Action of Buffers

A buffer resists changes in pH because it contains both an acid to neutralize added OH ions and a base to neutralize added H ions. The acid and base that make up the buffer, however, must not consume each other through a neutralization reaction. ooo (Section 4.3) These requirements are fulfilled by a weak acid-base conjugate pair, such as CH3COOH/CH3C00 or NH4 /NH3. The key is to have roughly equal concentrations of both the weak acid and its conjugate base. 

By choosing appropriate components and adjusting their relative concentrations, we can buffer a solution at virtually any pH. 

Which of these conjugate acid-base pairs will not function as a buffer C2H5COOH and C2H5C00, HCOa and C03 , or HNO3 and N03 Explain. 

To understand how a buffer works, let s consider one composed of a weak acid HA and one of its salts MA, where M could be Na , IC, or any other cation that does not react with water. The acid-dissociation equilibrium in this buffered solution involves both the acid and its conjugate base  

We see from this expression that [H ] and, thus, the pH are determined by two factors the value of K, for the weak-acid component of the buffer and the ratio of the concentrations of the conjugate add-base peur, [HA]/[A ]. 

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An excellent flash animation with audio to explain the action of buffering solutions is found at... 

The action of buffer systems to resist change follows Le Chatelier s principle, where the system acts to restore the disturbance in equilibrium. So if alkali is added to this system there is a build up on the right-hand side of the equation that shifts the equilibrium to the right, and in order to maintain balance, more acid (HA) needs to be dissociated. Conversely, if strong acid is added to the system there is a build up on the left-hand side of the equation that shifts the equilibrium to the left, and in order to maintain balance, more base (A ) needs to be protonated (i.e. becomes HA). It is important to note that in circumstances where very strong acids or bases are added to a buffer system it may not be able to restore equilibrium and there will be a large associated change in pH. 

Section 19.1 discusses the Brpnsted theory of acids and bases, which extends the concepts of add and base beyond aqueous solutions and also explains the acidic or basic nature of solutions of most salts. Dissociation constants, the equilibrium constants for the reactions of weak acids or bases with water, are introduced in Section 19.2. The concept of the ionization of covalent compounds is extended to water itself in Section 19.3, which also covers pH, a scale of acidity and basicity. Section 19.4 describes buffer solutions, which resist change in their acidity or basicity even when some strong acid or base is added. Both the preparation and the action of buffer solutions are explained. Section 19.5 discusses the equilibria of acids containing more than one ionizable hydrogen atom per molecule. 

Obviously, there must be systems in place that prevent the delicate pH balance of living systems from being fatally upset, such as in Figure 5.5.1. Every system is different, but the general principles for the action of buffers are the same. 

Buffers are defined as substances that resist changes in the pH of a system. All weak acids or bases, in the presence of their salts, form buffer systems. The action of buffers and their role in maintaining the pH of a solution can best be explained with the aid of the Henderson-Hasselbalch equation, which may be derived as follows. 

The mode of action of buffers on the Maillard reaction has been proposed to be (11) ... 

Human blood is held at nearly constant pH by the action of buffers, a main topic of this chapter. 

A pH of 6-0-6-5 is generally ensured by the buffering action of the ammonia produced by hydrolysis upon the hexamine salt.)... 

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. 

It is a consequence of the action of different pH values in the aeration cell that these cells do not arise in well-buffered media [4] and in fast-flowing waters [5-7]. The enforced uniform corrosion leads to the formation of homogeneous surface films in solutions containing Oj [7-9]. This process is encouraged by film-forming inhibitors (HCOj, phosphate, silicate, Ca and AP ) and disrupted by peptizing anions (CP, SO ") [10]. In pure salt water, no protective films are formed. In this case the corrosion rate is determined by oxygen diffusion [6,7,10]... 

These reactions have very large equilibrium constants, as we will see in Section 14.3, and so go virtually to completion. As a result, the added H+ or OH- ions are consumed and do not directly affect the pH. This is the principle of buffer action, which explains why a buffered solution is much more resistant to a change in pH than one that is unbuffered (Figure 14.1, p. 384). 

Before leaving the subject of buffer solutions, it is necessary to draw attention to a possible erroneous deduction from equation (21), namely that the hydrogen-ion concentration of a buffer solution is dependent only upon the ratio of the concentrations of acid and salt and upon Ka, and not upon the actual concentrations otherwise expressed, that the pH of such a buffer mixture should not change upon dilution with water. This is approximately although not strictly true. In deducing equation (18), concentrations have been substituted for activities, a step which is not entirely justifiable except in dilute solutions. The exact expression controlling buffer action is ... 

Buffer action 46 Buffer capacity 48 Buffer mixture universal, (T) 831 Buffer solutions 46, (T) 831 acetic acid-sodium acetate, 49 for EDTA titrations, 329 preparation of IUPAC standards, 569 Bumping of solutions 101 Buoyancy of air in weighing 77 Burette 84, 257 piston, 87 reader, 85 weight, 86... 

The active luciferase in the soluble luminescence system occurs in two molecular sizes, 130 kDa and 35kDa. The 130 kDa luciferase is the native form and occurs in extracts made at pH 8, and if luciferase is extracted with a pH 6 buffer, 130 kDa luciferase is converted into 35 kDa luciferase by the action of a protease (Krieger and Hastings, 1968 Fogel and Hastings, 1971 Krieger et al., 1974). The 130 kDa species is considered the naturally occurring form. 

Coordinated phosphate control charts assume either that all contribution to pH level is derived from phosphate or that the buffering action of phosphate is sufficient to overcome the presence of other alkaline species, such as amines. Neither assumption is true. This may lead operators to conclude perhaps that a higher than anticipated bulk water pH level (caused by the presence of amine) should be rectified by the addition of MSP. This action may lower localized Na P04 ratios below 2.2 1.0, producing acid phosphate corrosion (phosphate wastage) and resulting in tube thinning and ultimately tube failure. 

We can use these numbers to express the range of buffer action in terms of the pH of the solution. The Henderson-Hasselbalch equation shows us that,... 

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