1 / 68

Reactions in Aqueous Solution

Reactions in Aqueous Solution. Chapter 4. PbI 2. Reaction of lead nitrate with sodium Iodide. Types of Chemical Reactions 01. Precipitation Reactions: A process in which an insoluble solid (precipitate) drops out of the solution.

sook
Download Presentation

Reactions in Aqueous Solution

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Reactions in Aqueous Solution Chapter 4

  2. PbI2 Reaction of lead nitrate with sodium Iodide

  3. Types of Chemical Reactions 01 • Precipitation Reactions:A process in which an insoluble solid (precipitate) drops out of the solution. • Most precipitation reactions occur when the anions and cations of two ionic compounds change partners. Pb(NO3)2(aq) + 2 NaI(aq)  2 NaNO3(aq) + PbI2(s)

  4. Types of Chemical Reactions 02 • Acid–Base Neutralization:A process in which an acid reacts with a base to yield water plus an ionic compound called a salt. • The driving force of this reaction is the formation of the stable water molecule. HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)

  5. 4.4

  6. Types of Chemical Reactions 03 • Oxidation–Reduction (Redox) Reaction:A process in which one or more electrons are transferred between reaction partners. • The driving force of this reaction is the decrease in electrical potential. Mg(s) + I2(g)  MgI2(s)

  7. Types of Chemical Reactions 04 • Metathesis Reactions: These are reactions where two reactants just exchange parts. AX + BY AY + BX HNO3(aq) + KOH(aq) KNO3(aq) + HOH(l) BaCl2(aq) + K2SO4(aq)  BaSO4(s) + 2 KCl(aq)

  8. Electrolytes in Solution 01 • Why do ionic compounds conduct electricity when molecular ones generally do not?

  9. Electrolytes in Solution 02 • Electrolytes:Dissolve in water to produce ionic solutions. • Nonelectrolytes: Do not form ions when they dissolve in water.

  10. nonelectrolyte weak electrolyte strong electrolyte 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, results in a solution that does not conduct electricity. 4.1

  11. H2O NaCl (s)Na+ (aq) + Cl- (aq) CH3COOHCH3COO- (aq) + H+ (aq) Conduct electricity in solution? Cations (+) and Anions (-) Strong Electrolyte – 100% dissociation Weak Electrolyte – not completely dissociated 4.1

  12. Electrolytes in Solution 03 • Dissociation: • The process by which a compound splits up to form ions in the solution.

  13. d- d+ H2O Hydration is the process in which an ion is surrounded by water molecules arranged in a specific manner.

  14. Electrolytes in Solution 05

  15. PbI2 Reaction of lead (II) nitrate with sodium Iodide

  16. precipitate Pb(NO3)2(aq) + 2NaI (aq) PbI2(s) + 2NaNO3(aq) Pb2+ + 2NO3- + 2Na+ + 2I- PbI2 (s) + 2Na+ + 2NO3- Pb2+ + 2I- PbI2 (s) PbI2 Net Ionic Equations in Precipitation Reactions Precipitate – insoluble solid that separates from solution Formula unit equation ionic equation net ionic equation Na+ and NO3- are spectator ions 4.2

  17. AgNO3 (aq) + NaCl (aq) AgCl (s) + NaNO3 (aq) Write the net ionic equation for the reaction of silver nitrate with sodium chloride. Ag+ + NO3- + Na+ + Cl- AgCl (s) + Na+ + NO3- Ag+ + Cl- AgCl (s) Writing Net Ionic Equations • Write the balanced formula unit equation. • Write the ionic equation showing dissociation of strong electrolytes • Determine precipitate from solubility rules, leave them un-dissociated. • Cancel the spectator ions on both sides of the ionic equation 4.2

  18. Solubility Rules for Ionic Compounds

  19. Acids and Bases Regulation of blood pH Blood pH = 7.2-7.4

  20. Figure 4.7: Household acids and bases. Photo courtesy of American Color.

  21. Figure 4.8: Preparation of red cabbage juice as an acid-base indicator.Photo courtesy of James Scherer.

  22. Acids 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 and bicarbonates to produce carbon dioxide gas Bases Have a bitter taste. Feel slippery. Many soaps contain bases. 4.3

  23. Arrhenius acid is a substance that produces H+ (H3O+) in water Arrhenius base is a substance that produces OH- in water 4.3

  24. Arrhenius Acid:

  25. Arrhenius Base • Arrhenius Base: • A substance that dissociates in, or reacts with, water to form hydroxide ions (OH–).

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

  27. A Brønsted acid is a proton donorA Brønsted base is a proton acceptor • Dissociation of Water: This equilibrium gives us the ion product of water.Kw = Kc = [H+][OH–] = 1.0 x 10–14

  28. Acid–Base Concepts 05 • Lewis Acid: Electron pair acceptor. e.g.: Al3+, H+, BF3. • Lewis Base: Electron pair donor. e.g.: H2O, NH3, O2–. • Bond formed is called a coordinate bond.

  29. H F F F B F B N H F F H H N H H • • Lewis Acids and Bases + acid base No protons donated or accepted!

  30. Strong and Weak Acid-Bases

  31. Strong and Weak Acids HCl H+ + Cl- HNO3 H+ + NO3- CH3COOH H+ + CH3COO- H2SO4 H+ + HSO4- HPO42- H+ + PO43- H2PO4-- H+ + HPO42- H3PO4 H+ + H2PO4- HSO4- H+ + SO42- Monoprotic acids Strong electrolyte, strong acid Strong electrolyte, strong acid Weak electrolyte, weak acid Diprotic acids Strong electrolyte, strong acid Weak electrolyte, weak acid Triprotic acids Weak electrolyte, weak acid Weak electrolyte, weak acid Weak electrolyte, weak acid

  32. pH [H+] pH – A Measure of Acidity pH = -log [H+] Solution Is At 250C neutral [H+] = [OH-] [H+] = 1 x 10-7 pH = 7 [H+] > 1 x 10-7 pH < 7 acidic [H+] > [OH-] basic [H+] < [OH-] [H+] < 1 x 10-7 pH > 7

  33. pOH = -log [OH-] [H+][OH-] = Kw = 1.0 x 10-14 -log [H+] – log [OH-] = 14.00 pH + pOH = 14.00

  34. The OH- ion concentration of a blood sample is 2.5 x 10-7 M. What is the pH of the blood? The pH of rainwater collected in a certain region of the northeastern United States on a particular day was 4.82. What is the H+ ion concentration of the rainwater? pH = -log [H+] = 10-4.82 = 1.5 x 10-5M [H+] = 10-pH pH + pOH = 14.00 pOH = -log [OH-] = -log (2.5 x 10-7) = 6.60 pH = 14.00 – pOH = 14.00 – 6.60 = 7.40

  35. Oxidation–Reduction Reactions 02 Oxidation Is Loss (of electrons) AnodeOxidation Reducing Agent

  36. Oxidation–Reduction Reactions 03 Reduction Is Gain (of electrons) Cathode Reduction Oxidizing Agent

  37. 2Mg (s) + O2 (g) 2MgO (s) 2Mg 2Mg2+ + 4e- O2 + 4e- 2O2- 2Mg + O2 + 4e- 2Mg2+ + 2O2- + 4e- 2Mg + O2 2MgO Oxidation-Reduction Reactions (electron transfer reactions) Oxidation half-reaction (lose e-) Reduction half-reaction (gain e-) 4.4

  38. A + B C S + O2 SO2 C A + B 2KClO3 2KCl + 3O2 Types of Oxidation-Reduction Reactions Combination Reaction +4 -2 0 0 Decomposition Reaction +1 +5 -2 +1 -1 0

  39. NH3 + H+ NH4+ Classify the following reactions. Zn + 2HCl ZnCl2 + H2 Ca + F2 CaF2 Precipitation Ca+2(aq) + CO32-(aq) CaCO3(s) Acid-Base Redox (H2 Displacement) Redox (Combination)

  40. 4.4

  41. Cu (s) + 2AgNO3 (aq) Cu(NO3)2 (aq) + 2Ag (s) Zn (s) + CuSO4 (aq) ZnSO4 (aq) + Cu (s) Cu2+ + 2e- Cu Copper wire reacts with silver nitrate to form silver metal. What is the oxidizing agent in the reaction? Cu Cu2+ + 2e- Zn Zn2+ + 2e- Ag+ + 1e- Ag Zn is the reducing agent Zn is oxidized Cu2+is reduced Cu2+ is the oxidizing agent Ag+is reduced Ag+ is the oxidizing agent 4.4

  42. OxidationNumber Rules

  43. Oxidation Number Rules

  44. Oxidation numbers of all the elements in HCO3- ? HCO3- O = -2 H = +1 3x(-2) + 1 + ? = -1 C = +4 4.4

  45. Oxidation numbers of all the elements in the following ? IF7 F = -1 7x(-1) + ? = 0 I = +7 K2Cr2O7 NaIO3 O = -2 K = +1 O = -2 Na = +1 3x(-2) + 1 + ? = 0 7x(-2) + 2x(+1) + 2x(?) = 0 I = +5 Cr = +6 4.4

  46. 2Ag(s) + Cu2+(g) Cu(s) + 2Ag1+(g) 2Ag1+(aq) + Cu(s) Cu2+(aq) + 2Ag(s) The Activity Series of the Elements Which one of these reactions will occur?

  47. Activity Series of Elements 01

  48. The Activity Series of the Elements The elements that are higher up in the table are more likely to be oxidized. Thus, any element higher in the activity series will reduce the ion of any element lower in the activity series.

  49. Activity Series of Elements 02 • Activity series looks at the relative reactivity of a free metal with an aqueous cation. • Fe(s) + Cu2+(aq)  Fe2+(aq) + Cu(s) • Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s) • Cu(s) + 2 Ag+(aq)  2 Ag(s) + Cu2+(aq) • Mg(s) + 2 H+(aq) Mg2+(aq) + H2(g) Cu2+ (aq) + 2 Ag (s) X Cu (s) + 2 Ag+ (aq)

More Related