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Chapter 4 (Hill/Petrucci/McCreary/Perry Chemical Reactions in Aqueous Solutions This chapter deals with reactions that occur in aqueous solution …these solutions all use water as the solvent. We will look at some properties of these solutions and also look briefly at three different general types of reactions that occur in aqueous solutions. “water is such a good solvent for so many ionic and molecular substances that it has been called the universal solvent.” (Hill, p.125) Electrical Properties of Aqueous Solutions Nonelectrolytes do not conduct electricity – electrical conductance requires the presence of charged particles. Electrolytes do conduct electricity, in proportion to the concentrations of their ions in solution. Strong electrolyte : almost all molecules or neutral “units” present form ions in aqueous solution HCl(aq) + H2 O " H3 O+(aq) + Cl1-(aq) (strong electrolyte) 100% ionized (converted to ions) Weak electrolyte: relatively few molecules or neutral “units” present form ions in aqueous solution HC2 H3 O2 (l) + H2 O " H3 O+(aq) + C2 H3 O2 1- (aq) (weak electrolyte) <<100% ionized Carefully read pp. 125-128 in Hill! Molarities of Ions in Strong Electrolytes: to calculate the molarity of an ion in a solution of strong electrolyte, simply multiply the subscript for that ion in the compound by the given molarity of the electrolyte See Example 4.1, Hill, p.129 See Exercise 4.1A, Hill, p.129 Reactions of Acids in Aqueous Solution Recall: an acid is a proton donor in aqueous solution; a base is a proton acceptor in aqueous solution. Strong Acids Are Strong Electrolytes - Six strong acids that you should recognize: HCl hydrochloric acid HBr hydrobromic acid HI hydroiodic acid See also Table 4.1, p. 131, Hill HNO3 nitric acid H2 SO4 sulfuric acid HClO 4 perchloric acid Most of the other acids that you will encounter are weak acids that are weak electrolytes. Ionization of Acids in Aqueous Solution By convention, the chemical formulas for acids have their “ionizable” protons (H+ ions) at the “front” of the formula. Recall the strong acids: HCl HBr HI HNO3 H2 SO4 HClO4 hydrochloric acid hydrobromic acid hydroiodic acid nitric acid sulfuric acid perchloric acid 1 ionizable proton 1 ionizable proton 1 ionizable proton 1 ionizable proton 2 ionizable protons 1 ionizable proton Ionization of Acids in Aqueous Solution 0.10 M 0 0 HCl(aq) + H2 O " H3 O+(aq) + Cl1-(aq) 0.0 M 0.10 M 0.10 M " means 100% 100% ionization in water = strong acid 0.10 M 0 0 HC2 H3 O2 (aq) + H2 O D H3 O+(aq) + C2 H3 O21-(aq) 0.09 M 0.01 M 0.01 M at equilibrium D means <<100% ionization in water = weak acid Reactions of Bases in Aqueous Solution Recall: a base is a proton acceptor in aqueous solution. Strong Bases Are Strong Electrolytes Most strong bases are Group IA and Group IIA hydroxides: IA: IIA: LiOH, NaOH, KOH, RbOH, CsOH Mg(OH)2 , Ca(OH)2 , Sr(OH)2 , Ba(OH)2 , See also Table 4.1, p. 131, Hill Ionization of Bases in Aqueous Solution 0.10 M 0 0 + NaOH(aq) (+ H2 O) " Na (aq) + OH1-(aq) 0.0 M 0.10 M 0.10 M 100% ionization in water = strong base 0.10 M 0 0 + NH3 (aq) + H2 O(l) D NH4 (aq) + OH1-(aq) 0.09 M 0.01 M 0.01 M at equilibrium <<100% ionization in water = weak base Most common weak bases: NH3 and amines Reaction of Acids with Bases: Neutralization neutralization reaction: the reaction of ionizable H+ ions on acid molecules with OH1- or other anions (such as HCO3 1- or CO3 2-) on base “molecules” Example. We represent an acid-base reaction as a “molecular equation.” (no ions involved) NaOH(aq) + HCl(aq) " H2 O(l) + NaCl(aq) But, underlying reaction: H+ + OH1- " H2 O(l) Classically, acid-base reactions produce a salt and water. Molecular Equations to Ionic Equations ionic equation: all ionizable species written as ions, i.e. in their ionized or dissociated forms Example. If the molecular equation is NaOH(aq) + HCl(aq) " H2 O(l) + NaCl(aq) we must “break” up the aqueous ionizable species into their respective ions: NaOH(aq) ionizes to Na+(aq) + OH1-(aq) HCl(aq) ionizes in water to H3 O+(aq) + Cl1-(aq) Molecular Equation: NaOH(aq) + HCl(aq) " H2 O(l) + NaCl(aq) NaOH(aq): Na+(aq) + OH1-(aq) HCl(aq): H3 O+(aq) + Cl1-(aq) NaCl(aq): ionizes to Na+(aq) + Cl1-(aq) Corresponding Ionic Equation: Na+(aq) + OH1-(aq) + H3 O+(aq) + Cl1-(aq) " H2 O(l) + Na+(aq) + Cl1-(aq) ] Species in bold that appear on both sides are called spectator ions and “cancel” out. These species do not participate in the chemical reaction. Molecular Equations to Ionic Equations Molecular Equation: NaOH(aq) + HCl(aq) " H2 O(l) + NaCl(aq) Ionic Equation: Na+(aq) + OH1-(aq) + H3 O+(aq) + Cl1-(aq) " H2 O(l) + Na+(aq) + Cl1-(aq) Net Ionic Equation: H3 O+(aq) + OH1-(aq) " H2 O(l) See Example 4.2 and following Exercises 4.2A and 4.2B on Hill, p. 133 Acid-Base Reactions That Form Gases 1. Carbonates (compounds that contain CO3 2-) CaCO3 (s) + HCl(aq) " H2 O(l) + CaCl2 (aq) + CO2 (g)# 2. Sulfites (compounds that contain SO3 2-) K2 SO3 (s) + H2 SO4 (aq) " H2 O(l) + K2 SO4 (aq) + SO2 (g)# 3. Sulfides (compounds that contain S2-) Na2 S(aq) + 2 HCl(aq) " 2 NaCl(aq) + H2 S(g)# Acid-Base Reactions: Another Example When balancing acid-base equations that have hydroxyl bases, use the lowest common denominator for the number of ionizable protons and the number of OH1- ions per base unit. Use this number for the number of H2 O molecules formed. Example. (unbalanced) H3 PO4 (aq) + Ca(OH)2 (aq) " ??? 2 H3 PO4 (aq) + 3 Ca(OH)2 (aq) " 6 H2 O(l) + ?? 2 H3 PO4 (aq) + 3 Ca(OH)2 (aq) " 6 H2 O(l) + Ca2+(aq) + PO4 3-(aq) 2 H3 PO4 (aq) + 3 Ca(OH)2 (aq) " 6 H2 O(l) + Ca3 (PO4 )2 (aq) Reactions That Form Precipitates precipitate: a solid product formed from the reaction of two soluble ions (a cation and an anion); a precipitate is, by definition, insoluble (not soluble) in the solvent used Example of a Precipitation Reaction: Ba2+(aq) + SO4 2-(aq) " BaSO4 (s) ] The chemical equation above is the net ionic equation for the reaction between barium chloride and sodium sulfate: BaCl2 (aq) + Na2 SO4 (aq) " 2NaCl(aq) + BaSO4 (s) Can you get from this equation to the net ionic equation? Reaction of Ag+ with I1The reaction: AgNO3 (aq) + KI(aq) " " AgI(s) + KNO3 (aq) Net ionic reaction: Ag+(aq) + I1-(aq) " AgI(s) ]Silver iodide, AgI precipitates! Solubility Rules for Common Ionic Compounds 1.Group IA ions and NH4 + are almost always SOLUBLE when paired with NO3 1-, C2 H3 O21- and ClO 4 12. Most salts of Cl1-, Br1-, and I1- are SOLUBLE; exceptions are combinations of these anions with Pb2+, Ag+ or Hg2 2+. 3. Compounds containing SO4 2- are SOLUBLE except those with Sr2+, Ba2+, Pb2+, and Hg2 2+; CaSO4 is slightly soluble. 4. Compounds containing CO3 2-, OH1-, PO4 3- and S2- are INSOLUBLE except Group IA cations, NH4 +; combinations of OH1- and S2- with Ca2+, Sr2+, Ba2+ are slightly to moderately soluble. Memorize these solubility rules! See also Table 4.3 on p. 136 (Hill) See Example 4.4 p. 137 Exercise 4.4A, p. 137 (a) MgSO4 (aq) + KOH(aq) " ? (b) FeCl3 (aq) + Na2 S(aq) " ? (c) Sr(NO3 )2 (aq) + Na2 SO4 (aq) " ? Oxidation-Reduction Reactions (Redox Reactions) In addition to acid-base and precipitation reactions, there is a third type of reaction: oxidationreduction or redox reactions. Oxidation-reduction reactions are electron exchange reactions (electron = e1-) Oxidation - loss of 1 or more electrons by an ion or molecule Reduction - gain of 1 or more electrons by an ion/molecule Example of an Oxidation Reaction. Example of a Reduction Reaction. Fe0 " Fe3+ + 3 e1Cl2 + 2 e1- " 2 Cl1- Oxidation Numbers and the Oxidation Number Concept oxidation number: the charge on a monatomic ion or the nominal charge on an atom in a unit of a compound (oxidation number is also referred to as the oxidation state of an atom) Short List of Oxidation Number Rules. 1. The oxidation number of a neutral, uncharged atom is 0 2. Ions: IA metals = +1; IIA metals = +2 3. Hydrogen : H is usually +1; sometimes -1 in hydrides 4. Oxygen: O is usually -2; sometimes -1 in peroxides 5. The sum of all the oxidation numbers in a molecule or an ion is equal to the charge on the molecule (0) or ion. Examples. See Example 4.7, p. 141 (Hill) Exercise 4.7 A, p. 142 Assign known oxidation numbers and then set sum of the oxidation numbers equal to the charge and solve algebraically. Al2 O3 P4 HAsO 4 3- NaMnO 4 Oxidation Numbers in Nitrogen, Sulfur and Chlorine Species (see text Figure 4.12) Identifying Oxidation and Reduction Reactions To classify a reaction as an oxidation process or as a reduction process, first assign oxidation numbers to all atoms on both sides of the equation. 1. Oxidation. If the oxidation number for an element increases (becomes more positive) from reactant to product, the process is an oxidation process 2. Reduction. If the oxidation number for an element decreases (becomes more negative) from reactant to product, the process is a reduction process Identifying Oxidation and Reduction Reactions Example. The “thermite” reaction: 2 Al(s) + Fe2 O3 (s) " 2 Fe(l) + Al2 O3 (s) Here, Al0 " Al3+ (oxidation) and Fe3+ " Fe0 (reduction) We say that Al was oxidized to Al3+ and that Fe3+ was reduced to Fe0 The species oxidized (Al) is the reducing agent, and the species reduced (Fe3+) is the oxidizing agent. Two Other Examples. 4 HCl(aq) + O2 (g) " 2 Cl2 (aq) + 2 H2 O(l) Cl1- in HCl is oxidized to Cl0 in Cl2 , and O2 (O 0 ) is reduced to O2- in H2 O Ag(s) + H+(aq) + NO3 1- " Ag+(aq) + H2 O(l) + NO(g) Ag0 is oxidized to Ag+, and N+5 in NO3 1- is reduced to N+2 in NO Oxidation-Reduction Reactions (Redox Reactions) See Hill, Figure 4.14, pp. 146: HNO3 oxidizes Cu to Cu 2+, but HCl doesn’t … why? Activity Series of the Metals A metal will “displace from solution the ions of any metal that lie below it in the activity series. Example. metal = Mg and ion = Ni2+ Mg0 (s) + Ni2+(aq) " Mg2+(aq) + Ni0 (s) Read Hill, Section 4.5, pp. 148-150 Figure 14.15. Maryland uses the Breathalyzer to determine blood alcohol levels of drivers. Cr2 O7 2- + ethanol " Cr3+ Titrations We are interested in quantitatively determining the concentration of a chemical species (called the analyte) in a sample. The sample is placed in a flask or beaker, and a solution containing a known concentration of a chemical reagent (the titrant ) that will react with the analyte is added until no more analyte remains (the titration endpoint). The chemical reaction between the analyte and the titrant is known. We also know the concentration of the titrant solution and the volume of the solution required to just react with all of the analyte. A titration is carried out using a tube ( a buret) calibrated along its length, typically in 0.1 mL increments. The volume before titrant is measured and recorded; the volume after reaching the endpoint is then measured and recorded. The volume of titrant used is then: Vrequired = Vfinal - Vinitial Acid-Base Titrations See Examples and Exercises on pp. 153-154 (Hill) Precipitation Titrations See Examples and Exercises on pp. 155-156 (Hill) Redox Titrations See Examples and Exercises on pp. 156-157 (Hill)