1. AP Chapter 4 Aqueous Reactions and Solutions HW: 1 5 7 17 21 23 27 31 39 51 55 61 67 80 87 94 100 106 109

2. 4.1 – General Solution Properties Solution – Homogeneous mixture -Aqueous Solution – Solvent is water Solvent – One of larger amount Solute – One of smaller amount Electrolyte – Substance that conducts electricity when dissolved in water. Dissociates into ions (ex-ionic, acids, bases) Nonelectrolyte – Substance that does not conduct electricity when dissolved in water (most molecular compounds)

3. Electrolytes Open File “SALT WATER CONDUCTIVITY” in AP Power Point folder using Real Player http://www.youtube.com/watch?v=NdG3wK9kNcg

4. Ionic Compounds in Water Dissociates = Each ion separates from the solid structure and disperses through the solution. Solvation = Ions become surrounded by water molecules. Prevents the ions from recombining. When water is the solvent, AKA hydration.

5. Ionic Compounds in Water Dissociates = Each ion separates from the solid structure and disperses through the solution. Solvation = Ions become surrounded by water molecules. Prevents the ions from recombining.

6. Molecular Compounds in Water Ionization = Become ions, especially when dissolving in water. Used typically referring to molecular electrolytes. HCl(g)  H + (aq) + Cl - (aq) H2OH2O

7. © 2009, Prentice-Hall, Inc. Electrolytes A strong electrolyte dissociates completely when dissolved in water. A weak electrolyte only dissociates partially when dissolved in water.

8. © 2009, Prentice-Hall, Inc. Strong Electrolytes Are… Strong acids Strong bases Soluble ionic salts

9. Reversible Reaction – Reaction can proceed in both directions Chemical Equilibrium – No net change in reaction direction. Activity continues on a molecular level. General Information

10. Chapter 18 Chemical Equilibrium http://www.learner.org/resources/series61.html?pop =yes&pid=806#http://www.learner.org/resources/series61.html?pop =yes&pid=806# - Molecules in Action Starting at 15:30-23:52 (FYI: Co(H 2 O) 6 +2 + 4Cl - -> CoCl 4 + 6 H 2 O) Equilibrium

11. © 2009, Prentice-Hall, Inc. 4.2 - Precipitation Reactions Mixing of two ionic solutions that have ions that form compounds that are insoluble (as could be predicted by the solubility guidelines), forms a precipitate.

12. 4.2 – Precipitation Reactions Precipitation Rxn – Results in a precipitate Precipitate – Insoluble solid (typically formed from two solutions)

13. 4.2 – Precipitation Reactions Solubility – The maximum amount of solute that will dissolve in a given quantity of solvent at a specific temp Insoluble – A substance that does not dissolve. Specifically, the solubility must be less than 0.01 mol/L to be considered insoluble.

14. Solubility Rules 1. Group 1 & Ammonium cmpds are soluble 2. Nitrates, Bicarbonates, Chlorates are soluble 3. Halides are soluble except Ag, Hg 2 2+, Pb 4. Sulfates are soluble except Ag, Ca, Sr, Ba, Pb 5. Carbonates, phosphates, chromates, sulfides are insoluble

15. © 2009, Prentice-Hall, Inc. Double Displacement (Metathesis / Exchange) Reactions Metathesis comes from a Greek word that means “to transpose.” AgNO 3 (aq) + KCl (aq)  AgCl (s) + KNO 3 (aq)

16. © 2009, Prentice-Hall, Inc. Ionic Equations It is helpful to pay attention to exactly what species are present in a reaction mixture (i.e., solid, liquid, gas, aqueous solution). If we are to understand reactivity, we must be aware of just what is changing during the course of a reaction.

17. © 2009, Prentice-Hall, Inc. Molecular Equation The molecular equation lists the reactants and products in their molecular form. This is typically NOT a good representation of what is actually happening on an atomic level. AgNO 3 (aq) + KCl (aq)  AgCl (s) + KNO 3 (aq)

18. © 2009, Prentice-Hall, Inc. Complete Ionic / Ionic Equation In the ionic equation all strong electrolytes (strong acids, strong bases, and soluble ionic salts) are dissociated into their ions. This more accurately reflects the species that are found in the reaction mixture. Ag + (aq) + NO 3 - (aq) + K + (aq) + Cl - (aq)  AgCl (s) + K + (aq) + NO 3 - (aq)

19. © 2009, Prentice-Hall, Inc. Net Ionic Equation To form the net ionic equation, cross out anything that does not change from the left side of the equation to the right. Ag + (aq) + NO 3 - (aq) + K + (aq) + Cl - (aq)  AgCl (s) + K + (aq) + NO 3 - (aq)

20. © 2009, Prentice-Hall, Inc. Net Ionic Equation To form the net ionic equation, cross out anything that does not change from the left side of the equation to the right. The only things left in the equation are those things that change (i.e., react) during the course of the reaction. Ag + (aq) + Cl - (aq)  AgCl (s)

21. © 2009, Prentice-Hall, Inc. Net Ionic Equation To form the net ionic equation, cross out anything that does not change from the left side of the equation to the right. The only things left in the equation are those things that change (i.e., react) during the course of the reaction. Those things that didn’t change (and were deleted from the net ionic equation) are called spectator ions. Ag + (aq) + NO 3 - (aq) + K + (aq) + Cl - (aq)  AgCl (s) + K + (aq) + NO 3 - (aq)

22. © 2009, Prentice-Hall, Inc. Writing Net Ionic Equations 1.Write a balanced molecular equation. 2.Dissociate all strong electrolytes. 3.Cross out anything that remains unchanged from the left side to the right side of the equation. 4.Write the net ionic equation with the species that remain.

23. Practice

24. 4.3 - Acid – Base Reactions Properties of Acids: -Sour -Color of indicators -React with some metals (ex-Zn, Mg, Fe) -React with carbonates to form CO 2 -Electrolytes

25. 4.3 - Acid – Base Reactions Base Properties: -Bitter -Feel slippery -Change color of indicators -Electrolytes

26. © 2009, Prentice-Hall, Inc. Acids Arrhenius defined acids as substances that increase the concentration of H + when dissolved in water. Brønsted and Lowry defined them as proton donors.

27. © 2009, Prentice-Hall, Inc. Acids There are only seven strong acids: Hydrochloric (HCl) Hydrobromic (HBr) Hydroiodic (HI) Nitric (HNO 3 ) Sulfuric (H 2 SO 4 ) Chloric (HClO 3 ) Perchloric (HClO 4 ) Weak acid examples: Acetic Hydrofluoric Ammonium

28. © 2009, Prentice-Hall, Inc. Bases Arrhenius defined bases as substances that increase the concentration of OH − when dissolved in water. Brønsted and Lowry defined them as proton acceptors.

29. © 2009, Prentice-Hall, Inc. Bases The strong bases are the soluble metal salts of hydroxide ion: Alkali metals (Grp 1) Calcium Strontium Barium Weak Base - ammonia

30. Bronsted Definitions Bronsted Acid = Proton donor Bronsted Base = Proton acceptor HCl + H 2 O -> Cl - + H 3 O + Hydronium

31. 4.3 – Acid / Base Weak Acid = Have only partial ionization Weak Base = Have only partial ionization Hydronium = H 3 O + Salt = Any ionic compound

32. Acids Monoprotic – One proton -Hydrochloric, Nitric, Acetic Diprotic – Two protons -Sulfuric Can look at in two steps:

33. © 2009, Prentice-Hall, Inc. Acid-Base Reactions In an acid-base reaction, the acid donates a proton (H + ) to the base.

34. © 2009, Prentice-Hall, Inc. Neutralization Reactions Generally, when aqueous solutions of an acid and a base are combined, the products are a salt and water. CH 3 COOH (aq) + NaOH (aq)  CH 3 COONa (aq) + H 2 O (l)

35. © 2009, Prentice-Hall, Inc. Neutralization Reactions When a strong acid reacts with a strong base, the net ionic equation is… HCl (aq) + NaOH (aq)  NaCl (aq) + H 2 O (l)

36. © 2009, Prentice-Hall, Inc. Neutralization Reactions When a strong acid reacts with a strong base, the net ionic equation is… HCl (aq) + NaOH (aq)  NaCl (aq) + H 2 O (l) H + ( aq ) + Cl - ( aq ) + Na + ( aq ) + OH - ( aq )  Na + ( aq ) + Cl - ( aq ) + H 2 O ( l )

37. © 2009, Prentice-Hall, Inc. Neutralization Reactions When a strong acid reacts with a strong base, the net ionic equation is… HCl (aq) + NaOH (aq)  NaCl (aq) + H 2 O (l) H + ( aq ) + Cl - ( aq ) + Na + ( aq ) + OH - ( aq )  Na + ( aq ) + Cl - ( aq ) + H 2 O ( l ) H + (aq) + OH - (aq)  H 2 O (l)

38. © 2009, Prentice-Hall, Inc. Gas-Forming Reactions Some metathesis reactions do not give the product expected. In this reaction, the expected product (H 2 CO 3 ) decomposes to give a gaseous product (CO 2 ). CaCO 3 (s) + HCl (aq)  CaCl 2 (aq) + CO 2 (g) + H 2 O (l)

39. © 2009, Prentice-Hall, Inc. Gas-Forming Reactions When a carbonate or bicarbonate reacts with an acid, the products are a salt, carbon dioxide, and water. CaCO 3 (s) + HCl (aq)  CaCl 2 (aq) + CO 2 (g) + H 2 O (l) NaHCO 3 (aq) + HBr (aq)  NaBr (aq) + CO 2 (g) + H 2 O (l)

40. © 2009, Prentice-Hall, Inc. Gas-Forming Reactions Similarly, when a sulfite reacts with an acid, the products are a salt, sulfur dioxide, and water. SrSO 3 (s) + 2 HI (aq)  SrI 2 (aq) + SO 2 (g) + H 2 O (l)

41. © 2009, Prentice-Hall, Inc. Gas-Forming Reactions This reaction gives the predicted product, but you had better carry it out in the hood, or you will be very unpopular! But just as in the previous examples, a gas is formed as a product of this reaction. Na 2 S (aq) + H 2 SO 4 (aq)  Na 2 SO 4 (aq) + H 2 S (g)

42. 4.4 Oxidation-Reduction (Redox) Redox Reactions Half-Reaction Oxidation Reaction Reduction Reaction Oxidation Numbers Reducing Agent Oxidizing Agent

43. © 2009, Prentice-Hall, Inc. Oxidation-Reduction Reactions An oxidation occurs when an atom or ion loses electrons. A reduction occurs when an atom or ion gains electrons. One cannot occur without the other.

44. © 2009, Prentice-Hall, Inc. 4.4 – Redox Reactions Oxidation Numbers Elements in their elemental form have an oxidation number of 0. The oxidation number of a monatomic ion is the same as its charge.

45. © 2009, Prentice-Hall, Inc. 4.4 – Redox Reactions Oxidation Numbers Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions.  Oxygen has an oxidation number of −2, except in the peroxide ion in which it has an oxidation number of −1.  Hydrogen is −1 when bonded to a metal, +1 when bonded to a nonmetal.

46. © 2009, Prentice-Hall, Inc. 4.4 – Redox Reactions Oxidation Numbers Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions.  Fluorine always has an oxidation number of −1.  The other halogens have an oxidation number of −1 when they are negative; they can have positive oxidation numbers, however, most notably in oxyanions.

47. © 2009, Prentice-Hall, Inc. 4.4 – Redox Reactions Oxidation Numbers The sum of the oxidation numbers in a neutral compound is 0. The sum of the oxidation numbers in a polyatomic ion is the charge on the ion.

48. Types of Redox Reactions 1.Combination -Two or more reactants from a single product -Combustion of Magnesium metal

49. Types of Redox Reactions 2.Decomposition -Break down into 2 or more products -Decomposition of mercury (II) oxide -Elephant’s Toothpaste Dawn detergent 30% hydrogen peroxide (H 2 0 2 ) saturated solution of potassium iodide (KI)

50. Types of Redox Reactions 3.Single Displacement Reactions -An atom or ion in a compound is displaced by another a. -Hydrogen displacement: -Group 1, Ca, Sr, Ba can displace from water -Al & Fe can replace from steam -Many metals can replace from Acids (ex – Zn)

51. b. Metal Displacement -Use activity series -Elements high on the list can remove lower elements from a compound

52. © 2009, Prentice-Hall, Inc. Displacement Reactions Metal Displacement: In this reaction, silver ions oxidize copper metal. Cu (s) + 2 Ag + (aq)  Cu 2+ (aq) + 2 Ag (s)

53. © 2009, Prentice-Hall, Inc. Displacement Reactions The reverse reaction, however, does not occur. Cu 2+ (aq) + 2 Ag (s)  Cu (s) + 2 Ag + (aq) x

54. c. Halogen displacement Fluorine Gas > Cl > Br > I ex – Chlorine gas + Potassium bromide

55. 4. Disproportionation -One species is both oxidized and reduced ex – Hydrogen peroxide -> water + oxygen gas

56. Concentration Molarity

57. © 2009, Prentice-Hall, Inc. Mixing a Solution To create a solution of a known molarity, one weighs out a known mass (and, therefore, number of moles) of the solute. The solute is added to a volumetric flask, and solvent is added to the line on the neck of the flask. (best to start with some solvent in flask)

58. © 2009, Prentice-Hall, Inc. Dilution One can also dilute a more concentrated solution by – Using a pipet to deliver a volume of the solution to a new volumetric flask, and – Adding solvent to the line on the neck of the new flask.

59. © 2009, Prentice-Hall, Inc. 4.6 - Using Molarities in Stoichiometric Calculations

60. Titration The analytical technique in which one can calculate the concentration of a solute in a solution.

61. Acid Base Titration Understand Equivalence point vs. End Point -Equivalence Pt = -End Pt = *Need to pick the right indicator so the Equiv Pt = the End Pt

62. Redox Titration Titrate a reducing agent with an oxidizing agent Ex: dichromate -> chromium (III) ion Ex: permanganate -> manganese (II) ion