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) + H 2 O H 3 O + (aq) + Cl 1- (aq) (strong electrolyte) 100% ionized (converted to ions) Weak electrolyte: relatively few molecules or neutral units present form ions in aqueous solution HC 2 H 3 O 2 (l) + H 2 O H 3 O + (aq) + C 2 H 3 O 1-2 (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 HNO 3 nitric acid HBr hydrobromic acid H 2 SO 4 sulfuric acid HI hydroiodic acid HClO 4 perchloric acid See also Table 4.1, p. 131, Hill
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 hydrochloric acid 1 ionizable proton HBr hydrobromic acid 1 ionizable proton HI hydroiodic acid 1 ionizable proton HNO 3 nitric acid 1 ionizable proton H 2 SO 4 sulfuric acid 2 ionizable protons HClO 4 perchloric acid 1 ionizable proton Ionization of Acids in Aqueous Solution 0.10 M 0 0 HCl(aq) + H 2 O H 3 O + (aq) + Cl 1- (aq) means 100% 0.0 M 0.10 M 0.10 M 100% ionization in water = strong acid 0.10 M 0 0 HC 2 H 3 O 2 (aq) + H 2 O H 3 O + (aq) + C 2 H 3 O 2 1- (aq) 0.09 M 0.01 M 0.01 M at equilibrium 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: LiOH, NaOH, KOH, RbOH, CsOH IIA: 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) (+ H 2 O) Na + (aq) + OH 1- (aq) 0.0 M 0.10 M 0.10 M 100% ionization in water = strong base
0.10 M 0 0 NH 3 (aq) + H 2 O(l) NH 4 + (aq) + OH 1- (aq) 0.09 M 0.01 M 0.01 M at equilibrium <<100% ionization in water = weak base Most common weak bases: NH 3 and amines Reaction of Acids with Bases: Neutralization neutralization reaction: the reaction of ionizable H + ions on acid molecules with OH 1- or other anions (such as HCO 3 1- or CO 3 2- ) on base molecules Example. We represent an acid-base reaction as a molecular equation. (no ions involved) NaOH(aq) + HCl(aq) H 2 O(l) + NaCl(aq) But, underlying reaction: H + + OH 1- H 2 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) H 2 O(l) + NaCl(aq) we must break up the aqueous ionizable species into their respective ions: NaOH(aq) ionizes to Na + (aq) + OH 1- (aq) HCl(aq) ionizes in water to H 3 O + (aq) + Cl 1- (aq) Molecular Equation: NaOH(aq) + HCl(aq) H 2 O(l) + NaCl(aq) NaOH(aq): Na + (aq) + OH 1- (aq) HCl(aq): H 3 O + (aq) + Cl 1- (aq) NaCl(aq): ionizes to Na + (aq) + Cl 1- (aq) Corresponding Ionic Equation: Na + (aq) + OH 1- (aq) + H 3 O + (aq) + Cl 1- (aq) H 2 O(l) + Na + (aq) + Cl 1- (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) H 2 O(l) + NaCl(aq) Ionic Equation: Na + (aq) + OH 1- (aq) + H 3 O + (aq) + Cl 1- (aq) H 2 O(l) + Na + (aq) + Cl 1- (aq) Net Ionic Equation: H 3 O + (aq) + OH 1- (aq) H 2 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 CO 3 2- ) CaCO 3 (s) + HCl(aq) H 2 O(l) + CaCl 2 (aq) + CO 2 (g) 2. Sulfites (compounds that contain SO 3 2- ) K 2 SO 3 (s) + H 2 SO 4 (aq) H 2 O(l) + K 2 SO 4 (aq) + SO 2 (g) 3. Sulfides (compounds that contain S 2- ) Na 2 S(aq) + 2 HCl(aq) 2 NaCl(aq) + H 2 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 OH 1- ions per base unit. Use this number for the number of H 2 O molecules formed. Example. (unbalanced) H 3 PO 4 (aq) + Ca(OH) 2 (aq)??? 2 H 3 PO 4 (aq) + 3 Ca(OH) 2 (aq) 6 H 2 O(l) +?? 2 H 3 PO 4 (aq) + 3 Ca(OH) 2 (aq) 6 H 2 O(l) + Ca 2+ (aq) + PO 4 3- (aq) 2 H 3 PO 4 (aq) + 3 Ca(OH) 2 (aq) 6 H 2 O(l) + Ca 3 (PO 4 ) 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: Ba 2+ (aq) + SO 4 2- (aq) BaSO 4 (s) The chemical equation above is the net ionic equation for the reaction between barium chloride and sodium sulfate: BaCl 2 (aq) + Na 2 SO 4 (aq) 2NaCl(aq) + BaSO 4 (s)
Can you get from this equation to the net ionic equation? Reaction of Ag+ with I 1- The reaction: AgNO 3 (aq) + KI(aq) AgI(s) + KNO 3 (aq) Net ionic reaction: Ag+(aq) + I 1- (aq) AgI(s) Silver iodide, AgI precipitates! Solubility Rules for Common Ionic Compounds 1.Group IA ions and NH 4 + are almost always SOLUBLE when paired with NO 3 1-, C 2 H 3 O 2 1- and ClO 4 1-2. Most salts of Cl 1-, Br 1-, and I 1- are SOLUBLE; exceptions are combinations of these anions with Pb 2+, Ag + or Hg 2 2+. 3. Compounds containing SO 4 2- are SOLUBLE except those with Sr 2+, Ba 2+, Pb 2+, and Hg 2 2+ ; CaSO 4 is slightly soluble. 4. Compounds containing CO 3 2-, OH 1-, PO 4 3- and S 2- are INSOLUBLE except Group IA cations, NH 4 + ; combinations of OH 1- and S 2- with Ca 2+, Sr 2+, Ba 2+ 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) MgSO 4 (aq) + KOH(aq)? (b) FeCl 3 (aq) + Na 2 S(aq)? (c) Sr(NO 3 ) 2 (aq) + Na 2 SO 4 (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 = e 1- ) 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. Fe 0 Fe 3+ + 3 e 1- Example of a Reduction Reaction. Cl 2 + 2 e 1-2 Cl 1- 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. Al 2 O 3 P 4 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) + Fe 2 O 3 (s) 2 Fe(l) + Al 2 O 3 (s) Here, Al 0 Al 3+ (oxidation) and Fe 3+ Fe 0 (reduction) We say that Al was oxidized to Al 3+ and that Fe 3+ was reduced to Fe 0 The species oxidized (Al) is the reducing agent, and the species reduced (Fe 3+ ) is the oxidizing agent. Two Other Examples. 4 HCl(aq) + O 2 (g) 2 Cl 2 (aq) + 2 H 2 O(l) Cl 1- in HCl is oxidized to Cl 0 in Cl 2, and O 2 (O 0 ) is reduced to O 2- in H 2 O Ag(s) + H + (aq) + NO 3 1- Ag+(aq) + H 2 O(l) + NO(g) Ag 0 is oxidized to Ag +, and N +5 in NO 3 1- is reduced to N +2 in NO Oxidation-Reduction Reactions (Redox Reactions) See Hill, Figure 4.14, pp. 146: HNO 3 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+ Mg 0 (s) + Ni 2+ (aq) Mg 2+ (aq) + Ni 0 (s) Read Hill, Section 4.5, pp. 148-150 Figure 14.15. Maryland uses the Breathalyzer to determine blood alcohol levels of drivers. Cr 2 O 7 2- + ethanol Cr 3+ 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: V required = V final - V initial 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)