Solutes weak bases

Strong acids (the acids listed in Table J.l) are completely deprotonated in solution weak acids (most other acids) are not. Strong bases (the metal oxides and hydroxides listed in Table J.l) are completely protonated in solution. Weak bases (ammonia and its organic derivatives, the amines) are only partially protonated in solution. 

The greater the acidity of a solution, the lower its pH. Weak acids partially ionize to release a hydrogen ion, thus lowering the pH of the aqueous solution. Weak bases accept a hydrogen ion, increasing the pH. The extent of these processes is characteristic of each particular weak acid or base and is expressed as a disso-... 

The sodium ethanoate which is largely dissociated, serves as a source of ethanoate ions, which combine with any hydrogen ions which may be added to the solution to yield more of the acid. The addition of hydrogen ions has therefore much less effect on such a solution than it would have on water. In a similar manner, the solution of the salt of a strong acid and a weak base, in the presence of a weak base, has a pH that is insensitive to additions of alkali. 

Aqueous ammonia can also behave as a weak base giving hydroxide ions in solution. However, addition of aqueous ammonia to a solution of a cation which normally forms an insoluble hydroxide may not always precipitate the latter, because (a) the ammonia may form a complex ammine with the cation and (b) because the concentration of hydroxide ions available in aqueous ammonia may be insufficient to exceed the solubility product of the cation hydroxide. Effects (a) and (b) may operate simultaneously. The hydroxyl ion concentration of aqueous ammonia can be further reduced by the addition of ammonium chloride hence this mixture can be used to precipitate the hydroxides of, for example, aluminium and chrom-ium(III) but not nickel(II) or cobalt(II). 

Conversion of the salt of a weak base into the free base. Prepare a column of a strong base anion resin (such as Amberlite IRA-40o(OH) ) washed with distilled water as above. Drain off most of the water and then allow 100 ml. of A//2.Na.2C03 solution to pass through the column at 5 ml. per minute. Again wash the column with 200 ml. of distilled water. Dissolve 0-05 g. of aniline hydrochloride in 100 ml. of distilled water and pass the solution down the column. The effluent contains aniline in solution and free from all other ions. 

On the other hand, the two nitro groups make 2,4-dinitrophenylhydrazine a very weak base, and it has therefore to be used in reasonably concentrated acid solution. 

Aniline hydrochloride, being a salt formed from a very weak base and a strong acid, undergoes considerable hydrolysis in aqueous solution to aniline... 

Adenine is a weak base Which one of the three nitrogens designated by arrows in the struc tural formula shown is protonated in acidic solution" A resonance evaluation of the three protonated forms will tell you which one is the most stable... 

Strong and Weak Bases Just as the acidity of an aqueous solution is a measure of the concentration of the hydronium ion, H3O+, the basicity of an aqueous solution is a measure of the concentration of the hydroxide ion, OH . The most common example of a strong base is an alkali metal hydroxide, such as sodium hydroxide, which completely dissociates to produce the hydroxide ion. 

Besides equilibrium constant equations, two other types of equations are used in the systematic approach to solving equilibrium problems. The first of these is a mass balance equation, which is simply a statement of the conservation of matter. In a solution of a monoprotic weak acid, for example, the combined concentrations of the conjugate weak acid, HA, and the conjugate weak base, A , must equal the weak acid s initial concentration, Cha- ... 

Adding as little as 0.1 mb of concentrated HCl to a liter of H2O shifts the pH from 7.0 to 3.0. The same addition of HCl to a liter solution that is 0.1 M in both a weak acid and its conjugate weak base, however, results in only a negligible change in pH. Such solutions are called buffers, and their buffering action is a consequence of the relationship between pH and the relative concentrations of the conjugate weak acid/weak base pair. 

A solution containing a conjugate weak acid/weak base pair that is resistant to a change in pH when a strong acid or strong base is added. 

The same approach can be used to derive an equation for the distribution ratio when the solute is a molecular weak base, B, (Figure 7.24). The resulting distribution ratio is... 

Another important parameter that may affect a precipitate s solubility is the pH of the solution in which the precipitate forms. For example, hydroxide precipitates, such as Fe(OH)3, are more soluble at lower pH levels at which the concentration of OH is small. The effect of pH on solubility is not limited to hydroxide precipitates, but also affects precipitates containing basic or acidic ions. The solubility of Ca3(P04)2 is pH-dependent because phosphate is a weak base. The following four reactions, therefore, govern the solubility of Ca3(P04)2. 

Any solution containing comparable amounts of a weak acid, HA, and its conjugate weak base, A-, is a buffer. As we learned in Chapter 6, we can calculate the pH of a buffer using the Henderson-Hasselbalch equation. 

At the equivalence point, the moles of acetic acid initially present and the moles of NaOH added are identical. Since their reaction effectively proceeds to completion, the predominate ion in solution is CH3COO-, which is a weak base. To calculate the pH we first determine the concentration of CH3COO-. 

It has been shown that for most acid-base titrations the inflection point, which corresponds to the greatest slope in the titration curve, very nearly coincides with the equivalence point. The inflection point actually precedes the equivalence point, with the error approaching 0.1% for weak acids or weak bases with dissociation constants smaller than 10 , or for very dilute solutions. Equivalence points determined in this fashion are indicated on the titration curves in figure 9.8. 

The majority of titrations involving basic analytes, whether conducted in aqueous or nonaqueous solvents, use HCl, HCIO4, or H2SO4 as the titrant. Solutions of these titrants are usually prepared by diluting a commercially available concentrated stock solution and are stable for extended periods of time. Since the concentrations of concentrated acids are known only approximately,the titrant s concentration is determined by standardizing against one of the primary standard weak bases listed in Table 9.7. 

In this case, the components are mixed, the pH adjusted to about 6.0 with sodium hydroxide, and the solution appHed to the textile via a pad-dry-cure treatment. The combination of urea and formaldehyde given off from the THPC further strengthens the polymer and causes a limited amount of cross-linking to the fabric. The Na2HP04 not only acts as a catalyst, but also as an additional buffer for the system. Other weak bases also have been found to be effective. The presence of urea in any flame-retardant finish tends to reduce the amount of formaldehyde released during finishing. 

The absorption of sulfur dioxide in alkaline (even weakly alkaline) aqueous solutions affords sulfites, bisulfites, and metabisulfites. The chemistry of the interaction of sulfur dioxide with alkaline substances, either in solution, slurry, or soHd form, is also of great technological importance in connection with air pollution control and sulfur recovery (25,227,235—241). Even weak bases such as 2inc oxide absorb sulfur dioxide. A slurry of 2inc oxide in a smelter can be used to remove sulfur dioxide and the resultant product can be recycled to the roaster (242). 

---