Balancing Chemical Equations (Chapter 8)

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Equations MUST be balanced before performing any quantitative calculations.

Rules and Suggestions for Balancing Equations

1)     The same # and type of atom must be present on each side of the equation.

2)     Balancing is accomplished by adding coefficients.  NEVER change the subscripts.

3)     Coefficients must be in the smallest whole # ratio.

4)     Balancing is done by trial and error.

5)     Balance H’s and O’s last.

6)      Balance polyatomic ions as one unit.

Rules and Suggestions from Physical Science Text

Balancing chemical equations is mostly trial-and-error procedure. The key to success at balancing equations is to think it out one step-by-step while remembering the following:

1)     Atoms are neither lost nor gained nor do they change their identity in a chemical reaction. The same kind and number of atoms in the reactants must appear in the products, meaning atoms are conserved.

2)     A correct formula of a compound can not be changed by altering or placement of subscripts. Changing subscripts changes the identity of a compound and the meaning of the entire equation.

3)     A coefficient in front of a formula multiplies everything in the formula by that number.

There are also a few generalizations that can be helpful for success in balancing equations:

1. Use a one as a coefficient when you balance a particular atom instead of not writing a coefficient so that you will remember that the compound has a one in front of it.
2. Look first to formulas of compounds with the most atoms and try to balance the or compounds they were formed from or decomposed to.
3. Polyatomic ions that appear on both sides of the equation should be treated as independent units with a charge. That is, consider the polyatomic ions as a unit while taking an inventory rather than the individual atoms making up the polyatomic ion. This will save time, frustration, and simplify the procedure.

Try the following:

P4         +         O2          à              P4O10

S8         +          O2        à            SO2

Ca        +              O2          à          CaO

Fe          +             O2           à          Fe2O3

Combustion Reactions – Burning a fuel in oxygen – produces heat and light.

Example:   Alkane +    O2   à CO2    +     H2O   (products of complete combustion.)

C3H8            +       O2             à          CO2          +            H2O

C4H10            +       O2             à        CO2          +            H2O

Balancing Ionic Equations

CaCrO4            +             NaCl             à                Na2CrO4              +            CaCl2

Na2SO4           +               AlP              à               Na3P                  +              Al2(SO4)3

Mass Relationships in Chemical Reactions:   Stoichiometry (Chap 10)

Balance the following equation.

Show how mass is balanced.

P4              +                 Cl2                  à                   PCl3

Mass of A            à           Moles of A            à       Moles of B       à       Mass of B

(1 mole = Molar Mass in g)           (mole ratio)           (1 mole = Molar Mass in g)

What mass of oxygen, O2, is required to completely combust 454 g of propane, C3H8?  What masses of CO2 and H2O are produced?

C3H8            +       5 O2             à       3 CO2          +            4 H2O