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Reactions in Aqueous Solutions

Reactions in Aqueous Solutions. Chapter 7. Aqueous Solutions. Water is the dissolving medium, or solvent. Driving Forces for Chemical Reactions. 1. Formation of a precipitate (solid). 2. Formation of molecular substance (water). 3. Formation of a gas.

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Reactions in Aqueous Solutions

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  1. Reactions in Aqueous Solutions • Chapter 7

  2. Aqueous Solutions Water is the dissolving medium, or solvent.

  3. Driving Forces for Chemical Reactions • 1. Formation of a precipitate (solid). • 2. Formation of molecular substance (water). • 3. Formation of a gas. • 4. Transfer of electrons (Redox).

  4. Types of Solution Reactions • Precipitation reactions • AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq) • Acid-base reactions • NaOH(aq) + HCl(aq)  NaCl(aq) + H2O(l)

  5. Electrolytes • Strong - conduct current efficiently - many ions in solution. • NaCl, KNO3, HNO3, NaOH • Weak - conduct only a small current - few ions in solution, • HC2H3O2, aq. NH3, tap H2O • Non - no current flows - no ions in solution. • pure H2O, sugar solution, glycerol

  6. In what two ways can a solid ionic compound be made to conduct electricity? Dissolve it in water. Melt or fuse it. Ions must be free to move (mobile) in order to conduct electricity!

  7. Dissociation • ionic compounds • metal + nonmetal (Type I & II) • metal + polyatomic anion • Ammonium compounds • Acids • When ionic compounds dissolve in water the anions and cations are separated from each other; this is called dissociation • We know that ionic compounds dissociate when they dissolve in water because the solution conducts electricity

  8. K Cl K+ Cl- Cu SO4 Cu+2 SO42- Dissociation • potassium chloride dissociates in water into potassium cations and chloride anions • KCl(aq) = K+ (aq) + Cl- (aq) • copper(II) sulfate dissociates in water into copper(II) cations and sulfate anions • CuSO4(aq) = Cu+2(aq) + SO42-(aq)

  9. K+ SO4 K K SO42- K+ Dissociation • potassium sulfate dissociates in water into potassium cations and sulfate anions • K2SO4(aq) = 2 K+ (aq) + SO42-(aq)

  10. Ionic Compounds in Solution In aqueous solution, soluble ionic compounds exist in the form of ions. K2CrO4(aq) + Ba(NO3)2(aq) ----> BaCrO4(s) + 2KNO3(aq)

  11. Figure 7.1: The precipitation reaction that occurs when yellow potassium chormate, K2CrO4 (aq), is mixed with a colorless barium nitrate solution, Ba(NO3)2 (aq)

  12. Solubility • 1. A soluble solid readily dissolves in water-- designated with (aq). • 2. A slightly soluble solid only dissolves to a tiny extent in water--designated with (s). • 3. An insoluble solid does not dissolve to any appreciable extent in water--designated with (s).

  13. Simple Rules for Solubility • 1. Most nitrate (NO3), acetate (C2H3O2-), & chlorate (ClO3-) salts are soluble. • 2. Most alkali (group 1A) salts and NH4+are soluble. • 3. Most Cl, Br, and I salts are soluble (NOTAg+, Pb2+, Hg22+) • 4. Most sulfate salts are soluble (NOTBaSO4, PbSO4, HgSO4, CaSO4) • 5. Most OH salts are only slightly soluble (NaOH, KOH are soluble, Ba(OH)2, Ca(OH)2 are marginally soluble) • 6. Most S2, CO32, CrO42, PO43 salts are only slightly soluble.

  14. K+ Ag+ K+ Ag I I- NO3- NO3- Precipitation Reactions • In all precipitation reactions, the ions of one substance are exchanged with the ions of another substance when their aqueous solutions are mixed • At least one of the products formed is insoluble in water • KI(aq) + AgNO3(aq) KNO3(aq) + AgIs

  15. Describing Reactions in Solution • 1. Molecular equation (reactants and products as compounds) • AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq) • 2.Complete ionic equation (all strong electrolytes shown as ions) • Ag+(aq) + NO3(aq) + Na+(aq) + Cl(aq)  • AgCl(s) + Na+(aq) + NO3(aq)

  16. Describing Reactions in Solution (continued) • 3. Net ionic equation (show only components that actually react) • Ag+(aq) + Cl(aq)  AgCl(s) • Na+ and NO3 are spectator ions.

  17. Solutions of: AgNO3 Na2SO4

  18. If AgNO3 is mixed with Na2SO4 what ions are most abundant in the solution? AgNO3 Na2SO4 With what ions is the solution saturated?

  19. 2AgNO3(aq) + Na2SO4(aq) 2 NaNO3 (aq) + Ag2SO4(s) Molecular Equation 2Ag+(aq)+ 2NO3-(aq) + 2Na+(aq) + SO42-(aq) 2Na+(aq) + 2 NO3-(aq) + Ag2SO4(s) Overall Ionic Equation 2Ag+(aq)+ SO42-(aq) Ag2SO4(s) Net Ionic Equation Silver Sulfate Precipitate

  20. Acids • The nature of acids was discovered by Svante Arrhenius. • Acids are characterized by: • a sour taste (lemons -- citric acid). • producing H+ ions (protons) in aqueous solution. • turn blue litmus red.

  21. Acids • Strong acids - dissociate completely (nearly 100 %) to produce H+ in solution • HCl, H2SO4, HNO3, HBr, HI, & HClO4 • Weak acids - dissociate to a slight extent (approximately 1 %) to give H+ in solution • HC2H3O2, HCOOH, HNO2, & H2SO3

  22. Figure 7.5: When gaseous HCl is dissolved in water, each molecule dissociates to produce H+ and Cl- ions

  23. Bases • Bases are characterized by: • bitter taste (soap). • feel slippery. • produce hydroxide (OH-) ions in aqueous solution. • turn red litmus blue. • A basic solution is said to be alkaline since they often contain one of the alkali metals --Na, K, Li, etc.

  24. Bases • Strong bases - react completely with water to give OH ions. • sodium hydroxide • Weak bases - react only slightly with water to give OH ions. • ammonia

  25. Acid-Base Reactions Molecular Equation HCl(aq) + NaOH(aq) ----> NaCl(aq) + HOH(l) Complete Ionic Equation H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) -----> Na+(aq) + Cl-(aq) + HOH(l) Acid + Base ----> Salt + water Net Ionic Equation H+(aq) + OH-(aq) -----> HOH(l)

  26. Acid-Base Neutralization The net ionic equation for the neutralization of any strong acid and base will always be: H+(aq) + OH-(aq) -----> HOH(l)

  27. Salts • Salts are ionic compounds consisting of: • a. metal & nonmetal -- KCl • b. metal & polyatomic ion -- CuSO4 • c. polyatomic ion & nonmetal -- NH4Cl • d. two polyatomic ions -- (NH4)2SO4

  28. Oxidation-Reduction Reactions(Redox Reactions) • Redox reaction -- involves the transfer of electrons • loss and gain of electrons must be exactly equal. • loss and gain of electrons must be simultaneous.

  29. Oxidation • loss of electrons • metal atoms -- Na, Ca, & K • Nao ---> Na1+ + 1e- • Cao ---> Ca2+ + 2e- • nonmetal ions -- Cl-, S2-, & O2- • Cl1- ---> Clo +1e- • S2- ---> So + 2e-

  30. Reduction • gain of electrons. • nonmetal atoms. • Oo + 2e- ---> O2- • Fo + 1e- ---> F1- • metal ions. • K1+ + 1e- ---> Ko • Ba2+ + 2e- ---> Bao

  31. Figure 7.7: When powdered aluminum and iodine (shown in the foreground) are mixed (and a little water added), they react vigorously

  32. Oxidation & Reduction Half-Reactions • Always add electrons (negative) to the more positive side of the equation. • The charge on both sides of an equation must be equal. • Oxidation -- Nao ----> Na1+ + 1 e- • Reduction -- Cl2 + 2 e- ----> 2 Cl-

  33. OIL RIG • Oxidation Is Loss. • Reduction Is Gain.

  34. Redox Reactions • Metal-nonmetal reactions are always redox reactions. • Any reaction that has a free element (such as O2) as a reactant or product are redox. • All single replacement reactions are redox. • All combustion reactions are redox.

  35. Synthesis I (Composition) • A + X ----> AX ALWAYS REDOX • Element + Element -----> Binary Compound • Fe(s) + S(s) ----> FeS(s) • 4 Al(s) + 3 O2(g) ----> 2 Al2O3(s)

  36. Synthesis II (Composition) • A + X ----> AX NOT REDOX • Compound + Compound ----> Compound (3 or more elements) • Ammonia + Acid ----> Ammonium Salt • NH3(g) + HCl(g) ----> NH4Cl(s) • 2 NH3(aq) + H2SO4(aq) ----> (NH4)2SO4(aq)

  37. Synthesis II (Composition)Continued • Water + An Oxide • Rule # 1 --Water + Metal Oxide ----> Metal Hydroxide (Base) • HOH(l) + CaO(s) ----> Ca(OH)2(s) • HOH(l) + Na2O(s) ----> 2 NaOH(aq) • Rule # 2 --Water +Nonmetal Oxide ----> Acid • HOH(l) + SO3(g) ----> H2SO4(aq) • HOH(l) + N2O5(g) ----> 2 HNO3(aq)

  38. Single Replacement I • A + BX ----> AX + B ALWAYS REDOX • Element + Compound ----> Different Element + Different Compound • Metal Reactivities: • K > Na > Ca > Mg > Al > Zn > Cr > Fe > Cd > Co > Ni > Sn > Pb > H > Sb > Cu > Hg > Ag > Pt > Au

  39. Single Replacement I(Continued) • A is a metal. • A is more reactive than B • A + BX ----> AX + B • Cu(s) + 2 AgNO3(aq) ---->Cu(NO3)2(aq) + 2 Ag(s) • Zn(s) + 2 HCl(aq) ----> ZnCl2(aq) + H2(g) • Rule # 3 -- Active Metal + Water ----> Metal Hydroxide + Hydrogen • Ca(s) + 2 HOH(l) ----> Ca(OH)2(s) +H2(g)

  40. Single Replacement I(Continued) • A is a metal. • A is less reactive than B • A + BX ----> No Reaction (NR) • Cu(s) + HCl(aq) ----> NR

  41. Single Replacement II • Y + BX ----> BY + X ALWAYS REDOX • Element + Compound ----> Different Element + Different Compound • Nonmetal Reactivities: • F > O > Cl > Br > I

  42. Single Replacement II(Continued) • Y is a nonmetal. • Y is more reactive than X • Y + BX ----> BY + X • Cl2(g) + 2 KBr(aq) ----> 2 KCl(aq) + Br2(aq) • Y is less reactive than X • Y + BX ----> No Reaction • Cl2(g) + KF(aq) ----> NR

  43. Single Replacement(Continued) • Remember!! • Metals replace metals. • Mg(s) + 2 Ag(NO3)2(aq) ----> 2 Ag(s) + Mg(NO3)2(aq) • Nonmetals replace nonmetals. • Cl2(g) + 2 KBr(aq) ----> Br2(aq) + 2 KCl(aq)

  44. Figure 7.6: The thermite reaction gives off so much heat that the iron formed is molten

  45. Decomposition I • AX ----> A + X ALWAYS REDOX • AX is a binary compound. • Binary Compound ----> Element + Element • elect • 2 NaCl(l) ----> 2 Na(l) + Cl2(g) • elect • 2 HOH(l) ----> 2 H2(g) + O2(g)

  46. Decomposition II • AX -----> A + X MAY OR MAY NOT BE REDOX • AX is a ternary compound. • Rule # 4 -- Metal Chlorates ----> Metal Chlorides + Oxygen •  • 2 KClO3(s) ----> 2 KCl (s) + 3 O2(g) REDOX

  47. Decomposition II(Continued) • Rule # 5 -- Metal Carbonates ----> Metal Oxides + Carbon Dioxide •  • CuCO3(s) ----> CuO (s) + CO2(g) NOTREDOX

  48. Decomposition II(Continued) • Rule # 6 -- Metal Hydroxides ----> Metal Oxides + Water •  • Ca(OH)2(s) ----> CaO (s) + HOH(g) NOTREDOX

  49. Decomposition II(Continued) • Rule # 7 -- Acid ----> Nonmetal Oxide + Water •  • H2SO4(aq) ----> SO3(g) + HOH(g) NOTREDOX

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