General Chemistry Principles and Modern Applications Petrucci Harwood Herring 8 th Edition Chapter 5: Introduction to Reactions in Aqueous Solutions Philip Dutton University of Windsor, Canada N9B 3P4 Contents 5-1 The Nature of Aqueous Solutions 5-2 Precipitation Reactions 5-3 Acid-Base Reactions 5-4 Oxidation-Reduction: Some General Principles 5-5 Balancing Oxidation-Reduction Equations 5-6 Oxidizing and Reducing Agents 5-7 Stoichiometry of Reactions in Aqueous Solutions: Titrations Focus on Water Treatment Slide 1 of 43 Slide 2 of 43 5.1 The Nature of Aqueous Solutions Electrolytes Some solutes can dissociate into ions. Electric charge can be carried. Slide 3 of 43 Slide 4 of 43 Types of Electrolytes Strong electrolyte dissociates completely. Good electrical conduction. Weak electrolyte partially dissociates. Fair conductor of electricity. Non-electrolyte does not dissociate. Poor conductor of electricity. Representation of Electrolytes using Chemical Equations A strong electrolyte: MgCl 2 (s) Mg 2+ (aq) + 2 Cl - (aq) A weak electrolyte: CH 3 CO 2 H(aq) CH 3 CO 2 - (aq) + H + (aq) A non-electrolyte: CH 3 OH(aq) Slide 5 of 43 Slide 6 of 43
Notation for Concentration MgCl 2 (s) Mg 2+ (aq) + 2 Cl - (aq) Example 5-1 Calculating Ion concentrations in a Solution of a Strong Electolyte. What are the aluminum and sulfate ion concentrations in 0.0165 M Al 2 (SO 4 ) 3?. In 0.0050 M MgCl 2: Stoichiometry is important. [Mg 2+ ] = 0.0050 M [Cl - ] = 0.0100 M [MgCl 2 ] = 0 M Balanced Chemical Equation: Al 2 (SO 4 ) 3 (s) 2 Al 3+ (aq) + 3 SO 4 (aq) Slide 7 of 43 Slide 8 of 43 Example 5-1 5-2 Precipitation Reactions Aluminum Concentration: 0.0165 mol Al 2 mol Al [Al] = 2 (SO 4 ) 3 1 L 1 mol Al 2 (SO 4 ) 3 = 0.0330 M Al 3+ Soluble ions can combine to form an insoluble compound. Precipitation occurs. Sulfate Concentration: [SO 0.0165 mol Al 3 mol SO 4 ] = 2 (SO 4 ) 3 4 = 1 L 1 mol Al 2 (SO 4 ) 3 0.0495 M SO 4 Ag + (aq) + Cl - (aq) AgCl(s) Slide 9 of 43 Slide 10 of 43 Net Ionic Equation Solubility Rules Overall Precipitation Reaction: AgNO 3 (aq) +NaI (aq) AgI(s) + NaNO 3 (aq) Complete ionic equation: Spectator ions Ag + (aq) + NO 3- (aq) + Na + (aq) + I - (aq) AgI(s) + Na + (aq) + NO 3- (aq) Net ionic equation: Ag + (aq) + I - (aq) AgI(s) Compounds that are soluble: Alkali metal ion and ammonium ion salts Li +, Na +, K +, Rb +, Cs + NH + 4 Nitrates, perchlorates and acetates NO - 3 ClO - 4 CH 3 CO - 2 Slide 11 of 43 Slide 12 of 43
Solubility Rules Solubility Rules Compounds that are mostly soluble: Chlorides, bromides and iodides Cl -, Br -, I - Except those of Pb 2+, Ag +, and Hg 2 2+. Sulfates SO 4 Except those of Sr 2+, Ba 2+, Pb 2+ and Hg 2 2+. Ca(SO 4 ) is slightly soluble. Compounds that are insoluble: Hydroxides and sulfides HO -, S Except alkali metal and ammonium salts Sulfides of alkaline earths are soluble Hydroxides of Sr 2+ and Ca 2 + are slightly soluble. Carbonates and phosphates CO 3, PO 4 3- Except alkali metal and ammonium salts Slide 13 of 43 Slide 14 of 43 5-3 Acid-Base Reactions Latin acidus (sour) Sour taste Arabic al-qali (ashes of certain plants) Bitter taste Acids Acids provide H + in aqueous solution. Strong acids: HCl(aq) H + (aq) + Cl - (aq) Svante Arrhenius 1884 Acid-Base theory. Weak acids: CH 3 CO 2 H(aq) H + (aq) + CH 3 CO (aq) Slide 15 of 43 Slide 16 of 43 Bases Bases provide OH - in aqueous solution. Strong bases: NaOH(aq) H 2 O Na + (aq) + OH - (aq) Recognizing Acids and Bases. Acids have ionizable hydrogen ions. CH 3 CO 2 H or HC 2 H 3 O 2 Bases have OH - combined with a metal ion. KOH Weak bases: or are identified by chemical equations NH 3 (aq) + H 2 O(l) OH - (aq) + NH 4+ (aq) Na 2 CO 3 (s) + H 2 O(l) HCO 3- (aq) + 2 Na + (aq) + OH - (aq) Slide 17 of 43 Slide 18 of 43
More Acid-Base Reactions More Acid-Base Reactions Milk of magnesia Mg(OH) 2 Mg(OH) 2 (s) + 2 H + (aq) Mg 2+ (aq) + 2 H 2 O(l) Limestone and marble. CaCO 3 (s) + 2 H + (aq) Ca 2+ (aq) + H 2 CO 3 (aq) Mg(OH) 2 (s) + 2 CH 3 CO 2 H(aq) Mg 2+ (aq) + 2 CH 3 CO (aq) + 2 H 2 O(l) But: H 2 CO 3 (aq) H 2 O(l) + CO 2 (g) CaCO 3 (s) + 2 H + (aq) Ca 2+ (aq) + H 2 O(l) + CO 2 (g) Slide 19 of 43 Slide 20 of 43 Limestone and Marble Gas Forming Reactions Slide 21 of 43 Slide 22 of 43 5-4 Oxidation-Reduction: Some General Principles Hematite is converted to iron in a blast furnace. Fe 2 O 3 (s) + 3 CO(g) 2 Fe(l) + 3 CO 2 (g) Oxidation and reduction always occur together. Fe 3+ is reduced to metallic iron. CO(g) is oxidized to carbon dioxide. Oxidation State Changes Assign oxidation states: 3+ 2+ 0 4+ Fe 2 O 3 (s) + 3 CO(g) 2 Fe(l) + 3 CO 2 (g) Fe 3+ is reduced to metallic iron. CO(g) is oxidized to carbon dioxide. Slide 23 of 43 Slide 24 of 43
Oxidation and Reduction Zinc in Copper Sulfate Oxidation O.S. of some element increases in the reaction. Electrons are on the right of the equation Reduction O.S. of some element decreases in the reaction. Electrons are on the left of the equation. Zn(s) + Cu 2+ (aq) Zn 2+ (aq) + Cu(s) Slide 25 of 43 Slide 26 of 43 Half-Reactions Represent a reaction by two half-reactions. Oxidation: Zn(s) Zn 2+ (aq) + 2 e - Balancing Oxidation-Reduction Equations Few can be balanced by inspection. Systematic approach required. Reduction: Overall: Cu 2+ (aq) + 2 e- Cu(s) Cu 2+ (aq) + Zn(s) Cu(s) + Zn 2+ (aq) The Half-Reaction (Ion-Electron) Method Slide 27 of 43 Slide 28 of 43 Example 5-6 Example 5-6 Balancing the Equation for a Redox Reaction in Acidic Solution. The reaction described below is used to determine the sulfite ion concentration present in wastewater from a papermaking plant. Write the balanced equation for this reaction in acidic solution.. SO 3 (aq) + MnO 4- (aq) SO 4 (aq) + Mn 2+ (aq) Determine the oxidation states: 4+ 7+ 6+ 2+ SO 3 (aq) + MnO 4- (aq) SO 4 (aq) + Mn 2+ (aq) Write the half-reactions: SO 3 (aq) SO 4 (aq) + 2 e - (aq) 5 e - (aq) +MnO 4- (aq) Mn 2+ (aq) Balance atoms other than H and O: Already balanced for elements. Slide 29 of 43 Slide 30 of 43
Example 5-6 Balance O by adding H 2 O: H 2 O(l) + SO 3 (aq) SO 4 (aq) + 2 e - (aq) Example 5-6 Multiply the half-reactions to balance all e - : 5 H 2 O(l) + 5 SO 3 (aq) 5 SO 4 (aq) + 10 e - (aq) + 10 H + (aq) 5 e - (aq) +MnO 4- (aq) Mn 2+ (aq) + 4 H 2 O(l) Balance hydrogen by adding H + : H 2 O(l) + SO 3 (aq) SO 4 (aq) + 2 e - (aq) + 2 H + (aq) 8 H + (aq) + 5 e - (aq) +MnO 4- (aq) Mn 2+ (aq) + 4 H 2 O(l) 16 H + (aq) + 10 e - (aq) + 2 MnO 4- (aq) 2 Mn 2+ (aq) + 8 H 2 O(l) Add both equations and simplify: 5 SO 3 (aq) + 2 MnO 4- (aq) + 6H + (aq) 5 SO 4 (aq) + 2 Mn 2+ (aq) + 3 H 2 O(l) Check that the charges are balanced: Add e - if necessary. Check the balance! Slide 31 of 43 Slide 32 of 43 Balancing in Acid Write the equations for the half-reactions. Balance all atoms except H and O. Balance oxygen using H 2 O. Balance hydrogen using H +. Balance charge using e -. Equalize the number of electrons. Add the half reactions. Check the balance. Balancing in Basic Solution OH - appears instead of H +. Treat the equation as if it were in acid. Then add OH - to each side to neutralize H +. Remove H 2 O appearing on both sides of equation. Check the balance. Slide 33 of 43 Slide 34 of 43 5-6 Oxidizing and Reducing Agents. Redox An oxidizing agent (oxidant ): Contains an element whose oxidation state decreases in a redox reaction A reducing agent (reductant): Contains an element whose oxidation state increases in a redox reaction. Slide 35 of 43 Slide 36 of 43
Example 5-8 Identifying Oxidizing and Reducing Agents. Hydrogen peroxide, H 2 O 2, is a versatile chemical. Its uses include bleaching wood pulp and fabrics and substituting for chlorine in water purification. One reason for its versatility is that it can be either an oxidizing or a reducing agent. For the following reactions, identify whether hydrogen peroxide is an oxidizing or reducing agent. Example 5-8 H 2 O 2 (aq) + 2 Fe 2+ (aq) + 2 H + 2 H 2 O(l) + 2 Fe 3+ (aq) Iron is oxidized and peroxide is reduced. 5 H 2 O 2 (aq) + 2 MnO 4- (aq) + 6 H + 8 H 2 O(l) + 2 Mn 2+ (aq) + 5 O 2 (g) Manganese is reduced and peroxide is oxidized. Slide 37 of 43 Slide 38 of 43 5-7 Stoichiometry of Reactions in Aqueous Solutions: Titrations. Indicators Titration Carefully controlled addition of one solution to another. Equivalence Point Both reactants have reacted completely. Indicators Substances which change colour near an equivalence point. Slide 39 of 43 Slide 40 of 43 Example 5-10 Standardizing a Solution for Use in Redox Titrations. A piece of iron wire weighing 0.1568 g is converted to Fe 2+ (aq) and requires 26.42 ml of a KMnO 4 (aq) solution for its titration. What is the molarity of the KMnO 4 (aq)? 5 Fe 2+ (aq) + MnO 4- (aq) + 8 H + (aq) 4 H 2 O(l) + 5 Fe 3+ (aq) + Mn 2+ (aq) Example 5-10 5 Fe 2+ (aq) + MnO 4- (aq) + 8 H + (aq) 4 H 2 O(l) + 5 Fe 3+ (aq) + Mn 2+ (aq) Determine KMnO 4 consumed in the reaction: n H 2 O 2+ 1molFe 1molFe = 0.1568g Fe 55.847g Fe 1molFe 4 2+ 1molMnO 5molFe Determine the concentration: 1molKMnO 4 = 5.615 10 molkmno4 1molMnO 4 4 4 5.615 10 mol KMnO4 [ KMnO 4] = 0.02624 L = 0. 02140 M KMnO4 Slide 41 of 43 Slide 42 of 43
Chapter 5 Questions 1, 2, 3, 5, 6, 8, 14, 17, 19, 24, 27, 33, 37, 41, 43, 51, 53, 59, 68, 71, 82, 96. Slide 43 of 43