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Chapter 4 Part 2

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norman-white
27 diapositives · 560 Views · Nov 19, 2014

Chapter 4 Part 2. CHM 108 Suroviec Spring 2014. I. Solution Stoichiometry. According to the following reaction, how many moles of Fe(OH)2 can form from 175.0 mL of 0.227 M LiOH solution? Assume that there is excess FeCl2. FeCl2(aq) + 2 LiOH(aq) → Fe(OH)2(s) + 2 LiCl(aq)

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  1. Chapter 4Part 2 CHM 108 Suroviec Spring 2014

  2. I. Solution Stoichiometry • According to the following reaction, how many moles of Fe(OH)2 can form from 175.0 mL of 0.227 M LiOH solution? Assume that there is excess FeCl2. FeCl2(aq) + 2 LiOH(aq) → Fe(OH)2(s) + 2 LiCl(aq) • Determine the number of grams H2 formed when 250.0 mL of 0.743 M HCl solution reacts with 3.41 x 1023 atoms of Fe according to the following reaction. 2 HCl(aq) + Fe(s) → H2(g) + FeCl2(aq)

  3. II. Aqueous Solution and Solubility • Consider salt dissolving in water and sugar dissolving in water.

  4. A. Electrolyte • The way ionic compounds vs. molecular compounds dissolve in water shows the difference between types of solution.

  5. A. Electrolyte • Electrolytes – ions that act at charge carriers • Solutes that completely dissociate into ions are called strong electrolytes

  6. B. Solubility of Ionic Compounds • Most ionic compounds when dissolved in water the solute breaks into ions. • Not true for all ionic compounds

  7. Determine the insoluble compounds • AgCl • NaNO3 • PbCl2 • Ba(OH)2

  8. III. Precipitation Reactions • Precipitate: insoluble solid that separates from solution where no solid existed before reaction • Hard water contains Ca2+ and Mg2+ • Laundry detergent contains Na2CO3

  9. Examples • silver nitrate and potassium chloride • lead (II) nitrate and potassium chromate • potassium chromate and silver nitrate • sodium carbonate and copper (II) chloride • nickel (II) chloride and potassium hydroxide

  10. IV. Molecular and Ionic Equations A. Molecular Equations • Consider the following equation: CaCl2 (aq) + Na2SO4 (aq)  CaSO4 (s) + 2NaCl (aq)

  11. B. Ionic Equations • In these equations, see that some of the ions are present on both sides of the arrow

  12. Example Given: 2AgNO3(aq) + MgCl2(aq)  2AgCl (s) + Mg(NO3)2 (aq) What is the ionic equation? Net Ionic?

  13. Write NET ionic equations • AlCl3 (aq) + Na3PO4 (aq) • lead (II) nitrate and potassium chloride

  14. H2SO4 (aq) H+ (aq) + HSO4- (aq) HSO4- (aq) H+ (aq) + SO42- (aq) IV. Acid and Base Reactions • Bronstead definition of acid: proton donor • Bronstead definition of base: proton acceptor • Acids HCl (aq) + H2O (l) → H3O+ (aq) + Cl- (aq)

  15. B. Bases • Proton acceptors • Strong bases ionize completely to OH- • NaOH (s) → • Ca(OH)2 → • Weak bases ionize only partially • NH3 (aq) + H2O ⇌

  16. acid + base salt + water HCl (aq) + NaOH (aq) C. Reactions of Acids and Bases • Neutralization

  17. C. Reactions of Acids and Bases 2. Weak Acid/Base reactions CH3COOH + NaOH

  18. V. pH • Concentration scale for acids and bases Vinegar: [H+] = 1.610-3M Pure Water: [H+] = 1.010-7M Ammonia: [H+] = 1.010-11M pH = -log[H+] Determine the pH of the above. What is the trend of acids and bases?

  19. VI. Acid-Base Titrations • Commonly used to determine the concentration of a dissolved species or its percentage in a mixture • Titration • Measuring the volume of a standard solution (known concentration) needed to react with a measured quantity of a sample • Titrant (in the buret) • Analyte (in the Erlenmeyer flask)

  20. VI. Acid-Base Titrations • Equivalence point is where the number of moles of acid equals the number of moles of base • The endpoint is indicated by a color change in the acid-base indicator

  21. Example • What volume (in mL) of a 1.420 M NaOH solution is required to titrate 25.00 mL of a 4.50 M H2SO4 solution? • What volume (in mL) of 0.955 M HCl is required to titrate 2.152g of Na2CO3 to the equivalence point?

  22. VI. Redox Reactions • Oxidation Numbers • Needed when we are looking at reactions between 2 nonmetals. • The oxidation number of an atom in a compound is the “charge” that it would have it all shared electrons were assigned to the atom with higher electronegativity.

  23. III. Oxidation-Reduction Reactions • Short name: Redox reactions • Electron exchange • Oxidation is a loss of electrons • Reduction is a gain of electrons

  24. III. Redox Reactions Fe (s)  Fe? (aq) + 2e- 2H? (aq)+ 2e- H2 (g)

  25. Examples A. Fe3+ (aq) + H2 (g) Fe2+ (aq) + H+ (aq) B. Au (s) + F2 (aq) F- (aq) + Au3+(aq) • Break into 1/2 reactions • Mass balance 1/2 reactions • Combine and check for neutrality and check again for mass balance