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5.1 The Nature of Aqueous Solutions Water Inexpensive Can dissolve a vast number of substances Many substances dissociate into ions Aqueous solutions are found everywhere ◦ Seawater ◦ Living systems Slide 1 of 47 General Chemistry: Chapter 5 Prentice-Hall © 2007 Electrolytes Some solutes can dissociate into ions. Electric charge can be carried. Slide 2 of 47 General Chemistry: Chapter 5 Prentice-Hall © 2007 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. Slide 3 of 47 General Chemistry: Chapter 5 Prentice-Hall © 2007 Representation of Electrolytes using Chemical Equations A strong electrolyte: MgCl2(s) → Mg2+(aq) + 2 Cl-(aq) A weak electrolyte: → CH3CO2-(aq) + H+(aq) CH3CO2H(aq) ← A non-electrolyte: CH3OH(aq) Slide 4 of 47 General Chemistry: Chapter 5 Prentice-Hall © 2007 Three Types of Electrolytes Slide 5 of 47 General Chemistry: Chapter 5 Prentice-Hall © 2007 Solvation: The Hydrated Proton Slide 6 of 47 General Chemistry: Chapter 5 Prentice-Hall © 2007 Relative Concentrations in Solution MgCl2(s) → Mg2+(aq) + 2 Cl-(aq) In 0.0050 M MgCl2: Stoichiometry is important. [Mg2+] = 0.0050 M Slide 7 of 47 [Cl-] = 0.0100 M General Chemistry: Chapter 5 [MgCl2] = 0 M Prentice-Hall © 2007 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 Al2(SO4)3?. Write a Balanced Chemical Equation: Al2(SO4)3 (s) → 2 Al3+(aq) + 3 SO42-(aq) Identify the Stoichiometric Factors : 2 mol Al3+ 1 mol Al2(SO4)3 Slide 8 of 47 3 mol SO421 mol Al2(SO4)3 General Chemistry: Chapter 5 Prentice-Hall © 2007 EXAMPLE 7-3 Aluminum Concentration: [Al] = 0.0165 mol Al2(SO4)3 1L 2 mol Al3+ 1 mol Al2(SO4)3 mol Al3+ = 0.0330 1L Sulfate Concentration: 0.0165 mol Al2(SO4)3 3 mol SO42[SO42-] = = 1L 1 mol Al2(SO4)3 Slide 9 of 47 General Chemistry: Chapter 5 0.0495 M SO42- Prentice-Hall © 2007 5-2 Precipitation Reactions Soluble ions can combine to form an insoluble compound. Precipitation occurs. A test for the presence of chloride ion in water. Ag+(aq) + Cl-(aq) → AgCl(s) Slide 10 of 47 General Chemistry: Chapter 5 Prentice-Hall © 2007 Silver Nitrate and Sodium Iodide AgNO3(aq) NaI(aq) AgI(s) Na+(aq) NO3-(aq) Slide 11 of 47 General Chemistry: Chapter 5 Prentice-Hall © 2007 Net Ionic Equation Overall Precipitation Reaction: AgNO3(aq) +NaI(aq) → AgI(s) + NaNO3(aq) Complete ionic equation: Spectator ions Ag+(aq) + NO3-(aq) + Na+(aq) + I-(aq) → AgI(s) + Na+(aq) + NO3-(aq) Net ionic equation: Ag+(aq) + I-(aq) → AgI(s) Slide 12 of 47 General Chemistry: Chapter 5 Prentice-Hall © 2007 Solubility Rules Compounds that are soluble: Alkali metal ion and ammonium ion salts Li+, Na+, K+, Rb+, Cs+ NH4+ Nitrates, perchlorates and acetates NO3- Slide 13 of 47 ClO4- CH3CO2- General Chemistry: Chapter 5 Prentice-Hall © 2007 Solubility Rules Compounds that are mostly soluble: Chlorides, bromides and iodides Cl- Br- I- ◦ Except those of Pb2+, Ag+, and Hg22+. Sulfates SO42- ◦ Except those of Sr2+, Ba2+, Pb2+ and Hg22+. ◦ Ca(SO4) is slightly soluble. Slide 14 of 47 General Chemistry: Chapter 5 Prentice-Hall © 2007 5-3 Acid-Base Reactions Latin acidus (sour) Sour taste Arabic al-qali (ashes of certain plants) Bitter taste Svante Arrhenius 1884 Acid-Base theory. Slide 15 of 47 General Chemistry: Chapter 5 Prentice-Hall © 2007 Acids and Bases Some Definitions Arrhenius An acid is a substance that, when dissolved in water, ionizes and increases the concentration of hydrogen ions, H+. HCl → H+ + Cl A base is a substance that, when dissolved in water, increases the concentration of hydroxide ions, OH-. But, Not all bases contain OHNaOH → Na+ + OH- Brønsted-Lowry An acid is a proton donor. A base is a proton acceptor. e.g. ammonia, NH3, is a base. A Brønsted-Lowry acid… …must have an ionizable or removable (acidic) proton. A Brønsted-Lowry base… …must have a pair of nonbonding electrons. In an acid – base reaction, the acid donates a proton (H+) to the base. If a substance can be either… …it is amphiprotic. H2O HCO3 What Happens When an Acid Dissolves in Water? Water acts as a BrønstedLowry base and abstracts a proton (H+) from the acid. As a result, the conjugate base of the acid and a hydronium ion are formed. Self-ionization of water H2O + H2O ⇄ H3O+ (aq) + OH- (aq) Autoionization of Water As we have seen, water is amphoteric. In pure water, a few molecules act as bases and a few act as acids. H2O (l) + H2O (l) H3O+ (aq) + OH- (aq) This is referred to as autoionization. For ALL aqueous solutions: Kc = [H3O+] [OH-] At 25C, Kw = [H3O+]·[OH-] = 1.0·10-14 = Kw (at 25 ºC) [H3O+] = [OH-] the solution is…. [H3O+] > [OH-] the solution is…. [H3O+] < [OH-] the solution is…. Acids Acids provide H+ in aqueous solution. Strong acids completely ionize: HCl(aq) → H+(aq) + Cl-(aq) Weak acid ionization is not complete: CH3CO2H(aq) Slide 26 of 47 ← → H+(aq) + CH3CO2-(aq) General Chemistry: Chapter 5 Prentice-Hall © 2007 Acids Arrhenius defined acids as substances that increase the concentration of H+ when dissolved in water. Brønsted and Lowry defined them as proton donors. Bases Bases provide OH- in aqueous solution. Strong bases: NaOH(aq) → Na+(aq) + OH-(aq) H2O Weak bases: NH3(aq) + H2O(l) Slide 28 of 47 ← → OH-(aq) + NH4+(aq) General Chemistry: Chapter 5 Prentice-Hall © 2007 Acids There are only seven strong acids: • • • • • • • Hydrochloric (HCl) Hydrobromic (HBr) Hydroiodic (HI) Nitric (HNO3) Sulfuric (H2SO4) Chloric (HClO3) Perchloric (HClO4) Acid-Base Reactions In an acid-base reaction, the acid donates a proton (H+) to the base. Recognizing Acids and Bases. Acids have ionizable hydrogen ions. CH3CO2H or HC2H3O2 Bases have OH- combined with a metal ion. KOH or can be identified by chemical equations Na2CO3(s) + H2O(l)→ HCO3-(aq) + 2 Na+(aq) + OH-(aq) Slide 31 of 47 General Chemistry: Chapter 5 Prentice-Hall © 2007 More Acid-Base Reactions Milk of magnesia Mg(OH)2 Mg(OH)2(s) + 2 H+(aq) → Mg2+(aq) + 2 H2O(l) Mg(OH)2(s) + 2 CH3CO2H(aq) → Mg2+(aq) + 2 CH3CO2-(aq) + 2 H2O(l) Slide 32 of 47 General Chemistry: Chapter 5 Prentice-Hall © 2007 More Acid-Base Reactions Limestone and marble. CaCO3(s) + 2 H+(aq) → Ca2+(aq) + H2CO3(aq) But: H2CO3(aq) → H2O(l) + CO2(g) CaCO3(s) + 2 H+(aq) → Ca2+(aq) + H2O(l) + CO2(g) Slide 33 of 47 General Chemistry: Chapter 5 Prentice-Hall © 2007 Electrochemical Reactions In electrochemical reactions, electrons are transferred from one species to another. Mg (s) + HCl (aq) → Oxidation Numbers In order to keep track of what loses electrons and what gains them, we assign oxidation numbers. Oxidation and Reduction A species is oxidized when it loses electrons. Here, zinc loses two electrons to go from neutral zinc metal to the Zn2+ ion. Oxidation and Reduction A species is reduced when it gains electrons. Here, each of the H+ gains an electron, and they combine to form H2. Oxidation and Reduction What is reduced is the oxidizing agent. H+ oxidizes Zn by taking electrons from it. What is oxidized is the reducing agent. Zn reduces H+ by giving it electrons. Assigning Oxidation Numbers, ON 1. Elements in their elemental form have an oxidation number of 0. Mg Cl2 Pb 1. The oxidation number of a monatomic ion is the same as its charge. Zn2+ F- Na+ In combination with other elements: 3. Nonmetals tend to have negative oxidation numbers Oxygen has an oxidation number of −2, except in the peroxide ion, which has an oxidation number of −1. MgO Na2O HF Hydrogen is −1 when bonded to a metal and +1 when bonded to a nonmetal. NaH Assigning Oxidation Numbers Fluorine always has an oxidation number of −1. HF NaF The other halogens have an oxidation number of −1 when they are negative; they can have positive oxidation numbers, however, most notably in oxyanions. ZnBr2 CuCl2 Assigning Oxidation Numbers 4. The sum of the oxidation numbers in a neutral compound is 0. HCl H2O K2O 5. The sum of the oxidation numbers in a polyatomic ion is the charge on the ion. CO32- ClO4- Examples Determine the oxidation number of sulfur in the following compounds: H2S SO42S8 S2O32- Electrochemical Reactions In a spontaneous Oxidation-Reduction (redox) Reaction electrons are transferred and energy is released as the reaction proceeds: Zn strip dissolves… The blue color due to Cu2+ fades. Copper metal is deposited. Oxidation and Reduction Half-Reactions A reaction represented by two half-reactions. Oxidation: Zn(s) → Zn2+(aq) + 2 e- Reduction: Cu2+(aq) + 2 e- → Cu(s) Overall: Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq) Slide 46 of 47 General Chemistry: Chapter 5 Prentice-Hall © 2007 Corrosion and… 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. Slide 49 of 47 General Chemistry: Chapter 5 Prentice-Hall © 2007 Oxidation State Changes Assign oxidation states: D Fe2O3(s) + 3 CO(g) → 2 Fe(l) + 3 CO2(g) Fe3+ is CO(g) is Slide 50 of 47 to metallic iron. to carbon dioxide. General Chemistry: Chapter 5 Prentice-Hall © 2007 Balancing Oxidation-Reduction Equations Few can be balanced by inspection. Systematic approach required. The Half-Reaction (Ion-Electron) Method Slide 51 of 47 General Chemistry: Chapter 5 Prentice-Hall © 2007 Balancing in Acid Write the equations for the half-reactions. Balance all atoms except H and O. Balance oxygen using H2O. Balance hydrogen using H+. Balance charge using e-. Equalize the number of electrons. Add the half reactions. Check the balance. Slide 52 of 47 General Chemistry: Chapter 5 Prentice-Hall © 2007 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.. SO32-(aq) + MnO4-(aq) → SO42-(aq) + Mn2+(aq) Slide 53 of 47 General Chemistry: Chapter 5 Prentice-Hall © 2007 EXAMPLE 5-6 Determine the oxidation states: 4+ 6+ 7+ 2+ SO32-(aq) + MnO4-(aq) → SO42-(aq) + Mn2+(aq) Write the half-reactions: SO32-(aq) → SO42-(aq) + 2 e-(aq) 5 e-(aq) +MnO4-(aq) → Mn2+(aq) Balance atoms other than H and O: Already balanced for elements. Slide 54 of 47 General Chemistry: Chapter 5 Prentice-Hall © 2007 EXAMPLE 5-6 Balance O by adding H2O: H2O(l) + SO32-(aq) → SO42-(aq) + 2 e-(aq) 5 e-(aq) +MnO4-(aq) → Mn2+(aq) + 4 H2O(l) Balance hydrogen by adding H+: H2O(l) + SO32-(aq) → SO42-(aq) + 2 e-(aq) + 2 H+(aq) 8 H+(aq) + 5 e-(aq) +MnO4-(aq) → Mn2+(aq) + 4 H2O(l) Check that the charges are balanced: Slide 55 of 47 General Chemistry: Chapter 5 Add e- if necessary. Prentice-Hall © 2007 EXAMPLE 5-6 Multiply the half-reactions to balance all e-: 5 H2O(l) + 5 SO32-(aq) → 5 SO42-(aq) + 10 e-(aq) + 10 H+(aq) 16 H+(aq) + 10 e-(aq) + 2 MnO4-(aq) → 2 Mn2+(aq) + 8 H2O(l) 6 3 Add both equations and simplify: 5 SO32-(aq) + 2 MnO4-(aq) + 6H+(aq) → 5 SO42-(aq) + 2 Mn2+(aq) + 3 H2O(l) Check the balance! Slide 56 of 47 General Chemistry: Chapter 5 Prentice-Hall © 2007 EXAMPLE 7-3 3+ 2+ H2O2(aq) + 2 Fe2+(aq) + 2 H+ → 2 H2O(l) + 2 Fe3+(aq) Iron is oxidized and peroxide is reduced. 7+ 5 H2O2(aq) + 2 MnO4-(aq) + 6 H+ → 2+ 8 H2O(l) + 2 Mn2+(aq) + 5 O2(g) Manganese is reduced and peroxide is oxidized. Slide 57 of 47 General Chemistry: Chapter 5 Prentice-Hall © 2007