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Rules for Balancing Equations

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					                                  Rules for Balancing Equations

1.   If formulas are not given, write formulas of compounds correctly. Do not change the subscripts
     of the compounds once they are written correctly.
2.   Start with the largest subscript of the elements (Do not count the subscript of a polyatomic ion).
     Balance oxygen last and hydrogen next to last.
3.   Place a coefficient in front of the compound to balance subscripts. Remember – do not change
     subscripts.
4.   If a polyatomic ion remains in tact on both sides of the arrow, count it as one unit. You do not
     have to separate it into its elements.
5.   Count the number of atoms of each element on the reactant side and the product side of the
     equation to be sure they equal. If they do not equal, start over.

                                       Predicting Products

1.   Write formulas for reactants and then decide the type of reaction.
2.   Follow the rules for types of reactions to predict the products.
         a. Synthesis: Combine elements (check charges if metal is involved).
         b. Single Replacement: Element replaces an element in the compound. Balance charges.
         c. Double Replacement: Cations (metals) trade places. Balance charges.
         d. Decomposition: Follow the decompositions rules.
         e. Combustion: Carbohydrates combust in the presence of oxygen to form CO 2 and H2O.
3.   Do not bring subscripts across the arrow unless they are part of a polyatomic ion.
4.   Make certain all compounds are written correctly. Did you criss-cross the charges in the products.
5.   Now, add coefficients to balance the equation.

                                     Special Synthesis Reactions

1.   Nonmetallic oxide and water react to product an acid. Ex: CO2 + H2O  H2CO3
2.   Metallic oxides react with water to produce a base, a metallic hydroxide. Ex: CaO + H 2O 
     Ca(OH)2

                                       Decomposition Reactions

1.   Metallic carbonates, when heated, form metallic oxides and CO2.
     Ex: CaCO3  CaO + CO2
2.   Many metallic hydroxides, when heated, decompose into metallic oxides and water.
     Ex: Ca(OH)2  CaO + H2O
3.   Metallic chlorates, when heated, decompose into metallic chlorides and oxygen.
     Ex: 2KClO3  2 KCl + 3O2
4.   Some acids, when heated, decompose into nonmetallic oxides and water.
     Ex: H2SO3  H2) + SO2
5.   Some oxides, when heated, decompose.
     Ex: HgO  Hg + O2                          Ex: PbO2  PbO + O2
6.   Some decompositions reactions are produced by an electric current.
     Ex: H20  H2 + O2
                                                   Note

     Remember the diatomic elements: H2, N2, F2, O2, Cl2, Br2, I2,

				
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