Chapter 3: Stoichiometry: Calculations with Chemical Formulas and Equations

In Chapter 3 we are going dwell on the application of the law of conservation of mass by studying the quantitative relationship which exists between the reactants and products in chemical reactions. To communicate such quantitative information about chemical reactions, we will employ chemical equations as a symbolic way of representing a chemical reaction. The study of the quantitative relationships in chemical reactions is called stoichiometry.

We will look at six kinds of chemical reactions. You will need to able to recognize the different types of reactions and to write the correct products give the reactants. We will see examples of the different types of reactions in class. You must know these examples for the examination. You will also be expected to extend the information you receive in class to other examples.

The six different reaction types are;

1. Formation
   • A reaction between elements in their standard state to form a single product in its standard state. Such reactions may require heat or other form of energy to occur.
   • Fe_(s) + S_(s) -(heat)--> FeS_(s)
   • Al_(s) + Br_2(s) ---> Al₂Br_6(s)
2. Combustion
   • A reaction between a substance and oxygen. The common example is the combustion of a hydrocarbon. A hydrocarbon is a covalent (organic) compound which only contains carbon and hydrogen. The combustion of a hydrocarbon produces carbon dioxide (CO₂) and water (H₂O)
   • CH_4(g) + O_2(g) ---> CO_2(g) + H₂O_(g)
3. Decomposition
   • A reaction of a single substance to produce two or more new substances. Such reactions may require heat or other form of energy to occur.
   • NaN_3(s) ---> Na_(s) + N_2(g)
4. Single replacement
   • A reaction in which an element in a reactant compound is replaced by a second reacting element producing a new compound and an element.
   • Na_(s) + H₂O_(l) ---> NaOH_(aq) + H_2(g)
5. Double replacement
   • A reaction between two ionic compounds producing two new ionic compounds where the cations and anions have exchanged 'partners'.
   • NaCl_(aq) + AgNO_3(aq) ---> NaNO_3(aq) + AgCl_(s)
6. Neutralization
   • A reaction between an acid and a base to produce a salt and water.
   • HCl_(aq) + NaOH_(aq) ---> NaCl_(aq) + H₂O_(g)

Important Acids

Name of Acid	Formula of Acid
Hydrochloric Acid	HCl
Sulfuric Acid	H2SO4
Nitric Acid	HNO3
Phosphoric Acid	H3PO4

Important Bases

Name of Base	Formula of Base
Sodium hydroxide	NaOH
Potassium hydroxide	KOH
Barium hydroxide	Ba(OH)2
Ammonia	NH3

How to balance equations

One very important aspect of any chemical equation is that it be balanced. While the process of balancing chemical equations is straight forward, the real challenge is in the laboratory trying to identify the products of a chemical reaction. In this course our goal will generally be to simply balance a chemical equation. The primary rule that must be remembered when balancing equations is not to change subscripts in the formulas. Balancing equations is accomplished by changing coefficients. Coefficients are whole numbers immediately preceding the chemical formula for the substance in the equation. Never change a subscript or add or remove elements or compounds in an equation.

__Fe(s) + __S₈(s) ----> __FeS(s)

To balance this equation, that is, obtain equals numbers of atoms of each element in the products and the reactants, can be accomplished by placing a coefficient of 8 in front of iron(II) sulfide, and than a 8 in front of iron. Now the equation is balanced.

There are some simple rules for balancing equations:

• Things not to do when balancing an equation:
  1. Change a subscripted number
  2. Add or remove an element arbitrarily
  3. Add or remove a compound arbitrarily
• Things to do when balancing an equation:
  1. Change only the coefficients preceding an element or compound.
  2. Try to balance non-oxygen and non-hydrogen elements first.
  3. Balance substances in their elemental form last.

Previous page

Next page