Tables of Acids and Bases

Use Table 4.2 to predict whether the equilibrium for these reactions favors the reactants or the products. 

Sodium hydroxide and sodium ethoxide are used in many reactions where a moderate base will suffice. Carbonate and bicarbonate are often employed to remove acids in the workup of organic reactions. For many of these bases it does not usually matter whether the cation is sodium or potassium. 

Some care must be given to the choice of an acid or a base to use in a reaction. It must be strong enough to do the job but not so strong as to cause undesired reactions at other, less reactive functional groups. 

Finally, remember that pAys more negative than —2 (strong acids) and pA a s greater than 16 (very weak acids) can be measured only indirectly, so the values listed in these tables are only approximate. For this reason, different sources often list somewhat different values for these acidity constants. 

C S V -6.5 Benzenesulfonic acid (note similarity to sulfuric acid) 

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This IS a very useful relationship You should practice writing equations according to the Brpnsted-Lowry definitions of acids and bases and familiarize yourself with Table 1 7 which gives the s of various Br0n sted acids... 

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. 

From Table 10.1, the pKa for formic acid is 3.75. Assuming little change in the concentrations of acid and base due to the deprotonation of HCOOH, we write... 

Using Environmental Examples to Teach About Acids. Acid-base reactions are usually presented to secondary students as examples of aqueous equilibrium (2). In their study of acids and bases, students are expected to master the characteristic properties and reactions. They are taught to test the acidity of solutions, identify familiar acids and label them as strong or weak. The ionic dissociation of water, the pH scale and some common reactions of acids are also included in high school chemistry. All of these topics may be illustrated with examples related to acid deposition (5). A lesson plan is presented in Table I. 

C17-0038. Prepare a table listing the various types of acids and bases and the identifying features of each. 

Neutral molecules show a range of retention properties between those of acids and bases. Progesterone membrane retention is very high in all cases. Griseofulvin and carbamazepine retention steeply increase with phospholipid content. The patterns of retention follow the lipophilicity properties of the molecules, as indicated by octanol-water apparent partition coefficients (Table 7.4). 

As we have seen, the Lewis theory of acid-base interactions based on electron pair donation and acceptance applies to many types of species. As a result, the electronic theory of acids and bases pervades the whole of chemistry. Because the formation of metal complexes represents one type of Lewis acid-base interaction, it was in that area that evidence of the principle that species of similar electronic character interact best was first noted. As early as the 1950s, Ahrland, Chatt, and Davies had classified metals as belonging to class A if they formed more stable complexes with the first element in the periodic group or to class B if they formed more stable complexes with the heavier elements in that group. This means that metals are classified as A or B based on the electronic character of the donor atom they prefer to bond to. The donor strength of the ligands is determined by the stability of the complexes they form with metals. This behavior is summarized in the following table. 

In the case of the four-parameter equation, the enthalpies of interaction of a large number of acids and bases were determined calorimetrically in an inert solvent. With these values being known, a value of 1.00 was assigned for EA and CA for the Lewis acid iodine. The experimental enthalpies for the interaction of iodine with several molecular Lewis bases were fitted to the data to determine EB and CB values for the bases. Values were thus established for the four parameters for many acids and bases so that they can be used in Eq. (9.112) to calculate the enthalpies of the interactions. The agreement of the experimental and calculated enthalpies is excellent in most cases. However, the four-parameter approach is used primarily in conjunction with interactions between molecular species, although extensions of the approach to include interactions between charged species have been made. Table 9.7 gives the parameters for several acids and bases. 

We can prepare a buffer of almost any pH provided we know the pAa of the acid and such values are easily calculated from the Ka values in Table 6.5 and in most books of physical chemistry and Equation (6.50). We first choose a weak acid whose pKa is relatively close to the buffer pH we want. We then need to measure out accurately the volume of acid and base solutions, as dictated by Equation (6.50). 

In this chapter, you learned about solutions and how to use molarity to express the concentration of solutions. You also learned about electrolytes and nonelectrolytes. Using a set of solubility rules allows you to predict whether or not precipitation will occur if two solutions are mixed. You examined the properties of acids and bases and the neutralization reactions that occur between them. You then learned about redox reactions and how to use an activity table to predict redox reactions. You learned about writing net ionic equations. Finally, you learned how to use the technique of titrations to determine the concentration of an acid or base solution. 

As was pointed out in the previous chapter, biologically important metal ions and their ligands can be classified according to the hard-soft theory of acids and bases (Table 2.1). While there are exceptions, most metal ions bind to donor ligands as a function of preferences based on this concept, with hard acids (metal ions) binding preferentially to hard bases (ligands) and soft acids to soft bases. 

According to the hard and soft acids and bases (HSAB) principle, developed by Pearson in 1963232,233, Lewis acids and Lewis bases are divided into two groups hard and soft. Pearson correlated the hardness of acids and bases with their polarizability, whereby soft acids and bases are large and easily polarizable, and vice versa. A selected list of Lewis acids ordered according to their hardness in aqueous solution is presented in Table 18. The HSAB principle predicts strong association of like partners. Hard acid-soft base complexes mainly result from electrostatic interactions, while soft acid-soft base complexes are dominated by covalent interactions. 

Table 8.1 outlines properties of acids and bases that you have examined in previous courses. In this section, you will review two theories that help to explain these and other properties. As well, you will use your understanding of molecular structure to help you understand why acids and bases differ in strength. 

Table 8.1 Examples and Common Properties of Acids and Bases...

Table 7. A set of acids and bases in which E and C both increase in the series...

It was not until the last decade of the nineteenth century that chemists had an adequate theoretical description of acid and bases. Until then, most acids and bases were classified according to their general properties. Chemists knew acids and bases displayed contrary properties, and these were adequate for identifying many acids and bases. Some of these properties are listed in Table 13.1. 

Table 13.2 illustrates the presence of hydrogen in acids. It is also apparent that bases contain hydroxide ions, but the weak base ammonia seems to be an exception. Ammonia illustrates one of the shortcomings of the Arrhenius definition of acids and bases specifically, bases do not have to contain the hydroxide ion to produce hydroxide in aqueous solution. When ammonia dissolves in water, the reaction is represented by ... 

Lewis defined a base as an electron pair donor and an acid as an electron pair acceptor. Lewis electron pair donor was the same as Bronsted-Lowry s proton acceptor, and therefore, was an equivalent way of defining a base. Lewis acids were defined as a substance with an empty valence shell that could accommodate a pair of electrons. This definition broadened the Bronsted-Lowry definition of an acid. The three definitions of acids and bases are summarized in Table 13.3. 

For complete neutralization to take place, the proper amounts of acid and base must be present. The salt formed in the above reaction is NaCl. If the water were evaporated after completing the reaction, we would be left with common table salt. Sodium chloride is just one of hundreds of salts that form during neutralization reactions. While we commonly think of salt, NaCl, as a seasoning for food, in chemistry a salt is any ionic compound containing a metal cation and a nonmetal anion (excluding hydroxide and oxygen). Some examples of salts that result from neutralization reactions include potassium chloride (KCl), calcium fluoride (CaF ), ammonium nitrate (NH NOj), and sodium acetate (NaC2H302). 

Peruse Table 16-1 for a list of common acids and bases, noting that all the acids in the list contain a hydrogen at the beginnings of their formulas and that most of the bases contain a hydroxide. The Arrhenius definition of acids and bases is straightforward and works for many common acids and bases, but it s limited by its narrow definition of bases. 

You no doubt noticed that some of the bases in Table 16-1 don t contain a hydroxide ion, which means that the Arrhenius definition of acids and bases can t apply. When chemists realized that several substances behaved like bases but didn t contain a hydroxide ion, they reluctantly acknowledged that another determination method was needed. Independently proposed by Johannes Bronsted and Thomas Lowry in 1923 and therefore named cifter both of them, the Bronsted-Lowry method for determining acids and bases accounts for those pesky non-hydroxide-containing bases. 

Substances typical of acids and bases are, respectively, HCl and NaOH. Hydrogen chloride dissolves in water with practically complete dissociation into hydrated protons and hydrated chloride ions. Sodium hydroxide dissolves in water to give a solution containing hydrated sodium ions and hydrated hydroxide ions. Table 3.6 gives values of the mean ionic activity coefficients, y , at different concentrations and indicates the pH values and those expected if the activity coefficients are assumed to be unity. 

Electrical Conductance of Aqueous Solutions of Ammonia and Metal Hydroxides. Check the electrical conductance of 1 W solutions of sodium hydroxide, potassium hydroxide, and ammonia. Record the ammeter readings. Arrange the studied alkalies in a series according to their activity. Acquaint yourself with the degree of dissociation and the dissociation constants of acids and bases (see Appendix 1, Tables 9 and 10). Why is the term apparent degree of dissociation used to characterize the dissociation of strong electrolytes ... 

A selection of cases in which 2 has been found to be particularly efficacious is given in the Table. Additional examples are cited in references 1 and 5. Particularly noteworthy examples include the oxidation of acid- and base-sensitive systems, systems containing sulfur and selenium, and 1,3-diols to 1,3-dicarbonyl compounds. Use of chromium reagents in these latter cases often leads to fragmentation products. 

The (a) md IM symboBsm is arbitrary. Sometimes the symbols A and B are used. Neither should be confused wiih the A and B subgroups of the periodic table or A and B as generic representations of acids and bases. 

For preparation of acid and base stock solutions, see Tables A.2A.1 and A.2A.2 as well as individual recipes. 

We can use the information in Tables 10.1 and 10.2 to determine the relative strengths of acids and bases in solution. For example, suppose we need to decide which member of (a) HF or HI03 (b) NOz or CN is the stronger acid or base in water. We need to know that the higher the K, of a weak acid, the stronger the acid and the weaker its conjugate base. Similarly, the... 

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