Concentration of weak acids

Equation 6.44 is written in terms of the concentrations of CH3COOH and CH3COO- at equilibrium. A more useful relationship relates the buffer s pH to the initial concentrations of weak acid and weak base. A general buffer equation can be derived by considering the following reactions for a weak acid, HA, and the salt of its conjugate weak base, NaA. 

If the initial concentrations of weak acid and weak base are greater than [H3O+] and [OH ], the general equation simplifies to the Henderson-Hasselhalch equation. 

Percent ionization also depends on the concentration of weak acid, increasing as the acid is diluted (Figure 13.8). 

In most of the problems you will work, the approximation a — x a is valid, and you can solve for [H+] quite simply, as in Example 13.7, where x = 0.012a. Sometimes, though, you will find that the calculated [H+] is greater than 5% of the original concentration of weak acid. In that case, you can solve for x by using either the quadratic formula or the method of successive approximations. 

The general approach illustrated by Example 18.7 is widely used to determine equilibrium constants for solution reactions. The pH meter in particular can be used to determine acid or base equilibrium constants by measuring the pH of solutions containing known concentrations of weak acids or bases. Specific ion electrodes are readily adapted to the determination of solubility product constants. For example, a chloride ion electrode can be used to find [Cl-] in equilibrium with AgCl(s) and a known [Ag+]. From that information, Ksp of AgCl can be calculated. 

Buffer capacity also depends on the relative concentrations of weak acid and base. Broadly speaking, a buffer is found experimentally to have a high capacity for acid when the amount of base present is at least 10% of the amount of acid. Otherwise, the base is used up quickly as strong acid is added. Similarly, a buffer has a high capacity for base when the amount of acid present is at least 10% of the amount of base, because otherwise the acid is used up quickly as strong base is added. 

Thus, at equilibrium, the transbilayer concentration gradient of the weak acid reflects the inverse of the transbilayer concentration gradient of protons (Fig. 14). For example, a pH difference of 2 units (e.g., internal pH = 9 and external pH = 7) shoud lead to 100-fold higher concentration of weak acid within the vesicle as compared to the external concentration. 

Rate measurements are straightforward if the carbenes can be monitored directly. As a rule, the decay of carbene absorption is (pseudo) first-order, due to rearrangement and/or reaction with the solvent. In the presence of a quencher, the decay is accelerated (Eq. 1), and the rate constant kq is obtained from a plot of k0bs versus [Q], Curved plots were often observed with proton donors (HX) as quenchers, particularly for high concentrations of weakly acidic alcohols. Although these effects have been attributed to oligomerization of the alcohols,91 the interpretation of curved plots remains a matter of dispute.76 Therefore, the rate constants reported in Tables 2-4 are taken from linear (regions of) obs-HX plots, or refer to a specified concentration of HX. 

As base is added, a mixture of weak acid and conjugate base is formed. This is a buffer solution and can be treated as one in the calculations. Determine the moles of acid consumed from the moles of titrant added—that will be the moles of conjugate base formed. Then calculate the molar concentration of weak acid and conjugate base, taking into consideration the volume of titrant added. Finally, apply your buffer equations. 

Strong acids and bases (and strong electrolytes) dissociate completely in water. Therefore, you can use the concentrations of these compounds to determine the concentrations of the ions they form in aqueous solutions. You cannot, however, use the concentrations of weak acids, bases, and electrolytes in the same way. Their solutions contain some particles that have not dissociated into ions. Nevertheless, important changes in [HsOT and [OH ] take place because dissolved ions affect the dissociation of water. 

Table 3.1 Concentrations of weak acids and bases in natural waters...

Although cellulose acetate is not inherently a polyelectrolyte there are reports which indicate that it contains a low concentration of weak acid, presumably carboxylic, groups (1). Water absorbed by cellulose acetate membranes might be preferentially located, to some extent, in the region of these ionogenic groups and so assist in their dissociation. 

Buffer with equal concentrations of weak acid and its... 

Samples containing high concentrations of weak acids in the matrix can be treated to remove the... 

In addition to changing the pH of the water, the uptake and release of CO2 alter the buffer capacity of the water. The effect upon buffer capacity is the result of two factors (1) the dependence of buffer capacity on the hydrogen ion concentration, and (2) the dependence of buffer capacity on the total concentration of weak acid and conjugate base in solution (67, 68). The precipitation of CaCO in natural waters reduces the buffer capacity to a value lower than that predicted on the basis of pH change and respiratory or photosynthetic changes in COL content of the water. 

FIGURE 2.5 The effect of the total concentration of weak acid ( pKa = 4.8 )-conjugate buffer on the pH of the buffer. 

The hydrogen ion concentration of weak acids must be determined by considering the equilibrium concentrations of all ions in the equilibrium mixture. 

Logs to the base 10 are used to express the concentration of weak acid and alkaline solutions as a more easily understood whole number, pH (see Chapter 9). 

Occasions arise when it is useful to calculate the pH of a solution resulting from addition of a known concentration of weak acid or base to water. The pH may be determined by considering the dissociation of the acid. 

At what concentration of weak acid, HA (in terms of its K,), will the acid be 25% dissociated ... 

In Section 4.4 we saw that most soilwaters that feed rivers and groundwater have near-neutral pH, with I ICO, as the major anion. This results from the dissolution of C02 in water (see eqn. 4.7) and from the acid hydrolysis of silicates and carbonates. The total concentration of weak acid anions like HCOf in water is referred to as alkalinity. These anions are available to neutralize acidity (H+) in natural waters, consequently it is important to understand their chemical behaviour. 

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