Weak bases buffered solutions

In a weak base buffer solution, when the concentrations of the weak base and its conjugate acid are equal the pOH of the solution is equal to the pKb of the weak base. 

EXAMPLE 19-3 Weak Base/Salt of Weak Base Buffer Solution... 

Because the ionic product of water = [H ] [OH ] = 1.04 x 10" at 25°C, it follows that pH = 14 - pOH. Thus, a neutral solution (e.g., pure water at 25°C) in which [H j = [OH ] has a pH = pOH = 7. Acids show a lower pH and bases a higher pH than this neutral value of 7. The hydrogen ion concentrations can cover a wide range, from -1 g-ion/liter or more in acidic solutions to -lO" " g-ion/liter or less in alkaline solutions [53, p. 545]. Buffer action refers to the property of a solution in resisting change of pH upon addition of an acid or a base. Buffer solutions usually consist of a mixture of a weak acid and its salt (conjugate base) or of a weak base and its salt (conjugate acid). 

A buffer solution is a solution that resists a change in pH after addition of small amounts of an acid or a base. Buffer solutions require the presence of an acid to neutralize an added base and also the presence of a base to neutralize an added acid. These two components present in the buffer also must not neutralize each other A conjugate acid-base pair is present in buffers to fulfill these requirements. Buffers are prepared by mixing together a weak acid or base and a salt of the acid or base that provides the conjugate. 

Buffer Solutions A buffer solution contains a weak acid and a salt derived from the acid. To maintain a relatively constant pH, the acid and base components of the buffer solution react with added acid or base. Buffer solutions play an important role in many chemical and biological processes. 

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. 

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. 

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. 

Because they are weak acids or bases, the iadicators may affect the pH of the sample, especially ia the case of a poorly buffered solution. Variations in the ionic strength or solvent composition, or both, also can produce large uncertainties in pH measurements, presumably caused by changes in the equihbria of the indicator species. Specific chemical reactions also may occur between solutes in the sample and the indicator species to produce appreciable pH errors. Examples of such interferences include binding of the indicator forms by proteins and colloidal substances and direct reaction with sample components, eg, oxidising agents and heavy-metal ions. 

The characteristics of soluble sihcates relevant to various uses include the pH behavior of solutions, the rate of water loss from films, and dried film strength. The pH values of sihcate solutions are a function of composition and concentration. These solutions are alkaline, being composed of a salt of a strong base and a weak acid. The solutions exhibit up to twice the buffering action of other alkaline chemicals, eg, phosphate. An approximately linear empirical relationship exists between the modulus of sodium sihcate and the maximum solution pH for ratios of 2.0 to 4.0. 

Aqueous solutions of citric acid make excellent buffer systems when partially neutralized because citric acid is a weak acid and has three carboxyl groups, hence three p-K s. At 20°C pifj = 3.14, pi 2 4.77, and = 6.39 (2). The buffer range for citrate solutions is pH 2.5 to 6.5. Buffer systems can be made using a solution of citric acid and sodium citrate or by neutralizing a solution of citric acid with a base such as sodium hydroxide. In Table 4 stock solutions of 0.1 Af (0.33 N) citric acid are combined with 0.1 Af (0.33 N) sodium citrate to make a typical buffer solution. 

Buffers are solutions that tend to resist changes in their pH as acid or base is added. Typically, a buffer system is composed of a weak acid and its conjugate base. A solution of a weak acid that has a pH nearly equal to its by definition contains an amount of the conjugate base nearly equivalent to the weak acid. Note that in this region, the titration curve is relatively flat (Figure 2.15). Addition of H then has little effect because it is absorbed by the following reaction ... 

Textbooks of analytieal ehemistry should be eonsulted for further details eoneeming the ionization of weak aeids and bases and the theory of indieators, buffer solutions, and aeid-alkali titrations. " ... 

The concentration of the acid is usually of the order 0.05-0.2 mol L" Similar remarks apply to weak bases. It is clear that the greater the concentrations of acid and conjugate base in a buffer solution, the greater will be the buffer capacity. A quantitative measure of buffer capacity is given by the number of moles of strong base required to change the pH of 1 litre of the solution by 1 pH unit. 

The most important type of mixed solution is a buffer, a solution in which the pH resists change when small amounts of strong acids or bases are added. Buffers are used to calibrate pH meters, to culture bacteria, and to control the pH of solutions in which chemical reactions are taking place. They are also administered intravenously to hospital patients. Human blood plasma is buffered to pH = 7.4 the ocean is buffered to about pH = 8.4 by a complex buffering process that depends on the presence of hydrogen carbonates and silicates. A buffer consists of an aqueous solution of a weak acid and its conjugate base supplied as a salt, or a weak base and its conjugate acid supplied as a salt. Examples are a solution of acetic acid and sodium acetate and a solution of ammonia and ammonium chloride. 

The pH of a buffer solution is close to the pKa of the weak acid component when the acid and base have similar concentrations. 

Now consider the overall shape of the pH curve. The slow change in pH about halfway to the stoichiometric point indicates that the solution acts as a buffer in that region (see Fig. 11.3). At the halfwayr point of the titration, [HA] = [A ] and pH = pfCa. In fact, one way to prepare a buffer is to neutralize half the amount of weak acid present with strong base. The flatness of the curve near pH = pKa illustrates very clearly the ability of a buffer solution to stabilize the pH of the solution. Moreover, we can now see how to determine pKa plot the pH curve during a titration, identify the pH halfway to the stoichiometric point, and set pKa equal to that pH (Fig. 11.8). To obtain the pfCh of a weak base, we find pK3 in the same way but go on to use pKa -1- pfq, = pKw. The values recorded in Tables 10.1 and 10.2 were obtained in this way. 

The pH is governed by the major solute species present in solution. As strong base is added to a solution of a weak acid, a salt of the conjugate base of the weak acid is formed. This salt affects the pH and needs to be taken into account, as in a buffer solution. Table 11.2 outlines the regions encountered during a titration and the primary equilibrium to consider in each region. 

To protect a solution against pH variations, a major species in the solution must react with added hydronium ions, and another major species must react with added hydroxide ions. The conjugate base of a weak acid will react readily with hydronium ions, and the weak acid itself will react readily with hydroxide ions. This means that a buffer solution can be defined in terms of its composition. 

A buffer solution contains both a weak acid and its conjugate base as major species in solution. 

The analysis carried out in Example reveals one of the key features of buffer solutions The equilibrium concentrations of both the weak acid and its conjugate base are essentially the same as their initial concentrations. 

In the laboratory, chemists prepare buffer solutions in three different ways. Each results in a solution containing a weak acid and its conjugate base as major species. The most straightforward way to produce a buffer solution is by dissolving a salt of a weak acid in a solution of the same weak acid, as described in Example. ... 

A second way to prepare a buffer is by adding strong base to a solution of a weak acid. This produces a buffer solution if the number of moles of strong base is about half the number of moles of weak acid. As a simple example, if 1 L of 0.5 M NaOH is mixed with 1 L of 1.0 M CH3 CO2 H, hydroxide anions react quantitatively... 

A third approach to buffer solutions is to add strong acid to a solution of a weak base. This produces a buffer solution if the amount of strong acid is about half the amount of weak base. Continuing with our examples of acetic acid-acetate buffers, if a solution of hydrochloric acid is added to a solution of sodium acetate, then hydronium ions react quantitatively with acetate anions ... 

The question asks if this is a buffer solution. A buffer solution contains both a weak acid and its conjugate base as major species. Thus, to answer the question, we must calculate the concentrations of acetate anions and acetic acid in the solution. We use a compressed version of the seven-step method to obtain these concentrations. 

As long as the buffer solution contains acetic acid as a major species, a small amount of hydroxide ion added to the solution will be neutralized completely. Figure 18-1 shows two hydroxide ions added to a portion of a buffer solution. When a hydroxide ion collides with a molecule of weak acid, proton transfer forms a water molecule and the conjugate base of the weak acid. As long as there are more weak acid molecules in the solution than the number of added hydroxide ions, the proton transfer reaction goes virtually to completion. Weak acid molecules change into conjugate base anions as they mop up added hydroxide. 

When a strong base is added to a buffer solution, the weak acid H4 donates protons to hydroxide ions to form the conjugate base... 

The same molecular reasoning shows that a buffer solution can absorb added hydronium ions. Consider what happens when some hydronium ions are added to the acetic acid-acetate buffer solution described in Example. The hydronium ion is a strong acid, and the acetate anion is a weak base, so proton transfer from CH3 CO2 to H3 goes essentially to completion ... 

When a strong acid is added to a buffer solution, the conjugate base A accepts protons from hydronium ions to form the weak acid HA, preventing a large increase in hydronium ion concentration. (All water molecules except those produced in the proton transfer process are omitted for clarity.)... 

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