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Reactions in Aqueous Solution

Reactions in Aqueous Solution. Chapter 4. Copyright © The McGraw-Hill Companies, Inc.  Permission required for reproduction or display. Solution. Solvent. Solute. A solution is a homogenous mixture of 2 or more substances.

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Reactions in Aqueous Solution

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

  2. Solution Solvent Solute A solution is a homogenous mixture of 2 or more substances The solute is(are) the substance(s) present in the smaller amount(s) and is dissolved in the solvent The solvent is the substance present in the larger amount and dissolves the solute H2O Soft drink (l) Sugar, CO2 Air (g) N2 O2, Ar, CH4 Pb Sn Soft Solder (s) 4.1

  3. 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

  4. 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

  5. A reversible reaction. The reaction can occur in both directions. Ionization of acetic acid CH3COOHCH3COO- (aq) + H+ (aq) Acetic acid is a weak electrolyte because its ionization in water is incomplete. 4.1

  6. d- d+ H2O Hydration is the process in which an ion is surrounded by water molecules arranged in a specific manner. 4.1

  7. H2O C6H12O6 (s) C6H12O6 (aq) Nonelectrolyte does not conduct electricity? No cations (+) and anions (-) in solution 4.1

  8. precipitate Pb(NO3)2(aq) + 2NaI (aq) PbI2(s) + 2NaNO3(aq) Pb2+ + 2NO3- + 2Na+ + 2I- PbI2 (s) + 2Na+ + 2NO3- Pb2+ + 2I- PbI2 (s) PbI2 Precipitation Reactions Precipitate – insoluble solid that separates from solution molecular equation ionic equation net ionic equation Na+ and NO3- are spectator ions 4.2

  9. Pb2+ + 2I- PbI2 (s) Precipitation of Lead Iodide PbI2 4.2

  10. Solubility is the maximum amount of solute that will dissolve in a given quantity of solvent at a specific temperature. 4.2

  11. Write the net ionic equation for the reaction of silver nitrate with sodium chloride. Writing Net Ionic Equations • Write the balanced molecular equation. • Write the ionic equation showing the strong electrolytes completely dissociated into cations and anions. • Cancel the spectator ions on both sides of the ionic equation • Check that charges and number of atoms are balanced in the net ionic equation 4.2

  12. Practice Problem • Write the net ionic equation for the following reaction: • MgCl2(aq) + Na2S(aq)2 NaCl(aq) + MgS(s)

  13. 2HCl (aq) + Mg (s) MgCl2 (aq) + H2 (g) 2HCl (aq) + CaCO3 (s) CaCl2 (aq) + CO2 (g) + H2O (l) Acids Have a sour taste. Vinegar owes its taste to acetic acid. Citrus fruits contain citric acid. Cause color changes in plant dyes. React with certain metals to produce hydrogen gas. React with carbonates and bicarbonates to produce carbon dioxide gas Aqueous acid solutions conduct electricity. 4.3

  14. Bases Have a bitter taste. Feel slippery. Many soaps contain bases. Cause color changes in plant dyes. Aqueous base solutions conduct electricity. 4.3

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

  16. Hydronium ion, hydrated proton, H3O+ 4.3

  17. 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

  18. HCl H+ + Cl- HNO3 H+ + NO3- CH3COOH H+ + CH3COO- H2SO4 H+ + HSO4- HSO4- H+ + SO42- H3PO4 H+ + H2PO4- H2PO4- H+ + HPO42- HPO42- H+ + PO43- 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 4.3

  19. HI (aq) H+ (aq) + I- (aq) CH3COO- (aq) + H+ (aq) CH3COOH (aq) H2PO4- (aq) H+ (aq) + HPO42- (aq) H2PO4- (aq) + H+ (aq) H3PO4 (aq) Identify each of the following species as a Brønsted acid, base, or both. (a) HI, (b) CH3COO-, (c) H2PO4- 4.3

  20. acid + base salt + water HCl (aq) + NaOH (aq) NaCl (aq) + H2O H+ + Cl- + Na+ + OH- Na+ + Cl- + H2O H+ + OH- H2O Neutralization Reaction 4.3

  21. Oxidation number The charge the atom would have in a molecule (or an ionic compound) if electrons were completely transferred. • Free elements (uncombined state) and have an oxidation number of zero. Na, Be, K, Pb, H2, O2, P4 = 0 • In monatomic ions, the oxidation number is equal to the charge on the ion. Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2 • The oxidation number of oxygen isusually–2. In H2O2 and O22- it is –1. 4.4

  22. Oxidation numbers of all the elements in HCO3- ? • The oxidation number of hydrogen is +1except when it is bonded to metals in binary compounds. In these cases, its oxidation number is –1. • Group IA metals are +1, IIA metals are +2 and fluorine is always –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. 7. Oxidation numbers do not have to be integers. Oxidation number of oxygen in the superoxide ion, O2-, is -½. HCO3- 4.4

  23. The oxidation numbers of elements in their compounds 4.4

  24. Oxidation numbers of all the elements in the following ? IF7 NaIO3 K2Cr2O7 4.4

  25. Practice Problems • Find the oxidation number for Nitrogen in the following compounds: • HNO3 • Ca3N2 • Mg(NO2)2

  26. 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

  27. 4.4

  28. Steps for Determining Oxidation and Reduction • Assign oxidation numbers to each atom in the reaction. • Determine if an atom gains or loses electrons by comparing the oxidation number in the reactants to the oxidation number in the products. • If an atom loses electrons, then it is Oxidized OiL • If an atom gains electrons then it is Reduced RiG • Write each of the half-reactions.

  29. 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? Zn Zn2+ + 2e- Zn is the reducing agent Zn is oxidized Cu2+is reduced Cu2+ is the oxidizing agent 4.4

  30. A + B C 2Al + 3Br2 2AlBr3 C A + B 2KClO3 2KCl + 3O2 Types of Oxidation-Reduction Reactions Combination Reaction +3 -1 0 0 Decomposition Reaction +1 +5 -2 +1 -1 0 4.4

  31. A + O2 B S + O2 SO2 2Mg + O2 2MgO Types of Oxidation-Reduction Reactions Combustion Reaction +4 -2 0 0 +2 -2 0 0 4.4

  32. A + BC AC + B Sr + 2H2O Sr(OH)2 + H2 TiCl4 + 2Mg Ti + 2MgCl2 Cl2 + 2KBr 2KCl + Br2 Types of Oxidation-Reduction Reactions Displacement Reaction +1 +2 0 0 Hydrogen Displacement +4 0 0 +2 Metal Displacement -1 0 0 -1 Halogen Displacement 4.4

  33. M + BC AC + B Ca + 2H2O Ca(OH)2 + H2 Pb + 2H2O Pb(OH)2 + H2 The Activity Series for Metals Hydrogen Displacement Reaction M is metal BC is acid or H2O B is H2 4.4

  34. Cl2 + 2KBr 2KCl + Br2 I2 + 2KBr 2KI + Br2 The Activity Series for Halogens F2 > Cl2 > Br2 > I2 Halogen Displacement Reaction -1 0 0 -1 4.4

  35. Cl2 + 2OH- ClO- + Cl- + H2O Chlorine Chemistry Types of Oxidation-Reduction Reactions Disproportionation Reaction Element is simultaneously oxidized and reduced. +1 -1 0 4.4

  36. NH3+H2O NH4OH Classify the following reactions. Zn + 2HCl ZnCl2 + H2 Ca + F2 CaF2 CaCl2 + Na2CO32NaCl + CaCO3 4.4

  37. What mass of KI is required to make 500. mL of a 2.80 M KI solution? moles of solute M = molarity = liters of solution Solution Stoichiometry The concentration of a solution is the amount of solute present in a given quantity of solvent or solution. 4.5

  38. Practice Problems • What mass of MgCl2 would you need to make 250.0 ml of a 6.50 M solution? • How much water would be needed to make a 2.0M solution given 96.0 g of NaCl? • How many grams of FeO are needed to make 500.00 ml of a 1.6M solution?

  39. 4.5

  40. Moles of solute before dilution (1) Moles of solute after dilution (2) = Dilution Add Solvent = M2 V2 M1V1 Dilution is the procedure for preparing a less concentrated solution from a more concentrated solution. 4.5

  41. How would you prepare 60.0 mL of 0.200 M HNO3 from a stock solution of 4.00 M HNO3? 4.5

  42. Practice Problem • How would make 250.0 ml of 2.00M HCl from a 12.0M stock solution?

  43. Titrations In a titration a solution of accurately known concentration is added 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

  44. WRITE THE CHEMICAL EQUATION! What volume of a 1.420 M NaOH solution is Required to titrate 25.00 mL of a 4.50 M H2SO4 solution? H2SO4 + 2NaOH 2H2O + Na2SO4 4.7

  45. Practice Problem • What volume of 2.0 M Mg(OH)2 is needed to titrate 20.0 ml of 1.0M H3PO4?

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