Prentice-Hall © 2002General Chemistry: Chapter 5Slide 1 of 43 Chapter 5: Introduction to Reactions in Aqueous Solutions Philip Dutton University of Windsor, Canada Prentice-Hall © 2002 General Chemistry Principles and Modern Applications Petrucci Harwood Herring 8 th Edition

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 2 of 43 Contents 5-1The Nature of Aqueous Solutions 5-2Precipitation Reactions 5-3Acid-Base Reactions 5-4Oxidation-Reduction: Some General Principles 5-5Balancing Oxidation-Reduction Equations 5-6Oxidizing and Reducing Agents 5-7Stoichiometry of Reactions in Aqueous Solutions: Titrations Focus on Water Treatment

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 3 of The Nature of Aqueous Solutions

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 4 of 43 Electrolytes Some solutes can dissociate into ions. Electric charge can be carried.

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 5 of 43 Types of Electrolytes Weak electrolyte partially dissociates. –Fair conductor of electricity. Non-electrolyte does not dissociate. –Poor conductor of electricity. Strong electrolyte dissociates completely. –Good electrical conduction.

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 6 of 43 Representation of Electrolytes using Chemical Equations MgCl 2 (s) → Mg 2+ (aq) + 2 Cl - (aq) A strong electrolyte: A weak electrolyte: CH 3 CO 2 H(aq) ← CH 3 CO 2 - (aq) + H + (aq) → CH 3 OH(aq) A non-electrolyte:

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 7 of 43 MgCl 2 (s) → Mg 2+ (aq) + 2 Cl - (aq) [Mg 2+ ] = M [Cl - ] = M [MgCl 2 ] = 0 M Notation for Concentration In M MgCl 2 : Stoichiometry is important.

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 8 of 43 Example 5-1 Calculating Ion concentrations in a Solution of a Strong Electolyte. What are the aluminum and sulfate ion concentrations in M Al 2 (SO 4 ) 3 ?. Al 2 (SO 4 ) 3 (s) → 2 Al 3+ (aq) + 3 SO 4 2- (aq) Balanced Chemical Equation:

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 9 of 43 [Al] = × = 1 L 2 mol Al 3+ 1 mol Al 2 (SO 4 ) mol Al 2 (SO 4 ) M Al 3+ Example M SO 4 2- [SO 4 2- ] = × = 1 mol Al 2 (SO 4 ) 3 Sulfate Concentration: 1 L 3 mol SO mol Al 2 (SO 4 ) 3 Aluminum Concentration:

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 10 of Precipitation Reactions Soluble ions can combine to form an insoluble compound. Precipitation occurs. Ag + (aq) + Cl - (aq) → AgCl(s)

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 11 of 43 Ag + (aq) + NO 3 - (aq) + Na + (aq) + I - (aq) → AgI(s) + Na + (aq) + NO 3 - (aq) Spectator ions Ag + (aq) + NO 3 - (aq) + Na + (aq) + I - (aq) → AgI(s) + Na + (aq) + NO 3 - (aq) Net Ionic Equation AgNO 3 (aq) +NaI (aq) → AgI(s) + NaNO 3 (aq) Overall Precipitation Reaction: Complete ionic equation: Ag + (aq) + I - (aq) → AgI(s) Net ionic equation:

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 12 of 43 Solubility Rules Compounds that are soluble: Li +, Na +, K +, Rb +, Cs + NH 4 + NO 3 - ClO 4 - CH 3 CO 2 - –Alkali metal ion and ammonium ion salts –Nitrates, perchlorates and acetates

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 13 of 43 Solubility Rules –Chlorides, bromides and iodides Cl -, Br -, I - Except those of Pb 2+, Ag +, and Hg –Sulfates SO 4 2- Except those of Sr 2+, Ba 2+, Pb 2+ and Hg Ca(SO 4 ) is slightly soluble. Compounds that are mostly soluble:

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 14 of 43 Solubility Rules –Hydroxides and sulfides HO -, S 2- 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 2-, PO 4 3- Except alkali metal and ammonium salts Compounds that are insoluble:

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 15 of Acid-Base Reactions Latin acidus (sour) –Sour taste Arabic al-qali (ashes of certain plants) –Bitter taste Svante Arrhenius 1884 Acid-Base theory.

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 16 of 43 Acids Acids provide H + in aqueous solution. Strong acids: Weak acids: HCl(aq) H + (aq) + Cl - (aq) → → ← CH 3 CO 2 H(aq) H + (aq) + CH 3 CO 2 - (aq)

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 17 of 43 Bases Bases provide OH - in aqueous solution. Strong bases: Weak bases: → ← NH 3 (aq) + H 2 O(l) OH - (aq) + NH 4 + (aq) NaOH(aq) Na + (aq) + OH - (aq) → H2OH2O

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 18 of 43 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 or are identified by chemical equations Na 2 CO 3 (s) + H 2 O(l)→ HCO 3 - (aq) + 2 Na + (aq) + OH - (aq)

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 19 of 43 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) Mg(OH) 2 (s) + 2 CH 3 CO 2 H(aq) → Mg 2+ (aq) + 2 CH 3 CO 2 - (aq) + 2 H 2 O(l)

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 20 of 43 More Acid-Base Reactions Limestone and marble. CaCO 3 (s) + 2 H + (aq) → Ca 2+ (aq) + H 2 CO 3 (aq) 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) CaCO 3 (s) + 2 H + (aq) → Ca 2+ (aq) + H 2 CO 3 (aq) CaCO 3 (s) + 2 H + (aq) → Ca 2+ (aq) + H 2 O(l) + CO 2 (g)

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 21 of 43 Limestone and Marble

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 22 of 43 Fe 2 O 3 (s) + 3 CO(g) → 2 Fe(l) + 3 CO 2 (g)  Hematite is converted to iron in a blast furnace. Fe 2 O 3 (s) + 3 CO(g) → 2 Fe(l) + 3 CO 2 (g)   CO(g) is oxidized to carbon dioxide. Fe 3+ is reduced to metallic iron. 5-4 Oxidation-Reduction: Some General Principles Oxidation and reduction always occur together.

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 23 of 43 Oxidation State Changes Fe 2 O 3 (s) + 3 CO(g) → 2 Fe(l) + 3 CO 2 (g)  Assign oxidation states: CO(g) is oxidized to carbon dioxide. Fe 3+ is reduced to metallic iron.

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 24 of 43 Oxidation and Reduction 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.

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 25 of 43 Zinc in Copper Sulfate Zn(s) + Cu 2+ (aq) → Zn 2+ (aq) + Cu(s)

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 26 of 43 Half-Reactions Represent a reaction by two half-reactions. Oxidation: Reduction: Overall: Zn(s) → Zn 2+ (aq) + 2 e - Cu 2+ (aq) + 2 e- → Cu(s) Cu 2+ (aq) + Zn(s) → Cu(s) + Zn 2+ (aq)

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 27 of 43 Balancing Oxidation-Reduction Equations Few can be balanced by inspection. Systematic approach required. The Half-Reaction (Ion-Electron) Method

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 28 of 43 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 2- (aq) + MnO 4 - (aq) → SO 4 2- (aq) + Mn 2+ (aq)

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 29 of 43 Example 5-6 SO 3 2- (aq) + MnO 4 - (aq) → SO 4 2- (aq) + Mn 2+ (aq) Determine the oxidation states: SO 3 2- (aq) → SO 4 2- (aq) + 2 e - (aq) Write the half-reactions: 5 e - (aq) +MnO 4 - (aq) → Mn 2+ (aq) Balance atoms other than H and O: Already balanced for elements.

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 30 of 43 Example 5-6 Balance O by adding H 2 O: H 2 O(l) + SO 3 2- (aq) → SO 4 2- (aq) + 2 e - (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 2- (aq) → SO 4 2- (aq) + 2 e - (aq) + 2 H + (aq) 8 H + (aq) + 5 e - (aq) +MnO 4 - (aq) → Mn 2+ (aq) + 4 H 2 O(l) Check that the charges are balanced: Add e - if necessary.

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 31 of 43 Example 5-6 Multiply the half-reactions to balance all e - : 5 H 2 O(l) + 5 SO 3 2- (aq) → 5 SO 4 2- (aq) + 10 e - (aq) + 10 H + (aq) 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 2- (aq) + 2 MnO 4 - (aq) + 6H + (aq) → 5 SO 4 2- (aq) + 2 Mn 2+ (aq) + 3 H 2 O(l) Check the balance!

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 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.

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 33 of 43 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.

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 34 of Oxidizing and Reducing Agents. 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.

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 35 of 43 Redox

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 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.

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 37 of 43 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) 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. Manganese is reduced and peroxide is oxidized.

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 38 of Stoichiometry of Reactions in Aqueous Solutions: Titrations. 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.

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 39 of 43 Example 5-10 Standardizing a Solution for Use in Redox Titrations. A piece of iron wire weighing g is converted to Fe 2+ (aq) and requires 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)

Prentice-Hall © 2002General Chemistry: Chapter 5Slide 40 of 43 Chapter 5 Questions 1, 2, 3, 5, 6, 8, 14, 17, 19, 24, 27, 33, 37, 41, 43, 51, 53, 59, 68,