1. Reactions in Aqueous Solution Chapter 4 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

2. 4.1 A solution is a homogenous mixture of 2 or more substances The solute is(are) the substance(s) present in the smaller amount(s) The solvent is the substance present in the larger amount SolutionSolventSolute Soft drink (l) Air (g) Soft Solder (s) H2OH2O N2N2 Pb Sugar, CO 2 O 2, Ar, CH 4 Sn Examples :

3. An electrolyte is a substance that, when dissolved in water, results in a solution that can conduct electricity. A nonelectrolyte is a substance that, when dissolved in water, results in a solution that does not conduct electricity. nonelectrolyte weak electrolyte strong electrolyte 4.1

4. Strong Electrolyte – 100% dissociation NaCl (s) Na + (aq) + Cl - (aq) H2OH2O Weak Electrolyte – not completely dissociated CH 3 COOH CH 3 COO - (aq) + H + (aq) Conduct electricity in solution? Cations (+) and Anions (-) 4.1

5.  H2OH2O  A Polar molecule

6. Hydration is the process in which an ion is surrounded by water molecules arranged in a specific manner.   H2OH2O It’s a battle. Charge!

7.   H2OH2O   H2OH2O   H2OH2O   H2OH2O

8. Nonelectrolyte does not conduct electricity? No cations (+) and anions (-) in solution C 6 H 12 O 6 (s) C 6 H 12 O 6 (aq) No dissociation H2OH2O

9. Examples:

10. Precipitation Reactions Precipitate – insoluble solid that separates from solution 4.2 PbI 2

11. Memorization time…

12. 1. All alkali metals (Group 1A) compounds are soluble. Solubility Rules for Common Ionic Compounds In water at 25 o C 2. All ammonium (NH 4 + ) compounds are soluble. 3. All nitrate (NO 3 - ), chlorate (ClO 3 - ), and perchlorate (ClO 4 - ) compounds are soluble. 4. Most hydroxides (OH - ) are insoluble. The exceptions are barium hydroxide [Ba(OH) 2 ], which is very soluble, calcium hydroxide [Ca(OH) 2 ], which is slightly soluble, and previous examples. 5. Most compounds containing chlorides (Cl - ), bromides (Br - ), or iodides (I - ), are soluble. The exceptions are those containing Ag +, Hg 2 2+, and Pb 2+. 7. All carbonates (CO 3 2- ), phosphates (PO 4 3- ), and sulfides (S 2- ) are insoluble, excepting previous rules. 6. Most sulfates (SO 4 2- ) are soluble, excepting previous rules. Calcium sulfate [CaSO 4 ] and silver sulfate [Ag 2 SO 4 ] are slightly soluble. Barium sulfate [BaSO 4 ], mercury (II) sulfate [HgSO 4 ], and lead sulfate [PbSO 4 ] are insoluble.

14. Solubility Rules for Common Ionic Compounds In water at 25 0 C

15. Precipitation Reactions Precipitate – insoluble solid that separates from solution molecular equation ionic equation net ionic equation Pb 2+ + 2NO 3 - + 2Na + + 2I - PbI 2 (s) + 2Na + + 2NO 3 - Na + and NO 3 - are spectator ions PbI 2 Pb(NO 3 ) 2 (aq) + 2NaI (aq) PbI 2 (s) + 2NaNO 3 (aq) precipitate Pb 2+ + 2I - PbI 2 (s) 4.2

16. Writing Net Ionic Equations 1.Write the balanced molecular equation. 2.Write the ionic equation showing the strong electrolytes 3.Determine precipitate from solubility rules 4.Cancel the spectator ions on both sides of the ionic equation AgNO 3 (aq) + NaCl (aq) AgCl (s) + NaNO 3 (aq) Ag + + NO 3 - + Na + + Cl - AgCl (s) + Na + + NO 3 - Ag + + Cl - AgCl (s) 4.2 Write the net ionic equation for the reaction of silver nitrate with sodium chloride.

17. Acids (according to Arrhenius) Have a sour taste. Vinegar owes its taste to acetic acid. Citrus fruits contain citric acid. React with certain metals to produce hydrogen gas. React with carbonates (CO 3 2- ) and bicarbonates (HCO 3 - ) to produce carbon dioxide gas. 4.3 Are electrolytes. Cause certain plant dyes to change color. For example, they turn litmus from blue to red.

18. Have a bitter taste. Feel slippery (they make water insoluble organic molecules into water soluble molecules). Many soaps contain bases. 4.3 Bases (according to Arrhenius) Are electrolytes. Cause certain plant dyes to change color. For example, they turn litmus from red to blue.

19. Arrhenius acid is a substance that produces H + (H 3 O + ) in water Arrhenius base is a substance that produces OH - in water 4.3

20. A Brønsted acid is a proton donor A Brønsted base is a proton acceptor acidbaseacidbase 4.3 A Brønsted acid must contain at least one ionizable proton!

21. Monoprotic acids HCl H + + Cl - HNO 3 H + + NO 3 - CH 3 COOH H + + CH 3 COO - Strong electrolyte, strong acid Weak electrolyte, weak acid Diprotic acids H 2 SO 4 H + + HSO 4 - HSO 4 - H + + SO 4 2- Strong electrolyte, strong acid Weak electrolyte, weak acid Triprotic acids H 3 PO 4 H + + H 2 PO 4 - H 2 PO 4 - H + + HPO 4 2- HPO 4 2- H + + PO 4 3- Weak electrolyte, weak acid 4.3

22. Neutralization Reaction acid + base salt + water HCl (aq) + NaOH (aq) NaCl (aq) + H 2 O H + + Cl - + Na + + OH - Na + + Cl - + H 2 O H + + OH - H 2 O 4.3 net ionic equation for strong acid / strong base reaction

23. Neutralization Reaction acid + base salt + water HCl (aq) + NH 3 (aq) NH 4 Cl(aq) H + + Cl - + NH 3 (aq) NH 4 + + Cl - H + + NH 3 (aq) NH 4 + 4.3 net ionic equation for strong acid / weak base reaction

24. Oxidation-Reduction Reactions (electron transfer reactions) 2Mg (s) + O 2 (g) 2MgO (s) 2Mg 2Mg 2+ + 4e - O 2 + 4e - 2O 2- Oxidation half-reaction (lose e - ) Reduction half-reaction (gain e - ) 2Mg + O 2 + 4e - 2Mg 2+ + 2O 2- + 4e - 2Mg + O 2 2MgO 4.4

25. Fig. 4.9a-c Iron (Fe) reacting with hydrochloric acid (HCl)

26. Fe (s) + 2HCl (aq) FeCl 2 (aq) + H 2 (g) Fe is oxidizedFe Fe 2+ + 2e - H + is reduced2H + + 2e - H 2 Fe is the reducing agent H + is the oxidizing agent 4.4

27. The Activity Series for Metals M + BC AC + B Displacement Reaction M is metal BC is acid or H 2 O B is H 2 Ca + 2H 2 O Ca(OH) 2 + H 2 Pb + 2H 2 O Pb(OH) 2 + H 2 4.4

29. Zn (s) + CuSO 4 (aq) ZnSO 4 (aq) + Cu (s) Zn is oxidizedZn Zn 2+ + 2e - Cu 2+ is reducedCu 2+ + 2e - Cu Zn is the reducing agent Cu 2+ is the oxidizing agent 4.4 Copper wire reacts with silver nitrate to form silver metal. What is the oxidizing agent in the reaction? Cu (s) + 2AgNO 3 (aq) Cu(NO 3 ) 2 (aq) + 2Ag (s) Cu Cu 2+ + 2e - Ag + + 1e - AgAg + is reducedAg + is the oxidizing agent Cu is oxidizedCu is the reducing agent

30. Oxidation number The charge the atom would have in a molecule (or an ionic compound) if electrons were completely transferred. 1.Free elements (uncombined state) have an oxidation number of zero. Na, Be, K, Pb, H 2, O 2, P 4 = 0 2.In monatomic ions, the oxidation number is equal to the charge on the ion. Li +, Li = +1; Fe 3+, Fe = +3; O 2-, O = -2 3.The oxidation number of oxygen is usually –2. In H 2 O 2 and O 2 2- it is –1. 4.4

31. 4.The oxidation number of hydrogen is +1 except when it is bonded to metals in binary compounds. In these cases, its oxidation number is –1. 6. The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on the molecule or ion. 5.Group IA metals are +1, IIA metals are +2 and fluorine is always –1. HCO 3 - O = -2H = +1 3x(-2) + 1 + ? = -1 C = +4 Oxidation numbers of all the elements in HCO 3 - ? 4.4

32. Fig. 4.10

33. NaIO 3 Na = +1 O = -2 3x(-2) + 1 + ? = 0 I = +5 IF 3 F = -1 3x(-1) + ? = 0 I = +3 K 2 Cr 2 O 7 O = -2K = +1 7x(-2) + 2x(+1) + 2x(?) = 0 Cr = +6 Oxidation numbers of all the elements in the following ? 4.4

34. Types of Oxidation-Reduction Reactions Combination Reaction A + B C S + O 2 SO 2 Decomposition Reaction 2KClO 3 2KCl + 3O 2 C A + B 00 +4-2 +1+5-2+10 4.4

35. Displacement Reaction A + BC AC + B Sr + 2H 2 O Sr(OH) 2 + H 2 TiCl 4 + 2Mg Ti + 2MgCl 2 Cl 2 + 2KBr 2KCl + Br 2 Hydrogen Displacement Metal Displacement Halogen Displacement Types of Oxidation-Reduction Reactions 4.4 0 +1+20 0+40+2 0 0

36. The Activity Series for Metals M + BC AC + B Displacement Reaction M is metal BC is acid or H 2 O B is H 2 Ca + 2H 2 O Ca(OH) 2 + H 2 Pb + 2H 2 O Pb(OH) 2 + H 2 4.4

37. Ca 2+ + CO 3 2- CaCO 3 NH 3 + H + NH 4 + Zn + 2HCl ZnCl 2 + H 2 Ca + F 2 CaF 2 Precipitation Acid-Base Redox (H 2 Displacement) Redox (Combination) Classify the following reactions. 4.4

38. Solution Stoichiometry The concentration of a solution is the amount of solute present in a given quantity of solvent or solution. M = molarity = moles of solute liters of solution What mass of KI is required to make 500. mL of a 2.80 M KI solution? volume KImoles KIgrams KI M KI 500. mL= 232 g KI 166 g KI 1 mol KI x 2.80 mol KI 1 L soln x 1 L 1000 mL x 4.5

40. Dilution is the procedure for preparing a less concentrated solution from a more concentrated solution. Dilution Add Solvent Moles of solute before dilution (i) Moles of solute after dilution (f) = MiViMiVi MfVfMfVf = 4.5

41. How would you prepare 60.0 mL of 0.200 M HNO 3 from a stock solution of 4.00 M HNO 3 ? M i V i = M f V f M i = 4.00 M f = 0.200V f = 0.0600 L V i = ? L 4.5 V i = MfVfMfVf MiMi = 0.200 x 0.0600 4.00 = 0.00300 L = 3.00 mL 3.00 mL of acid + 57.0 mL of water= 60.0 mL of solution

42. Titrations In a titration a solution of accurately known concentration is gradually added to another solution of unknown concentration until the chemical reaction between the two solutions is complete. Equivalence point – the point at which the reaction is complete Indicator – substance that changes color at (or near) the equivalence point Slowly add base to unknown acid UNTIL the indicator changes color 4.7

43. What volume of a 1.420 M NaOH solution is Required to titrate 25.00 mL of a 4.50 M H 2 SO 4 solution? 4.7 WRITE THE CHEMICAL EQUATION! volume acidmoles acidmoles basevolume base H 2 SO 4 + 2NaOH 2H 2 O + Na 2 SO 4 4.50 mol H 2 SO 4 1000 mL soln x 2 mol NaOH 1 mol H 2 SO 4 x 1000 ml soln 1.420 mol NaOH x 25.00 mL = 158 mL M acid rx coef. M base

44. What volume of a 0.800 M HCl solution is required to titrate 40.00 mL of a 1.600 M NH 3 solution? 4.7 HCl + NH 3 NH 4 Cl 1.600 mol NH 3 1000 mL soln x 1 mol HCl 1 mol NH 3 x 1000 ml soln 0.800 mol HCl x 40.00 mL = 80.0 mL

45. What volume of a 1.200 M Ba(OH) 2 solution is required to titrate 90.00 mL of a 3.600 M H 3 PO 4 solution? 4.7 3Ba(OH) 2 + 2H 3 PO 4 6H 2 O + Ba 3 (PO 4 ) 2 3.600 mol H 3 PO 4 1000 mL soln x 3 mol Ba 2 (OH) 2 mol H 3 PO 4 x 1000 ml soln 1.200 mol Ba 2 (OH) x 90.00 mL = 405.0 mL