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Lecture #26 What’s on the Final?. Chemistry 142 B Autumn Quarter, 2004 J. B. Callis, Instructor. General. Date as Announced in Syllabus 44 Multiple-Choice Questions - Chapters 1-8 22 questions from chapters 1-5 22 questions from chapters 6-8

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lecture 26 what s on the final

Lecture #26What’s on the Final?

Chemistry 142 B

Autumn Quarter, 2004

J. B. Callis, Instructor

general
General
  • Date as Announced in Syllabus
  • 44 Multiple-Choice Questions - Chapters 1-8
  • 22 questions from chapters 1-5
  • 22 questions from chapters 6-8
  • Practice exam available on the class web site – key on class bulletin board.
chapter 2 dalton s atomic hypothesis and beyond
Chapter 2Dalton’s Atomic Hypothesisand Beyond
  • Law of multiple proportions as explained by Dalton
  • Naming compounds
problem 26 1 multiple proportions
Problem 26-1: Multiple Proportions

A chemist studies three oxides of iodine, and finds their % oxygen as follows: 20.14 %, 23.97% and 22.10 %. (a) Calculate the mass of oxygen per gram of iodine in each compound. Express the result as a ratio. Then form the ratio of ratios by dividing by the ratio of the first compound. (b) How do the numbers in part (a) support Dalton’s atomic theory?

chapter 3 stoichiometry
Chapter 3 - Stoichiometry
  • Convert between moles, mass and number of molecules.
  • Determine empirical formula from elemental composition.
  • Balance a chemical reaction using the atom balance equations.
  • Solve stoichiometric problems involving a limiting reagent using the reactant ratio method.
slide7

Problem 26-2: Bromoform is 94.85% bromine, 0.40% hydrogen and 4.75% carbon by mass. Determine its empirical formula.

slide8

Problem 26-3: Balance the following chemical equation:

x NaCl + y SO2 + z H2O + w O2 => u Na2SO4 + v HCl

slide9

Problem 26-4: Sulfuric acid (H2SO4) forms in the following reaction:

2 SO2 + O2 + 2 H2O => 2 H2SO4

Suppose 400. g SO2, 175. g O2 and 125. g H2O are mixed and the reaction proceeds until one of the reactants is used up. Identify the limiting reactant and determine what masses of the other reactants remain.

chapter 4 chemical reactions
Chapter 4Chemical Reactions
  • Classify a given reaction and identify products such as precipitates.
  • Write the net ionic reaction.
  • Determine the oxidation number of an element in a compound.
  • Balance a redox equation using the atom balance and charge balance equations.
  • Perform stoichiometric calculations on solutions
slide11

Problem 26-5: Balance the following chemical equation:

x H+ + y H2O2 + z Fe2+ => u Fe3+ + v H2O

chapter 5 the gas phase
Chapter 5 – The Gas Phase
  • Use the ideal gas law to compute density from other independent variables, e.g. pressure, temperature.
  • Work with partial pressures.
  • Calculate the amount of gas collected over water.
  • Use the ideal gas law in stoichiometric problems.
slide13

Problem 26-6: A 20.6 L sample of air is collected in Greenland at –20. oC at a pressure of 1.01 atm and forced into a 1.05 L bottle for shipment to Europe for analysis. Compute the pressure inside the bottle as it is opened in the laboratory at 21. oC.

chapter 6 chemical equilibrium
Chapter 6 – Chemical Equilibrium
  • Write the equilibrium constant for a balanced chemical reaction.
  • Solve equilibrium problems given starting concentrations and K.
  • Apply Le Chatelier’s principle to determine direction of shift of reaction in response to a change in reaction conditions.
slide15

Problem 26-7: At 5000 K, even the nitrogen molecule (N2) breaks down (into 2 atoms of nitrogen). At this temperature, when the total pressure of nitrogen is 1.00 atm, N2(g) is a = 0.65% dissociated at equilibrium: N2(g) = 2 N(g). Compute the equilibrium constant at 5000K.

chapter 7 acid base equilibria
Chapter 7 – Acid/Base Equilibria
  • Understand the pH scale
  • Calculate the pH of strong, weak and very weak acids and bases
slide18

Problem 26-8: What is the pH of 0.15 M methylammonium bromide, CH3NH3Br? (Kb of CH3NH2 = 4.4 x 10-4

chapter 8 applications of aqueous equilibria
Chapter 8 – Applications of Aqueous Equilibria
  • Calculate the pH of buffer solutions
  • Give recipes for making up buffers
  • Calculate pH at different points in a titration curve
  • Perform solubility calculations
  • Understand the basic ideas of complex ion equilibria
slide20

Problem 9: You have at your disposal an ample quantity of a solution of 0.0500 M NaOH and 500 mL of a solution of 0.100 M formic acid (HCOOH, Ka = 1.77 x 10-4). How much of the NaOH solution should be added to the acid solution to produce a buffer of pH 4.00?

Ans: Use the base to produce a sufficient amount of the formate ion to provide a buffer of the desired pH. Do the problem at the Henderson-Hasselbalch level, ignoring the ionization of water. Allow for dilution of the original acid solution with the base, just as in a titration.

The relevant reaction is:

slide23

Problem 10: A saturated solution of Mg(OH)2 at 25oC is prepared by equilibrating solid Mg(OH)2 with water. Concentrated NaOH is then added until the solubility of Mg(OH)2 is 0.001 times that in H2O alone. (Ignore the change in volume resulting from the addition of NaOH.) The solubility product Ksp of Mg(OH)2 is 1.2 x 10-11 at 25oC. Calculate the concentration of hydroxide ion in the solution after the addition of the NaOH.

Ans: First calculate the solubility of Mg(OH)2 in water. Then calculate the concentration of [Mg2+] after addition of OH- put this into the mass action expression and solve for [OH-].

slide24

Problem 10 (cont.):

Addition of OH- shifts the equilibrium to the left and the concentration of Mg2+ must diminish to maintain the mass-action expression at a constant value. After the addition of base, the new concentration of Mg2+ is

answers to questions from lecture 26
Answers to Questions from Lecture 26
  • 1, 1.25, 1.125
  • CHBr3
  • 4 NaCl + 2 SO2 + 2 H2O + O2 => 2 Na2SO4 + 4 HCl
  • 2 H++H2O2 + 2 Fe2+=> 2 Fe3+ + 2 H2O
  • 23.0 atm
  • K = 1.67 x 10-4
  • pH = 5.73
  • 639 mL
  • 9.1 x 10-3