons 0-1e (does not come from shell, but from 10n 11p + 0-1e) 131 I 131 Xe + 0 e 53 54 -1 Gamma () decay = high-energy photon (very short wavelength) (is usually not shown) Positron emission = 11p 10n + 01e (= antiparticle to electron 0-1e) 11 C 11 B + 0 e 6 5 1 Electron capture from inner shell = conversion of 81 Rb + 0 e 81 Kr 37 -1 36 1 p 1 + 0-1e 10n 162. Stability of Nuclei All nuclei with more than 83 protons are unstable. Magic numbers (2, 8, 20, 50, 82, 126) of protons or neutrons are more stable than adjacent nuclei. Nuclei with even numbers of protons or neutrons are more stable than those with uneven numbers. Nuclei with equal numbers of protons and neutrons are more stable than those with different numbers. 163. Radioactive Series Certain nuclei keep disintegrating until a proton number of less than 84 is reached. Example: 23892U (…) (…) 20882Pb 164. Transmutation Formation of new nuclei Target nuclei are bombarded with highly accelerated p + or n 51 165. Half-life Time needed for a radioisotope to decay to ½ of its mass (first-order process) Application: radiocarbon dating (= determination of ages of minerals and organic matter, e.g., wood) Useful formulas: ln [A]t = ln [A]0 – kt 1 [A]0 k = — · ln —— t [A]t ln 2 t1/2 = —— k t1/2 [A]0 t = —— · ln —— ln 2 [A]t 166. Mass-energy Conversions Masses of nuclei are lower than the masses of individual nucleons. Difference: mass converted to energy E = mc2 (= “mass defect”) Mass defect is converted into binding energy. Notes: 52 BEFORE THE TEST Form study groups, but with no more than three people at a time. Get a review book. Then use it. And use it. And use it. Do as many sample AP exam questions as you possibly can. Bother your teacher with questions. After all, it’s his job. Stop studying the day before. Marathon runners don’t run the day before a race. You won’t forget anything if you don’t study during the last 24 hours. Work out. Go for a long walk or see a really bad movie. Have a nice dinner. Distract yourself. Get plenty of sleep and have breakfast. If you are tired or hungry, you won’t do well. DURING THE TEST Read the questions carefully. Select essential from superfluous information. Scan the questions before you answer them. Rate their difficulty from easy to medium to hard. Do the easy ones first. Elimininate wrong answers in the Multiple Choice section. If you don’t know the right answer, guess as soon as you have eliminated two answers as wrong. Do not ramble. The graders don’t want to know what else you know – keep your answers brief. Stick to basic concepts. Don’t waste your time. If a question doesn’t ask you to explain something, don’t. Keep an eye on the time. Do not expect to be able to answer every question completely. Don’t fret over the really hard ones you couldn’t answer. 53 APPENDIX A: EQUATIONS IN THE AP CHEMISTRY EXAMINATIONS As of May 2007, the equations in this section need to be balanced. Balancing can be done either by trial-and-error as in the past, or by a more sophisticated method that you will need to apply to redox equations. Omit spectator ions, i.e., ions that stay in solution during the reaction and form no precipitates (see worksheet “Solubility Rules”). Strong acids that are mostly dissociated (HCl, HNO3, H2SO4) should be written in their ionized form, e.g., HCl = H+ + Cl-. Weak acids that are mostly undissociated (HC2H3O2 and similar organic acids) must be written as undissociated molecules. Each question consists of two parts: One equation that must be correctly balanced. A question about that reaction. This question may involve numbers of electrons transferred, oxidation numbers, observable changes during reactions, formulas/names of precipitates, etc. Since this change is a very recent one, my suggestions as how to approach this new format can be only tentative. I will therefore include my previous strategies. 54 1) ACID-BASE REACTIONS: There are no changes of oxidation states in acid-base reactions. a) acid and base solutions: H2SO4 (aq) + Ca(OH)2 (aq) CaSO4 (s) + 2 H2O should be written as: H+ + SO42- + Ca2+ + 2 OH- CaSO4 + 2 H2O Some acids are unstable: HCO3- / CO32- + H+ (H2CO3)* H2O + CO2 (Don’t include H2CO3!) HSO3- / SO32- + H+ (H2SO3)* H2O + SO2 (Don’t include H2SO3!) A star superscript after a compound in parentheses (…)* is often used to indicate instable compounds or intermediates. b) solution and anhydride: N2O5 + 2 OH- 2 NO3- + H2O or 6 H+ + Cr2O3 2 Cr3+ + 3 H2O c) anhydrides: P2O3 + 3 MgO Mg3(PO3)2 d) amphoteric hydroxides: Some hydroxides (Sn4+, Zn2+, Al3+, Cr3+) dissolve in acids as well as excess base. Al(OH)3 + 3 H+ Al3+ + 3 H2O Al(OH)3 + 3 OH- (xs.) Al(OH)63- (also: AlO33- + 3 H2O e) Lewis acid-base reactions: Compounds with lone electron pairs react with electron-deficient compounds. BF3 + NH3 BF3 • NH3 (F3B—NH3) f) hydrolysis of nonmetal halides or oxyhalides: PCl5 + 4 H2O 8 H+ + 5 Cl- + PO43- Yields two acids. SO2Cl2 + 2 H2O 4 H+ + SO42- + 2 Cl- g) hydrolysis of certain transition metal halides: UF6 + 2 H2O UO2F2 + 4 H+ + 4 F- Yields oxyhalides and acid. 55 2) DOUBLE REPLACEMENT: compound+ compound compound + compound AB + CD AD + BC 2 AgNO3(aq) + BaCl2(aq) 2 AgCl(s) + Ba(NO3)2(aq) (Ag+ + Cl- AgCl) Driving force: Formation of precipitate. 3) SINGLE REPLACEMENT: element + compound compound + element a) cationic single replacement: A + BC AC + B (A, B = metals or hydrogen) H2 + CuO H2O + Cu 2 Al + Cr2O3 Al2O3 + 2 Cr Mg + 2 H+ Mg2+ + H2 b) anionic single replacement: A + BC AB + C (A, C = nonmetals) F2 + 2 I- 2 F- + I2 Cl2 + 2 Br- 2 Cl- + Br2 Driving force: Check Table of Standard Reduction Potentials, often hydrogen is formed. 56 4) REDOX REACTIONS: a) oxidizing agents: oxygen: halogens: Elements in (unusually) high oxidation states. O2 (0), peroxides O22- (-1) O2- X2 (0) e.g. Cl2 XO- (+1) e.g. ClOXO2- (+3) e.g. ClO2XO3 (+5) e.g. ClO3XO4- (+7) e.g. ClO4(oxidation strength increases with higher oxidation number) chromates and dichromates: CrO42-, Cr2O72- Cr3+ manganese dioxide and (per)manganates: MnO2 Mn2+ MnO42-, MnO4- MnO2 (basic) MnO42-, MnO4- Mn2+ (acidic) oxyacids (esp. when concentrated), common: HNO3, H2SO4 and HClO4 HNO3 NO H2SO4 SO2 HClO4 Cl- b) reducing agents: alkali and alkaline earth metals Na, K, Ca Na+, K+, Ca2+ aluminum Al Al3+ transition metals in low oxidation states Fe2+ Fe3+ hydrogen H2 H2O carbon monoxide CO CO2 sulfite and thiosulfate ion S2O32- / SO32- SO42- Cu+ Cu2+ 57 c) miscellaneous: Halogens often oxidize and reduce themselves in water (= disproportionation), e.g. Cl2 + H2O H+ + Cl- + ClOHalogens in low AND high oxidation states reach one oxidation state (= comproportionation), e.g. I- + IO3- + H+ I2 + H2O 5) (THERMAL) DECOMPOSITIONS: a) Carbonates (esp. alkaline earth): CaCO3 CaO + CO2 b) Hydrogencarbonates: 2 NaHCO3 Na2O + H2O + CO2 c) Azides (esp. Na, Ag): decompose explosively and form nitrides and nitrogen, e.g. 3 NaN3 Na3N + 4 N2 Driving force: Entropy gain (gas!) and enthalpy gain by formation of NN bond 6) LIGAND EXCHANGE REACTIONS: Molecules or anions with lone pairs can have a higher affinity to hydrated metal ions than water and subsequently replace one or all water molecules. (examples: NH3, CN-, CO) Cu(H2O)42+ + 4 NH3 Cu(NH3)4+ + 4 H2O Fe(H2O)63+ + 6 CN- Fe(CN)63- + 6 H2O AgCl + 2 CN- Ag(CN)2- + ClAmmine ligands (NH3) can be removed by acids: Ag(NH3)2+ + 2 H+ Ag+ + 2 NH4+ 58 7) ORGANIC REACTIONS: a) Addition to multiple bonds: Alkynes and alkenes add halogens, acids and water. HCCH + Cl2 HClC=CClH H2C=CH2 + HBr H3C—CH2Br H2C=CH—CH3 + H2O H3C—CHOH—CH3 b) Substitution of single bonds: Alkanes and aromatic hydrocarbons substitute halogens for one or more hydrogen. CH4 + 4 Cl2 CCl4 + 4 HCl C6H6 + Br2 C6H5Br + HBr c) Ether formation: alcohol + alcohol ether + water (catalysis: H+) C2H5OH + C2H5OH C2H5—O—C2H5 + H2O d) Ester formation: alcohol + acid ester + water (catalysis: H+) CH3CH2COOH + CH3CH2CH2CH2OH CH3CH2—COO—CH2CH2CH2CH3 + H2O 59 EXAMPLES FROM AP EXAMS: 2007 (example): (i) A strip of magnesium is added to a solution of silver (I) nitrate. Mg + 2 Ag+ Mg2+ + 2 Ag (ii) Which substance is oxidized in the reaction? Mg is oxidized. 2007 (a): (i) A solution of sodium hydroxide is added to a solution of lead (II) nitrate. Pb2+ + 2 OH- Pb(OH)2 (ii) If 1.0 L volumes of 1.0 M solutions of sodium hydroxide and lead (II) nhitrate are mixed together, how many moles of product(s) will be produced. Assume the reaction goes to completion. 0.5 moles of lead (II) hydroxide Pb(OH)2 will be formed. (Since OH- is used up at twice the rate, it is the limiting reactant.) 2007 (b): (i) Excess nitric acid is added to solid sodium carbonate. 2 H+ + CaCO3 Ca2+ + CO2 + H2O (ii) Briefly explain why statues made of marble (calcium carbonate) displayed outdoors in urban areas are deteriorating. The air in urban areas contains nonmetal oxides such as SO2, SO3, and NOx. These react as acid anhydrides and create acid rain which has a deteriorating effect on calcium carbonate. 2007 (c) (i) A solution containing silver (I) ion (an oxidizing agent) is mixed with a solution containing iron (II) ion (a reducing agent). Ag+ + Fe2+ Ag + Fe3+ (ii) If the contents of the reaction mixture described above are filtered, what substance(s), if any, would remain on the filter paper. Silver metal (Ag). 60 1974/2 A sample of pure 2-butene is treated with hydrogen bromide gas. Answer: 2-butene is an alkene and undergoes addition to the double bond. CH3—CH=CH—CH3 + HBr CH3—CH2—CHBr—CH3 1988/1 A solution of potassium iodide is added to an acidified solution of potassium dichromate. Answer: a) potassium K+ is a spectator ion omit! b) Cr2O72- is an oxidizing agent iodide is oxidized, dichromate reduced I- + H+ + Cr2O72- I2 + Cr3+ + H2O 1989/2 Solutions of silver nitrate and lithium bromide are mixed. Answer: a) lithium Li+ and nitrate NO3- are spectator ions omit! b) silver bromide is insoluble. Ag+ + Br- AgBr 1990/5 Hydrogen sulfide gas is bubbled through a solution of potassium hydroxide. Answer: a) potassium K+ is a spectator ion omit! b) H2S is an acid, OH- a base H2S + OH- S2- + H2O 1993/4 Excess chlorine gas is passed over hot iron filings. Answer: a) chlorine acts as an oxidizing agent, iron as a reducing agent b) synthesis reaction: metal + nonmetal salt Fe + Cl2 FeCl3 61 Please use “Equations in the AP Examinations” and Appendices D and E in your textbook. For each of the following reactions, (i) write a balanced equation for the reaction (ii) answer the question about the reaction. In part (i), coefficients should be in terms of lowest whole numbers. Assume that solutions are aqueous unless otherwise indicated. Represent substances in solutions as ions if they are extensively ionized. Omit formulas for ions or molecules that are unchanged by the reaction. 1) 2) 3) 4) 5) 6) 7) 8) 9) 10) 11) 12) 13) 14) 15) —I— (i) Solid lithium hydride is added to water. (ii) How does the pH of the reaction mixture change? (i) Excess ammonia is added to a solution of zinc sulfate. (ii) What particles act as LEWIS acids and bases? (i) Sodium acetate is acidified with dilute hydrochloric acid. (ii) How does the pH of the reaction mixture change? (i) A solution of potassium dichromate is added to an acidified solution of sodium iodide. (ii) What is total number of electrons transferred? (i) Solid barium carbonate is added to dilute nitric acid. (ii) What is a simple analytic test for one of the products? — II — (i) Solid aluminum sulfide is added to water. (ii) Why is the gaseous product toxic? (i) Silver metal is added to 6M nitric acid. (ii) What acts as an oxidizing agent? (i) A solution of tin(II)nitrate is dropped onto a piece of zinc metal. (ii) What observation can be made about the piece of zinc? (i) Solid sodium dichromate is added to an acidified solution of iron(II) sulfate. (ii) What is the total number of electrons transferred? (i) Liquid titanium tetrachloride is given into excess water. (ii) How does the pH of the reaction mixture change? — III — (i) Solid ammonium chloride is added to a concentrated solution of sodium hydroxide. (ii) What ion(s), if any, act as spectator ions? (i) Solid sodium hydrogencarbonate is added to concentrated sodium hydroxide. (ii) Sodium hydrogencarbonate reacts with hydrochloric acid as well. Name this behavior. (i) Solid calcium sulfite is acidified. (ii) Name one unstable reaction intermediate. (i) Liquid acetic acid and methanol are combined and warmed. (ii) What specific organic reaction type occurs here? (i) Aqueous solutions of sodium iodate and sodium iodide are combined, then acidified. (ii) What specific type of redox reaction occurs here? 62 16) 17) 18) 19) 20) 21) 22) 23) 24) 25) 26) 27) 28) 29) 30) 31) 32) 33) 34) 35) — IV — (i) Solid calcium phosphate is given to excess hydrobromic acid. (ii) What phosphorus-containing product would be formed if hydrobromic acid and calcium phosphate reacted in a mole ratio of 1:1? (i) A piece of magnesium is dropped into hydrochloric acid. (ii) How many electrons are transferred in this reaction? (i) A stream of carbon dioxide gas is bubbled through a suspension of barium carbonate. (ii) What observation can be made regarding the appearance of the reaction mixture? (i) Potassium permanganate solution is added to concentrated hydrochloric acid. (ii) How does the color of the reaction mixture change? (i) Ammonia gas is bubbled through a suspension of zinc hydroxide. (ii) What observation can be made regarding the appearance of the reaction mixture? —V— (i) Hydrogen sulfide gas reacts with moist lead(II)chloride. (ii) What color change could be observed? (i) Moist sodium hydroxide reacts with carbon dioxide gas. (ii) How and why does the pH of the reaction mixture change? (i) An aqueous solution of ammonia is added to solid silver iodide. (ii) Name this type of reaction. (i) Rubidium permanganate is added to hydroiodic acid. (ii) How many electrons are transferred in this reaction? (i) Liquid phosporus trichloride is poured into an excess of sodium hydroxide. (iii) What are the oxidation states of phosphorus in reactant(s) and product(s)? — VI — (i) 30% hydrogen peroxide is heated. (ii) What are the oxidation states of oxygen in reactant(s) and product(s)? (i) Liquid mercury is added to concentrated nitric acid. (ii) What substance acts as a reducing agent? (i) Solid iron(III)sulfate is added to a sodium iodide solution. (ii) What substance could be used to detect one of the products in small concentrations? (i) Hydrogen sulfide is bubbled through a solution of silver nitrate. (ii) Name this specific type of reaction. (i) Concentrated (15 M) ammonia solution is added in excess to a solution of copper(II)nitrate. (ii) Describe any observable color changes. — VII — (i) Magnesium metal is added to dilute nitric acid, giving as one of the products a compound in which the oxidation number of nitrogen is -3. (ii) Name said product. (i) Excess water is added to solid calcium hydride. (ii) Why must calcium hydride be stored in a well-sealed container? (i) Silver acetate solution is dropped into a solution of sodium phosphate. (ii) If the contents of the reaction mixture would be filtered, what substance(s) would remain on the filter paper? (i) Solid sodium cyanide is added to water. (ii) Why is sodium cyanide toxic? (i) Solid potassium hydride is added to anhydrous ethyl alcohol. (ii) Name the solid reaction product. 63 36) 37) 38) 39) 40) 41) 42) 43) 44) 45) 46) 47) 48) 49) 50) 51) 52) 53) 54) 55) — VIII — (i) Lithium metal is burned in air. (ii) Given sufficiently high temperatures, what byproduct could be formed? (i) Aluminum metal is added to a solution of copper(II)chloride. (ii) What, if any, color change would be observable? (i) Manganese(II)nitrate is mixed with sodium hydroxide solution. (ii) Name any spectator ions. (i) Equal volumes of dilute equimolar solutions of sodium carbonate and hydrochloric acid are mixed. (ii) What ratio of sodium carbonate to hydrochloric acid would lead to the formation of a gas? (i) Solid sodium carbide is added to an excess of water. (ii) How does the pH of the reaction mixture change? — IX — (i) An excess of sodium hydroxide solution is added to a solution of aluminum chloride. (ii) Describe any visible changes during the reaction. (i) Solid ammonium carbonate is heated. (ii) Which of the products has the lowest boiling point? (i) Phosphorus pentachloride is added to an excess of potassium hydroxide solution. (ii) Name the type of reaction that the phosphorus pentachloride undergoes. (i) Potassium chlorate is heated in the presence of manganese dioxide. (ii) What function does the manganese dioxide have? (i) Chlorine gas is bubbled into a solution of sodium hydroxide. (ii) What oxidation state(s) change(s)? —X— (i) Magnesium carbonate is heated strongly. (ii) Which of the product is an acid anhydride? (i) Methanol and propanoic acid are mixed. (ii) What general organic reaction takes place here? (i) Barium acetate is added to dilute sulfuric acid. (ii) What, if any, spectator ions occur here? (i) Ammonia gas is bubbled through a suspension of silver chloride. (ii) Describe any visible changes of the reaction mixture. (i) Hot iron(III)oxide is reacted with carbon monoxide. (ii) How many electrons are transferred during the reaction? — XI — (i) Silane (SiH4) is burned in air. (ii) What type of solid is the product containing silicon? (i) Cesium metal is heated with sulfur. (ii) Name this type of reaction. (i) Sodium phosphate is dissolved in water. (ii) How does the pH of the reaction mixture change? (i) Ethyl ethanoate is boiled with sodium hydroxide. (ii) What type of organic compound is ethyl ethanoate? (i) Solutions of sodium phosphate and calcium chloride are combined. (ii) Name any spectator ion(s). 64 56) 57) 58) 59) 60) 61) 62) 63) 64) 65) 66) 67) 68) 69) 70) 71) 72) 73) 74) 75) — XII — (i) A suspension of aluminum hydroxide is added to hydroiodic acid. (ii) How would the reaction be different if hydrochloric acid were used? (i) Hydrogen peroxide is added to an acidified solution of sodium iodide. (ii) How many electrons are transferred during this reaction? (i) Hydrogen peroxide is added to an acidified solution of potassium dichromate. (ii) What are the oxidation states of chromium in this reaction? (i) Calcium is added to dilute hydrochloric acid. (ii) What specific inorganic reaction type occurs here? (i) Aluminum nitrate is dissolved in water. (ii) Which substance acts as LEWIS acid here? — XIII — (i) Phosphorus tribromide is added to water. (ii) Name the product containing phosporus. (i) Solutions of sulfuric acid and lithium hydroxide are mixed. (ii) What specific analytical method would use this type of reaction? (i) Nitric acid is added to a solution of tetraamminecadmium(II). (ii) What spectator ion(s), if any, occur in this reaction? (i) 2-butene reacts with hydrobromic acid. (ii) What functional group of 2-butene is active in this reaction? (i) Hydrogensulfide is bubbled through a solution of lead(II)nitrate. (ii) What color change could be observed? — XIV — (i) Solutions of iron(III)nitrate and sodium thiocyanate are combined. (ii) What color change could be observed? (i) Ammonia gas is bubbled through a suspension of zinc hydroxide. (ii) How many electrons are transferred during this reaction? (i) Sulfur dioxide gas is bubbled through an acidified solution of potassium permanganate. (ii) How many electrons are transferred during this reaction? (i) Aluminum is added to excess sodium hydroxide solution. (ii) Name the ionic product of this reaction. (i) A solution of tin(II)nitrate is added to an acidified solution of potassium dichromate. (ii) How many electrons are transferred during this reaction? — XV — (i) Equal volumes of equimolar sodium hydroxide and sodium dihydrogenphosphate are mixed. (ii) What phosphorus-containing product would be formed if the ratio of hydroxide to dihydrogenphosphate were 2:1? (i) Excess sodium hydroxide is added to a solution of calcium hydrogencarbonate. (ii) Name any spectator ions. (i) Acetylene (ethyne) reacts with an an excess of chlorine. (ii) What product would be formed if the reactants were used in a ratio of 1:1? (i) Solid silver nitrate is added to a solution of sodium chromate. (ii) Name any spectator ions. (i) An excess of solid potassium hydroxide is added to a suspension of zinc hydroxide. (ii) What changes in the reaction mixture can be observed? 65 76) 77) 78) 79) 80) 81) 82) 83) 84) 85) 86) 87) 88) 89) 90) 91) 92) 93) 94) 95) — XVI — (i) Benzene reacts with bromine. (ii) Why does benzene not react as if it had double bonds? (i) A direct electric current is passed through a dilute solution of sulfuric acid. (ii) What gas is formed at the cathode? (i) Ammonia and oxygen are heated in the presence of a catalyst. (ii) How does the oxidation number for nitrogen change? (i) Liquid bromine is added to a solution of sodium iodide. (ii) What specific type of inorganic reaction occurs here? (i) Dilute hydrochloric acid is added to solid calcium oxide. (ii) Name the spectator ion. — XVII — (i) Sodium hydride is added to water. (ii) Compare the relative reactivity of potassium hydride with sodium hydride. (i) Ethene reacts with water in the presence of a catalyst. (ii) Name the product. (i) Solutions of barium hydroxide and iron(III)sulfate are combined. (ii) If equal moles of reactants were combined, which would be the limiting reactant? (i) Ammonia gas is bubbled through a solution of copper(II)nitrate. (ii) What color change could be observed? (i) Lead is added to a dilute solution of sulfuric acid. (ii) How many electrons are transferred during this reaction? — XVIII — (i) Lead is added to a hot solution of concentrated sulfuric acid. (ii) What acts as the oxidizing agent? (i) Calcium oxide is exposed to an atmosphere of carbon dioxide. (ii) Which reactant is an acid anhydride? (i) Nitrogen(V)oxide is bubbled through water. (ii) How does the pH of this reaction change? (i) Solid sodium hydrogencarbonate is added to water. (ii) How does the pH of this reaction change? (i) Solid sodium hydroxide and solid ammonium chloride are mixed and heated. (ii) Name the gaseous product. — XIX — (i) Methyl iodide (iodomethane) is heated with a solution of sodium hydroxide. (ii) Compare the reactivity of methyl iodide with the reactivity of methyl chloride. (i) Solid barium hydroxide and solid ammonium sulfate are mixed and heated. (ii) What product has the highest melting point? (i) A solution of sodium oxalate is added to an acidified solution of potassium permanganate. (ii) How many electrons are transferred during this reaction? (i) Equal volumes of equimolar hydrochloric acid are added to a solution of sodium hydrogenphosphate. (iii) What would be the stoichiometric ratio of the reactants? (i) Ethanol burns in air. (ii) Name a reagent that would detect the heavier of the products. 66 96) 97) 98) 99) 100) 101) 102) 103) 104) 105) 106) 107) 108) 109) 110) 111) 112) 113) 114) 115) — XX — (i) A strip of magnesium is put in a solution of iron(III)nitrate. (ii) What is reduced in this reaction? (i) Hydrogen peroxide is heated. (ii) What changes of oxidation state(s) occur here? (i) Iron filings are sprinkled into a solution of iron(III)chloride. (ii) What substance is oxidized? (i) Chlorine is bubbled through a solution of sodium bromide. (ii) What reaction would occur if a solution of sodium fluoride were used? (i) Solid lithium oxide is added to water. (ii) Name another metal whose oxide would react similarly. — XXI — (i) Methane reacts with an excess of chlorine gas. (ii) What type of organic reaction occurs here? (i) Hydrogen sulfide gas is bubbled into excess sodium hydroxide solution. (ii) Name any spectator ion(s). (i) Solutions of ammonia and carbon dioxide are mixed. (ii) Name the acid that is formed as an intermediate. (i) Solid copper(II)sulfide is added to a dilute solution of nitric acid. (ii) How would the appearance of the reaction mixture change? (i) Zinc is added to a solution of copper(II)sulfate. (ii) Describe any visible changes to the zind metal. — XXII — (i) Solutions of potassium hydroxide and ammonium sulfate are mixed. (ii) Name any spectator (ion)s. (i) Ethanol reacts with methanoic acid. (ii) Name this specific organic reaction. (i) Dilute sulfuric acid is added to solid sodium fluoride. (ii) Name the nonionic product. (i) An electric current is passed through a solution of copper(II)sulfate. (ii) What reaction occurs at the anode? (i) Solid sodium acetate is mixed with dilute hydrochloric acid. (ii) Name any spectator ion(s). — XXIII — (i) Formic acid is reacted with acidified potassium dichromate. (ii) How would the color of the reaction mixture change? (i) Sulfur dioxide gas is passed over solid calcium oxide. (ii) Name the product. (i) Ethene reacts with liquid bromine. (ii) Name the product. (i) Hydrochloric acid is given to a solution of dimercury(I)nitrate. (ii) Name any spectator ion(s). (i) Concentrated hydrobromic acid is heated with manganese dioxide. (ii) How many electrons are transferred in this reaction? 67 116) 117) 118) 119) 120) 121) 122) 123) 124) 125) 126) 127) 128) 129) 130) 131) 132) 133) 134) 135) — XXIV — (i) Bromine is added to a dilute solution of sodium hydroxide. (ii) How does the pH change during this reaction? (i) Acetic acid is added to solid sodium hydrogencarbonate. (ii) What gaseous product is formed? (i) Solutions of sodium sulfide and zinc nitrate are combined. (ii) Name this type of inorganic reaction. (i) A basic solution of potassium permanganate is added to a solution of sodium sulfite. (ii) How does the oxidation state of manganese change? (i) A stoichiometric amount of sulfuric acid is added to a solution of lithium carbonate. (ii) What product would be formed if the mole ratio of reactant were 1:1? — XXV — (i) Hydrogen sulfide gas is bubbled through a solution of nickel(II)nitrate. (ii) How would the pH of the solution change? (i) Magnesium is added to dilute nitric acid. (ii) What acts as an oxidizing agent? (i) Lead is added to a solution of silver nitrate. (ii) What observable change would occur to the piece of lead? (i) Borontrifluoride and ammonia gas are mixed. (ii) What are the formal charges of B and N in the product? (i) Propanol combusts in air. (ii) If 25 g of propanol were used, what volume of products at STP would be formed, assuming that all products are gaseous? — XXVI — (i) A solution of barium chloride is added to solid silver nitrate. (ii) Name any spectator ion(s). (i) Hydrogen gas is passed over hot iron(II)oxide. (ii) What acts as an oxidizing agent? (i) Hydrogen peroxide is mixed with an acidified solution of sodium bromide. (ii) How many electrons are transferred during this reaction? (i) Phosphorusoxytrichloride is added to an excess of potassium hydroxide solution. (ii) Name this type of reaction. (i) Rubidium is added to water. (ii) How would the pH of the reaction change? — XXVII — (i) Solutions of iron(III)sulfate and tin(II)chloride are mixed. (ii) How many electrons are transferred during this reaction? (i) An excess of lauric acid CH3(CH2)10COOH reacts with glycerol. (iii) Name the type of organic molecule formed. (i) Solid sodium sulfite is added to an acidified solution of sodium permanganate. (ii) What acts as a reducing agent? (i) Sodium sulfite solution is added to hydrochloric acid. (ii) Name the unstable intermediate in this reaction. (i) Solid barium methanoate is added to dilute nitric acid. (ii) Name any spectator ion(s). 68 136) 137) 138) 139) 140) 141) 142) 143) 144) 145) 146) 147) 148) 149) 150) — XXVIII — (i) Lithium reacts with nitrogen. (ii) How many electrons are transferred during this reaction? (i) Iron filings are boiled with a solution of iron(III)nitrate. (ii) What substance is reduced? (i) Water is added to solid sodium oxide. (ii) What color would added litmus show? (i) Water is added to solid sodium peroxide. (ii) How would the pH of the reaction mixture change? (i) Chlorine gas is bubbled through a solution of sodium iodide. (ii) What different product would be formed if sodium bromide were used? — XXIX — (i) Ammonia is bubbled through dilute acetic acid. (ii) Name any spectator ion(s). (i) 1-Pentene reacts with hydrogen in the presence of a catalyst. (ii) Name the product. (i) Magnesium oxide is exposed to sulfur trioxide gas. (ii) What species acts a base anhydride? (i) Calcium is put in water. (ii) Name Group 2 element that would show higher reactivity. (i) Silver is dropped into a concentrated solution of nitric acid. (ii) What acts as the oxidizing agent? — XXX — (i) Equal volumes of equimolar solutions of potassium hydroxide and sodium dihydrogenphospate are mixed. (ii) Name the product if stoichiometric amounts would be used. (i) Excess dilute hydrochloric acid is added to a solution of sodium phosphate. (ii) Name any spectator ion(s). (i) Solutions of silver nitrate and lithium bromide are combined. (ii) Name this type of inorganic reaction. (i) Carbon disulfide is burned in an excess of fluorine. (ii) What structure would you predict for the sulfur-containing product? (i) Butyl ethanoate is boiled with potassium hydroxide. (ii) Name the alcohol that is formed. Notes: 69 HHS, AP Chemistry, 2012/2013 AP Equations (Answer Key) Please use “Equations in the AP Examinations” and Appendices D and E in your textbook. The answer key might not be free of errors. Symbol Reaction Type SY synthesis DC decomposition DR double replacement/precipitation AS anionic single replacement CS cationic single replacement CO combustion RO redox (other) AB acid/base LE ligand exchange HY hydrolysis DH dehydration OA OS organic (addition to double or triple bonds) organic (substitution to alkanes or aromatic compounds) ET ether reactions ES ester reactions Examples Li + O2 Li2O Si + F2 SiF4 NaHCO3 Na2O + CO2 + H2O CaSO4 CaO + SO3 OH- + Fe3+ Fe(OH)3 Ca2+ + PO43- Ca3(PO4)2 Cl2 + Br- Cl- + Br2 F2 + Cl- F- + Cl2 Ca + Au3+ Ca2+ + Au Al + Cu+ Al3+ + Cu CH3OH + O2 CO2 + H2O C8H18 + O2 CO2 + H2O CrO42- + H2O2 + H+ Cr3+ + O2 + H2O Sn4+ + Fe2+ Sn2+ + Fe3+ MgO + H+ Mg2+ + H2O SO3 + OH- SO42- + H2O Cu(OH)2 + NH3 Cu(NH3)42+ + OHAgCl + CN- Ag(CN)2- + ClNa3N + H2O NH3 + Na+ + OHPCl5 + H2O H+ + PO43- + ClFe(OH)3 Fe2O3 + H2O H2SO3 H2O + SO2 C2H4 + Cl2 C2H4Cl2 C2H2 + Cl2 C2H2Cl2 CH4 + Br2 CH3Br + HBr C6H6 + Br2 C6H5 Br + HBr CH3OCH2CH3 + H2O CH3OH + CH3CH2OH CH3COOCH2CH3 + H2O CH3COOH + CH3CH2OH 70 —I— 1) (i) Solid lithium hydride is added to water. (ii) How does the pH of the reaction mixture change? LiH + H2O Li+ + OH- + H2 AB The pH would increase. 2) (i) Excess ammonia is added to a solution of zinc sulfate. (ii) What particles act as LEWIS acids and bases? 4 NH3 + Zn2+ Zn(NH3)42+ LE LEWIS acid: Zn2+; LEWIS base: NH3 3) (i) Sodium acetate is acidified with dilute hydrochloric acid. (ii) How does the pH of the reaction mixture change? CH3COO- + H+ CH3COOH AB The pH would increase. 4) (i) A solution of potassium dichromate is added to an acidified solution of sodium iodide. (ii) What is total number of electrons transferred? Cr2O72- + 14 H+ + 6 I- 2 Cr3+ + 3 I2 + 7 H2O RO Six electrons are transferred. 5) (i) Solid barium carbonate is added to dilute nitric acid. (ii) What is a simple analytic test for one of the products? BaCO3 + 2 H+ Ba2+ + H2O + CO2 AB, DC CO2 gas in calcium hydroxide solution precipitates CaCO3. 71 — II — 6) (i) Solid aluminum sulfide is added to water. (ii) Why is the gaseous product toxic? Al2S3 + 6 H2O 2 Al(OH)3 + 3 H2S AB/DR 7) H2S is an acid and dissociates, forming sulfide ions S2- that can precipitate Fe2+ in hemoglobin, thus inactivating it. (i) Silver metal is added to 6M nitric acid. (ii) What acts as an oxidizing agent? 3 Ag + 4 H+ + NO3- 3 Ag+ + NO + 2 H2O RO The nitrate ion NO3- acts an oxidizing agent. 8) (i) A solution of tin(II)nitrate is dropped onto a piece of zinc metal. (ii) What observation can be made about the piece of zinc? Sn2+ + Zn Sn4+ + Zn2+ RO The zinc metal corrodes. 9) (i) Solid sodium dichromate is added to an acidified solution of iron(II) sulfate. (ii) What is the total number of electrons transferred? Na2Cr2O7 + 14 H+ + 6 Fe2+ 2 Na+ + 2 Cr3+ + 6 Fe3+ + 7 H2O RO Six electrons are transferred. 10) (i) Liquid titanium tetrachloride is given into excess water. (ii) How does the pH of the reaction mixture change? TiCl4 + 2 H2O TiO2 + 4 H+ + 4 ClHY The pH decreases. 72 — III — 11) (i) Solid ammonium chloride is added to a concentrated solution of sodium hydroxide. (ii) What ion(s), if any, act as spectator ions? NH4Cl + OH- NH3 + H2O + ClAB 12) Sodium ion Na+ is a spectator ion. (i) Solid sodium hydrogencarbonate is added to concentrated sodium hydroxide. (ii) Sodium hydrogencarbonate reacts with hydrochloric acid as well. Name this behavior. NaHCO3 + OH- Na+ + CO32- + H2O AB Amphoterism 13) (i) Solid calcium sulfite is acidified. (ii) Name one unstable reaction intermediate. CaSO3 + 2 H+ Ca2+ + H2O + SO2 AB, DC Sulfurous acid H2SO3 14) (i) Liquid acetic acid and methanol are combined and warmed. (ii) What specific organic reaction type occurs here? CH3COOH + CH3OH CH3COOCH3 + H2O ES, DH Esterification 15) (i) Aqueous solutions of sodium iodate and sodium iodide are combined, then acidified. (ii) What specific type of redox reaction occurs here? IO3- + 5 I- + 6 H+ 3 I2 + 3 H2O RO Comproportionation 73 — IV — 16) (i) Solid calcium phosphate is given to excess hydrobromic acid. (ii) What phosphorus-containing product would be formed if hydrobromic acid and calcium phosphate reacted in a mole ratio of 1:1? Ca3(PO4)2 + 6 H+ 3 Ca2+ + 2 H3PO4 AB 17) Hydrogenphosphate ion HPO42- (i) A piece of magnesium is dropped into hydrochloric acid. (ii) How many electrons are transferred in this reaction? Mg + 2 H+ Mg2+ + H2 CS Two electrons are transferred. 18) (i) A stream of carbon dioxide gas is bubbled through a suspension of barium carbonate. (ii) What observation can be made regarding the appearance of the reaction mixture? CO2 + H2O + BaCO3 Ba2+ + 2 HCO3AB The precipitate dissolves. 19) (i) Potassium permanganate solution is added to concentrated hydrochloric acid. (ii) How does the color of the reaction mixture change? 2 MnO4- + 16 H+ + 10 Cl- 2 Mn2+ + 5 Cl2 + 8 H2O RO The solution color changes from violet to colorless (faint pink). 20) (i) Ammonia gas is bubbled through a suspension of zinc hydroxide. (ii) What observation can be made regarding the appearance of the reaction mixture? 4 NH3 + Zn(OH)2 Zn(NH3)42+ + 2 OHLE The precipitate dissolves. 74 —V— 21) (i) Hydrogen sulfide gas reacts with moist lead(II)chloride. (ii) What color change could be observed? H2S + Pb2+ PbS + 2 H+ AB, DR The reaction mixture turns black. 22) (i) Moist sodium hydroxide reacts with carbon dioxide gas. (ii) How and why does the pH of the reaction mixture change? 2 OH- + CO2 H2O + CO32- (or: OH- + CO2 HCO3-) AB Carbon dioxide is an acid anhydride, therefore the pH decreases. 23) (i) An aqueous solution of ammonia is added to solid silver chloride. (ii) Name this type of reaction. 2 NH3 + AgCl Ag(NH3)2+ + ClLE Ligand-exchange reaction 24) (i) Rubidium permanganate is added to hydroiodic acid. (ii) How many electrons are transferred in this reaction? 2 MnO4- + 16 H+ + 10 I- 2 Mn2+ + 5 I2 + 8 H2O RO Ten electrons are transferred. 25) (i) Liquid phosporus trichloride is poured into an excess of sodium hydroxide. (iii) What are the oxidation states of phosphorus in reactant(s) and product(s)? PCl3 + 6 OH- PO33- + 3 Cl- + 3 H2O HY, AB The oxidation state is +3 in both reactant and product. 75 — VI — 26) (i) 30% hydrogen peroxide is heated. (ii) What are the oxidation states of oxygen in reactant(s) and product(s)? 2 H2O2 2 H2O + O2 RO/DC Oxidation states of oxygen: H2O2: -1; H2O: -2; O2: 0 27) (i) Liquid mercury is added to concentrated nitric acid. (ii) What substance acts as a reducing agent? 3 Hg + 8 H+ + 2 NO3- 3 Hg2+ + 2 NO + 4 H2O RO Mercury Hg acts as reducing agent. 28) (i) Solid iron(III)sulfate is added to a sodium iodide solution. (ii) What substance could be used to detect one of the products in small concentrations? Fe2(SO4)3 + 2 I- 2 Fe2+ + I2 + 3 SO42RO 29) Iodine I2 can be detected by starch; it forms a dark violet solution. (i) Hydrogen sulfide is bubbled through a solution of silver nitrate. (ii) Name this specific type of reaction. H2S + Ag+ Ag2S + 2 H+ AB, DR Double replacement 30) (i) Concentrated (15 M) ammonia solution is added in excess to a solution of copper(II)nitrate. (ii) Describe any observable color changes. 4 NH3 + Cu2+ Cu(NH3)42+ LE The blue color intensifies during the reaction. 76 — VII — 31) (i) Magnesium metal is added to dilute nitric acid, giving as one of the products a compound in which the oxidation number of nitrogen is -3. (ii) Name said product. 4 Mg + 10 H+ + NO3- 4 Mg2+ + NH4+ + 3 H2O RO 32) Ammonium ion NH4+ (i) Excess water is added to solid calcium hydride. (ii) Why must calcium hydride be stored in a well-sealed container? 2 H2O + CaH2 Ca2+ + 2 OH- + H2 AB 33) It could react with moisture in the air; the hydrogen could burst the container. (i) Silver acetate solution is dropped into a solution of sodium phosphate. (ii) If the contents of the reaction mixture would be filtered, what substance(s) would remain on the filter paper? 3 Ag+ + PO43- Ag3PO4 DR Silver phosphate 34) (i) Solid sodium cyanide is added to water. (ii) Why is sodium cyanide toxic? NaCN + H2O HCN + Na+ + OHAB 35) Cyanide ions remove Fe2+ from hemoglobin as Fe(CN)64-, therefore inactivating it. (i) Solid potassium hydride is added to anhydrous ethyl alcohol. (ii) Name the solid reaction product. KH + CH3CH2OH CH3CH2OK + H2 AB Potassium ethanoate 77 — VIII — 36) (i) Lithium metal is burned in air. (ii) Given sufficiently high temperatures, what byproduct could be formed? 4 Li + O2 2 Li2O SY Lithium nitride Li3N 37) (i) Aluminum metal is added to a solution of copper(II)chloride. (ii) What, if any, color change would be observable? 2 Al + 3 Cu2+ 2 Al3+ + 3 Cu CS The blue color of Cu2+ would fade over time; Al3+ is colorless. 38) (i) Manganese(II)nitrate is mixed with sodium hydroxide solution. (ii) Name any spectator ions. Mn2+ + 2 OH- Mn(OH)2 DR Nitrate NO3- and sodium ion Na+ 39) (i) Equal volumes of dilute equimolar solutions of sodium carbonate and hydrochloric acid are mixed. (ii) What ratio of sodium carbonate to hydrochloric acid would lead to the formation of a gas? CO32- + H+ HCO3- AB, DC (sodium carbonate) : (hydrochloric acid) = 2:1 40) (i) Solid sodium carbide is added to an excess of water. (ii) How does the pH of the reaction mixture change? Na2C2 + 2 H2O 2 Na+ + 2 OH- + C2H2 AB The pH increases. 78 — IX — 41) (i) An excess of sodium hydroxide solution is added to a solution of aluminum chloride. (ii) Describe any visible changes during the reaction. 6 OH- + Al3+ Al(OH)63AB 42) An initial precipitate of Al(OH)3 would dissolve when further OHis added. (i) Solid ammonium carbonate is heated. (ii) Which of the products has the lowest boiling point? (NH4)2CO3 2 NH3 + H2O + CO2 DC Carbon dioxide CO2 43) (i) Phosphorus pentachloride is added to an excess of potassium hydroxide solution. (ii) Name the type of reaction that the phosphorus pentachloride undergoes. PCl5 + 8 OH- PO43- + 5 Cl- + 4 H2O HY, AB Hydrolysis 44) (i) Potassium chlorate is heated in the presence of manganese dioxide. (ii) What function does the manganese dioxide have? 2 KClO3 2 KCl + 3 O2 DC Manganese dioxide acts as a catalyst. 45) (i) Chlorine gas is bubbled into a solution of sodium hydroxide. (ii) What oxidation state(s) change(s)? Cl2 + 2 OH- Cl- + OCl- + H2O RO, AB Chlorine: 0 -1 and +1 79 —X— 46) (i) Magnesium carbonate is heated strongly. (ii) Which of the product is an acid anhydride? MgCO3 MgO + CO2 DC Carbon dioxide CO2 is an acid anhydride. 47) (i) Methanol and propanoic acid are mixed. (ii) What general organic reaction takes place here? CH3OH + CH3CH2COOH CH3CH2COOCH3 + H2O ES, DH Dehydration synthesis (condensation) 48) (i) Barium acetate is added to dilute sulfuric acid. (ii) What, if any, spectator ions occur here? Ba2+ + 2 CH3COO- + 2 H+ + SO42- 2 CH3COOH + BaSO4 AB, DR There are no spectator ions. 49) (i) Ammonia gas is bubbled through a suspension of silver chloride. (ii) Describe any visible changes of the reaction mixture. 2 NH3 + AgCl Ag(NH3)2+ + ClLE The precipitate of AgCl would dissolve. 50) (i) Hot iron(III)oxide is reacted with carbon monoxide. (ii) How many electrons are transferred during the reaction? Fe2O3 + 3 CO 2 Fe + 3 CO2 RO Six electrons are transferred. 80 — XI — 51) (i) Silane (SiH4) is burned in air. (ii) What type of solid is the product containing silicon? SiH4 + 2 O2 SiO2 + 2 H2O RO, CO Silicon dioxide SiO2 is a network solid. 52) (i) Cesium metal is heated with sulfur. (ii) Name this type of reaction. 16 Cs + S8 8 Cs2S SY Synthesis 53) (i) Sodium phosphate is dissolved in water. (ii) How does the pH of the reaction mixture change? Na3PO4 + H2O HPO42- + 3 Na+ + OHAB The pH increases. 54) (i) Ethyl ethanoate is boiled with sodium hydroxide. (ii) What type of organic compound is ethyl ethanoate? CH3COOCH2CH3 + OH- CH3COO- + CH3CH2OH ES, HY Ethyl ethanoate is an ester. 55) (i) Solutions of sodium phosphate and calcium chloride are combined. (ii) Name any spectator ion(s). 2 PO43- + 3 Ca2+ Ca3(PO4)2 DR Sodium Na+ and chloride ion Cl- 81 — XII — 56) (i) A suspension of aluminum hydroxide is added to hydroiodic acid. (ii) How would the reaction be different if hydrochloric acid were used? Al(OH)3 + 3 H+ Al3+ + 3 H2O AB The reaction would be the same. 57) (i) Hydrogen peroxide is added to an acidified solution of sodium iodide. (ii) How many electrons are transferred during this reaction? H2O2 + 2 H+ + 2 I- 2 H2O + I2 RO Two electrons are transferred. 58) (i) Hydrogen peroxide is added to an acidified solution of potassium dichromate. (ii) What are the oxidation states of chromium in this reaction? 3 H2O2 + 8 H+ + Cr2O72- 3 O2 + 7 H2O + 2 Cr3+ RO The oxidation states of chromium are +6 (Cr2O72-) and +3 (Cr3+). 59) (i) Calcium is added to dilute hydrochloric acid. (ii) What specific inorganic reaction type occurs here? Ca + 2 H+ Ca2+ + H2 CS Cationic single replacement 60) (i) Aluminum nitrate is dissolved in water. (ii) Which substance acts as LEWIS acid here? Al(H2O)63+ [Al(H2O)5(OH)]- + H+ AB Aluminum ion Al3+ 82 — XIII — 61) (i) Phosphorus tribromide is added to water. (ii) Name the product containing phosporus. PBr3 + 3 H2O PO33- + 3 Br- + 6 H+ HY Phosphite ion PO33- 62) (i) Solutions of sulfuric acid and lithium hydroxide are mixed. (ii) What specific analytical method would use this type of reaction? H+ + OH- H2O AB Acid-base titration 63) (i) Nitric acid is added to a solution of tetraamminecadmium(II). (ii) What spectator ion(s), if any, occur in this reaction? 4 H+ + Cd(NH3)42+ Cd2+ + 4 NH4+ AB, LE Nitrate ion NO3- 64) (i) 2-butene reacts with hydrobromic acid. (ii) What functional group of 2-butene is active in this reaction? CH3CH=CHCH3 + H+ + Br- CH3CHBr—CH2CH3 OA The functional group is the C=C double bond. 65) (i) Hydrogensulfide is bubbled through a solution of lead(II)nitrate. (ii) What color change could be observed? H2S + Pb2+ 2 H+ + PbS AB, DR The solution would turn black. 83 — XIV — 66) (i) Solutions of iron(III)nitrate and sodium thiocyanate are combined. (ii) What color change could be observed? Fe(H2O)63+ + SCN- [Fe(SCN)(H2O)5]2+ + H2O LE/DR The solution would turn red. 67) (i) Ammonia gas is bubbled through a suspension of zinc hydroxide. (ii) How many electrons are transferred during this reaction? 4 NH3 + Zn(OH)2 Zn(NH3)42+ + 2 OHLE No electrons would be transferred. 68) (i) Sulfur dioxide gas is bubbled through an acidified solution of potassium permanganate. (ii) How many electrons are transferred during this reaction? 5 SO2 + 2 H2O + 2 MnO4- 5 SO42- + 4 H+ + 2 Mn2+ RO Ten electrons are transferred. 69) (i) Aluminum is added to excess sodium hydroxide solution. (ii) Name the ionic product of this reaction. 2 Al + 6 OH- + 6 H2O 2 Al(OH)63- + 3 H2 RO, LE Hexahydroxoaluminate ion 70) (i) A solution of tin(II)nitrate is added to an acidified solution of potassium dichromate. (ii) How many electrons are transferred during this reaction? 3 Sn2+ + 14 H+ + Cr2O72- 3 Sn4+ + 2 Cr3+ + 7 H2O RO Six electrons are transferred. 84 — XV — 71) (i) Equal volumes of equimolar sodium hydroxide and sodium dihydrogenphosphate are mixed. (ii) What phosphorus-containing product would be formed if the ratio of hydroxide to dihydrogenphosphate were 2:1? OH- + H2PO4- HPO42- + H2O AB 72) The phosphate ion PO43- (i) Excess sodium hydroxide is added to a solution of calcium hydrogencarbonate. (ii) Name any spectator ions. OH- + HCO3- + Ca2+ H2O + CaCO3 AB, DR 73) Sodium ion Na+ (i) Acetylene (ethyne) reacts with an excess of chlorine. (ii) What product would be formed if the reactants were used in a ratio of 1:1? HCCH + Cl2 CHCl2—CHCl2 OA 74) 1,2-dichloroethene ClCH2—CH2Cl (i) Solid silver nitrate is added to a solution of sodium chromate. (ii) Name any spectator ions. DR 2 AgNO3 + CrO42- Ag2CrO4 + 2 NO3Sodium ion Na+ 75) (i) An excess of solid potassium hydroxide is added to a suspension of zinc hydroxide. (ii) What changes in the reaction mixture can be observed? 2 KOH + Zn(OH)2 2 K+ + Zn(OH)42LE The precipitate of zinc hydroxide would dissolve. 85 — XVI — 76) (i) Benzene reacts with bromine. (ii) Why does benzene not react as if it had double bonds? C6H6 + Br2 C6H5Br + HBr OS 77) Benzene does not react as if it had double bond due to resonance stabilization. (i) A direct electric current is passed through a dilute solution of sulfuric acid. (ii) What gas is formed at the cathode? 2 H2O 2 H2 + O2 RO Hydrogen gas is formed at the cathode. 78) (i) Ammonia and oxygen are heated in the presence of a catalyst. (ii) How does the oxidation number for nitrogen change? 4 NH3 + 5 O2 4 NO + 6 H2O RO The oxidation states of nitrogen are -3 (NH3) and +2 (NO). 79) (i) Liquid bromine is added to a solution of sodium iodide. (ii) What specific type of inorganic reaction occurs here? Br2 + 2 I- 2 Br- + I2 AS Anionic single replacement 80) (i) Dilute hydrochloric acid is added to solid calcium oxide. (ii) Name the spectator ion. 2 H+ + CaO Ca2+ + H2O AB Chloride Cl- 86 — XVII — 81) (i) Sodium hydride is added to water. (ii) Compare the relative reactivity of potassium hydride with sodium hydride. NaH + H2O Na+ + OH- + H2 AB Potassium hydride is more reactive than sodium hydride. 82) (i) Ethene reacts with water in the presence of a catalyst. (ii) Name the product. CH2=CH2 + H2O CH3—CH2OH OA Ethanol 83) (i) Solutions of barium hydroxide and iron(III)sulfate are combined. (ii) If equal moles of reactants were combined, which would be the limiting reactant? 3 Ba2+ + 6 OH- + 2 Fe3+ + 3 SO42- 3 BaSO4 + 2 Fe(OH)3 DR The limiting reactant would be barium hydroxide Ba(OH)2. 84) (i) Ammonia gas is bubbled through a solution of copper(II)nitrate. (ii) What color change could be observed? 4 NH3 + Cu2+ Cu(NH3)42+ LE The blue color intensifies during the reaction. 85) (i) Lead is added to a dilute solution of sulfuric acid. (ii) How many electrons are transferred during this reaction? Pb + 2 H+ + SO42- PbSO4 + H2 CS, DR Two electrons are transferred. 87 — XVIII — 86) (i) Lead is added to a hot solution of concentrated sulfuric acid. (ii) What acts as the oxidizing agent? Pb + 4 H+ + 2 SO42- PbSO4 + SO2 + 2 H2O RO, DR The oxidizing agent is the sulfate ion SO42-. 87) (i) Calcium oxide is exposed to an atmosphere of carbon dioxide. (ii) Which reactant is an acid anhydride? CaO + CO2 CaCO3 SY The acid anhydride is carbon dioxide CO2. 88) (i) Nitrogen(V)oxide is bubbled through water. (ii) How does the pH of this reaction change? N2O5 + H2O 2 H+ + 2 NO3AB The pH decreases. 89) (i) Solid sodium hydrogencarbonate is added to water. (ii) How does the pH of this reaction change? NaHCO3 Na+ + CO2 + OHAB The pH increases. 90) (i) Solid sodium hydroxide and solid ammonium chloride are mixed and heated. (ii) Name the gaseous product. NaOH + NH4Cl NaCl + NH3 + H2O AB Ammonia NH3 88 — XIX — 91) (i) Iodomethane is heated with a solution of sodium hydroxide. (ii) Compare the reactivity of iodomethane with chloromethane. CH3I + OH- CH3OH +IOS Chloromethane is less reactive than iodomethane. 92) (i) Solid barium hydroxide and solid ammonium sulfate are mixed and heated. (ii) What product has the highest melting point? Ba2+ + 2 OH- + 2 NH4+ + SO42- BaSO4 + 2 NH3 + 2 H2O DR, AB Barium sulfate BaSO4 has the highest melting point. 93) (i) A solution of sodium oxalate is added to an acidified solution of potassium permanganate. (ii) How many electrons are transferred during this reaction? 5 -OOC—COO- + 16 H+ + 2 MnO4- 10 CO2 + 2 Mn2+ + 8 H2O RO Ten electrons are transferred. 94) (i) Equal volumes of equimolar hydrochloric acid are added to a solution of sodium hydrogenphosphate. (iii) What would be the stoichiometric ratio of the reactants? H+ + HPO42- H2PO4AB 95) The stoichiometric ratio would be (hydrochloric acid) : (hydrogenphosphate) = 2:1 (i) Ethanol burns in air. (ii) Name a reagent that would detect the heavier of the products. CH3CH2OH + O2 CO2 + H2O CO Carbon dioxide would precipitate calcium carbonate out of a solution of calcium hydroxide Ca(OH)2 — XX — 89 96) (i) A strip of magnesium is put in a solution of iron(III)nitrate. (ii) What is reduced in this reaction? Mg + 2 Fe3+ Mg2+ + 2 Fe2+ CS/RO The iron (III) ion would be reduced. 97) (i) Hydrogen peroxide is heated. (ii) What changes of oxidation state(s) occur here? 2 H2O2 2 H2O + O2 DC Oxidation states of oxygen: -1 (H2O2) to -2 (H2O) and 0 (O2) 98) (i) Iron filings are sprinkled into a solution of iron(III)chloride. (ii) What substance is oxidized? Fe + 2 Fe3+ 3 Fe2+ RO The iron Fe is oxidized. 99) (i) Chlorine is bubbled through a solution of sodium bromide. (ii) What reaction would occur if a solution of sodium fluoride were used? Cl2 + 2 Br- 2 Cl- + Br2 AS No reaction would occur. 100) (i) Solid lithium oxide is added to water. (ii) Name another metal whose oxide would react similarly. Li2O + H2O 2 Li+ + 2 OHAB Any other Group 1 metal (Na, K, Rb, Cs) 90 — XXI — 101) (i) Methane reacts with an excess of chlorine gas. (ii) What type of organic reaction occurs here? CH4 + 4 Cl2 CCl4 + 4 HCl OS (Aliphatic) Substitution 102) (i) Hydrogen sulfide gas is bubbled into excess sodium hydroxide solution. (ii) Name any spectator ion(s). H2S + 2 OH- S2- + 2 H2O AB Sodium ion Na+ 103) (i) Solutions of ammonia and carbon dioxide are mixed. (ii) Name the acid that is formed as an intermediate. 2 NH3 + CO2 + H2O 2 NH4+ + CO32AB/SY Carbonic acid H2CO3 104) (i) Solid copper(II)sulfide is added to a dilute solution of nitric acid. (ii) How would the appearance of the reaction mixture change? CuS + 2 H+ Cu2+ + H2S AB The black CuS precipitate would disappear and the solution would turn blue. 105) (i) Zinc is added to a solution of copper(II)sulfate. (ii) Describe any visible changes to the zinc metal. Zn + Cu2+ Zn2+ + Cu CS The zinc would corrode. 91 — XXII — 106) (i) Solutions of potassium hydroxide and ammonium sulfate are mixed. (ii) Name any spectator (ion)s. OH- + NH4+ NH3 + H2O AB Potassium ion K+ and sulfate SO42- 107) (i) Ethanol reacts with methanoic acid. (ii) Name this specific organic reaction. CH3CH2OH + HCOOH HCOOCH2CH3 + H2O ES, DH Esterification 108) (i) Dilute sulfuric acid is added to solid sodium fluoride. (ii) Name the nonionic product. 2 H+ + SO42- + CaF2 2 HF + CaSO4 AB, DR Hydrogen fluoride 109) (i) An electric current is passed through a solution of copper(II)sulfate. (ii) What reaction occurs at the anode? 2 Cu2+ + 2 OH- 2 Cu + O2 + 2 H+ RO The oxidation at the anode is 2OH- O2 + 4 e- + 2 H+ 110) (i) Solid sodium acetate is mixed with dilute hydrochloric acid. (ii) Name any spectator ion(s). NaCH3COO + H+ Na+ + CH3COOH AB Chloride Cl- 92 — XXIII — 111) (i) Formic acid is reacted with acidified potassium dichromate. (ii) How would the color of the reaction mixture change? 3 HCOOH + 8 H+ + Cr2O72- 3 CO2 + 2 Cr3+ + 7 H2O RO The color would change from orange to green. 112) (i) Sulfur dioxide gas is passed over solid calcium oxide. (ii) Name the product. SO2 + CaO CaSO3 SY Calcium sulfite 113) (i) Ethene reacts with liquid bromine. (ii) Name the product. CH2=CH2 + Br2 CH2Br—CH2Br OA 1,2-dibromoethane 114) (i) Hydrochloric acid is given to a solution of dimercury(I)nitrate. (ii) Name any spectator ion(s). Cl- + Hg22+ Hg2Cl2 DR Hydrogen ion H+ and nitrate NO3- 115) (i) Concentrated hydrobromic acid is heated with manganese dioxide. (ii) How many electrons are transferred in this reaction? 4 H+ + 2 Br- + MnO2 Br2 + Mn2+ + 2 H2O RO Two electrons are transferred. 93 — XXIV — 116) (i) Bromine is added to a dilute solution of sodium hydroxide. (ii) How does the pH change during this reaction? Br2 + 2 OH- Br- + OBr- + H2O RO, AB The pH decreases. 117) (i) Acetic acid is added to solid sodium hydrogencarbonate. (ii) What gaseous product is formed? CH3COOH + NaHCO3 Na+ + CH3COO- + CO2 + H2O AB, DC Carbon dioxide CO2 118) (i) Solutions of sodium sulfide and zinc nitrate are combined. (ii) Name this type of inorganic reaction. S2- + Zn2+ ZnS DR Double replacement 119) (i) A basic solution of potassium permanganate is added to a solution of sodium sulfite. (ii) How does the oxidation state of manganese change? H2O + 2 MnO4- + 3 SO32- 2 MnO2 + 2 OH- + 3 SO42RO Oxidation states of manganese: +7 (MnO4-) and +4 (MnO2) 120) (i) A stoichiometric amount of sulfuric acid is added to a solution of lithium carbonate. (ii) What product would be formed if the mole ratio of reactant were 1:1? 2 H+ + CO3- H2O + CO2 AB, DC Hydrogencarbonate HCO3- 94 — XXV — 121) (i) Hydrogen sulfide gas is bubbled through a solution of nickel(II)nitrate. (ii) How would the pH of the solution change? H2S + Ni2+ NiS + 2 H+ AB, DR The pH would decrease. 122) (i) Magnesium is added to dilute nitric acid. (ii) What acts as an oxidizing agent? Mg + 2 H+ Mg2+ + H2 RO Hydrogen ion H+ 123) (i) Lead is added to a solution of silver nitrate. (ii) What observable change would occur to the piece of lead? Pb + 2 Ag+ Pb2+ + 2 Ag CS The lead would corrode. 124) (i) Borontrifluoride and ammonia gas are mixed. (ii) What are the formal charges of B and N in the product? BF3 + NH3 F3B—NH3 (addition) The formal charges would be +1 (N) and -1 (B). 125) (i) Propanol combusts in air. (ii) If 25.0 g of propanol were used, what volume of products at STP would be formed, assuming that all products are gaseous? 2 CH3CH2CH2OH + 9 O2 6 CO2 + 8 H2O CO The total product volume would be 65.3 liters. 95 — XXVI — 126) (i) A solution of barium chloride is added to solid silver nitrate. (ii) Name any spectator ion(s). Cl- + Ag+ AgCl DR Barium ion Ba2+ and nitrate NO3- 127) (i) Hydrogen gas is passed over hot iron(II)oxide. (ii) What acts as an oxidizing agent? H2 + FeO H2O + Fe CS The oxidizing agent is Fe2+ in FeO. 128) (i) Hydrogen peroxide is mixed with an acidified solution of sodium bromide. (ii) How many electrons are transferred during this reaction? H2O2 + 2 H+ + 2 Br- 2 H2O + Br2 RO Two electrons are transferred. 129) (i) Phosphorusoxytrichloride is added to an excess of potassium hydroxide solution. (ii) Name this type of reaction. POCl3 + 6 OH- PO43- + 3 Cl- + 3 H2O HY, AB Hydrolysis and subsequent deprotonation 130) (i) Rubidium is added to water. (ii) How would the pH of the reaction change? 2 Rb + 2 H2O 2 Rb+ + 2 OH- + H2 RO The pH would increase. 96 — XXVII — 131) (i) Solutions of iron(III)sulfate and tin(II)chloride are mixed. (ii) How many electrons are transferred during this reaction? 2 Fe3+ + Sn2+ 2 Fe2+ + Sn4+ RO Two electrons are transferred. 132) (i) An excess of lauric acid CH3(CH2)10COOH reacts with glycerol. (iii) Name the type of organic molecule formed. 3 CH3(CH2)10COOH + C3H5(OH)3 C3H5[OOC(CH2)10CH3]3 + 3 H2O ES, DH Triglyceride 133) (i) Solid sodium sulfite is added to an acidified solution of sodium permanganate. (ii) What acts as a reducing agent? 5 Na2SO3 + 6 H+ + 2 MnO4- 10 Na+ + 5 SO42- + 2 Mn2+ + 3 H2O RO The sulfite ion SO32- is the reducing agent. 134) (i) Sodium sulfite solution is added to hydrochloric acid. (ii) Name the unstable intermediate in this reaction. SO32- + 2 H+ H2O + SO2 AB, DC Sulfurous acid H2SO3 135) (i) Solid barium methanoate is added to dilute nitric acid. (ii) Name any spectator ion(s). Ba(CH3O)2 + 2 H+ Ba2+ + CH3OH AB Nitrate NO3- 97 — XXVIII — 136) (i) Lithium reacts with nitrogen. (ii) How many electrons are transferred during this reaction? 6 Li + N2 2 Li3N SY Six electrons are transferred. 137) (i) Iron filings are boiled with a solution of iron(III)nitrate. (ii) What substance is reduced? Fe + 2 Fe3+ 3 Fe2+ RO The iron (III) ion Fe3+ is reduced. 138) (i) Water is added to solid sodium oxide. (ii) What color would added litmus show? H2O + Na2O 2 Na+ + 2 OHAB Litmus would turn blue. 139) (i) Water is added to solid sodium peroxide. (ii) How would the pH of the reaction mixture change? 2 H2O + Na2O2 2 Na+ + 2 OH- + H2O2 AB The pH would increase. 140) (i) Chlorine gas is bubbled through a solution of sodium iodide. (ii) What different product would be formed if sodium bromide were used? Cl2 + 2 I- 2 Cl- + I2 AS Bromine Br2 98 — XXIX — 141) (i) Ammonia is bubbled through dilute acetic acid. (ii) Name any spectator ion(s). CH3COOH + NH3 CH3COO- + NH4+ AB There are no spectator ions. 142) (i) 1-Pentene reacts with hydrogen in the presence of a catalyst. (ii) Name the product. CH2=CHCH2CH2CH3 + H2 CH3CH2CH2CH2CH3 OA Pentane 143) (i) Magnesium oxide is exposed to sulfur trioxide gas. (ii) What species acts a base anhydride? MgO + SO3 MgSO4 AB The base anhydride is magnesium oxide MgO. 144) (i) Calcium is put in water. (ii) Name a Group 2 element that would show higher reactivity. Ca + 2 H2O Ca2+ + 2 OH- + H2 RO Strontium Sr, barium Ba, or radium Ra 145) (i) Silver is dropped into a concentrated solution of nitric acid. (ii) What acts as the oxidizing agent? 3 Ag + 4 H+ + NO3- 3 Ag+ + NO + 2 H2O RO The oxidizing agent is the nitrate ion NO3-. 99 — XXX — 146) (i) Equal volumes of equimolar solutions of potassium hydroxide and sodium dihydrogenphospate are mixed. (ii) Name the product if stoichiometric amounts would be used. OH- + H2PO4- HPO42- + H2O AB Phosphoric acid H3PO4 147) (i) Excess dilute hydrochloric acid is added to a solution of sodium phosphate. (ii) Name any spectator ion(s). H+ + PO43- H3PO4 AB Chloride Cl- and sodium ion Na+ 148) (i) Solutions of silver nitrate and lithium bromide are combined. (ii) Name this type of inorganic reaction. Ag+ + Br- AgBr DR Double replacement 149) (i) Carbon disulfide is burned in an excess of fluorine. (ii) What structure would you predict for the sulfur-containing product? CS2 + 8 F2 CF4 + 2 SF6 RO Sulfur hexafluoride SF6 has an octahedral structure. 150) (i) Butyl ethanoate is boiled with potassium hydroxide. (ii) Name the alcohol that is formed. CH3COOCH2CH2CH2CH3 + OH- CH3COO- + CH3CH2CH2CH2OH ES, HY 1-butanol 100 APPENDIX B: SOLUBILITY RULES FOR IONIC COMPOUNDS The following cations are considered: 1) 2) 3) 4) 5) 6) 7) 8) alkali metals (Li+, Na+, etc.) and NH4+ (similar reactivity) alkaline earth metals (Be2+, Mg2+, Ca2+, etc.) Al3+ Sn2+, Pb2+ Cr3+ Mn2+ Fe2+, Fe3+, Co2+, Ni2+ Cu2+, Ag+, Zn2+, Cd2+, Hg22+ (= "Hg+"), Hg2+ There is no completely insoluble compound; they are often classified as "mostly soluble" or "mostly (nearly) insoluble." For brevity I will use “soluble” and “insoluble.” Soluble salts with the following anions: NO3-, C2H3O2-, ClO3-, ClO4-, FCl- (except: Ag+, Hg22+, Pb2+) Br- (except: Ag+, Hg22+, Hg2+, Pb2+) I- (except: Ag+, Hg22+, Hg2+, Pb2+) SO42- (except: Ca2+, Sr2+, Ba2+, Hg22+, Hg2+, Pb2+) Insoluble salts with the following anions S2- (except: alkali and alkaline earth , NH4+) CO32- (except: alkali, NH4+) SO32- (except: alkali, NH4+) PO43- (except: alkali, NH4+) OH- (except: alkali, Ba2+, Sr2+, Ca2+) NH4+ decomposes: NH4+ + OH- NH3 + H2O O2- (except: alkali, Ba2+, Sr2+, Ca2+) NH4+ decomposes: NH4+ + O2- NH3 + OHReaction with water: O2- + H2O 2 OH- 101 Notes: 102 APPENDIX C: NOMENCLATURE OF INORGANIC COMPOUNDS 1) Binary Acids: Binary acids contain only one hydrogen and one more element. The names of acids and corresponding ions are derived from the element name. Formula Acid: “hydro” + (element) + “ic” Anion: (element) + “ide” HCl hydrochloric acid Cl- chloride HBr hydrobromic acid Br- bromide H2S hydrosulfuric acid S2(HS- sulfide hydrogensulfide) analogous names: F- fluoride, N3- nitride, P3- phosphide, O2- oxide, etc. 2) Ternary Acids: Ternary acids contain hydrogen and two more elements, one of them usually being oxygen ( oxyacids). The amount of oxygen present (i.e., the oxidation state of the third element) is denoted by various pre- and suffixes. Group 17 Formula Name Formula Name HClO4 HClO3 HClO2 HClO perchloric acid chloric acid chlorous acid hypochlorous acid ClO4ClO3ClO2ClO- perchlorate chlorate chlorite hypochlorite (analogous: Br, I) Fluorine forms only hypofluorous acid HOF. 103 Group 16 Formula Name Formula Name H2SO4 sulfuric acid SO42(HSO4- sulfate hydrogensulfate) H2SO3 sulfurous acid SO32(HSO3- sulfite hydrogensulfite) analogous: Se, Te Group 15 Formula Name Formula Name H3PO4 H3PO3 phosphoric acid phosphorous acid PO43PO33- phosphate phosphite H3AsO5 H3AsO4 perarsenic acid arsenic acid AsO53AsO53- perarsenate arsenate CO32BO33SiO44- carbonate borate silicate Some elements form only one type of oxyacid: H2CO3 H3BO3 H4SiO4 carbonic acid boric acid silicic acid In general, salts of elements in the same group follow the same pattern (formula and name). This is not true for carbon (no H4CO4) and nitrogen (no H3NO4). And some ions do not follow these “rules” at all: MnO4RuO43) permanganate perruthenate but but MnO42RuO42- manganate ruthenate Molecular or ionic compounds Using prefixes (Appendix E) or oxidation states, the nomenclature is similar for both. The cationic (less electronegative) particle/ion is written first. Notes: 104 APPENDIX D: NOMENCLATURE OF ORGANIC COMPOUNDS Alkanes C—C CnH2n+2 CH4 methane C2H6 ethane C3H8 propane C4H10 butane C5H12 pentane (cont. with Greek prefixes) Alkenes C==C CnH2n C2H4 ethene C3H6 propene etc. Alkynes CC CnH2n-2 C2H2 ethyne (acetylene) C3H4 propyne etc. Nomenclature: 1) find longest continuous chain name for skeleton 2) branches = alkyl groups (methyl, ethyl, propyl, etc.) or other functional groups 3) position alkyl groups to minimize numbers indicating their position 4) more than one substituent of a kind: di, tri, tetra, penta, etc. The nature of functional groups is given by using specific pre- and suffixes: Compound halides alcohols aldehydes ketones ethers acid esters amines amides Notes: Functional group -F, -Cl, -Br, -I R-OH R-CHO R1-CO-R2 R1-O-R2 R-COOH R1-COO-R2 R-NH2 R-CO-NH2 Pre-/suffix “fluoro-“, “chloro-“, “bromo-“, “iodo-“ “-ol” “-al” “-one” “(Rest 1)(Rest 2) ether” (carbon chain) + “-oic acid” (Rest 2) (Rest 1) + “oate” “amino-“ “-amide” 105 APPENDIX E: MISCELLANEOUS SI units: Basic units are by definition and use an arbitrary standard. Basic quantity Name Symbol/Unit Mass kilogram kg Length meter m Time second s Electric current Ampere A Luminous intensity Candela Cd Amount of matter Mole mol Derived quantity Name Symbol/Unit Force Newton 1 N = 1 kg m s-2 work, energy Joule 1 J = 1 N m = 1 kg m2 s-2 Pressure Pascal 1 Pa = 1 N m-2 = 1 kg m-1 s-2 Density –– g cm-3 106 Constants and other assorted data Constant Atomic mass unit Symbol amu Numerical Value 1.66055 x 10-24 g Application Stoichiometry Avogadro’s Number NA Boltzmann’s constant k 1.381 x 10-23 J mol-1 K-1 Thermodynamics, entropy Gas Constant R 8.314 J mol-1 K-1 0.082 l atm mol-1 K-1 Thermodynamics, gas problems Faraday constant F 96,000 C mol-1 Electrochemistry 6.0221367 x 10-23 10-31 mol-1 electron mass me 9.110 x proton mass mp 1.673 x 10-27 kg neutron mass mn 1.675 x 10-27 kg Planck’s constant h 6.626 x 10-34 J s-1 108 Stoichiometry kg m s-1 Quantum mechanics speed of light c 2.9979 x Quantum mechanics Standard Pressure 101.3 kPa = 1 atm Thermodynamics, gas problems Standard Temperature 273 K = 0 oC Thermodynamics, gas problems 107 Prefixes Fractions and multiples of 10 Giga G Mega M Kilo k Hecto h deci d milli m micro nano n pico p Numerical prefixes from 1 to 10 Mono1 Di2 Tri3 Tetra4 Penta5 Hexa6 Hepta7 Octa8 Nona9 Deca10 Logarithms log 0 log 1 log 2 log 3 log 4 log 5 log 6 log 7 log 8 log 9 = = = = = = = = = = (not defined) 0 0.3 0.5 0.6 0.7 0.8 0.85 0.9 0.95 Logarithmic rules log (ab) = log a + log b log (a/b) = log a – log b log (ab) = b log a 109 106 103 102 10-1 10-3 10-6 10-9 10-12 1,000,000,000 1,000,000 1,000 100 0.1 0.001 0.000 001 0.000 000 001 0.000 000 000 001 108 Notes: 109 APPENDIX F: HELPFUL WEBSITES Collegeboard — Advanced Placement Courses (general information) http://apcentral.collegeboard.com/ Common misconceptions in chemistry http://www.princeton.edu/~lehmann/BadChemistry.html Various chemistry links (might be too much) http://www.liv.ac.uk/Chemistry/Links/links.html History of chemistry (biographies of chemists) http://www.woodrow.org/teachers/ci/1992/ Nomenclature of oxyacids http://chem01.usca.sc.edu/chemistry/genchem/nomen.htm Results that any search engine gives you for the search term “AP Chemistry” Notes: