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

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Presentation on theme: "Reactions in Aqueous Solution Chapter 4 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display."— Presentation transcript:

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

2 2 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 SolutionSolventSolute Soft drink (l) Air (g) Soft Solder (s) H2OH2O N2N2 Pb Sugar, CO 2 O 2, Ar, CH 4 Sn aqueous solutions of KMnO 4

3 3 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. nonelectrolyte weak electrolyte strong electrolyte

4 4 Strong Electrolyte – 100% dissociation NaCl (s) Na + (aq) + Cl - (aq) H2OH2O Weak Electrolyte – not completely dissociated CH 3 COOH CH 3 COO - (aq) + H + (aq) Conduct electricity in solution? Cations (+) and Anions (-)

5 5 Ionization of acetic acid CH 3 COOH CH 3 COO - (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.

6 6 Hydration is the process in which an ion is surrounded by water molecules arranged in a specific manner.   H2OH2O

7 7 Nonelectrolyte does not conduct electricity? No cations (+) and anions (-) in solution C 6 H 12 O 6 (s) C 6 H 12 O 6 (aq) H2OH2O

8 8 Precipitation Reactions Precipitate – insoluble solid that separates from solution molecular equation ionic equation net ionic equation Pb NO Na + + 2I - PbI 2 (s) + 2Na + + 2NO 3 - Na + and NO 3 - are spectator ions PbI 2 Pb(NO 3 ) 2 (aq) + 2NaI (aq) PbI 2 (s) + 2NaNO 3 (aq) precipitate Pb I - PbI 2 (s)

9 9 Solubility is the maximum amount of solute that will dissolve in a given quantity of solvent at a specific temperature.

10 10 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, Hg 2 Cl 2, PbCl 2 ; AgBr, HgBr 2, Hg 2 Br 2, PbBr 2 ; AgI, HgI 2, Hg 2 I 2, PbI 2 4.Most sulfates are soluble. Exceptions: CaSO 4, SrSO 4, BaSO 4, Ag 2 SO 4, Hg 2 SO 4, PbSO 4

11 11 5.Most carbonates are insoluble. Exceptions: Group IA carbonates and (NH 4 ) 2 SO 4 6.Most phosphates are insoluble. Exceptions: Group IA phosphates and (NH 4 ) 3 PO 4 7.Most sulfides are insoluble. Exceptions: Group IA sulfides and (NH 4 ) 2 S 8.Most hydroxides are insoluble. Exceptions: Group IA hydroxides, Ca(OH) 2, Sr(OH) 2, Ba(OH) 2

12 12 Decide whether the following reaction occurs. If it does, write the molecular, ionic, and net ionic equations. KBr + MgSO 4  1.Determine the product formulas: K + and SO 4 2− make K 2 SO 4 Mg 2+ and Br − make MgBr 2 2.Determine whether the products are soluble: K 2 SO 4 is soluble MgBr 2 is soluble KBr + MgSO 4  no reaction

13 13 Examples of Insoluble Compounds CdS PbSNi(OH) 2 Al(OH) 3

14 14 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 AgNO 3 (aq) + NaCl (aq) AgCl (s) + NaNO 3 (aq) Ag + + NO Na + + Cl - AgCl (s) + Na + + NO 3 - Ag + + Cl - AgCl (s) Write the net ionic equation for the reaction of silver nitrate with sodium chloride.

15 15 Properties of 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 Cause color changes in plant dyes. 2HCl (aq) + Mg (s) MgCl 2 (aq) + H 2 (g) 2HCl (aq) + CaCO 3 (s) CaCl 2 (aq) + CO 2 (g) + H 2 O (l) Aqueous acid solutions conduct electricity.

16 16 Have a bitter taste. Feel slippery. Many soaps contain bases. Properties of Bases Cause color changes in plant dyes. Aqueous base solutions conduct electricity. Examples:

17 17 Arrhenius acid is a substance that produces H + (H 3 O + ) in water Arrhenius base is a substance that produces OH - in water

18 18 Hydronium ion, hydrated proton, H 3 O +

19 19 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.

20 20 A Brønsted acid is a proton donor A Brønsted base is a proton acceptor acidbaseacidbase A Brønsted acid must contain at least one ionizable proton!

21 21 4 | 21 Household Acids and Bases

22 22 Monoprotic acids HCl H + + Cl - HNO 3 H + + NO 3 - CH 3 COOH H + + CH 3 COO - Strong electrolyte, strong acid Weak electrolyte, weak acid Diprotic acids H 2 SO 4 H + + HSO 4 - HSO 4 - H + + SO 4 2- Strong electrolyte, strong acid Weak electrolyte, weak acid Triprotic acids H 3 PO 4 H + + H 2 PO 4 - H 2 PO 4 - H + + HPO 4 2- HPO 4 2- H + + PO 4 3- Weak electrolyte, weak acid Polyprotic Acid An acid that results in two or more acidic hydrogens per molecule.

23 23

24 24 Identify each of the following species as a Brønsted acid, base, or both. (a) HI, (b) CH 3 COO -, (c) H 2 PO 4 - HI (aq) H + (aq) + I - (aq)Brønsted acid CH 3 COO - (aq) + H + (aq) CH 3 COOH (aq)Brønsted base H 2 PO 4 - (aq) H + (aq) + HPO 4 2- (aq) H 2 PO 4 - (aq) + H + (aq) H 3 PO 4 (aq) Brønsted acid Brønsted base

25 25 Neutralization Reaction acid + base salt + water HCl (aq) + NaOH (aq) NaCl (aq) + H 2 O H + + Cl - + Na + + OH - Na + + Cl - + H 2 O H + + OH - H 2 O

26 26 Neutralization Reaction Involving a Weak Electrolyte weak acid + base salt + water HCN (aq) + NaOH (aq) NaCN (aq) + H 2 O HCN + Na + + OH - Na + + CN - + H 2 O HCN + OH - CN - + H 2 O

27 27 Neutralization Reaction Producing a Gas acid + base salt + water + CO 2 2HCl (aq) + Na 2 CO 3 (aq) 2NaCl (aq) + H 2 O +CO 2 2H + + 2Cl - + 2Na + + CO Na + + 2Cl - + H 2 O + CO 2 2H + + CO 3 2- H 2 O + CO 2

28 28 4 | 28 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.

29 29 Oxidation-Reduction Reactions (electron transfer reactions) 2Mg 2Mg e - O 2 + 4e - 2O 2- Oxidation half-reaction (lose e - ) Reduction half-reaction (gain e - ) 2Mg + O 2 + 4e - 2Mg O e - 2Mg + O 2 2MgO

30 30 Zn (s) + CuSO 4 (aq) ZnSO 4 (aq) + Cu (s) Zn is oxidizedZn Zn e - Cu 2+ is reducedCu e - Cu Zn is the reducing agent Cu 2+ is the oxidizing agent Copper wire reacts with silver nitrate to form silver metal. What is the oxidizing agent in the reaction? Cu (s) + 2AgNO 3 (aq) Cu(NO 3 ) 2 (aq) + 2Ag (s) Cu Cu e - Ag + + 1e - AgAg + is reducedAg + is the oxidizing agent

31 31 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, H 2, O 2, P 4 = 0 2.In monatomic ions, the oxidation number is equal to the charge on the ion. Li +, Li = +1; Fe 3+, Fe = +3; O 2-, O = -2 3.The oxidation number of oxygen is usually –2. In H 2 O 2 and O 2 2- it is –1. 4.4

32 32 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. 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. 5.Group IA metals are +1, IIA metals are +2 and fluorine is always –1. HCO 3 - O = – 2H = +1 3x( – 2) ? = – 1 C = +4 What are the oxidation numbers of all the elements in HCO 3 - ? 7. Oxidation numbers do not have to be integers. Oxidation number of oxygen in the superoxide ion, O 2 -, is –½.

33 33 The Oxidation Numbers of Elements in their Compounds

34 34 KIO 3 K = +1O = -2 3x(-2) ? = 0 I = +5 IF 7 F = -1 7x(-1) + ? = 0 I = +7 Na 2 Cr 2 O 7 O = -2 Na = +1 7x(-2) + 2x(+1) + 2x(?) = 0 Cr = +6 What are the oxidation numbers of all the elements in each of these compounds? NaIO 3 IF 7 K 2 Cr 2 O 7

35 35 Potassium permanganate, KMnO 4, is a purple−colored compound; potassium manganate, K 2 MnO 4, is a green−colored compound. Obtain the oxidation numbers of the manganese in these compounds. KMnO 1(+1) + 1(oxidation number of Mn) + 4(−2) = (oxidation number of Mn) + (−8) = 0 (−7) + (oxidation number of Mn) = 0 Oxidation number of Mn = +7

36 36 KMnO 2(+1) + 1(oxidation number of Mn) + 4(−2) = (oxidation number of Mn) + (−8) = 0 (−6) + (oxidation number of Mn) = 0 Oxidation number of Mn = +6 In KMnO 4, the oxidation number of Mn is +7. In K 2 MnO 4, the oxidation number of Mn is +6.

37 37 What is the oxidation number of Cr in dichromate, Cr 2 O 7 2− ? CrO 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

38 38 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).

39 39 Oxidizing Agent A species that oxidizes another species; it is itself reduced. Reducing Agent A species that reduces another species; it is itself oxidized.

40 40 Types of Oxidation-Reduction Reactions Combination Reaction A + B C 2Al + 3Br 2 2AlBr A reaction in which two substances combine to form a third substance.

41 41 For example:2Na(s) + Cl 2 (g)  2NaCl(s)

42 42 Decomposition Reaction 2KClO 3 2KCl + 3O 2 C A + B A reaction in which a single compound reacts to give two or more substances. For example: 2HgO(s)  2Hg(l) + O 2 (g)

43 43 Types of Oxidation-Reduction Reactions Combustion Reaction A + O 2 B S + O 2 SO Mg + O 2 2MgO A reaction in which a substance reacts with oxygen, usually with the rapid release of heat to produce a flame.

44 44 4 | 44 For example: 4Fe(s) + 3O 2 (g)  2Fe 2 O 3 (s) Combustion Reaction

45 45 Displacement Reaction A + BC AC + B Sr + 2H 2 O Sr(OH) 2 + H 2 TiCl 4 + 2Mg Ti + 2MgCl 2 Cl 2 + 2KBr 2KCl + Br 2 Hydrogen Displacement Metal Displacement Halogen Displacement Types of Oxidation-Reduction Reactions A reaction in which an element reacts with a compound, displacing another element from it.

46 46 4 | 46 Displacement Reaction For example: Zn(s) + 2HCl(aq)  H 2 (g) + ZnCl 2 (aq)

47 47 The same element is simultaneously oxidized and reduced. Example: Disproportionation Reaction Cl 2 + 2OH - ClO - + Cl - + H 2 O Types of Oxidation-Reduction Reactions 0 +1 oxidized reduced

48 48 Ca 2+ + CO 3 2- CaCO 3 NH 3 + H + NH 4 + Zn + 2HCl ZnCl 2 + H 2 Ca + F 2 CaF 2 Precipitation Acid-Base Redox (H 2 Displacement) Redox (Combination) Classify each of the following reactions. Mg(s)  Mg 2+ (aq) + 2e − (oxidation) Fe 3+ (aq) + 3e −  Fe(s) (reduction)

49 49 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 What mass of KI is required to make 500. mL of a 2.80 M KI solution? volume of KI solutionmoles KIgrams KI M KI 500. mL= 232 g KI 166 g KI 1 mol KI x 2.80 mol KI 1 L soln x 1 L 1000 mL x Molality= amount of substance in mol of solute/mass in Kg of the solvent

50 50 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.

51 51 Preparing a Solution of Known Concentration

52 52 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) = MiViMiVi MfVfMfVf =

53 53 Diluting a solution quantitatively requires specific glassware. The photo at the right shows a volumetric flask used in dilution.

54 54 You place a 1.52−g of potassium dichromate, K 2 Cr 2 O 7, 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 K 2 Cr 2 O 7 in the solution? Molar mass of K 2 Cr 2 O 7 is 294 g M

55 55 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)?

56 56 How would you prepare 30.0 mL of M HNO 3 from a stock solution of 2.00 M HNO 3 ? M i V i = M f V f M i = 2.00 M M f = MV f = L V i = ? L V i = MfVfMfVf MiMi = M x L 2.00 M = L = 1.50 mL Dilute 1.50 mL of acid with water to a total volume of 30.0 mL.


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