This chapter is really a continuation of chapter 3. The same basic steps for solving stoichiometry problems are used: change a quantity to moles, use the mole ratio to change from moles of a given substance to moles of another substance, then change back to another quantity. The real difference is that in chapter 3 the "quantity" was mass; in chapter 4 it is often the volume of a solution of known concentration. But the important point is this: whether the reaction is an acid/base reaction, a gas-phase reaction, or a redox reaction, the procedures for solving stoichiometry problems are basically the same.

Most chemical reactions take place in aqueous solution, so this chapter is quite important. Remember that water is such a good solvent because it is a polar molecule. Figure 4.2 shows how water dissolves an ionic substance by attracting the cations and anions of the salt so strongly that the attractive forces in the ionic crystal are broken. The resulting ions are hydrated, as shown in the figure.

Strong and Weak Electrolytes

Electrolytes are substances whose aqueous solutions conduct electricity. It is the presence of ions which can move through the solution that allows the conduction of the electricity. Sometimes the ions were already there in the ionic salt, and other times they result from the reaction of a substance with the solvent, water. When an ionic substance is broken apart into the solvated cations and anions, the process is properly called dissociation, since the ions, which were already present in the salt, simply separated (dissociated). However, when the ions are formed by the reaction of the substance with water, and didn't exist in the original substance, then the process is properly called ionization, since it involves the formation of ions. You should be aware, however, that many chemists use these two terms interchangeably.

Strong electrolytes are substances which are completely dissociated in water or substances for which the ionization reaction goes to completion. In other words, after the substance has dissolved in water, the original substance no longer exists at any appreciable concentration. The types of substances which are typically classified as strong electrolytes are:

1. soluble salts, e.g. NaCl, MgCl₂.
2. strong inorganic acids (the common ones are: hydrochloric acid, hydrobromic acid, hydroiodic acid, sulfuric acid, nitric acid, chloric acid, and perchloric acid; KNOW THIS LIST!)
3. strong inorganic bases, e.g. NaOH.

Notice that the reaction between NaCl and water is a dissociation reaction, while the reaction between HCl(g) and water is an ionization reaction:

Weak electrolytes only ionize partially in water. In other words, the ionization reaction lies mostly to the left, producing only a small number of ions (relatively speaking) in solution. The types of substances which behave as weak electrolytes are:

1. Weak inorganic acids (inorganic acids which are not listed above), e.g. carbonic acid, nitrous acid, hypochlorous acid, etc.
2. Organic acids (carboxylic acids , e.g. acetic acid)
3. Ammonia, NH₃
4. Organic amines, e.g. methylamine.

Nonelectrolytes do not form ions in solution. Typical nonelectrolytes which will be encountered in this course are sugars and alcohols.

The Composition of Solutions

Although several other ways of expressing the concentration of a solution will be discussed in Chapter 11 (Properties of Solutions), the concentration unit which is most commonly used is presented in this chapter: molarity (abbreviation: M). The molarity of a solution is defined as the number of moles of solute dissolved in a liter of solution (or, equivalently, millimoles of solute per milliliter.) As a result of this definition, one of the most important results for this chapter is

Notice that this relationship is often used to calculate moles of a substance to perform stoichiometric calculations. Just as you divided grams of a substance by the molar mass to get moles in chapter 3 , you'll multiply the volume of a solution by its molarity to get moles in this chapter.

Study Figure 4.10 carefully, but note that there is a mistake. Step (e) should be to mix the solution thoroughly after diluting to the mark. You should be able to interconvert between moles, volume, and molarity, and do problems like sample exercise 4.6.

Dilution

Dilution problems are easy because they rely on a simple principle: when a solution is diluted, more solvent (water) is added to it. But the number of moles of solute is the same in the diluted and concentrated solution. Because of this fact, the following equation applies:

where C and D represent concentrated and dilute, respectively. DO NOT USE THIS EQUATION TO PERFORM STOICHIOMETRIC CALCULATIONS FOR SOLUTIONS. USE IT ONLY FOR DILUTION CALCULATIONS. When you use this equation, the units for volume aren't important so long as they are the same for the concentrated and dilute solutions.

Types of Chemical Reactions

The main types of reactions described in the chapter are precipitation reactions, acid/base reactions, and redox reactions.

In order to predict the products of precipitation reactions, it is necessary to know the solubility rules (Table 4.1). KNOW THESE RULES THOROUGHLY!

When describing reactions in solution, three types of reactions are commonly written:

1. The "molecular" equation lists the whole formula for reactants and products, not taking into account the form of the solute in solution. For example, NaCl doesn't exist as NaCl in water solution; it exists as sodium and chloride ions. Nevertheless, the formula NaCl would be used in a molecular equation. Note that the term "molecular" equation is a bit misleading, since ionic substances don't exist as molecules.
2. The ionic equation shows all strong electrolytes, reactants and products, as separated ions. Thus, NaCl would be shown as Na⁺ and Cl^- ions in an ionic equation. Sucrose (a sugar), however, would be shown as the molecular substance C₁₂H₂₂O₁₁, since it is a nonelectrolyte.
3. The net ionic equation results after common terms are subtracted from the ionic equation. These common terms do not participate in the reaction, and thus are called spectator ions.

Oxidation/Reduction Reactions

This very important class of reactions is based on the idea that formally, an atom can gain (or lose) electrons from (or to) another atom. Oxidation is the loss of electrons and reduction is the gain of electrons. A bookkeeping method for keeping track of electrons is the concept of oxidation number (sometimes called oxidation states). The rules for assigning oxidation states are listed in Table 4.2. You should know this thoroughly. An additional exception to rule 3 is the superoxide ion, O₂^-, in which each oxygen has an oxidation state of -1/2.

Remember that oxidation states are arbitrary charges, not necessarily real charges. The only time that an oxidation number is a real charge is for monatomic ions, like Cl^- (oxidation number -1) or aluminum ion, Al³⁺ (oxidation number +3). In hydrogen chloride, the oxidation numbers of +1 for the hydrogen and -1 for the chlorine are clearly not real; HCl is a covalent compound which doesn't contain ions.

Note that charges on ions are written as n+ or n-, whereas oxidation numbers are written as +n or -n.

Although some simple redox reactions can be balanced by inspection (this was the method used to balance the reactions in chapter 3, as well as the precipitation and acid/base reactions in this chapter), most redox reactions must be balanced systematically. Make sure you understand the half-reaction method of balancing redox reactions. This is described in section 4.10. Practice, practice, and practice is the name of the game for this technique. This is why it is very important that you do all homework problems for this chapter. Remember also that if you can't consistently assign oxidation numbers correctly, you won't be able to balance redox reactions. And without a balanced reaction, of course, you won't be able to perform stoichiometric calculations.

Last modified June 16, 1997