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Writing Chemical Equations Worksheet - DOC

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Writing Chemical Equations Worksheet - DOC Powered By Docstoc
					C. Morgan
AP Chem
May 2003


                                           AP TOPIC: Equation Writing
    Introduction
Question 4 in Section B, Part II of the AP Chemistry examination requires the candidates to write five chemical equations from a
choice of eight sentences that describe chemical reactions in words.
In general, it is worth noting the following points.
      Learn inorganic nomenclature rules. Without knowledge of these rules this question is almost impossible.
      Where appropriate all compounds that produce ions in solution should be written in their ionic form and spectator ions
         should be ignored, i.e. where appropriate, write the net-ionic equation.
      There are no reactions in the question that do not occur.
      There is no requirement to balance the equations you write, however, it may be helpful to do so.
      It is possible to get partial credit by writing reactants carefully and guessing products.

Double Replacement (Metathesis).
These reactions involve two reactants forming two products. They usually involve acids, bases and salts.
     Acids

    Formulae begin with H and have a hydrogen ion (H +) that they donate in reactions.

    Organic acids often have ionizable hydrogen ions written at the end of their formulae and include the -COOH or -CO2H
    group. E.g. ethanoic acid CH3COOH. Organic acids are weak.

    Learn the strong acids. HCl, HBr, HI and acids where the number of oxygen’s present exceeds the number of hydrogen’s by
    two or more, e.g. HNO3, H2SO4.

    Sulfuric acid, H2SO4, is more accurately represented as only losing one H + ion, i.e. as H+ + HSO4-, rather than 2H+ + SO42-,
    since the second dissociation of the HSO4- ions is only very slight in comparison to the loss of one H + from H2SO4 to form H+
    and HSO4-.

    Carbonic acid, H2CO3, is weak not very stable and is better represented as CO2 + H2O. If it were to be represented as an acid
    it should be shown as a monoprotic acid, forming H + and HCO3-.

    Sulfurous acid, H2SO3, is generally thought of as being a monoprotic acid forming H + and HSO3-, or alternatively, if a
    product, as a weak acid and not dissociating at all.

     Bases
    Formulae end in OH (hydroxide ions, OH- present) except for ammonia and organic bases that contain Nitrogen that has a
    lone pair of electrons allowing it to act as a base. Do not confuse organic alcohols that also have formulae ending in OH as
    hydroxide ions. An alcohol can be recognized as having a carbon chain preceding the –OH group, e.g. C2H5OH or
    CH3CH2CH2OH, for ethanol and propan-1-ol respectively.

    Learn the strong bases. Confined to Group I & II hydroxides and ammonium hydroxide, i.e. LiOH, NaOH, KOH, Ca(OH)2,
    Sr(OH)2, Ba(OH)2 and NH4OH.
    Ammonium hydroxide, NH4OH, is not very stable and is better represented as NH3 + H2O.

     Salts
    A salt is a compound where the hydrogen ion(s) in an acid have been replaced by metal ions or the ammonium (NH 4+) ion.
    Learn the solubility rules for salts.

    ACID + BASE  SALT + WATER (NEUTRALIZATION)
    ACID + CARBONATE  SALT + WATER + CARBON DIOXIDE
    ACID + METAL  SALT + HYDROGEN

    Equimolar means the reactants react in a 1:1 ratio and is often a clue to only some of the ionizable hydrogen’s in an acid
    being replaced.


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    If one of the products in a reaction is a salt that is thought of as being slightly soluble (commonly Ca(OH) 2, Sr(OH)2 and
    Ba(OH)2) credit may be given for writing either the dissociated or undissociated form.

Examples.

Solutions of sodium phosphate and calcium chloride are                A dilute solution of lithium hydroxide is mixed with dilute
mixed.                                                                hydrobromic acid.
Ca2+ + PO43-  Ca3(PO4)2                                              H+ + OH-  H2O

                                                                      Hydrogen sulfide gas is bubbled through a solution of silver
                                                                      nitrate.
                                                                      H2S + Ag+  H+ + Ag2S

SIMPLE REDOX (a simple change in oxidation state).
A reaction where electrons are transferred from one species to another. These reactions fall into a number of different sub-
categories.
     Single displacement
    Reactions where an element reacts with a compound. Common examples include a more reactive metal displacing another
    less reactive one from solution (See Standard Electrode potential table for clues), halogen displacement and hydrogen
    displacement from both acids and water.

     Combustion
    A reaction with oxygen. Common examples are organic compounds (hydrocarbons and carbohydrates) forming CO 2 and H2O
    and metals burning in air to produce the corresponding oxide.

     Combination
    A reaction where elements and/or compounds combine together. Usually fall into one of four types.

    Element (metal) + Element (non-metal)  compound
    Non-metal oxide + water  acid
    Metal oxide + water  base
    Metal oxide + non metal oxide  salt

     Decomposition
    The reverse of combination, where the compound decomposes to give elements, or a combination of elements and new
    compounds Clue, only one reactant usually being heated.

    Base  metal oxide + water
    Acid  non-metal oxide + water
    Salt containing oxygen (often carbonate)  metal oxide + non-metal oxide (often CO2)

Examples.

Magnesium metal is added to a solution of Iron (III)
chloride                                                              Sulfur trioxide is added to water
Mg + Fe3+  Fe2+ + Mg2+                                               SO3 + H2O  H+ + HSO4-

                                                                      Solid magnesium carbonate is heated
                                                                      MgCO3  MgO + CO2

NON-SIMPLE REDOX (a non-simple change in oxidation state).
In these reactions an oxidizing agent will cause the oxidation of another compound and in the process, itself will be reduced. The
opposite is true of reducing agents.
      If a reaction takes place in acid solution it means H + ions are reactants and water will be one of the products. Water is
          often formed when the oxidizing agent contains oxygen atoms.

       If a reaction takes place in basic solution it means OH- ions are a reactant and water is one of the products.

       Elements in their highest oxidation states can only be reduced and elements in their lowest oxidation states can only be
        oxidized.

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         Learn as many of the common oxidizing agents and reducing agents (and what they turn into) as possible.

         Disproportionation. Simultaneous oxidation and reduction of one species.
          E.g. 3NO2 + H2O  2H+ + 2NO3- + NO

                                                        Change & #
                                                                         Reducing                                   Change & # of
           Oxidizing agent         Turns into           of electrons                           Turns into
                                                                          agent                                     electrons lost
                                                           gained
                  MnO4-                                   +7  +2         Halide ions          Free Halogens                -1  0
                                        Mn2+
                 (in acid)                                  (+5)        (F-, Cl-, Br-, I-)    (F2, Cl2, Br2, I2)            (-1) x 2

                  MnO4-                                   +7  +4      Free metals (e.g.                               0  Various
                                        MnO2                                                    Metal cations
           (in neutral or basic)                            (+3)             Zn)                                        (- Various)

                  MnO2                                    +4  +2                                                          +4  +6
                                        Mn2+                             SO32- or SO2               SO42-
                 (in acid)                                  (+2)                                                             (-2)

                  Cr2O72-                                 +6  +3                                                          +3  +5
                                         Cr3+                                NO2-                   NO3-
                 (in acid)                                (+3) x 2                                                           (-2)

                  HNO3                                    +5  +4        Free Halogens         Hypohalite ions              0  +1
                                         NO2
                 (Conc.)                                    (+1)        (in dilute basic)    (XO-) + halide (X-)            (-1) x 2

                  HNO3                                    +5  +2        Free Halogens                                      0  +5
                                         NO                                                  Halite ions (XO3-)
                 (Dilute)                                   (+3)        (in Conc. basic)                                    (-5) x 2
                                                                        Lower charged
                  H2SO4                                   +6  +4                              Highly charged
                                         SO2                           metal cations (e.g                                   Various
              (hot & Conc.)                                 (+2)                                metal cations
                                                                          Sn2+, Fe2+)
           Highly charged metal    Lower charged                              H2O2                                          -1  0
                                                          Various                                 O2 + H2O
                  cations           metal cations                      (in basic solution)                                  (-1) x 2

              Free Halogens          Halide ions           0  -1                                                          +2  +2.5
                                                                             S2O32-                S4O62-
             (F2, Cl2, Br2, I2)    (F-, Cl-, Br-, I-)     (+1) x 2                                                          (-0.5) x 2

                   H2O2                                   -1  -2                             H+ (usually in the            0  +1
                                         H2O                                   H2
                 (in acid)                                (+1) x 2                             form of water)               (-1) x 2


Examples.
Sulfur dioxide gas is bubbled through a solution of acidified             Hydrogen peroxide is added to an acidified solution of
potassium manganate (VII) (permanganate)                                  potassium dichromate (VI)
SO2 + MnO4- + H+  H2O + SO42- + Mn2+                                     H2O2 + H+ + Cr2O72-  Cr3+ + O2 + H2O

                                                                          Concentrated hydrochloric acid is added to solid manganese
                                                                          (IV) oxide
                                                                          H+ + Cl- + MnO2  Mn2+ + Cl2 + H2O

HYDROLYSIS (reaction with water)
A reaction with water.

         Anion of weak acid + water  weak acid + hydroxide ion

         Cation of weak base + water  weak base + hydrogen ion

         Non-metal halide + water  (H+ ion + more electronegative element) + other elements combined

        PCl3 + H2O  HCl + H3PO3
        PCl5 + H2O  HCl + H3PO4
        SCl4 + H2O  HCl + H2SO3

        The oxidation number of the phosphorous and sulfur in each case does not change

         Hydrides of group I release hydrogen gas

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Examples.

Lithium hydride is added to water                                    Phosphorous tribromide is added to water
LiH + H2O  Li+ + OH- + H2                                           PBr3 + H2O  H+ + Br- + H3PO3
                                                                     Gaseous hydrogen chloride is bubbled into water
                                                                     HCl + H2O  H3O+ + Cl-

ORGANIC
Reactions where organic compounds (those involving carbon chains) react in many different ways.
    Combustion or oxidation. Products are usually CO2 + H2O

        Addition. Compounds with multiple bonds have a halogen, hydrogen halide or hydrogen added

        Substitution. Commonly hydrogen atoms are displaced (replaced) by another atom

        Esterification. Carboxylic acid + alcohol  ester + water

        Learn rules for nomenclature, learn functional groups

Examples.

Methane is mixed with an excess of bromine gas
CH4 + Br2  CBr4 + HBr                                               Ethanoic acid is refluxed with ethanol and a catalyst for
                                                                     several hours
Propanol is burned completely in air                                 CH3COOH + C2H5OH  CH3CO2C2H5 + H2O
CH3CH2CH2OH + O2  CO2 + H2O

COMPLEX ION FORMATION
These reactions involve a transition metal in aqueous solution reacting with (commonly) ammonia ligands or hydroxide ligands.
The breaking down of complexes can be achieved by reacting the complex ion with acid.

        Aluminum and zinc will act like transition metal ions and form complexes

        The word excess often appears in these questions

        Learn nomenclature of complexes and common reactions

        Most first row transition metal complexes exist in solution as hexaaqua ions

Examples.

A solution of ammonium thiocyanate is added to a solution of iron (III) chloride
[Fe(H2O)6]3+ + SCN-  [Fe(SCN)(H2O)5]2+

Excess dilute nitric acid is added to a solution containing tetraamminecadmium (II) ions
H+ + [Cd(NH3)4]2+  Cd2+ + NH4+

Aluminum pellets are treated with excess sodium hydroxide
Al + OH-  Al(OH)4-




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ion of iron (III) chlo ride
[Fe(H2 O)6 ]3+ + SCN-  [Fe(SCN)(H2 O)5 ]2+

Excess dilute nitric acid is added to a solution containing tetraamminecadmiu m (II) ions
H+ + [Cd(NH3 )4 ]2+  Cd 2+ + NH4 +

Aluminu m pellets are treated with excess sodium hydro xide
Al + OH-  Al(OH)4 -




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