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  • 1. Lab 5 Single Replacement Reactions
  • 2. Procedures
    • Four different metals were added to hydrochloric acid solution
    • Magnesium reacted immediately with rapid bubbling until the metal completely dissolved.
    • Zinc reacted with slow bubbling.
    • Aluminum reacted much slower. Only reacting with the 3.0 MHCl.
    • Copper did not react at all, regardless of the length of time.
  • 3. Single Replacement Reactions
    • Single Replacement Reactions occur when one element replaces another in a compound.
    • A metal can replace a metal (+) OR a nonmetal can replace a nonmetal (-).
    • element + compound  product + product
      • A + BC  AC + B (if A is a metal) OR
      • A + BC  BA + C (if A is a nonmetal)
    • (remember the cation always goes first!)
    • When H 2 O splits into ions, it splits into
    • H + and OH - (not H+ and O -2 !!)
  • 4. Single Replacement Reactions
    • The metal only replaces hydrogen when the metal is more active than hydrogen.
    • Magnesium , zinc and aluminum must be more active than hydrogen, since they replace hydrogen in the acid
    • Copper must be less active than the hydrogen, since it does react with the hydrogen in the acid.
  • 5. Analysis of the Single Replacement Reaction 1 p 1 p 1 e- 1 e- 12 p 17 p 12 e- 18 e- 1 p 17 p 0 e- 18 e- 12 p 12 e- H atoms sharing a pair of e- in nonpolar covalent bonds Mg 2+ & Cl - ions in solution by ion-dipole attractions H+ and Cl- ions in solution by ion-dipole attractions Atoms held in fixed positions by metallic bonds + H 2 (g) MgCl 2 (aq) HCl (aq) Mg(s) + PRODUCTS REACTANTS
  • 6. Oxidation and Reduction
    • Oxidation Half-reaction
      • Mg atoms  Mg 2+ ions
      • 12 p 12 p
      • 12 e- 10 e-
      • 0 +2 charge
      • Each magnesium atom lost two electrons
      • Mg  Mg 2+ + 2 e-
      • Notice the reaction is balanced for protons and electrons on both sides.
  • 7. Oxidation and Reduction
    • Reduction Half-reaction
      • hydrogen ions  hydrogen molecules
      • 1 p 1 p
      • 0 e- 1 e-
      • +1 0 charge
      • Each hydrogen ion gained an e- to form a hydrogen atom, then two hydrogen atoms get together to share a pair of electrons
    • 2e- + 2 H + (aq)  H 2 (g)
    • Notice the reaction is balanced for protons and electrons on both sides.
  • 8. LEO the lion says GER ! ose lectrons xidation ain lectrons Eduction GER! Oxidation and Reduction
  • 9. Equations to Summarize the Reaction
    • Molecular equation
    • Mg(s) + 2 HCl (aq)  MgCl 2 (aq) + H 2 (g)
    • Total Ionic Equation
    • Mg + 2H + (aq) + Cl - (aq)  Mg 2+ (aq) + Cl - (aq) + H 2 (g)
    • Net Ionic Equation
    • Mg + 2 H + (aq)  Mg 2+ (aq) + H 2 (g)
  • 10. Single Replacement Reactions
  • 11. Activity Series of Metals Each element on the list replaces from a compound any of the elements below it. The larger the interval between elements, the more vigorous the reaction. The first five elements (lithium - sodium) are known as very active metals and they react with cold water to produce the hydroxide and hydrogen gas. The next four metals (magnesium - chromium) are considered active metals and they will react with very hot water or steam to form the oxide and hydrogen gas. The oxides of all of these first metals resist reduction by H2. The next six metals (iron - lead) replace hydrogen from HCl and dil. sulfuric and nitric acids. Their oxides undergo reduction by heating with H2, carbon, and carbon monoxide. The metals lithium - copper, can combine directly with oxygen to form the oxide. The last five metals (mercury - gold) are often found free in nature, their oxides decompose with mild heating, and they form oxides only indirectly. lithium potassium strontium calcium sodium ------------------------------- magnesium aluminum zinc Chromium -------------------------------- iron cadmium cobalt nickel tin Lead -------------------------------- HYDROGEN antimony arsenic bismuth Copper -------------------------------- mercury silver palladium Platinum gold