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

Describing Chemical Reactions A chemical reaction is the process by which one or more substances are changed into one or more different substances (NH 4 ) 2 Cr 2 O 7(s)  N 2(g) + Cr 2 O 3(s) + 4H 2 O (g) “The reactant ammonium dichromate yields the products nitrogen, chromium (III) oxide and water” A CHEMICAL EQUATION represents, with symbols and formulas, the identifies and relative amounts of the reactants and products in a chemical equation reactants products 2

Indications of a Chemical Reaction 1. Evolution of heat and light is strong evidence that a chemical reaction has taken place! But, the evolution of heat or light by itself is not necessarily a sign of a chemical change since many physical changes also release either heat or light. 2. Production of gas! (aka bubbles when two substances are mixed) 3. Formation of precipitate! A solid that is produced as a result of a chemical reaction in solution and that separates from the solution is known as a precipitate 4. Color Change! 3

Characteristics of Chemical Equations 1. The equation must represent all reactants and products. 2. The equation must contain the correct formulas for the reactants and products 3. The law of conservation of mass MUST be satisfied!! Law of conservation of mass – atoms are neither created nor destroyed in ordinary chemical reactions To equalized numbers of atoms, coefficients are added in front of the formulas where necessary 4

Types of Chemical Reactions A + B ↔ C C ↔ A + B A + BC ↔ AC + B AB + CD ↔ AD + CB Acid + Base ↔ salt + water Change of oxidation state Hydrocarbon + O 2 ↔ CO 2 + H 2 O 5 Synthesis (Combination) Decomposition Single Replacement Precipitation Reactions (Double Replacement Reactions) Neutralization Reactions (Acid/Base) Redox Reactions Combustion

PRECIPITATION REACTIONS Double Replacement Reactions AB + CD ↔ AD + BC 6

7 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

8 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. nonelectrolyte weak electrolyte strong electrolyte

9 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 (-)

10 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.

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

12 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

Precipitation Reactions 13 Precipitate – insoluble solid that separates from solution molecular equation ionic equation net ionic equation Pb NO K + + 2I - PbI 2 (s) + 2K + + 2NO 3 - K + and NO 3 - are spectator ions PbI 2 Pb(NO 3 ) 2 (aq) + 2KI (aq) PbI 2 (s) + 2KNO 3 (aq) precipitate Pb I - PbI 2 (s)

Precipitation of Lead Iodide 14 PbI 2 Pb I - PbI 2 (s)

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

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

Problem Characterize the following compounds as (a) soluble or (b) insoluble in water: 1. CaCO 3 2. ZnSO 4 3. Hg(NO 3 ) 2 4. HgSO 4 5. NH 4 ClO 4

Writing Net Ionic Equations 18 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 + (aq) + NO 3 - (aq) + Na + (aq) + Cl - (aq) AgCl (s) + Na + (aq) + NO 3 - (aq) Ag + (aq) + Cl - (aq) AgCl (s) Write the net ionic equation for the reaction of silver nitrate with sodium chloride.reaction

19 PredictPredict what happens when a potassium hydroxide solution is mixed with a solution of sodium chloride. Write a net ionic equation for the reaction.

20 What do we expect to see if we add copper (II) sulfate to sodium hydroxide? Write the molecular equation, ionic equation, and net ionic equation.

21 What do we expect to see if we add copper (II) sulfate to sodium hydroxide? Write the molecular equation, ionic equation, and net ionic equation.

22 Predict what happens when a potassium phosphate solution is mixed with a solution of calcium nitrate. Write a net ionic equation for the reaction. EXTRA PRACTICE

23 PredictPredict what happens when a silver nitrate solution is mixed with a solution of potassium hydroxide. Write a net ionic equation for the reaction.

Types of Chemical Reactions A + B ↔ C C ↔ A + B A + BC ↔ AC + B AB + CD ↔ AD + BC Acid + Base ↔ salt + water Change of oxidation state Hydrocarbon + O 2 ↔ CO 2 + H 2 O 24 Synthesis (Combination) Decomposition Single Replacement Precipitation Reactions (Double Replacement Reactions) Neutralization Reactions (Acid/Base) Redox Reactions Combustion

NEUTRALIZATION REACTIONS Acid/Base Reactions Acid + Base ↔ Salt + Water 25

Properties of Acids 26 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.

27 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:

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

29 Hydronium ion, hydrated proton, H 3 O +

30 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!

31 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

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33 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

Problem Identify each of the following as either a (a) Brønsted acid, (b) Brønsted base, or (c) both. 1. PO ClO NH HCO 3 -

Neutralization Reaction 35 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

Neutralization Reaction Involving a Weak Electrolyte 36 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

Neutralization Reaction Producing a Gas 37 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

Types of Chemical Reactions A + B ↔ C C ↔ A + B A + BC ↔ AC + B AB + CD ↔ AD + BC Acid + Base ↔ salt + water Change of oxidation state Hydrocarbon + O 2 ↔ CO 2 + H 2 O 38 Synthesis (Combination) Decomposition Single Replacement Precipitation Reactions (Double Replacement Reactions) Neutralization Reactions (Acid/Base) Redox Reactions Combustion

OXIDATION REDUCTION REACTIONS Redox Reactions Synthesis Reactions:A + B ↔ C Decomposition Reactions:C ↔ A + B Single Replacement Reactions: A + BC ↔ AC + B Combustion Reactions:hydrocarbon + O 2 ↔ CO 2 + H 2 O 39

Oxidation-Reduction Reactions 40 (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

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42 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?

Oxidation number 43 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

44 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.

45 The Oxidation Numbers of Elements in their Compounds

46 NaIO 3 Na = +1O = -2 3x(-2) ? = 0 I = +5 IF 7 F = -1 7x(-1) + ? = 0 I = +7 K 2 Cr 2 O 7 O = -2K = +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

Problem Give the oxidation number for the underlined atoms or in each of the following species: a) Mg 3 N 2 b) CsO 2 c) CaC 2 d) CO 3 2- e) C 2 O 4 2- f) ZnO 2 2- g) NaBH 4 h) WO 4 2-

Types of Oxidation-Reduction Reactions 48 Synthesis/Combination Reaction A + B C 2Al + 3Br 2 2AlBr 3 Decomposition Reaction 2KClO 3 2KCl + 3O 2 C A + B

Types of Oxidation-Reduction Reactions 49 Combustion Reaction A + O 2 B S + O 2 SO Mg + O 2 2MgO

Types of Oxidation-Reduction Reactions 50 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

51 The Activity Series for Metals M + BC MC + B Hydrogen Displacement Reaction M is metal BC is acid or H 2 O B is H 2 Ca + 2H 2 O Ca(OH) 2 + H 2 Pb + 2H 2 O Pb(OH) 2 + H 2

Problem Which of the following metals can react with water to produce H 2 (g)? a) Au b) Li c) Hg d) Ca e) Pt

53 The Activity Series for Halogens Halogen Displacement Reaction Cl 2 + 2KBr 2KCl + Br F 2 > Cl 2 > Br 2 > I 2 I 2 + 2KBr 2KI + Br 2

Types of Oxidation-Reduction Reactions 54 The same element is simultaneously oxidized and reduced. Example: Disproportionation Reaction Cl 2 + 2OH - ClO - + Cl - + H 2 O 0 +1 oxidized reduced

55 Ca 2+ + CO 3 2- CaCO 3 NH 3 + H + NH 4 + Zn + 2HCl ZnCl 2 + H 2 Ca + F 2 CaF 2 Classify each of the following reactions.

Problem Predict the outcome of the reactions represented by the following equations by using the activity series, and balance the equations. Cu (s) + HCl (aq)  I 2 (g) + NaBr (aq)  Mg (s) + CuSO 4 (aq)  Cl 2 (g) + KBr (aq) 

Problem Classify the following redox reactions by type: P Cl 2  4PCl 5 2NO  N 2 + O 2 Cl 2 + 2KI  I 2 + 2KCl

Solution Stoichiometry Molarity Dilutions Gravimetric Analysis Titrations 58

59 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 5.00 x 10 2 mL of a 2.80 M KI solution? volume of KI solutionmoles KIgrams KI M KI 5.00x10 2 mL = 232 g KI 166 g KI 1 mol KI x 2.80 mol KI 1 L soln x 1 L 1000 mL x

Problem Calculate the mass in grams of sodium nitrate required to prepare 2.50 x 10 2 mL of a M solution.

61 Preparing a Solution of Known Concentration

62 DilutionDilution 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 =

63 How would you prepare 60.0 mL of M HNO 3 from a stock solution of 4.00 M HNO 3 ? M i V i = M f V f M i = 4.00 M M f = MV f = L V i = ? L V i = MfVfMfVf MiMi = M x L 4.00 M = L = 3.00 mL Dilute 3.00 mL of acid with water to a total volume of 60.0 mL.

Problem Water is added to 25.0 mL of a M KNO 3 solution until the volume of the solution is exactly 500 mL. What is the concentration of the final solution?

Problem You have 505 mL of a M HCl solution and you want to dilute it to exactly M. How much water should you add? (assume that the volumes are additive.)

66 Gravimetric Analysis 1.Dissolve unknown substance in water 2.React unknown with known substance to form a precipitate 3.Filter and dry precipitate 4.Weigh precipitate 5.Use chemical formula and mass of precipitate to determine amount of unknown ion

67 A g sample of an ionic compound containing chloride ions and an unknown metal is dissolved in water, and treated with excess AgNO 3. If g of AgCl precipitate forms, what is the percent by mass of Cl in the original compound?

Problem A sample of g of an unknown compound containing barium ions (Ba 2+ ) is dissolved in water and treated with an excess of Na 2 SO 4. If the mass of the BaSO 4 precipitate formed is g, what is the percent by mass of Ba in the original unknown compound?

69 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

70 Titrations can be used in the analysis of Acid-base reactions Redox reactions H 2 SO 4 + 2NaOH 2H 2 O + Na 2 SO 4 5Fe 2+ + MnO H + Mn Fe H 2 O

71 What volume of a M NaOH solution is required to titrate mL of a 4.50 M H 2 SO 4 solution? WRITE THE CHEMICAL EQUATION! volume acidmoles redmoles basevolume base H 2 SO 4 + 2NaOH 2H 2 O + Na 2 SO mol H 2 SO mL soln x 2 mol NaOH 1 mol H 2 SO 4 x 1000 ml soln mol NaOH x mL = 158 mL M acid rxn coef. M base

Problem Calculate the concentration (in molarity) of a NaOH solution if 25.0 mL of the solution are needed to neutralize 17.4 mL of a M HCl solution.

73 WRITE THE CHEMICAL EQUATION! volume redmoles redmoles oxidM oxid mol KMnO 4 1 L x 5 mol Fe 2+ 1 mol KMnO 4 x L Fe 2+ x L = M M red rxn coef. V oxid 5Fe 2+ + MnO H + Mn Fe H 2 O mL of M KMnO 4 solution is needed to oxidize mL of an acidic FeSO 4 solution. What is the molarity of the iron solution? mL = L25.00 mL = L

Problem The SO 2 present in air is mainly responsible for the acid rain phenomenon. Its concentration can be determined by titrating against a standard permanganate solution as follows: 5SO 2 + 2MnO H 2 O  5SO Mn H + Calculate the number of grams of SO 2 in a sample of air if 7.37 mL of M KMnO 4 solution are required for the titration.