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Reactions in Aqueous Solution Chapter 4 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display. 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 Solution Solvent Solute Soft drink (l) H2O Sugar, CO2 Air (g) N2 O2, Ar, CH4 Soft Solder (s) Pb Sn aqueous solutions of KMnO4 2 An electrolyte is a substance that, when dissolved in water, results in a solution that can conduct electricity. weak electrolyte is a substance that, when dissolved in water, results in a solution that can conduct small electricity A nonelectrolyte is a substance that, when dissolved, results in a solution that does not conduct electricity. 3 Conduct electricity in solution? Cations (+) and Anions (-) Strong Electrolyte – 100% dissociation NaCl (s) H 2O Na+ (aq) + Cl- (aq) Weak Electrolyte – not completely dissociated CH3COOH CH3COO- (aq) + H+ (aq) 4 Ionization of acetic acid CH3COOH CH3COO- (aq) + H+ (aq) A reversible reaction. The reaction can occur in both directions. Acetic acid is a weak electrolyte because its ionization in water is incomplete. 5 Hydration is the process in which an ion is surrounded by water molecules arranged in a specific manner. d- d+ H2O 6 Nonelectrolyte does not conduct electricity? No cations (+) and anions (-) in solution C6H12O6 (s) H 2O C6H12O6 (aq) 7 Precipitation Reactions Precipitate – insoluble solid that separates from solution precipitate Pb(NO3)2 (aq) + 2NaI (aq) PbI2 (s) + 2NaNO3 (aq) molecular equation Pb2+ + 2NO3- + 2Na+ + 2I- PbI2 (s) + 2Na+ + 2NO3- ionic equation Pb2+ + 2IPbI2 PbI2 (s) net ionic equation Na+ and NO3- are spectator ions 8 Solubility is the maximum amount of solute that will dissolve in a given quantity of solvent at a specific temperature. 9 Solubility Rules 1. Group IA and ammonium compounds are soluble. 2. Acetates and nitrates are soluble. 3. Most chlorides, bromides, and iodides are soluble. Exceptions: AgCl, Hg2Cl2, PbCl2; AgBr, HgBr2, Hg2Br2, PbBr2; AgI, HgI2, Hg2I2, PbI2 4. Most sulfates are soluble. Exceptions: CaSO4, SrSO4, BaSO4, Ag2SO4, Hg2SO4, PbSO4 10 5. Most carbonates are insoluble. Exceptions: Group IA carbonates and (NH4)2SO4 6. Most phosphates are insoluble. Exceptions: Group IA phosphates and (NH4)3PO4 7. Most sulfides are insoluble. Exceptions: Group IA sulfides and (NH4)2S 8. Most hydroxides are insoluble. Exceptions: Group IA hydroxides, Ca(OH)2, Sr(OH)2, Ba(OH)2 11 Decide whether the following reaction occurs. If it does, write the molecular, ionic, and net ionic equations. KBr + MgSO4 1. Determine the product formulas: K+ and SO42− make K2SO4 Mg2+ and Br − make MgBr2 2. Determine whether the products are soluble: K2SO4 is soluble MgBr2 is soluble KBr + MgSO4 no reaction 12 Examples of Insoluble Compounds CdS PbS Ni(OH)2 Al(OH)3 13 Writing Net Ionic Equations 1. Write the balanced molecular equation. 2. Write the ionic equation showing the strong electrolytes completely dissociated into cations and anions. 3. Cancel the spectator ions on both sides of the ionic equation 4. Check that charges and number of atoms are balanced in the net ionic equation Write the net ionic equation for the reaction of silver nitrate with sodium chloride. AgNO3 (aq) + NaCl (aq) AgCl (s) + NaNO3 (aq) Ag+ + NO3- + Na+ + Cl- AgCl (s) + Na+ + NO3- Ag+ + Cl- AgCl (s) 14 Properties of 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. 2HCl (aq) + Mg (s) MgCl2 (aq) + H2 (g) React with carbonates and bicarbonates to produce carbon dioxide gas 2HCl (aq) + CaCO3 (s) CaCl2 (aq) + CO2 (g) + H2O (l) Aqueous acid solutions conduct electricity. 15 Properties of Bases Have a bitter taste. Feel slippery. Many soaps contain bases. Cause color changes in plant dyes. Aqueous base solutions conduct electricity. Examples: 16 Arrhenius acid is a substance that produces H+ (H3O+) in water Arrhenius base is a substance that produces OH- in water 17 Hydronium ion, hydrated proton, H3O+ 18 In Figure A, a solution of HCl (a strong acid) illustrated on a molecular/ionic level, shows the acid as all ions. In Figure B, a solution of HF (a weak acid) also illustrated on a molecular/ionic level, shows mostly molecules with very few ions. 19 A Brønsted acid is a proton donor A Brønsted base is a proton acceptor base acid acid base A Brønsted acid must contain at least one ionizable proton! 20 Household Acids and Bases 21 4 | 21 Monoprotic acids HCl H+ + Cl- HNO3 Strong electrolyte, strong acid H+ + NO3H+ + CH3COO- CH3COOH Strong electrolyte, strong acid Weak electrolyte, weak acid Polyprotic Acid An acid that results in two or more acidic hydrogens per molecule. Diprotic acids H2SO4 H+ + HSO4- Strong electrolyte, strong acid HSO4- H+ + SO42- Weak electrolyte, weak acid Triprotic acids H3PO4 H2PO4HPO42- H+ + H2PO4H+ + HPO42H+ + PO43- Weak electrolyte, weak acid Weak electrolyte, weak acid Weak electrolyte, weak acid 22 23 Identify each of the following species as a Brønsted acid, base, or both. (a) HI, (b) CH3COO-, (c) H2PO4- HI (aq) H+ (aq) + I- (aq) CH3COO- (aq) + H+ (aq) H2PO4- (aq) Brønsted acid CH3COOH (aq) H+ (aq) + HPO42- (aq) H2PO4- (aq) + H+ (aq) H3PO4 (aq) Brønsted base Brønsted acid Brønsted base 24 Neutralization Reaction acid + base HCl (aq) + NaOH (aq) H+ + Cl- + Na+ + OH- H+ + OH- salt + water NaCl (aq) + H2O Na+ + Cl- + H2O H2O 25 Neutralization Reaction Involving a Weak Electrolyte weak acid + base HCN (aq) + NaOH (aq) HCN + Na+ + OH- HCN + OH- salt + water NaCN (aq) + H2O Na+ + CN- + H2O CN- + H2O 26 Neutralization Reaction Producing a Gas acid + base 2HCl (aq) + Na2CO3 (aq) 2H+ + 2Cl- + 2Na+ + CO32- 2H+ + CO32- salt + water + CO2 2NaCl (aq) + H2O +CO2 2Na+ + 2Cl- + H2O + CO2 H2O + CO2 27 Acid−Base Reaction with Gas Formation Some salts, when treated with an acid, produce a gas. Typically sulfides, sulfites, and carbonates behave in this way producing hydrogen sulfide, sulfur trioxide, and carbon dioxide, respectively. The photo to the right shows baking soda (sodium hydrogen carbonate) reacting with acetic acid in vinegar to give bubbles of carbon dioxide. 28 4 | 28 Oxidation-Reduction Reactions (electron transfer reactions) 2Mg O2 + 4e- 2Mg2+ + 4e- Oxidation half-reaction (lose e-) 2O2Reduction half-reaction (gain e-) 2Mg + O2 + 4e2Mg2+ + 2O2- + 4e29 2Mg + O2 2MgO Zn (s) + CuSO4 (aq) ZnSO4 (aq) + Cu (s) Zn2+ + 2e- Zn is oxidized Zn Cu2+ + 2e- Zn is the reducing agent Cu Cu2+ is reduced Cu2+ is the oxidizing agent Copper wire reacts with silver nitrate to form silver metal. What is the oxidizing agent in the reaction? Cu (s) + 2AgNO3 (aq) Cu Ag+ + 1e- Cu(NO3)2 (aq) + 2Ag (s) Cu2+ + 2eAg Ag+ is reduced Ag+ is the oxidizing 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, H2, O2, P4 = 0 2. In monatomic ions, the oxidation number is equal to the charge on the ion. Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2 3. The oxidation number of oxygen is usually –2. In H2O2 and O22- it is –1. 31 4.4 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. 5. 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 What are the oxidation numbers of all the elements in HCO3- ? O = –2 H = +1 3x(–2) + 1 + ? = –1 C = +4 32 The Oxidation Numbers of Elements in their Compounds 33 What are the oxidation numbers of all the elements in each of these compounds? NaIO3 IF7 K2Cr2O7 KIO3 K = +1 O = -2 3x(-2) + 1 + ? = 0 I = +5 Na2Cr2O7 IF7 F = -1 7x(-1) + ? = 0 I = +7 O = -2 Na = +1 7x(-2) + 2x(+1) + 2x(?) = 0 Cr = +6 34 Potassium permanganate, KMnO4, is a purple−colored compound; potassium manganate, K2MnO4, is a green−colored compound. Obtain the oxidation numbers of the manganese in these compounds. K Mn O 1(+1) + 1(oxidation number of Mn) + 4(−2) = 0 1 + 1(oxidation number of Mn) + (−8) = 0 (−7) + (oxidation number of Mn) = 0 Oxidation number of Mn = +7 35 K Mn O 2(+1) + 1(oxidation number of Mn) + 4(−2) = 0 2 + 1(oxidation number of Mn) + (−8) = 0 (−6) + (oxidation number of Mn) = 0 Oxidation number of Mn = +6 In KMnO4, the oxidation number of Mn is +7. In K2MnO4, the oxidation number of Mn is +6. 36 What is the oxidation number of Cr in dichromate, Cr2O72−? Cr O 2(oxidation number of Cr) + 7(−2) = −2 2(oxidation number of Cr) + (−14) = −2 2(oxidation number of Cr) = +12 Oxidation number of Cr = +6 37 Oxidation The half−reaction in which there is a loss of electrons by a species (or an increase in oxidation number). Reduction The half−reaction in which there is a gain of electrons by a species (or a decrease in oxidation number). 38 Oxidizing Agent A species that oxidizes another species; it is itself reduced. Reducing Agent A species that reduces another species; it is itself oxidized. 39 Types of Oxidation-Reduction Reactions Combination Reaction A reaction in which two substances combine to form a third substance. A+B 0 0 2Al + 3Br2 C +3 -1 2AlBr3 40 For example: 2Na(s) + Cl2(g) 2NaCl(s) 41 Decomposition Reaction C +1 +5 -2 2KClO3 A+B +1 -1 0 2KCl + 3O2 A reaction in which a single compound reacts to give two or more substances. For example: 2HgO(s) 2Hg(l) + O2(g) 42 Types of Oxidation-Reduction Reactions Combustion Reaction A reaction in which a substance reacts with oxygen, usually with the rapid release of heat to produce a flame. A + O2 B 0 0 S + O2 0 0 2Mg + O2 +4 -2 SO2 +2 -2 2MgO 43 Combustion Reaction For example: 4Fe(s) + 3O2(g) 2Fe2O3(s) 44 4 | 44 Types of Oxidation-Reduction Reactions Displacement Reaction A reaction in which an element reacts with a compound, displacing another element from it. A + BC 0 +1 +2 Sr + 2H2O +4 0 TiCl4 + 2Mg 0 AC + B -1 Cl2 + 2KBr 0 Sr(OH)2 + H2 0 Hydrogen Displacement +2 Ti + 2MgCl2 -1 Metal Displacement 0 2KCl + Br2 Halogen Displacement 45 Displacement Reaction For example: Zn(s) + 2HCl(aq) H2(g) + ZnCl2(aq) 46 4 | 46 Types of Oxidation-Reduction Reactions Disproportionation Reaction The same element is simultaneously oxidized and reduced. Example: reduced +1 0 Cl2 + 2OH- -1 ClO- + Cl- + H2O oxidized 47 Classify each of the following reactions. Ca2+ + CO32NH3 + H+ Zn + 2HCl Ca + F2 CaCO3 NH4+ ZnCl2 + H2 CaF2 Precipitation Acid-Base Redox (H2 Displacement) Redox (Combination) Mg(s) Mg2+(aq) + 2e− (oxidation) Fe3+(aq) + 3e− Fe(s) (reduction) 48 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 Molality= amount of substance in mol of solute/mass in Kg of the solvent What mass of KI is required to make 500. mL of a 2.80 M KI solution? M KI M KI volume of KI solution 500. mL x 1L 1000 mL moles KI x 2.80 mol KI 1 L soln x 166 g KI 1 mol KI grams KI = 232 g KI 49 To prepare a solution, add the measured amount of solute to a volumetric flask, then add water to bring the solution to the mark on the flask. 50 Preparing a Solution of Known Concentration 51 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) MiVi = MfVf 52 Diluting a solution quantitatively requires specific glassware. The photo at the right shows a volumetric flask used in dilution. 53 You place a 1.52−g of potassium dichromate, K2Cr2O7, into a 50.0−mL volumetric flask. You then add water to bring the solution up to the mark on the neck of the flask. What is the molarity of K2Cr2O7 in the solution? Molar mass of K2Cr2O7 is 294 g. 1 mol 1.52 g 294 g 0.103 M -3 50.0 10 L 54 A saturated stock solution of NaCl is 6.00 M. How much of this stock solution is needed to prepare 1.00−L of physiological saline soluiton (0.154 M)? M iVi M fVf M fVf Vi Mi (0.154 M )(1.00 L) Vi 6.00 M Vi 0.0257 L or 25.7 mL 55 How would you prepare 30.0 mL of 0.100 M HNO3 from a stock solution of 2.00 M HNO3? MiVi = MfVf Mi = 2.00 M Mf = 0.100 M Vf = 0.0300 L Vi = MfVf Mi Vi = ? L = 0.100 M x 0.0300 L = 0.00150 L = 1.50 mL 2.00 M Dilute 1.50 mL of acid with water to a total volume of 30.0 mL. 56