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ap chem seminar notes odds and ends: fractional reaction orders exist when: For example: find the order of the reaction for A] CH4 b] Cl2 In the mechanism: St 1 St 2 St 3 St 4 St 5 Cl2 ↔ 2 Cl CH4 + Cl CH3 + HCl CH3 + Cl2 CH3Cl + Cl CH3Cl + Cl CH2Cl2 + H H + Cl HCl fast equilibrium slow fast fast fast What this shows is that we can solve for the concentration of an ________________ by assuming that an _________________ is established in the __________ step. There is another way to classify chemical reactions. Reaction type 1: Synthesis Description General formula: Example Reaction type 2: Decomposition Description General formula: Example Note: hydroxides from a metal oxide and water when decomposed Carbonates form a metal oxide and ________ when decomposed. Reaction type 3: Displacement Description General formula: 3 types of displacement: Examples Reaction type 4: double displacement Description General formula: Example Note: When H2CO3 forms from an acid and a carbonate compound, the H2CO3 decomposes into: ___________ and ______________. Reaction type 5: combustion Description General formula: Example Reaction type 6: Complex ion equilibria: (see chart, p. 649 Brown LeMay) TTP 1: calculate the concentration of the silver ion present in solution at equilibrium when concentrated ammonia is added to a 0.010 M solution of silver nitrate to give an equilibrium concentration of ammonia = 0.20 M. Neglect the small volume change that occurs on addition of ammonia. TTP 2: (#108 p. 786 Zumdahl) A solution is formed by mixing 50.0 mL of 10.0 M NaX with 50.0 mL of 2.0 x 10-3 M CuNO3. Assume that Cu(I) forms complex ions with X- as follows: Cu+ + X- ↔ CuX K1 = 1.0 x 102 CuX + X- ↔ CuX2- K2 = 1.0 x 104 CuX2- + X- ↔ CuX3-2 K3 = 1.0 x 103 Calculate the following concentrations at equilibrium: a] CuX3-2 b] CuX2- c] Cu+ Predict the products for: (syn) 1] Al + Cl2 (comb) 2] C3H8 + O2 (dec) 3] CuBr (dis) 4] AlCl3 + F2 (dec) 5] Mg(OH)2 (dec) 6] MgCO3 (dis) 7] AlCl3 + K (double dis) 8] AlCl3 + AgNO3 (dis) 9] Al + HF (dec) 10] NaCl Buffered solution: chapter 15 Zumdahl A buffered solution is one that: Buffers contain: Derivation of Henderson-Hasselbach equation: TTP 1: A buffered solution contains 0.10 M acetic acid and 0.25 M sodium acetate. A] Find it’s pH b] find the pH after 0.010 moles of gaseous HCl is added to 1.0 L of the buffer solution c] find the pH after 0.010 moles of solid NaOH is added to 1.0 L of the buffer solution Buffer capacity: Selecting an ideal buffer: Titrations of strong acids and strong bases: TTP 2: calculate the pH when the following quantities of 0.100 M NaOH have been added to 50.0 mL of 0.100 M HCl solution (p. 633 Brown-LeMay): a] 49.00 mL B] 49.90 mL c] 50.10 mL d] 51.00 mL Titration of strong base with a weak acid: Since the ______________ _______ of a weak acid affect the pH, these problems are a bit more complicated. TTP 3: Calculate the pH in the titration of acetic acid by NaOH after 30.0 mL of 0.100 M NaOH has been added to 50.0 mL of 0.100 M acetic acid (p. 636 Brown-LeMay) Titration curves for weak acid with strong base Titration curve for weak base with a strong acid: Note: at the half-way point of the titration: Note: the pH is NOT ____ at the equivalence point for any strong acid/base with weak acid/base titration. For weak acid, strong base titration, the pH at the equivalence point is _________ 7. For a weak base, strong acid titration, the pH at the equivalence point is ___________ 7. Titrations of polyprotic acids: See curve of 25.0 mL of 0.10 M Na2CO3 with 0.10 M HCl (p. 641 Brown-LeMay) Acid-base indicators (p. 756 Zumdahl) The color change of an indicator takes place when: [In-] / [HIn] = TTP #1 Bromthymol blue, an indicator with a Ka value of 1.0 x 10-7, is yellow in its HIn form and blue in its In- form. Suppose we put a few drops of this indicator in a strongly acidic solution. If the solution is then titrated with NaOH, at what pH will the indicator color change first be visible? (p. 753 Zum) Chapter 6 – Thermochemistry Nature of Energy: State functions: _________ is a state function, but ________ and ________ are not state functions. First Law: Internal energy: Enthalpy and calorimetry Specific heat capacity Molar heat capacity Hess’s Law: Standard Enthalpies of formation TTP 1: For the reaction: CH4(g) + 2O2(g) CO2(g) +2H2O(g) H = -891 kJ Calculate the enthalpy change for each of the following: A] 1.00 g of methane is burned in excess oxygen B] 1.00 x 103 L of methane gas at 740. torr and 25oC is burned in excess oxygen. (p. 284 #36 Z) TTP 2: It takes 78.2 J to raise the temperature of 45.6 g lead by 13.3oC. Calculate the specific heat capacity and molar heat capacity of lead. (p. 284 #39 Z) TTP 3: In a coffe-cup calorimeter, 50.0 mL of 0.100 M AgNO3 and 50.0 mL of 0.100 M HCl are mixed to yield the reaction: Ag+ + Cl- AgCl The two solutions were initially at 22.60oC and the final temperature is 23.40oC. Calculate the heat that accompanies this reaction in kJ/mol of AgCl formed. Assume that the combined solution has a mass of 100.0 g and has a specific heat capacity of 4.18 J/oCg. (p. 284 #45) TTP 4: Given the following data: S + 3/2 O2 SO3 H = -395.2 kJ 2SO2 + O2 2SO3 H = -198.2 kJ calculate H for the reaction: S + O2 SO2 (p. 285 #53 Z) TTP 5: Using appendix 4, find Ho for C2H5OH(l) + 3O2(g) 2CO2(g) + 3H2O(g) Chapter 10 – Liquids and Solids Intermolecular Forces: Dipole-Dipole forces Typically these are about _____ percent as strong as __________ or ________ bonds. Hydrogen bonding: (see p. 454 Z) London Dispersion forces: Solids: two big categories: _______________ and ____________ . Smallest repeating unit of a lattice of a crystalline solid is called the ______ _____. p. 459. Bragg equation: p. 463 table 10.3 Metallic crystals pack together as either: The ccp structure has a unit cell that is a: Note: each corner of a unit cell is _____ of an atom. There are ____ corners of a unit cell. The center of each face contains _____ of an atom. There are ____ faces of a unit cell If there is a body-centered cubic unit cell, there is ___ atom in the center. See p. 465 Z. Bonding model for metals: Network atomic solids: Molecular solids: These have _______ forces between atoms in the molecule, but _______ forces between molecules. Examples: Ionic solids: Typically, the larger ions are the ________ and they are packed in either an hcp or ccp arrangement. The smaller __________ fit into the holes among the closest packed anions. There are three holes: In order of size: Table 10.7 p.483 Z. Vapor pressure and changes in state Eq. 10.4 p. 486 Z. Changes of state: Phase Diagrams: TTP 1: Identify the most important type of interparticle forces present in the solids of each of the following: (p. 504 # 35 Z.) A] Ar B] HCl c] HF d] CaCl2 e] CH4 f[ CO g] NaNO3 TTP 2] Predict which substance in each of the following pairs would have the greater intermolecular attraction: (p. 504 #37 Z.) A] CO2 or OCS b] PF3 or PF5 c] SF2 or SF6 d. SO3 or SO2 TTP 3] Iridium (Ir) has a face-centered cubic unit cell with an edge length of 383.3 pm. Calculate the density of solid Ir. (p. 505 #51 Z.) TTP #4] What type of solid will each of the following substances form? (p. 506 #71 Z.) A] CO2 b] SiO2 c] Si d] CH4 e] Ru f] I2 G] KBr h] H2O i] NaOH j] U k] CaCO3 l] PH3 TTP #5] How much energy does it take to convert 0.500 kg ice at -20.oC to steam at 250oC? The specific heat of ice is 2.1 J/oCg ; water is 4.2 J/oCg; steam is 2.0 J/oCg; Hvap = 40.7 kj/mol; Hfus = 6.02 kj/mol Chapter 16: Spontaneity, Entropy, and Free Energy Spontaneous process: Entropy: Entropy and the second law of thermo: The effect of temperature on spontaneity: Free Energy: The third law of thermo: The dependence of Free Energy on Pressure: Free energy and equilibrium: The temperature dependence of K TTP 1] For mercury, the enthalpy of vaporization is 58.51 kj/mol and the entropy of vaporization is 92.92 j/Kmol. What is the normal boiling point of Hg? TTP 2] For ammonia, the enthalpy of fusion is 5.65 kj/mol and the entropy of fusion is 28.9 j/Kmol. A] Will NH3(s) spontaneously melt at 200. K? b] What is the approximate melting point of ammonia? TTP 3] Predict the sign of So for each of the following changes: A] AgCl(s) Ag+ + Clb] 2H2(g) + O2(g) 2H2O(l) C] H2O(l) H2O(g) TTP 4] Predict the sign of So and then calculate So for each of the following reactions: A] 2H2S(g) + SO2(g) 3Srhombic(s) + 2H2O(g) B] 2SO3(g) 2so2(g) + O2(g) C] Fe2O3(s) + 3H2(g) 2Fe(s) + 3H2O(g) TTP #5] For the reaction at 298 K, 2NO2(g) ↔ N2O4(g) The values of Ho and So are -58.03 kj and -176.6 j/K respectively. What is the value of Go at 298 K? Assuming that Ho and So do not depend on temperature, at what temperature is Go = 0? Is Go negative above or below this temperature? TTP #6] For the reaction: SF4(g) + F2(g) SF6(g) The value of Go is -374 kj. Use this value and data from appendix 4 to calculate the value of Go for SF4(g) TTP 7] Consider the autoionization of water at 25oC, H2O(l) H+ + OH- ; Kw = 1.00 x 10-14 A] calculate Go for this reaction at 25oC b] at 40.oC, Kw = 2.92 x 10-14 . Calculate Go at 40.oC. p. 829 #’s 30, 31, 33, 33, 37, 46, 51, 61. Chapter 17 – electrochemistry Galvanic Cells: In a galvanic cell, _________________ energy is converted into ______________ Energy. In a galvanic cell, the emf (cell potential - €cell) is________ than zero. For example, if copper and zinc are used, the diagram for a galvanic cell is: Line notation Cell potential, electrical work, and free energy Faraday(F) G = -nF€o Concentration cells Nernst equation Ion-selective electrodes: Calculations of Equilibrium constants for redox reactions Batteries Fuel cells: Corrosion Electrolysis: TTP 1] Sketch the galvanic cells based on the following overall reactions. Show the direction of electron flow and identify the anode and cathode. Assume all concentrations are 1.0 M and all pressures are 1.0 atm. Determine the value of €o. Write the line notation. Find the value of K. (Z #’s 25, 27, 31, 69) A] Cu+2 + Mg(s) ↔Mg+2 + Cu(s) B] Cr+3 + Cl2 ↔ Cr2O7-2 + Cl- TTP 2] Place in order of increasing strength as oxidizing agents (all under standard conditions) Cd+2, IO3- , K+, H2O, AuCl4- , I2 TTP 3] Place in order of increasing strength as reducing agents (all under standard conditions) Cr+3 , H2, Zn, Li, F- , Fe+2 TTP 4] A galvanic cell is based on the following half-reactions at 25oC: Ag+ + e- Ag H2O2 + 2H+ + 2e- 2H2O Predict whether €cell is larger or smaller than €ocell for the following cases: A] [Ag+] = 1.0 M , [H2O2] = 2.0 M , [H+] = 2.0 M B] [Ag+] = 2.0 M , [H2O2] = 1.0 M , [H+] = 1.0 x 10-7 M C] for each case, find €cell (Z. #’s 55, 57) TTP 5] What volume of F2 gas, at 25oC and 1.00 atm, is produced when molten KF is electrolyzed by a current of 10.0 A for 2.00 hours? What mass of potassium metal is produced? At which electrode does each reaction occur? (Z #85) TTP 6] Electrolysis of a molten metal chloride, (MCl3) using a current of 6.50 A for 1397 seconds deposits 1.41 g of the metal at the cathode. What is the metal? Chapter 18 – The representative elements: Groups 1A – 4A The oxides of the 2A elements form bases in water, except: Carbon dioxide is very different from the oxide of Silicon because: Fluorine, F2, has a smaller electron affinity that Cl2 because: Lithium is a stronger reducing agent than the elements below it because: Lithium reacts more slowly with water than the elements below it b/c: The bonds aluminum forms with nonmetals are significantly___________________, causing it to have properties of both ____________ and bases. Substances that can act as both an acid and a base are said to be: Reaction: Carbon has 3 allotropes: Chapter 19 – The Representative Elements: Groups 5A – 8A All the 5A can form 5 covalent bonds except: This is because: Nitrogen fixation: chart p. 925 Reactions of nitrogen: Chemistry of phosphorus Chemistry of Sulfur: Why is HF a weak acid, but HCl is a strong acid? Disproportionation: Chapter 20 – Transition metals and Coordination Chemistry Ions of elements #’s 21-30: Table 20.5 Table 20.6 Coordination number: Table 20.12 Table 20.13 Rules for naming coordination compounds: (p. 979) The crystal field model: Strong field and weak field cases: Chapter 21: Nuclear chemistry Figure 21.1 Table 21.2 Chapter 22: Organic chemistry Alkanes Saturated vs. unsaturated Nomenclature: Alkenes and alkynes: Aromatic hydrocarbons: Common functional groups: table 22.5 Alchohols: Aldehydes and ketones: Carboxylic acids and esters amines polymers