Solutions of Acids or Bases Containing a Common Ion

Before we consider the other types of aqueous equilibria, we will deal with acid-base equilibria in more detail. 

In Chapter 7 we were concerned with calculating the equilibrium concentrations of species (particularly H+ ions) in solutions containing an acid or a base. In this section we discuss solutions that contain not only the weak acid HA but also its salt NaA. Although this case appears to be a new type of problem, it can be handled rather easily by using the procedures developed in Chapter 7. 

Suppose we have a solution containing the weak acid hydrofluoric acid (HF, Ka = 7.2 X 10-4) and its salt sodium fluoride (NaF). Recall that when a salt dissolves in water, it breaks up completely into its ions—it is a strong electrolyte  

Since hydrofluoric acid is a weak acid and only slightly dissociated, the major species in the solution are HF, Na+, F , and H20. The common ion in this solution is F, since it is produced by both hydrofluoric acid and sodium fluoride. What effect does the presence of the dissolved sodium fluoride have on the dissociation equilibrium of hydrofluoric acid  

To answer this question we compare the extent of dissociation of hydrofluoric acid in two different solutions, the first containing 1.0 M HF and the second containing 1.0 M HF and 1.0 M NaF. According to Le Chatelier s principle the dissociation equilibrium for HF 

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Solutions of Acids or Bases Containing a Common Ion Equilibrium Calculations... 

Buffered solutions are simply solutions of weak acids or bases containing a common ion. The pH calculations for buffered solutions require exactly the same procedures previously introduced in Chapter 7. This is not a new type of problem. 

The procedures for finding the pFI of a solution containing a weak acid or base plus a common ion are very similar to the procedures, which we covered in Chapter 14, for solutions containing the acids or bases alone. For example, in the case of a weak acid, the only important difference is that the initial concentration of the anion is not zero in a solution that also contains the salt NaA. Example 15.1 illustrates a typical example using the same general approach we developed in Chapter 14. 

The most important application of acid-base solutions containing a common ion is buffering. Thus, a buffer solution will marntam a relatively constant pH even when acidic or basic solutions are added to it. The most important practical example of a buffered solution is human blood, which can absorb the acids and bases produced by biological reactions without changing its pH. The normal pH of human blood is 7.4. A constant pH for blood is vital, because cells can only survive this narrow pH range around 7.4. 

Simple general chemistry principles can be used to help maximize precipitation and product yield. The first is the common ion effect, formally defined as making a weak acid or weak base weaker by the addition of a salt that contains a common ion, a direct consequence of Le Chatlier s principle. The common ion effect can also be used to understand metal complex precipitation. Consider a saturated solution of [Co(NH3)6]C13, equation (1.20). 

The most important appiication of acid-base soiutions containing a common ion is for buffering. A buffered solution is one that resists a change in itsp/f when either hydroxide ions or protons are added. The most important practicai exampie of a buffered soiu-tion is our biood, which can absorb the acids and bases produced in bioiogic reactions without changing its pH. A constant pH for biood is vitai because ceiis can survive oniy in a very narrow pH range. 

The Arrhenius theory accounts for the properties of many common acids and bases, but it has important limitations. For one thing, the Arrhenius theory is restricted to aqueous solutions for another, it doesn t account for the basicity of substances like ammonia (NH3) that don t contain OH groups. In 1923, a more general theory of acids and bases was proposed independently by the Danish chemist Johannes Bronsted and the English chemist Thomas Lowry. According to the Bronsted-Lowry theory, an acid is any substance (molecule or ion) that can transfer a proton (H + ion) to another substance, and a base is any substance that can accept a proton. In short, acids are proton donors, bases are proton acceptors, and acid-base reactions are proton-transfer reactions ... 

Our discussion of acid-base ionization and salt hydrolysis in Chapter 15 was limited to solutions containing a single solute, hi this section we will consider the acid-base properties of a solution with two dissolved solutes that contain the same ion (cation or anion), called the common ion. 

Let us consider the interesting case of solutions which contain weak electrolytes (acids or bases), and in addition contain a salt having a common ion. For example, a solution might be made up to the following specifications ... 

In Chapter 16, we examined the equilibrium concentrations of ions in solutions containing a weak acid or a weak base. We now consider solutions that contain a weak acid, such as acetic acid (CH3COOH), and a soluble salt of that acid, such as sodium acetate (CH3COONa). Notice that these solutions contain two substances that share a common ion, CH3COO . It is instructive to view these solutions from the perspective of Le Chatelier s principle. 000 (Section 15.7)... 

Weak acids or bases undergo only partial ionization, so that their solutions contain intact molecules as well as dissociated ions. When we write equations for the dissolution of weak electrolytes like these, we use a two-way arrow that emphasizes that the reaction does not proceed completely from left to right. For common weak acids, it is relatively easy to write the needed ionization equation. Many weak acids contain the —COOH functional group, and the H atom from that group tends to be lost in solution. Acetic acid is a good example. 

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