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CHEMICAL REACTIONS Chapter 3

CHEMICAL REACTIONS Chapter 3. Reactants: Zn + I 2. Product: ZnI 2. Chemical Equations. Depict the kind of reactants and products and their relative amounts in a reaction. 4 Al(s) + 3 O 2 (g)  2 Al 2 O 3 (s) The numbers in the front are called stoichiometric coefficients

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CHEMICAL REACTIONS Chapter 3

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  1. CHEMICAL REACTIONSChapter 3 Reactants: Zn + I2 Product: ZnI2

  2. Chemical Equations Depict the kind of reactants and products and their relative amounts in a reaction. 4 Al(s) + 3 O2(g)  2 Al2O3(s) The numbers in the front are called stoichiometric coefficients The letters (s), (g), and (s) are the physical states of compounds.

  3. Reaction of Phosphorus with Cl2 Notice the stoichiometric coefficients and the physical states of the reactants and products.

  4. Reaction of Iron with Cl2 Notice the stoichiometric coefficients and the physical states of the reactants and products.

  5. Chemical Equations 4 Al(s) + 3 O2(g) 2 Al2O3(s) This equation means 4 Al atoms + 3 O2 molecules  2 “molecules” of Al2O3 4 moles of Al + 3 moles of O2 2 moles of Al2O3

  6. Chemical Equations PLAY MOVIE • Because the same atoms are present in a reaction at the beginning and at the end, the amount of matter in a system does not change. • The Law of the Conservation of Matter 2HgO(s) 2 Hg(liq) + O2(g)

  7. Lavoisier, 1788 Chemical Equations Because of the principle of the conservation of matter, an equation must be balanced. It must have the same number of atoms of the same kind on both sides.

  8. PLAY MOVIE Balancing Equations ___ Al(s) + ___ Br2(s)  ___ Al2Br6(s)

  9. Balancing Equations PLAY MOVIE ____C3H8(g) + _____ O2(g)  _____CO2(g) + _____ H2O(g) ____B4H10(g) + _____ O2(g)  ___ B2O3(g) + _____ H2O(g)

  10. Chemical Equilibrium • Chemical reactions are reversible. • Ammonia can be produced from the elements in the Haber process N2(g) + 3 H2(g)  2 NH3(g) • But NH3 can also be decomposed to the elements 2 NH3(g)  N2(g) + 3 H2(g) • In a process to make NH3, the reaction can come eventually to equilbrium. N2(g) + 3 H2(g)  2 NH3(g) • Double arrows indicate equilibrium

  11. Reaction Reversibility Stalactites and stalagmites in caves depend on a reversible chemical reaction Ca2+(aq) + 2 HCO3–(aq)  CaCO3(s) + CO2(g) + H2O(s)

  12. Reaction Reversibility

  13. Chemical Equilibrium Once equilibrium is achieved, reaction continues, but there is no net change in amounts of products or reactants.

  14. K+(aq) + MnO4-(aq) Reactions in Aqueous Solution Many reactions involve ionic compounds, especially reactions in water — aqueous solutions. KMnO4 in water PLAY MOVIE PLAY MOVIE

  15. An Ionic Compound, CuCl2, in Water

  16. Aqueous Solutions How do we know ions are present in aqueous solutions? The solutions conduct electricity! They are called ELECTROLYTES HCl, CuCl2, and NaCl are strong electrolytes. They dissociate completely (or nearly so) into ions.

  17. Aqueous Solutions HCl, CuCl2, and NaCl are strong electrolytes. They dissociate completely (or nearly so) into ions. PLAY MOVIE

  18. Aqueous Solutions Acetic acid ionizes only to a small extent, so it is a weak electrolyte. CH3CO2H(aq)CH3CO2-(aq) + H+(aq) PLAY MOVIE

  19. Aqueous Solutions Acetic acid ionizes only to a small extent, so it is a weak electrolyte. CH3CO2H(aq)CH3CO2-(aq) + H+(aq)

  20. Aqueous Solutions Some compounds dissolve in water but do not conduct electricity. They are called nonelectrolytes. Examples include: sugar ethanol ethylene glycol

  21. Water Solubility of Ionic Compounds If one ion from the “Soluble Compd.” list is present in a compound, the compound is water soluble.

  22. Iron pyrite, a sulfide Orpiment, arsenic sulfide Azurite, a copper carbonate Water Solubility of Ionic Compounds Common minerals are often formed with anions that lead to insolubility: sulfide fluoride carbonate oxide

  23. Chemical Reactions in Water Pb(NO3) 2(aq) + 2 KI(aq)  PbI2(s) + 2 KNO3 (aq) We will look at EXCHANGEREACTIONS The anions exchange places between cations. PLAY MOVIE

  24. Precipitation Reactions The “driving force” is the formation of an insoluble compound — a precipitate. Pb(NO3)2(aq) + 2 KI(aq)  2 KNO3(aq) + PbI2(s) BaCl2(aq) + Na2SO4(aq)  BaSO4(s) + 2 NaCl(aq)

  25. Net Ionic Equations PLAY MOVIE Pb(NO3)2(aq) + K2CrO4(aq)  PbCrO4(s) + 2 KNO3(aq) This is the “complete equation” Because Pb(NO3)2 and K2CrO4 are strong electrolytes we should write Pb2+(aq) + 2 NO3-(aq) + 2 K+(aq) + CrO42-(aq)  PbCrO4(s) + 2 K+(aq) + 2 NO3-(aq) This is the “ionic equation” Question: do we need to include the K+ and NO3- ions?

  26. Net Ionic Equations Ionic equation: Pb2+(aq) + 2 NO3-(aq) + 2 K+(aq) + CrO42-(aq)  PbCrO4(s) + 2 K+(aq) + 2 NO3-(aq) The NO3- and K+ ions are SPECTATOR IONS — they do not participate. Could have used Na+ instead of K+. We leave the spectator ions out — Pb2+(aq) + CrO42-(aq)  PbCrO4(s) to give the NET IONIC EQUATION

  27. ACIDS An acid  H3O+ in water PLAY MOVIE

  28. The Nature of Acids PLAY MOVIE

  29. HNO3 ACIDS An acid  H3O+ in water Some strongacids are HCl hydrochloric H2SO4 sulfuric HClO4 perchloric HNO3 nitric

  30. Acetic acid Weak Acids WEAK ACIDS = weak electrolytes CH3CO2H acetic acid H2CO3 carbonic acid H3PO4 phosphoric acid HF hydrofluoric acid PLAY MOVIE

  31. ACIDS Nonmetal oxides can be acids CO2(aq) + H2O(s)  H2CO3(aq) SO3(aq) + H2O(s)  H2SO4(aq) and can come from burning coal and oil.

  32. BASESTable 3.2 Base  OH- in water NaOH(aq)  Na+(aq) + OH-(aq) NaOH is a strong base PLAY MOVIE

  33. Ammonia, NH3An Important Base PLAY MOVIE

  34. BASES Metal oxides are bases CaO(s) + H2O(s)  Ca(OH)2(aq) Metals from Groups 1A and 2A CaO in water. Indicator shows solution is basic.

  35. Know the strong acids & bases!

  36. Acid-Base Reactions • The “driving force” is the formation of water. NaOH(aq) + HCl(aq)  NaCl(aq) + H2O(liq) • Net ionic equation OH-(aq) + H3O+(aq)  2 H2O(l) • This applies to ALL reactions of STRONG acids and bases. PLAY MOVIE

  37. See Active Figure 3.14

  38. Acid-Base Reactions • A-B reactions are sometimes called NEUTRALIZATIONSbecause the solution is neither acidic nor basic at the end. • The other product of the A-B reaction is a SALT, MX. HX + MOH MX + H2O Mn+ comes from base &Xn- comes from acid This is one way to make compounds!

  39. Gas-Forming Reactions This is primarily the chemistry of metal carbonates. CO2 and water  H2CO3 H2CO3(aq) + Ca2+ 2 H+(aq) + CaCO3(s) (limestone) Adding acid reverses this reaction. MCO3 + acid  CO2 + salt

  40. Gas-Forming Reactions CaCO3(s) + H2SO4(aq)  2 CaSO4(s) + H2CO3(aq) Carbonic acid is unstable and forms CO2 & H2O H2CO3(aq) CO2 + water (Antacid tablet has citric acid + NaHCO3) PLAY MOVIE

  41. Oxidation-Reduction ReactionsSection 3.9 Thermite reaction Fe2O3(s) + 2 Al(s)  2 Fe(s) + Al2O3(s)

  42. EXCHANGE: Precipitation Reactions EXCHANGE Gas-Forming Reactions EXCHANGE Acid-Base Reactions REACTIONS REDOX REACTIONS

  43. REDOX REACTIONS REDOX = reduction & oxidation O2(g) + 2 H2(g) 2 H2O(s)

  44. REDOX REACTIONS REDOX = reduction & oxidation Corrosion of aluminum 2 Al(s) + 3 Cu2+(aq)  2 Al3+(aq) + 3 Cu(s)

  45. REDOX REACTIONS Cu(s) + 2 Ag+(aq)  Cu2+(aq) + 2 Ag(s) In all reactions IF something has been oxidized then something has also been reduced PLAY MOVIE

  46. REDOX REACTIONS Cu(s) + 2 Ag+(aq)  Cu2+(aq) + 2 Ag(s)

  47. Fuels Why Study Redox Reactions Batteries Corrosion Manufacturing metals

  48. REDOX REACTIONS Redox reactions are characterized byELECTRON TRANSFER between an electron donor and electron acceptor. Transfer leads to— 1. increase in oxidation number of some element = OXIDATION • decrease in oxidation number of some element = REDUCTION O I L R I G

  49. O I L R I G • Oxidation • It • Loses (electrons) • Reduction • It • Gains (electrons)

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