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Reaction Predictions

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					Reaction Predictions
       Most Commonly Used
        Cations and Anions
 Hydrogen H+           •Hydroxide OHˉ
 Sodium Na+            •Chloride Clˉ
 Potassium K+          •Sulfide Sˉ²
 Calcium Ca+²
                        •Bicarbonate HCOзˉ
                        •Carbonate COзˉ²
 Magnesium Mg+²
                        •Sulfate SO4ˉ²
 Iron (Ferrous) Fe+²
                        •Phosphate PO4ˉ³
 Iron (Ferric) Fe+³
      Cations/ Anions, contd.
 You can figure out the charge of an ion by
  using the periodic table. For Example:
 Alkali metals such as Lithium can easily
  lose an electron to become stable (just like
  a Noble gas) so taking away an electron
  give Lithium a +1 charge.
 On the other hand Halogens can easily
  accept an electron to become stable.
  Accepting an electron gives halogens a -1
  charge.
                   Practice
   What is the oxidation state of Oxide?

   What is the oxidation state of Iodide?

   What is the oxidation state of a Calcium
    ion?

   What is the oxidation state of a Lithium
    ion?
         Answers
   -2

   -1

   +2

   +1
         Net Ionic Equation
 To create a net ionic equation, you break
  apart all ionic molecules in a balanced
  molecular equation into their ions if they
  are soluble.
 If there are spectator ions, ions that
  appear on both sides of the equation, they
  cancel each other.
            Net Ionic Example
   Silver nitrate is mixed with potassium
    chromate
     2AgNO3 + K2CrO4 → Ag2CrO4 + 2KNO3
       Molecular Equation
     2Ag+ + 2NO3ˉ + 2K+ + CrO4-2 → Ag2CrO4 +
      2K+ + 2NO3-2
             Complete ionic equation
     2Ag+ + CrO4-2 → Ag2CrO4
             Net Ionic Equation
                 Solubility Rules
   NO3
       -     all nitrates are soluble

             -
    CH3COO or C2H3O2
                         -
    all acetates are soluble except AgCH3COO
   ClO3
         -     all chlorates are soluble

      -
    Cl all chlorides are soluble except AgCl, Hg2Cl2,
    PbCl2

        -
    Br all bromides are soluble except AgBr, PbBr2,
    Hg2Br2, and HgBr2

    -
    I all iodides are soluble except AgI, Hg2I2, HgI,
    and PbI2
      Solubility Rules, contd.
 SO4¯² all sulfates are soluble except
  BaSO4, PbSO4, Hg2SO4, CaSO4,
  AgSO4 and SrSO4
 Alkali metal, cations, and NH4 – all are
  soluble
 H+    all common inorganic acids and low
  molecular mass organic acids are soluble
       (In)Soubility Rules, contd.

        -
    CO3 ²     all carbonates are insoluble except those of
    alkali metals and NH4

          -
    CrO4 ² all chromates are insoluble except those of
    alkali metals, NH4, CaCrO4, and SrCO4
   OH
        -     all hydroxides are insoluble except those of the
    alkali metals, NH4, Ba(OH)2, Sr(OH)2, and Ca(OH)2
   PO4 ³
         -    all phosphates are insoluble except those of
    alkali metals and NH4
   SO3
         -² all sulfites are insoluble except those of alkali
    metals and NH4

      -
    S ²       all sulfides are insoluble except those of alkali
    metals and NH4
                       Synthesis
   Synthesis occurs when two or more reactants combine
    to form a single product. There are several common
    types of synthesis reaction.

   You know it happens when you have:
    -A metal combines with a nonmetal to form a bianary
    salt.
          -A piece of lithium metal is dropped into a container
    of nitrogen gas.
          6Li+ N2  2Li3N

    -Metal oxide and water forms a base (metallic
    hydroxide)
        -Solid sodium oxide is added to water.
        Na2O + H2O 2NaOH
          Synthesis, contd.
 Nonmetal oxide and water forms acids.
  Nonmetal retains its oxidation number.
     -Carbon dioxide is burned in water.
     CO2 + H2O  H2CO3
 Metallic oxides and nonmetallic oxides
  form salts.
     -Solid sodium oxide is added to carbon
  dioxide.
     Na2O + CO2  Na2CO2
            Decomposition
 Occurs when a single reactant is broken
  down into two or more products.
 The reactions react to form basic
  compounds or elements.
 When a compound is heated or
  electrolyzed, it means that it is broken up
  into its ions.
 AB A+B
    Examples of Decomposition
 A sample of magnesium carbonate is
  heated.
     MgCO3  MgO + CO2
 Molten sodium chloride is electrolyzed.
     2NaCl  2Na + Cl2
 A sample of ammonium carbonate is
  heated.
     (NH4)2CO3  2NH3 + H2O + CO2
        Single Replacement
 Reactions that involve an element
  replacing one part of a compound. The
  products include the displace element and
  a new compound. An element can only
  replace another element that is less active
  than itself. (Look a activity series/ AP
  packet)
 A +BX B+AX
     Single Replacement Rules
1.   Active metals replace less active metals
     from the less active metals‟ compounds
     in aqueous solutions
     ex. 3Mg+ 2FeCl3—> 2Fe + 3MgCl2
2.   Active metals replace hydrogen in water
     ex. 2Na + 2H2O—> H2 + 2NaOH
3.   Active metals replace hydrogen in acids
     ex. 2Li + 2HCl —> H2 + 2LiCl
   Single Replacement Rules,
             contd.
4. Active nonmetals replace less active
  nonmetals from their compounds in
  aqueous solutions
  ex. Cl2 + 2KI —> I2 + 2KCl
5. If a less reactive element is combined
  with a more reactive element in compound
  form, there will be no reaction
  ex. Cl2 + KF —> no reaction*
* On the AP test reactions will ALWAYS
  have products; it will never be “no
  reaction.”
    Activity Series (Single Replacement)

   Metals
       Li, Ca, Na, Mg, Al, Zn, Fe, Pb, [H2], Cu, Ag, Pt

   Nonmetals
       F2, Cl2, Br2, I2,

    More active                 Less Active
            Double Replacement
   Two compounds react to form two new compounds. No
    changes in oxidation numbers occur.
   Each cation pairs up with the anion in the other
    compound.
   The “driving force” in these reactions is the removal of at
    least one pair of ions from solution.
   This removal of ions happens with the formation of a
    precipitate, gas, or molecular species.
   When a double replacement reaction doesn‟t go to
    completion, it is a reversible reaction (no ions have been
    removed).
   AX+ BY  AY+ BX
       How do you know a double
      replacement reaction occurs?
   The reactants will contain a(n):
    -gas
    -insoluble precipitate
    -molecular species

*Remember– on the AP test the reaction will
  always occur
Common Gases Released (Dbl. Repl.)

 H2S Any sulfide plus any acid forms
      H2S and a salt.
 CO2 Any carbonate plus any acid form
      CO3, water, and a salt.
 SO2 Any sulfite plus any acid form SO2,
      water, and a salt.
 NH3 Any ammonium plus a soluble
      hydroxide form NH3, water, and a
           salt.
    Acid/ Base Reactions (Dbl.
              Repl.)
 An acid and a base will react and form
  water and a salt.
 Hydrochloric acid is added to sodium
  hydroxide.
     HCl + NaOH  NaCl + H2O
              Hydrolysis (Dbl. Repl.)
   It is the reverse of neutralization and results when a salt plus a water molecule yields
    an acid plus a base.
   Salt + water  acid + base

   Key things to know about hydrolysis reactions:
      Salts of a strong acid plus a weak base will hydrolyze into an acidic solution.
   NH4+ +Cl- +H2O → H+ +Cl- + (NH)4OH

        Salts of a weak acid and a strong base will always hydrolyze to give a basic
         solution.
   K+   +F- +H2O → K+ +OH- +HF

      Salts of a strong acid and a strong base will never undergo hydrolysis and
       therefore make a neutral solution.
   Na+ +Cl- +H2O → Na+ +OH- +H+ Cl-

       Salts of a weak acid plus salts of a weak base may hydrolyze as an acid, base,
        or a neutral solution; the final result depends on the Ka‟s and Kb‟s of the acids
        and bases formed during the hydrolysis process.
   Disclaimer!! The spectator ions were not removed 
 Examples of Dbl. Replacement
 Solutions of potassium bromide and silver
  nitrate are mixed.
      KBr + AgNO3  AgBr + KNO3
 A solution of sodium sulfate is added to a
  solution of hydrochloric acid.
      Na2SO3 + 2HCl  2NaCl + H2SO3
    Hydrolysis Sample Problems
   Try these:
     An aqueous solution of manganese (II) sulfate
      undergoes hydrolysis.
     Ammonium fluoride and water are mixed
      together.
      Hydrolysis answers
 MnSO4 + 2H2O → H2SO4 + Mn(OH)2
 NH4F + H2O → HF + NH4OH
    Combustion (Organic Reacs.)
 An organic compound reacts with O2 to
  form water and carbon dioxide.
 If something is burned there is a
  combustion reaction.
 Methanol is burned in oxygen gas.
      2CH3OH + 3O2  4H2O + 2CO2
    Addition (Organic Reacs.)
 A halogen or hydrogen is added to an
  alkene or alkyne, breaking apart the
  double or triple bonds and forming single
  bonds.
 Fluorine is added to ethene
     F2 + CH2=CH2  CH2F-CH2F
    Substitution (Organic Reacs.)
 An atom attached to a carbon is removed
  and something else takes its place.
 Bromine is added to methane
     Br2 + CH4  CH3Br + HBr
         Oxidizing Agents (Redox
                 Reacs.)
Common Oxidizing Agents                  Products Formed
                                         Mn+²
   MnO4¯ in acidic solution             Mn+²
   MnO2 in acidic solution              MnO2(s)
   MnO4¯ in neutral or basic solution   Cr+³
   Cr2O7ˉ² in acidic solution           NO2
   HNO3, concentrated                   NO
   HNO3, dilute                         SO2
   H2SO4, hot, concentrated             Metallous ions (lower oxidation #)
   Metallic ions (higher oxidation #)   Halide ions
   Free halogens                        NaOH
   Na2O2                                Clˉ
   HClO4                                CO2
   C2O4ˉ²                               O2
   H 2 O2
      Reduction Agents (Redo
              Reacs.)
Common Reducing           Products Formed
  Agents                  Free halogen
 Halide ions
                          Metal ions
 Free metals
                          Sulfate ions
 Sulfite ions or SO2
                          Nitrate ions
 Nitrite ions
                          Hypohalite ions
 Free halogens, dilute
  basic solution          Halite ions
 Free halogens,
                          Metallic ions
  concentrated basic      (higher oxidation #)
  solution
 Metallous ions (lower
  oxidation #)
     Electrolysis (Redox Reacs.)
   An electrolysis reaction is a reaction in which a non-
    spontaneous redox reaction is brought about by the
    passage of current under sufficient external electrical
    potential. The devices in which electrolysis reactions
    occur are called electrolytic cells.
   In theory, E° values (Standard Reduction Potentials) can
    be used to predict which element will plate out at a
    particular electrode when various solutions are
    combined.
   (B&L text)
    Rules for Predicting Cathode
       Reactions (Reduction)
 When a direct electric current is passed
  through a water solution of an electrolyte,
  two possible reduction processes may
  occur at the cathode.
 The cation may be reduced to the
  corresponding metal.
     Mn+ + ne-  M(s) (reaction 1)
       n = (charge of cation)
 Water molecule may be reduced to
  elementary hydrogen
     2H2O + 2eˉ  H2 + 2OHˉ (reaction 2)
      Rules for Predicting Cathode
           Reactions, contd.
   For salts containing transition metal cations,
    which are relatively easy to reduced compared
    to water, reaction #1 will occur at the cathode
    (and the transition metal will plate out).
Mn+ + ne-  M(s)
   If the cation is representative metal, the water
    molecules will be easier to reduce compared to
    the cation, and reaction #2 will occur at the
    cathode, producing hydrogen gas and hydrogen
    ions.
2H2O + 2eˉ  H2 + 2OHˉ
        Rules for Predicting Anode
           Reaction (oxidation)
   The oxidation process that occurs at the anode of an
    electrolytic cell operating in aqueous solution may be
    one of two oxidation processes.
   The anion may be oxidized to the corresponding
    nonmetal.
        - 2Xˉ  X2 + 2eˉ (reaction 1)
   Water molecules may be oxidized to elementary oxygen.
        - HOH  ½ O2 + 2H+ + 2eˉ (reaction 2)
      Rules for Predicting Anode
          Reactions, contd.
   For salts containing iodide, bromide, or chloride
    ions, it is usually easier to oxidize these
    nonmetals rather than water. It will be found that
    the nonmetal is formed at the anode.
   When the anion present is any other ion that is
    more difficult to oxidize than water, Reaction #2
    will occur at the anode producing elementary
    oxygen and aqueous hydrogen ions.
 Example Electrolysis Reactions
1.  Copper (II) chloride in water
     Cu+2 + 2Clˉ  Cu + Cl2
2. Copper (II) sulfate in water
     Cu+2 + HOH  Cu + ½ O2 + 2H+
3. Sodium chloride in water
     2Clˉ + 2HOH  H2 + Cl2 + 2OHˉ
4. Sodium sulfate in water
     2HOH  2H2 + O2
Metals w/ Multiple Oxidation Levels
         (Redox Reacs.)
   These metals can change their oxidation state in a redox reaction
      Antimony (III) or (V)
      Bismuth (III) or (IV)
      Cerium (III) or (IV)
      Chromium (II) or (III)
      Cobalt (II) or (III)
      Copper (I) or (II)
      Gallium (I) or (II) or (III)
      Germanium (II) or (IV)
      Gold (I) or (III)
      Iron (II) or (III)
      Lead (II) or (IV)
      Mercury (I) or (II)
      Nickel (II) or (III)
      Thallium (I) or (III)
      Thorium (II) or (IV)
      Tin (II) or (IV)
   Tin (II) sulfate is added to iron (III) sulfate
                      SnSO4 + Fe2(SO4)3  Sn(SO4)2 + 2FeSO4
          Complex Ion Reactions
   Nomenclature is on pages 23-27 of The Ultimate
    Chemical Equations Handbook
   There are a lot of very complicated types of these
    reactions, but, for all intensive purposes and for the AP
    test, you only need to be familiar with those reactions
    pertaining to ammonia and water.
   In a complex ion reaction, ligands will attach to a
    transition metal ion.
   There will usually be twice as many ligands as the
    metals oxidation number
 Complex Ion Reactions, contd.
 These reactions usually occur in a
  concentrated solution of the ligand.
 Copper chloride (II) is added to a
  concentrated solution of ammonia
       Cu2+ +NH3  [Cu(NH3)4]2+
       Common Reaction Terms
   Electrolysis: Electricity is run through a compound,
    resulting in a change of oxidation states.
   Hydrolysis: The reaction of a salt with water to form
    molecular species. Salts of a strong acid + a weak base
    will always hydrolyze to give an acidic solution.
   Neutralization: Acid and base react to form a salt and
    water.
   Catalyst: A molecule that speeds that speeds a reaction
    but that does not appear in the reaction.
   Oxidation number: the charge that it would have if all the
    ligands (atoms that donate electrons) were removed
    along with the electron pairs that were shared with the
    central atom
    Common Reaction Terms,
           contd.
 Precipitate: an insoluble substance formed
  by the reaction of two aqueous
  substances.
 Anode: the electrode where oxidation
  occurs        an ox
 Cathode: the electrode where reduction
  occurs        red cat
 By: Will Lambert, Adam Robinson,
  Michelle Klassen, and Tori Waldron
 (APChem „06-‟07)

				
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