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Introduction to Chem II

Introduction to Chem II. Instructors Course Objectives Course Topics Laboratory Exercises Course Website Today’s Agenda Syllabus. Course Objectives. Review some familiar topics Investigate some of these topics at a more in-depth level Model sound pedagogy

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Introduction to Chem II

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  1. Introduction to Chem II • Instructors • Course Objectives • Course Topics • Laboratory Exercises • Course Website • Today’s Agenda • Syllabus

  2. Course Objectives • Review some familiar topics • Investigate some of these topics at a more in-depth level • Model sound pedagogy • Obtain hand-on practice with Venier Data Collection • Show some effective demonstrations

  3. Course Topics • Stoichiometry • Calorimetry • Equilibrium • Solubility • Acid-base chemistry • Redox chemistry • Thermochemistry

  4. 5 Lab Exercises • A calorimetry experiment using a temperature probe • Solubility using a Ca ion selective electrode • Equilibrium constant using a Colorimeter • Acid-base titration • Ag Ion Indicator electrode

  5. Course website http://alpha.chem.umb.edu/chemistry/bpschemII/ Syllabus Lab experiments Course notes Homework solutions

  6. Today’s Agenda • Take a 2 hr exam • Paperwork, surveys • Lunch • Lecture; g/mol, Classification of reactions, Stoichiometry, LR, Energetics of Reactions • Lab Lecture; Calorimetry • Lab Experiment 1 • Early start on HW

  7. Investigating Stoichiometry using Calorimetry Experiment 1

  8. Heat of Reaction - DH • At constant pressure – most lab experiments • aA + bB → products DH/mol A • . DH = q (heat produced or absorbed)

  9. Calorimetry • Method of measuring the heat of reaction • Calorimeter-coffee cup • q = cmDT • c is the specific heat [J/(g ºC)] of solution • m = mass of solution • .DT is change in temperature • .DT is directly proportional to the heat of reaction

  10. The experiment • Mix reactants in different molar ratios • Predict the stoichiometry of the reaction from the ratio that gives the maximum temperature increase

  11. Example of the Experiment • 1 to 1, A + B → products • Mixing molar ratios • Constant total volume - cmDT

  12. 1:1 Stoichiometry (mol ratio A/B)

  13. Example 2 • 2 to 1, 2A + B → products

  14. 2:1 Stoichiometry (mol ratio A/B)

  15. Example 3 • 3 to 1, 3A + B → products

  16. 3:1 Stoichiometry (mol ratio A/B)

  17. Determining the DHm • . DH = cmDT = (4.4 J/gC)*(50 g)*(36) = 7920 J • mol A reacted = 18.8 mmol A • .DHm = DH/(mol A reacted) = (7920)/(.0188 mol) = 421276 J/mol = 421 kJ/mol

  18. Products • Thiosulfate is a classic reducing agent • 2S2O32-↔ S4O62- + 2e- • Cl- is the product of the reduction of OCl- • Write a balanced redox equation • Step 1: determine half reactions. • Step 2 Make the reduction half reaction and oxidation half reaction have the same number of electrons by multiply reactions by common denominator • Step 3: Add reactions

  19. OCl- + H2O + 2 e- ↔ Cl- + 2OH- • 2S2O32-↔ S4O62-+ 2e- • ________________________ • OCl- + H2O + 2S2O32-→ Cl- + 2OH- + S4O62-

  20. Solubility of CaSO4 Experiment 2

  21. Goals • Determine the solubility of CaSO4 in three different solution • Saturated CaSO4 in H2O • Saturated CaSO4 in 0.10 M KNO3 • Saturated CaSO4 in 0.10 M Na2SO4 • Compare and rationalize the results

  22. Major concepts • Solubility Product Constants and saturated solution • LeChatlier’s principle and the common ion effect • Effect of ionic strength and ion activities on Ksp • Ion Selective Electrodes

  23. Ksp of CaSO4 • CaSO4(s)↔ Ca2+ + SO42- • Ksp(CaSO4) = [Ca2+][SO42-] = 2.4∙10-5

  24. Saturated solution in water • Add several grams of CaSO4 to 1 L of water • Shake and mix for weeks • Allow CaSO4 that did not dissociate to settle to bottom • Ksp(CaSO4) = [Ca2+][SO42-] = 2.4∙10-5 = x2 [Ca2+] = 5.0∙10-3 M

  25. Saturated solution in 0.10 M Na2SO4 • Add several grams of CaSO4 to 1 L of 0.10 M Na2SO4 • Common Ion effect • Ksp(CaSO4) = [Ca2+][SO42-] = 2.4∙10-5 = x(x+0.10) Assume x <<< 0.10 x = 2.4∙10-4 M [Ca2+] = 2.4∙10-4 M

  26. Saturated solution in 0.10 M KNO3 • Activities • Ksp(CaSO4) = ACa2+ASO42- = [Ca2+]gCa2+[SO42-]gSO42- = 2.4∙10-5 • Activity coefficient (g) is dependent on the ionic strength of the solution, and the size and charge of the ion. It is a number between 0 and 1. At very low ionic strength, g approaches 1

  27. Ionic strength • A measure of the concentration of ions in solution m = ½ ∑ cizi2 Sat. solution in 0.10 M KNO3 m = ½ ([K+](+1)2 + [NO3-](-1)2 + [Ca2+](2+)2 + [SO42-](-2)2) = 0.12 M gCa2+@m=0.12 =

  28. Take home message • The common ion effect decreases the solubility by over an order of magnitude • At high ionic strengths, solubility increases slightly ( by a factor of 1.5 -5).

  29. Ion Selective Electrode • A probe that consists of two reference electrodes connected electrically through a specific type of salt bridge through the solution being measured. • The salt bridge is a membrane that specifically binds the ion of interest • A junction potential develops at this membrane that is proportional to the concentration of the ion of interest

  30. voltmeter Cathode Reference electrode Ag/AgCl, sat. KCl Anode Reference electrode Ag/AgCl, sat. KCl solution Ion selective membrane Engineer this whole set-up in one probe pH meter Ca2+ selective electrode

  31. Response of Ca2+ Selective Electrode • Ecell = constant + 29.58 logACa2+

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