Solutions of acids and bases

The complete results, up to the addition of 200 mL of alkali, are collected in Table 10.3 this also includes the figures for 0.1 M and 0.01 M solutions of acid and base respectively. The additions of alkali have been extended in all three cases to 200 mL it is evident that the range from 200 to 100 mL and beyond represents the reverse titration of 100 mL of alkali with the acid in the presence of the non-hydrolysed sodium chloride solution. The data in the table are presented graphically in Fig. 10.2. 

Examples through illustrate the two main types of equilibrium calculations as they apply to solutions of acids and bases. Notice that the techniques are the same as those introduced in Chapter 16 and applied to weak acids in Examples and. We can calculate values of equilibrium constants from a knowledge of concentrations at equilibrium (Examples and), and we can calculate equilibrium concentrations from a knowledge of equilibrium constants and initial concentrations (Examples, and ). 

It is a typical feature of aqueous electrolyte solutions that one can, within wide limits, change the solute concentrations and hence the conductivities themselves. Pure water has a very low value of o it is about 5 pS/m at room temperature after careful purification of the water. In the most highly conducting solutions (i.e., concentrated solutions of acids and bases), values of 80 S/m can be attained at the same temperature values seven orders of magnitude higher than those found for pure water. 

When solutions of acids and bases are sufficiently dilute or when other electrolytes are present, the activity, rather the concentration of hydrogen ions, should be substituted in the pH equation. 

Because the concentration of the hydronium ion, HsO+ (or H+ as a shorthand notation), can vary tremendously in solutions of acids and bases, a scale to represent the acidity of a solution was developed. This is the pH scale, which relates to the [H30+] ... 

When using dilute solutions of acids and bases, you can assume their density is close to the density of water. Therefore, you can easily measure the volume of the solutions and calculate their mass. 

In this section, you compared strong and weak acids and bases using your understanding of chemical equilibrium, and you solved problems involving their concentrations and pH. Then you considered the effect on pH of buffer solutions solutions that contain a mixture of acid ions and base ions. In the next section, you will compare pH changes that occur when solutions of acids and bases with different strengths react together. 

Furthermore, the advent of FT NMR allows the routine study of these nuclei, as well as the use of relatively dilute solutions of acids and bases, thus reducing the importance of concentration effects. 

Salts therefore, are prepared (1) from solutions of acids and bases by neutralization and separation by evaporation and crystallization (2) from solutions of two salts by precipitation where the solubility of the salt formed is slight (e.g., silver nitrate solution plus sodium chloride solution yields silver chloride precipitate [almost all as sulid], and sodium nitrate present in solution as sodium cations and nitrate anions [recoverable as sodium nitrate, solid by separation of silver chlondc and subsequent evaporation of the solution]) (3) from fusion of a basic oxide (or its suitable compound—sodium carbonate above) and an acidic oxide (or its suitable compound—ammonium phosphate), since ammonium and hydroxyl are volatilized as ammonia and water. Thus, sodium ammonium hydrogen phosphate... 

The rest of this chapter is a variation on a theme the use of equilibrium constants to calculate the equilibrium composition of solutions of acids and bases. We begin with solutions of acids, bases, and salts, explore the contribution of the autoprotolysis of the solvent to the pH, which is significant in very dilute solutions, and see how to handle the complications of acids that can donate more than one proton. Although the applications are varied, the techniques are all very similar and are based on the material in Chapter 9. 

What is it that makes an acid an acid and a base a base We first raised those questions in Section 4.5, and we now take a closer look at some of the concepts that chemists have developed to describe the chemical behavior of acids and bases. We ll also apply the principles of chemical equilibrium discussed in Chapter 13 to determine the concentrations of the substances present in aqueous solutions of acids and bases. An enormous amount of chemistry can be understood in terms of acid-base reactions, perhaps the most important reaction type in all of chemistry. 

Since a chemist should be familiar with the taste of hydrogen and hydroxyl ions, we have set the very bad precedent here of giving directions to taste the dilute solutions of acids and bases. In no other case should any laboratory chemical be taken in the mouth. 

Hence all aqueous solutions of acids and bases have the same theoretical decomposition voltage when during electrolysis hydrogen and oxygen are liberated. 

Another property that can be used to distinguish acids from bases is their conductivity in solution. As you can see in Figure 10.1, aqueous solutions of acids and bases conduct electricity. This is evidence that ions are present in acidic and basic solutions. Some of these solutions, such as hydrochloric acid and sodium hydroxide (a base), cause the bulb to glow brightly. Most acidic and basic solutions, however, cause the bulb to glow dimly. 

Aqueous solutions of acids and bases can be tested using a conductivity tester. The brightness of the bulb is a clue to the concentration of ions in the solution. Which of these solutions have higher concentrations of ions Which have lower concentrations ... 

In Figure 10.1, you saw evidence that ions are present in solutions of acids and bases. When hydrogen chloride dissolves in water, for example, it dissociates (breaks apart) into hydrogen ions and chloride ions. 

In this chapter we have encountered many different situations involving aqueous solutions of acids and bases, and in the next chapter we will encounter still more. In solving for the equilibrium concentrations in these aqueous solutions, you may be tempted to create a pigeonhole for each possible situation and to memorize the procedures necessary to deal with each particular situation. This approach is just not practical and usually leads to frustration Too many pigeonholes are required, because there seems to be an infinite number of cases. But you can handle any case successfully by taking a systematic, patient, and thoughtful approach. When analyzing an acid-base equilibrium problem, do not ask yourself how a memorized solution can be used to solve the problem. Instead, ask yourself this question What are the major species in the solution, and how does each behave chemically ... 

When the method just described was applied to the study of aqueous solutions of acids and bases, it was found, provided smooth platinum electrodes were used as anode and cathode, that the decomposition voltage was almost constant, irrespective of the nature of the electrolyte. The results obtained in a number of instances are quoted in Table LXXVI it is evident that the decomposition voltages of these solutions... 

The recognition that solutions of acid and bases (and also solutions of salts) contain ions rather than just neutral molecules is due to the Swedish chemist Svante Arrhenius (1859-1927), in 1887. Arrhenius pointed out that the electrical conductivity of these solutions could be understood as due to the motion of the ions through the solution between the electrodes. 

Silicones are resistant to water between 0°C and 100°C. Steam, at higher temperatures in prolonged attack, destroys the system. Normally, silicone rubbers are resistant to salt solution and to dilute solutions of acids and bases. However, they are not resistant to organic solvents. In this case a reversible swelling takes place. 

In dilute solutions of acids and bases and in pure water, the activities of H and OH" may be considered to be the same as their concentrations. 

PROPERTIES OF COMMERCIAL CONCENTRATED SOLUTIONS OF ACIDS AND BASES... 

The litmus in litmus paper is one of the dyes commonly used to distinguish solutions of acids and bases, as shown in Figure 19-1. Aqueous solutions of acids cause blue litmus paper to turn pink. Aqueous solutions of bases cause red litmus paper to turn blue. With this information you can now identify the two groups of household products you used in the DISCOVERY LAB. 

Water not only serves as the solvent in solutions of acids and bases, it also plays a role in the formation of the ions. In aqueous solutions of acids and bases, water sometimes acts as an acid and sometimes as a base. You can think of the self-ionization of water as an example of water assuming the role of an acid and a base in the same reaction. 

Our theme throughout this chapter is to manipulate equilibria to control the solubilities of ionic solids. In the first section we describe general features of the equilibria that govern dissolution and precipitation. In the remaining sections we explore quantitative aspects of these equilibria, including the effects of additional solutes, of acids and bases, and of ligands that can bind to metal ions to form complex ions. 

Percent concentration is the simplest concentration unit. The amount of solute is compared to the amount of solution in order to measure concentration. This concentration unit is generally used for concentrated solutions of acids and bases. The percentage of solute can be expressed by mass or volume. 

Most of the chemicals in a laboratory are toxic some are very toxic, and some—such as concentrated solutions of acids and bases—are highly coito-sive. Avoid contact between these liquids and the skin. In the event of such contact, immediately flood the affected area with copious quantities of water. If a corrosive solution is spilled on clothing, remove the garment immediately. Time is of the essence do not be concerned about modesty. 

The solutions of acids and bases that are sold commercially are too concentrated for most laboratory uses. We often dilute these solutions before we use them. We must know the molar concentration of a stock solution before it is diluted. This can be calculated from the specific gravity and the percentage data given on the label of the bottle. 

Because one mole of an acid does not necessarily neutralize one mole of a base, some chemists prefer a method of expressing concentration other than molarity to retain a one-to-one relationship. Concentrations of solutions of acids and bases are frequently expressed as normality (N). The normality of a solution is defined as the number of equivalent weights, or simply equivalents (eq), of solute per liter of solution. Normality may be represented symbolically as... 

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