Equilibria Aqueous

Identify the equilibrium relations among the solute species. 

Use the relations in number 2 above to determine the concentrations of the solute species. 

Examples Acetic acid (CH3COOH) and sodium acetate (NaCH3COO)  

Consists of a weak base and its conjugate acid provided as a salt Examples Ammonia (NH3) and ammonium chloride (NH4CI) sodium dihydrogenphosphate (NaH2P04) and sodium hydrogenphosphate (Na2HP04) 

The pH of a buffer solution is given approximately by the Henderson-Hasselbalch equation 

1 (a) When solid sodium acetate is added to an acetic acid solution, the 

NlijCaq) + H,0(1) NH/(aq) + OH (aq) shifts to the left to relieve the stress imposed by the increased [NH ] (Le Chatelier s principle). Because [OH ] decreases, [H3O ] increases and pH decreases. 

(aq) and SO (aq) are conjugate acid and base therefore, the pH calculation is most easily performed with the Henderson-Hasselbalch equation  

After adding HNOj [see part (a) of this exercise]  

15 In a solution containing HClO(aq) and CIO (aq), the following equilibrium occurs  

Solids can dissolve in water in two ways (1) with their molecules intact (e.g., when sugar dissolves in water individual sugar molecules pass from the solid to the liquid phase but the sugar molecule does not break up), and (2) by their molecules breaking up into positively and negatively charged ions. For example, common salt (NaCl) dissolves in the latter way, which can be represented by 

Aqueous solutions containing charged ions are electrically conducting and are called electrolytes. Aqueous ions are individual species, the properties of which are independent of their source. For example, chloride ions from NaCl are just the same as chloride ions from hydrochloric acid (HCl) or any other electrolyte containing chlorine. 

In Chapter 1 we discussed the basic principles of chemical equilibrium. We will now apply these principles to ionic equilibria in aqueous solutions. 

OH (hydroxide) Alkali ions, H, NH4 , Sr +, Ba All others Soluble Low solubility 

As a starting point, let us consider the dissolution of cuprous chloride (CuCl) in water 

Use the relations in number 2 above, along with equilibrium constants, to determine the concentrations of the solute species, often by using an equilibrium table. 

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Why Do We Need to Know This Material The techniques described in this chapter provide rhe tools that we need to analyze and control the concentrations of ions in solution. A great deal of chemistry is carried out in solution, and so this material is fundamental to understanding chemistry. The ionic compounds released into waterways by individuals, industry, and agriculture can impair the quality of our water supplies. However, these hazardous ions can be identified and removed if we add the right reagents. Aqueous equilibria govern the stabilization of the pH in blood, seawater, and other solutions encountered in biology, medicine, and the environment. 

Up to this point, we have focused on aqueous equilibria involving proton transfer. Now we apply the same principles to the equilibrium that exists between a solid salt and its dissolved ions in a saturated solution. We can use the equilibrium constant for the dissolution of a substance to predict the solubility of a salt and to control precipitate formation. These methods are used in the laboratory to separate and analyze mixtures of salts. They also have important practical applications in municipal wastewater treatment, the extraction of minerals from seawater, the formation and loss of bones and teeth, and the global carbon cycle. 

Most aqueous equilibria fall into three broad categories proton transfer, solubility, or complexation. The nature of the major species in the solution determines which category of equilibrium we need to consider. 

The major species in an aqueous solution determine which categories of equilibria are important for that solution. Each major species present in the solution must be examined in light of these general categories. Are any of the major species weak acids or weak bases Are there ions present that combine to form an insoluble salt Do any of the major species participate in more than one equilibrium Any chemical reaction can approach equilibrium from either direction. Consequently, there are six different t q)es of aqueous equilibria in which major species are reactants ... 

The types of aqueous equilibria described in this section have been given special names, and it is essential that you be able to recognize them. Keep in mind, however, that the principles described in the previous sections apply to all chemical equilibria. Chemists categorize equilibria for convenience, but they treat all equilibria the same way. Our Chemistry and the Environment Box explores the roles of these equilibria in a spectacular natural process, the formation of limestone caverns. 

Limestone caverns are among nature s most spectacular displays. These caves occur in many parts of the world. Examples are Carlsbad Caverns In New Mexico, Jeita Caves in Lebanon, the Blue Grotto in Italy, and the Jenolan Caves In Australia. Wherever they occur, the chemistry of their formation involves the aqueous equilibria of limestone, which Is calcium carbonate. Three such equilibria, linked to one another by Le Chatelier s principle, play essential roles In cave dynamics. 

This chapter describes several Important applications of aqueous equilibria. We begin with a discussion of buffer chemistry, followed by a description of acid and base titration reactions. Then we change our focus to examine the solubility equilibria of inorganic salts. The chapter concludes with a discussion of the equilibria of complex Ions. 

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