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Chapter 4 Reactions in Aqueous Solution

Chapter 4 Reactions in Aqueous Solution. HCl( aq ) + H 2 O( aq ). HCl( aq ). HA( aq ). H 3 O 1+ ( aq ) + Cl 1- ( aq ). H 1+ ( aq ) + A 1- ( aq ). H 1+ ( aq ) + Cl 1- ( aq ). Acids, Bases, and Neutralization Reactions.

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Chapter 4 Reactions in Aqueous Solution

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  1. Chapter 4Reactions in Aqueous Solution

  2. HCl(aq) + H2O(aq) HCl(aq) HA(aq) H3O1+(aq) + Cl1-(aq) H1+(aq) + A1-(aq) H1+(aq) + Cl1-(aq) Acids, Bases, and Neutralization Reactions Acid (Arrhenius): A substance that dissociates in water to produce hydrogen ions, H1+: In water, acids produce hydronium ions, H3O1+:

  3. NH3(aq) + H2O(aq) NaOH(aq) MOH(aq) M1+(aq) + OH1-(aq) Na1+(aq) + OH1-(aq) NH41+(aq) + OH1-(aq) Acids, Bases, and Neutralization Reactions Base (Arrhenius): A substance that dissociates in water to produce hydroxide ions, OH1-: Ammonia, commonly called “ammonium hydroxide” is a base:

  4. Acids, Bases, and Neutralization Reactions Strong acids and strong bases are strong electrolytes. Weak acids and weak bases are weak electrolytes.

  5. HA + MOH HA + MOH MA + H2O MA + HOH Acids, Bases, and Neutralization Reactions These acid-base neutralization reactions are double-replacement reactions just like the precipitation reactions: or Acid Base Salt Water

  6. Acids, Bases, and Neutralization Reactions Write the molecular, ionic, and net ionic equations for the reaction of aqueous HBr and aqueous Ba(OH)2. 1. Write the chemical formulas of the products (use proper ionic rules for the salt). HBr(aq) + Ba(OH)2(aq) H2O + BaBr2 Acid Base Water Salt

  7. Acids, Bases, and Neutralization Reactions Write the molecular, ionic, and net ionic equations for the reaction of aqueous NaOH and aqueous HF. 1. Write the chemical formulas of the products (use proper ionic rules for the salt). HF(aq) + NaOH(aq) H2O + NaF Acid Base Water Salt

  8. Examples • Predict the product and write a molecular equation, ionic equation and net ionic equation for the following reactions • K2CO3(aq) + NiCl2(aq)  • HNO3(aq) + LiOH(aq)  • HCN(aq) + Mg(OH)2

  9. 2Fe2O3(s) + 3C(s) 4Fe(s) + 3O2(g) 2Fe2O3(s) 4Fe(s) + 3CO2(g) Oxidation-Reduction (Redox) Reactions Rusting of iron: an oxidation of Fe Manufacture of iron: a reduction of Fe

  10. Oxidation-Reduction (Redox) Reactions Oxidation Number (State): A value which indicates whether an atom is neutral, electron-rich, or electron-poor. Rules for Assigning Oxidation Numbers An atom in its elemental state has an oxidation number of 0. Na H2 Br2 S Ne Oxidation number 0

  11. Oxidation-Reduction (Redox) Reactions A monatomic ion has an oxidation number identical to its charge. Na1+ +1 Ca2+ +2 Al3+ +3 Cl1- -1 O2- -2

  12. 1- H O H H Ca O H H Oxidation-Reduction (Redox) Reactions • An atom in a polyatomic ion or in a molecular compound usually has the same oxidation number it would have if it were a monatomic ion. • Hydrogen can be either +1 or -1. -1 +2 -1 +1 -2 • Oxygen usually has an oxidation number of -2. H O O H +1 -2 +1 +1 -1 -1 +1

  13. Cl O Cl Oxidation-Reduction (Redox) Reactions • Halogens usually have an oxidation number of -1. 3. H Cl +1 -1 +1 -2 +1

  14. Oxidation-Reduction (Redox) Reactions The sum of the oxidation numbers is 0 for a neutral compound and is equal to the net charge for a polyatomic ion. 2(+1) + x + 3(-2) = 0 (net charge) H2SO3 x = +4 +1 x -2 Cr2O72- 2(x) + 7(-2) = -2 (net charge) x = +6 x -2

  15. Example • Determine the oxidation number for each atom in the following compounds/molecules • CO2 • CCl4 • CoSO4 • K2O2

  16. reduction 0 -2 4Fe(s) + 3O2(g) 2Fe2 O3 (s) 0 +3 oxidation Identifying Redox Reactions Reduction: gaining one or more electron increasing in oxidation number  Oxidizing agent Oxidation: losing one or more electrons decreasing in oxidation number  Reducing agent

  17. Identifying Redox Reactions Oxidizing Agent • Causes oxidation • Gains one or more electrons • Undergoes reduction • Oxidation number of atom decreases Reducing Agent Causes reduction Loses one or more electrons Undergoes oxidation Oxidation number of atom increases

  18. Example Identify each of the following as 1) oxidation or 2) reduction. __A. Sn(s) Sn4+(aq) + 4e− __B. Fe3+(aq)+ 1e−Fe2+(aq) __C. Cl2(g) + 2e−2Cl-(aq)

  19. Writing Oxidation and Reduction Reactions Write the separate half oxidation and reduction reactions for the following equation. 2Cs(s) + F2(g) 2CsF(s) • 3 Na(l) + AlCl3(l) 3  NaCl(l) + Al(l)

  20. 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?

  21. Fe(s) + Cu2+(aq) Fe2+(aq) + Cu(s) The Activity Series of the Elements

  22. The Activity Series of the Elements 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.

  23. 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?

  24. Example • Predict whether the following redox reactions will occurred or not. If so, predict the products • Zn(s) + FeCl2(aq)  • Ni(s) + Mg(NO3)2(aq) 

  25. 5H2C2O4(aq) + 2MnO41-(aq) + 6H1+(aq) 10CO2(g) + 2Mn2+(aq) + 8H2O(l) Redox Titrations Titration: A procedure for determining the concentration of a solution by allowing a carefully measured volume to react with a solution of another substance (the standard solution) whose concentration is known. If the unknown concentration is the potassium permanganate solution, MnO41-, it can be slowly added to a known amount of oxalic acid, H2C2O4, until a faint purple color persists.

  26. 5H2C2O4(aq) + 2MnO41-(aq) + 6H1+(aq) 10CO2(g) + 2Mn2+(aq) + 8H2O(l) Redox Titrations A solution is prepared with 0.2585 g of oxalic acid, H2C2O4. 22.35 mL of an unknown solution of potassium permanganate are needed to titrate the solution. What is the concentration of the potassium permanganate solution?

  27. Calculation Set up Mass of H2C2O4 Moles of H2C2O4 Moles of KMnO4 Molarity of KMnO4 Mole Ratio Molar Mass of H2C2O4 Molarity of KMnO4

  28. Example • A 0.0484M standard solution of potassium permanganate was titrated against 25.00mL of an iron (II) sulfate solution. The equivalence point, as indicated by a faint pink color, was reached when 15.50mL of potassium permanganate solution had been added. Calculate the concentration of the iron (II) sulfate solution

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