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Chapter 5

Chapter 5. Reactions in Aqueous Solutions. Chapter goals. Understand the nature of ionic substances dissolved in water. Recognize common acids and bases and understand their behavior in aqueous solution. Recognize and write equations for the common types of reactions in aqueous solution.

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Chapter 5

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  1. Chapter 5 Reactions in Aqueous Solutions

  2. Chapter goals • Understand the nature of ionic substances dissolved in water. • Recognize common acids and bases and understand their behavior in aqueous solution. • Recognize and write equations for the common types of reactions in aqueous solution. • Recognize common oxidizing and reducing agents and identify common oxidation-reduction reactions. • Define and use the molarity in solution stoichiometry.

  3. Solution • homogeneous mixture • can be gas, liquid, or solid • solvent: component present in highest proportion • exception - water • solute: component(s) in solution other than solvent • We will mostly study aqueous solutions: human body is 2/3 water.

  4. Examples • mixture of 35% naphthalene and 65% benzene • solvent - benzene • solute – naphthalene • mixture of 10% ethanol, 40% methanol, and 50% propanol • solvent - propanol • solute - ethanol and methanol

  5. Examples • mixture of 40% ethanol, 40% methanol, and 20% butanol • solvent - ethanol/methanol • mixed solvent • solute – butanol • mixture of 40% ethanol, 50% propanol, and 10% water • solvent - water • solute - ethanol and propanol

  6. Next… We will focus on compounds that produce ions in aqueous solutions. They are named electrolytes and may be salts, acids, or bases.

  7. Salts Salts: ionic compounds made of cations other than H+ and anions other than OH− or O2−, O22− • NaCl: Na+ & Cl− • K2SO4: K+ & SO42− • FeBr3: Fe3+ & Br− • Zn3(PO4)2: Zn2+ & PO43− • Ca(HCO3)2: Ca2+ & HCO3−

  8. Electrolyte • substance that dissolves to produce an electrically conducting medium • forms ions in solution (dissociates/ionizes) • examples • soluble ionic compounds • H2O • KBr(s)  K+(aq)+ Br–(aq) • H2O • Acids, HCl(g)  H+(aq) + Cl–(aq) • bases, NH3 + H2O NH4+ + OH– • double arrow means equilibrium

  9. Nonelectrolytes • do not form ions in solution • do not form electrically conducting media upon dissolution • Examples: molecular compounds (alcohols, sugars & acetone) • H2O • CH3OH(l) CH3OH(aq) N.D. • Glucose C6H12O6(s)  C6H12O6(aq) N.D. • Sucrose C12H22O11(s)  C12H22O11(aq) N.D. • N.D. = no dissociation/ionization

  10. Types of Electrolytes • Strong: dissociate ~100% • most ionic compounds (soluble salts), strong acids, and strong bases H2O KBr(s)  K+(aq)+ Br–(aq) HCl(g)  H+(aq)+ Cl–(aq) NaOH(s)  Na+(aq) + OH−(aq) Weak: insoluble salts, weak acids and bases, water (H2O), and certain gases (e.g. CO2) H2O dissociate only slightly in water HF(g) H+(aq)+ F–(aq) Also acetic acid, CH3COOH NH3 + H2O NH4+(aq) + OH–(aq)

  11. Solubility of Ionic compounds in Water: Solubility Rules Soluble Compounds • 1. alkali metal salts (Li+, Na+, K+, Rb+…, ) except potassium perchlorate • 2. ammonium (NH4+) salts • 3. all nitrates(NO3−), chlorates (ClO3−), perchlorates (ClO4−), and acetates (C2H3O2−), except silver acetate and potassium perchlorate • 4. all Cl−, Br−, and l− are soluble except for Ag+, Pb2+, and Hg22+ salts • 5. all SO42− are soluble except for Pb2+, Sr2+, and Ba2+ salts

  12. Solubility of Ionic compounds in Water: Rules Insoluble or slightly soluble Compounds • 6. metal oxides (O2−) except those of the alkali metals, Ca2+, Sr2+, and Ba2+ • 7. hydroxides (OH−) except those of the alkali metals, Ba2+, Sr2+, and NH4+. Calcium hydroxide is slightly soluble • 8. carbonates, phosphates, sulfides, and sulfites except those of the alkali metals and the ammonium ion (NH4+) • 9. for salts of Cr2O72−, P3−, CrO42−, C2O42−, • assume they are insoluble except for IA metals and NH4+ salts

  13. Precipitation Reactions:A Driving Force in Chemical Reactions • formation of insoluble solid (precipitate, ppt) is a common reaction in aqueous solutions: • reactants are generally water-soluble ionic compounds • once substances dissolve in water they dissociate to give the appropriate cations and anions • if the cation of one compound forms an insoluble compound with the anion of another, precipitation will occur

  14. Precipitation Reaction:A Double Replacement (Metathesis) Reaction • Both ionic compounds trade partner ions __________ | | AB(aq) + CD(aq)  AD(s) + CB(aq) |_______| AD is an insoluble or slightly soluble salt A+, B−, C+, and D− are ions AgNO3(aq) + NaCl(aq) AgCl(s) + NaNO3(aq) weak electrolyte (unionized precipitate)

  15. Precipitation Reaction:A Double Replacement (Metathesis) ReactionA (solid) precipitate is formed.

  16. Example: complete and balance the equation (NH4)3PO4(aq) + MgSO4(aq) MgPO4(s) + NH4SO4(aq) we will write the right subscripts later Using the solubility rules, predict if at least one product is going to be insoluble in water. According to rule 8, MgPO4 (subscripts not right) is not soluble in water. Ions are Mg2+, PO43−, NH4+, and SO42−; subscripts (NH4)3PO4(aq) + MgSO4(aq)  Mg3(PO4)2(s) + (NH4)2SO4(aq) balancing 2(NH4)3PO4(aq) + 3MgSO4(aq)  Mg3(PO4)2(s) + 3(NH4)2SO4(aq)

  17. Example: complete and balance theequation • Na2SO4(aq) + BaBr2(aq)  Na2SO4(aq) + BaBr2(aq) BaSO4(s) + NaBr Na2SO4 + BaBr2BaSO4(s) + 2NaBr(aq) driving force = formation of insoluble barium sulfate (precipitate) • Os(NO3)5(aq) + Rb2S(aq)  Os(NO3)5 + Rb2S  Os2S5(s) + RbNO3(aq) 2 Os(NO3)5 + 5 Rb2S  Os2S5(s) + 10 RbNO3 driving force = form. of insoluble Os5+ sulfide (precipitate)

  18. Net Ionic Equations: Spectator Ions The equation AgNO3(aq) + NaCl(aq) AgCl(s) + NaNO3(aq) is not quite correct, because three salts are dissociated in ions while AgCl is a precipitate. Ag+(aq) + NO3−(aq) + Na+(aq) + Cl−(aq) AgCl(s) + Na+(aq) + NO3−(aq) before reaction after reaction Na+ and NO3− are present on both sides of equation, i.e., before and after reaction. They are called spectator ions; they do not participate in net reaction; they can be removed from the equation, but they remain in the solution. Ag+(aq) + Cl−(aq) AgCl(s) is the net ionic equation

  19. Net Ionic Equations: Spectator Ions For two previous examples: 2(NH4)3PO4(aq) + 3MgSO4(aq)  Mg3(PO4)2(s) + 3(NH4)2SO4(aq) 6NH4+(aq) + 2PO43−(aq) + 3Mg2+(aq) + 3SO42−(aq) Mg3(PO4)2(s) + beforereaction 6NH4+(aq) + 3SO42−(aq) after reaction 3Mg2+(aq) + 2PO43−(aq) Mg3(PO4)2(s) is the net equation spectator ions are eliminated from the equation ======================= Na2SO4(aq) + BaBr2(aq) BaSO4(s) + 2NaBr(aq) 2Na+(aq) + SO42−(aq) + Ba2+(aq) + 2Br−(aq) BaSO4(s) + 2Na+(aq) + 2Br− Ba2+(aq) + SO42−(aq)  BaSO4(s) net ionic equation

  20. Net Ionic Equations: Spectator Ions For the metathesis reaction 2 KF(aq) + Pb(NO3)2(aq)  PbF2(s) + 2 KNO3(aq) • formula unit equation spectator ions are eliminated 2K+ + 2F– + Pb2+ + 2 NO3– PbF2 + 2K+ + 2NO3– ionic equation PbF2 is the precipitate 2F–(aq) + Pb2+(aq)  PbF2(s) net ionic equation

  21. Net Ionic Equations: Spectator Ions NH4Cl(aq) + KNO3(aq)  NH4NO3(aq) + KCl(aq) NH4+ + Cl– + K+ + NO3– NH4+ + NO3– + K+ + Cl– all ions are spectators; all can be cancelled no net ionic equation no driving force for reaction N.R. (there is no reaction) The two salts are just dissolved in water.

  22. Acids and Bases Acid • Arrhenius definition • substance that ionizes in water to produce H+, hydrogen ion, and hence increases the concentration of this ion • HCl(aq)  H+(aq) + Cl–(aq) • Brønsted-Lowry definition • substance capable of donating H+ • HCl + H2O  H3O+ + Cl–(aq)

  23. Acids and Bases Base • Arrhenius definition • substance that increases the concentration of OH– in aqueous solution • KOH(aq)  K+(aq) + OH–(aq) • NH3 + H2O NH4+ + OH– • Brønsted/Lowry definition • substance capable of accepting H+ • KOH(aq)  K+(aq) + OH–(aq) • OH– + H+ H2O (OH– from NaOH accepts H+) • NH3 + H2O NH4+ + OH–(NH3 accepts H+)

  24. Water can act as both an acid and a base: it is an amphoteric substance HClO4 + H2O  H3O+ + ClO4– • acid base • (accepts H+ from HClO4) • NH3 + H2O NH4+ + OH– base acid (donates H+ to NH3)

  25. Strong Acids dissociate ~100% • HCl, HBr, HI (no HF) hydro…ic acid • HNO3 nitric acid • HClO3 chloric acid (moderate) • HClO4perchloric acid • H2SO4 (first proton) sulfuric acid H2SO4(aq)  H+(aq) + HSO4−(aq) 2ndweak:HSO4−(aq) H+(aq) + SO42−(aq)

  26. Weak Acids • dissociate <100% • most other acids HF hydrofluoric acid HCN hydrocyanic acid HNO2 nitrous acid CH3CO2H acetic acid H2CO3 carbonic acid (both protons) H3PO4 phosphoric acid (all protons) H2SO3 sulfurous acid (both protons) oxalic acid H2C2O4(aq) H+(aq) + HC2O4−(aq)

  27. Strong Bases dissociate ~100% • alkali metal hydroxides LiOH, NaOH, KOH, RbOH name: lithium hydroxide • hydroxide of • CaCa(OH)2 calcium hydroxide • Ba Ba(OH)2 • SrSr(OH)2 • Ammonia, NH3, is a weak base

  28. Neutralization Reactions • acid + OH-ctg. base  salt + water (a double replacement reaction) HF(aq) + KOH(aq)  KF(aq) + H2O HF(aq) + K+(aq) + OH–(aq)  K+(aq) + F–(aq) + H2O HF(aq) + OH–(aq)  F–(aq) + H2O net ionic spectator ions are eliminated from equation HF is a weak acid and HCl is a strong acid • acid + non-OH-ctg base  salt HCl(aq) + NH3(aq)  NH4Cl(aq) H+(aq) + Cl–(aq) + NH3(aq)  NH4+(aq) + Cl–(aq) H+(aq) + NH3(aq)  NH4+(aq) net ionic equation

  29. Neutralization Reactions • Strong acid + strong base  salt + water (a double replacement reaction) HClO3(aq) + NaOH(aq)  NaClO3(aq) + H2O chloric acid H+(aq) + ClO3−(aq) + Na+(aq) + OH–(aq)  Na+(aq) + ClO3–(aq) + H2O H+(aq) + OH–(aq)  H2O net ionic equation spectator ions, ClO3− + Na+, are eliminated from equation

  30. Formation of a Weak Acid or Base as a Driving Force(another double replacement reaction) • HNO3(aq) + KCN(aq)  HCN(aq) + KNO3(aq) H+(aq) + NO3–(aq) + K+(aq) + CN–(aq)  HCN (aq) + K+(aq) + NO3–(aq) H+(aq) + CN–(aq)  HCN(aq) (a weak acid) • NH4Cl + NaOH(aq)  NH4OH + NaCl(aq) NH4+(aq) + Cl–(aq) + Na+(aq) + OH–(aq)  NH4OH Na+(aq) + Cl–(aq) NH4+(aq) + OH–(aq)  NH4OH (a weak base) NH4OH is NH3 in water, i.e., NH3 + H2O

  31. When no Weak Electrolytes are Formed • HNO3(aq) + KCl(aq)  HCl(aq) + KNO3(aq) H+(aq) + NO3–(aq) + K+(aq) + Cl–(aq)  H+(aq) + Cl–(aq) + K+(aq) + NO3–(aq) There is no net reaction: N.R. No driving force All ions are spectators. • BaCl2(aq) + 2NaOH(aq)  Ba(OH)2(aq) + 2NaCl(aq) Ba2+(aq) + 2Cl–(aq) + 2Na+(aq) + 2OH–(aq)  Ba2+(aq) + 2OH–(aq) + 2Na+(aq) + 2Cl–(aq) There is no net reaction: N.R. No driving force

  32. Gas Forming Reactions (a Driving Force) Some of the weak acids and bases that are formed at double replacement reactions decompose to form a gas and water CO2 Na2CO3(aq) + 2HCl(aq) H2CO3(aq) + 2NaCl(aq) H2CO3(aq)  H2O + CO2(g) Na2CO3(aq) + 2HCl(aq) H2O + CO2(g) + 2NaCl(aq) SO2 Na2SO3(aq) + 2HCl(aq) H2SO3(aq) + 2NaCl(aq) H2SO3(aq)  H2O + SO2(g) Na2SO3(aq) + 2HCl(aq) H2O + SO2(g) + 2NaCl(aq)

  33. Redox (Oxidation-Reduction) Reactions • involve transfer of electron(s) • oxidation: loss of electron(s) • reduction: gain of electron(s) • some can be identified when an uncombined element is a reactant or a product • eg. Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s) • Zn  Zn2+ • Zn(s)  Zn2+(aq) + 2 e–, oxidation • Cu2+ Cu • Cu2+(aq) + 2e– Cu(s), reduction

  34. Single Displacement Reactions • Zn(s) + CuCl2(aq)  Cu(s) + ZnCl2(aq) • Zn oxidized to Zn2+ • Cu2+ reduced to Cu • occurs because zinc is more active than copper • Cl2(g) + CuBr2(aq)  Br2(l)+ CuCl2(aq) • Br oxidized from Br– to Br2 • Cl reduced from Cl2 to Cl– • Cl is more active than Br

  35. Oxidation Numbers • also an accounting tool • very useful • oxidation numbers of all atoms in substance add up to charge on substance e.g. (charge of species) zero for Al2(SO4)3 and H3PO4 +1 for NH4+ –2 for Cr2O72–

  36. Assigning Oxidation Numbers, ON • ON = 0 for all atoms in any substance in most elemental form, Na(s), Zn(s), Hg(l) H2(g), Cl2(g), I2(s), O2(g), C(s), P4(s), S8(s) • ON = charge for monatomic ions (Na+, S2−) • ON = –1 for F in all compounds • ON = –2 for O in compounds, usually • exceptions: peroxide, O22–, ON = –1 • superoxide, O2–, ON = –1/2 • ON = +1 for H in compounds, usually • exception: ON = –1 in metallic hydrides

  37. Assigning Oxidation Numbers, ON • ON = +1 for alkali metals in compounds • ON = +2 for alkaline earth metals in compounds • ON = +3 for Al in compounds • ON = −1 for Cl, Br, and I (iodine) in binary compounds except for those with oxygen (in these cases they variable positive) • F is always −1

  38. Assign ON to Each Atom in the Following Substances ? –1 WCl6 x + 6(–1) = 0 x –6 = 0 x = +6

  39. +1 –2 ? Na2S2O3 +2 + 2x – 6 = 0 2x = 6 – 2 x = +2 ON of S

  40. +1 –2 ? Na2S4O8 +2 + 4x –16 = 0 4x = 16 – 2 +14 +7 x = —— = —— ON of S 4 2

  41. ? –2 Cr2O72– 2x –14 = –2 2x = 14 – 2 +12 x = —— = +6 ON of Cr 2

  42. +1 –2 H2C2O4 +2 + 2x – 8 = 0 2x = 8 – 2 +6 x = —— = +3 ON of C 2

  43. ? –1 MoBr5+ x – 5 = +1 x = 5 + 1 x = 6 ON of Mo = +6

  44. Oxidizing and Reducing Agents In every redox reaction there is (at least) a Reducing agent (the one that is oxidized) and an oxidizing agent (the one that is reduced) ON increases ON decreases The species is The species is oxidized reduced +7 +6 +5 +4 +3 +2 +1 0 −1 −2 −3 −4 −5 −6 −7

  45. Activity (Electromotive) Series for metals

  46. Activity increases Li K Ba Ca Na Mg Al Mn Zn Cr Fe Co Ni Sn Pb (H2) Cu Ag Hg Pt Au Activity decreases

  47. Examples • Complete and balance each of the following chemical equations

  48. Li K Ba Ca Na Mg Al Mn Zn Cr Fe Co Ni Sn Pb H2 Cu Hg Ag Pt Au A free and chemically active metal displacing a less active metal from a compound • Mg + FeCl3 • Mg + FeCl3 Fe + MgCl2 • 3Mg(s) + 2FeCl3(aq)  2Fe(s) + 3MgCl2(aq) Mg  Mg2+, oxidized; Mg reducing agent Fe3+ Fe, reduced; Fe3+ oxidizing agent • Sn + CrF3 Sn is less reactive than Cr Sn + CrF3No Reaction • Pb(s) + Au(ClO3)3(aq)  • Pb(s) + Au(ClO3)3(aq)  Au(s) + Pb(ClO3)2(aq) 3Pb(s) + 2Au(ClO3)3(aq)  2Au(s) + 3Pb(ClO3)2(aq) Pb is oxidized Au+3 is reduced

  49. Li K Ba Ca Na Mg Al Mn Zn Cr Fe Co Ni Sn Pb H2 Cu Hg Ag Pt Au A free and chemically active metal displacing a less active metal from a compound • Zn + CrBr3 • Zn + CrBr3 Cr + ZnBr2 3Zn(s) + 2CrBr3(aq)  2Cr(s) + 3 ZnBr2(aq) • Zn oxidized to Zn2+; Zn reducing agent • Cr3+ reduced to Cr; Cr3+ oxidizing agent • Ag(s) + Hg(NO3)2 No Reaction • Ag is less reactive than Hg

  50. Li K Ba Ca Na Mg Al Mn Zn Cr Fe Co Ni Sn Pb H2 Cu Hg Ag Pt Au A free and chemically active metal displacing Hydrogen from acids or water • Fe + HBr  • Fe + HBr  H2 + FeBr3 • 0 +1 0 +3 • 2Fe + 6HBr  3H2 + 2 FeBr3 Fe oxidized to Fe3+; Fe reducing agent H+ reduced to H2; H+ oxidizing agent • Cu + HBr  Cu less active than H2 Cu + HBr  No Reaction

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