Buffer capacity concentrations

Buffer pH when working with buffer, always ensure that you are working within 1 pH unit of the buffer pKa to ensure maximum buffering capacity. Concentration is not of vital importance, assuming that sufficient buffer is present. To avoid buffer salts coming out of solution, it is advisable to work at as low a concentration as possible and to use low levels of organic modifier in the mobile phase. 

Phosphoric Acid. This acid is the primary acidulant in cola beverages. Phosphoric acid is stronger than most organic acids and weaker than other mineral acids. The dibasic properties of phosphoric acid provide minor buffering capacity in the beverage. Food-grade phosphoric acid is commercially available in concentrations of 75%, 80%, and 85% and is one of the most economical acidulants. 

FIGURE 2.15 A buffer system consists of a weak acid, HA, and its conjugate base, A. The pH varies only slightly in the region of the titration curve where [HA] = [A ]. The unshaded box denotes this area of greatest buffering capacity. Buffer action when HA and A are both available in sufficient concentration, the solution can absorb input of either H or OH, and pH is maintained essentially constant. 

A solution containing equal concentrations of acid and its salt, or a half-neutralised solution of the acid, has the maximum buffer capacity . Other mixtures also possess considerable buffer capacity, but the pH will differ slightly from that of the half-neutralised acid. Thus in a quarter-neutralised solution of acid, [Acid] = 3 [Salt] ... 

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. 

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. 

The susceptibility of the sulfamates to hydrolysis is intermediate with respect to procedures commonly used for extraction and manipulation of extracts. Quantitative hydrolysis of the pure sulfamate toxins can be accomplished (9) by heating at 100 C for 5 min in the presence of not less than 0.1 M free acid (pH 1 or below). Milder conditions appear insufficient (10). Figure 9 summarizes results from two separate experiments in which samples of nontoxic clam flesh, enriched with constant amounts of saxitoxin Cl (4), were acidified to differing final concentrations of HCl and heated for 5 min at 100 C. The difference between 0.1 M HCl, which would be sufficient for hydrolysis of the pure toxin, and the HCl concentration required to attain plateau toxicity, probably reflects the buffer capacity of... 

Buffer capacity is determined by the amounts of weak acid and conjugate base present in the solution. If enough H3 O is added to react completely with the conjugate base, the buffer is destroyed. Likewise, the buffer is destroyed if enough OH is added to consume all of the weak acid. Consequently, buffer capacity depends on the overall concentration as well as the volume of the buffer solution. A buffer solution whose overall concentration is 0.50 M has five times the capacity as an equal volume of a buffer solution whose overall concentration is 0.10 M. Two liters of 0.10 M buffer solution has twice the capacity as one liter of the same buffer solution. Example includes a calculation involving buffer capacity. 

During the lifetime of a root, considerable depletion of the available mineral nutrients (MN) in the rhizosphere is to be expected. This, in turn, will affect the equilibrium between available and unavailable forms of MN. For example, dissolution of insoluble calcium or iron phosphates may occur, clay-fixed ammonium or potassium may be released, and nonlabile forms of P associated with clay and sesquioxide surfaces may enter soil solution (10). Any or all of these conversions to available forms will act to buffer the soil solution concentrations and reduce the intensity of the depletion curves around the root. However, because they occur relatively slowly (e.g., over hours, days, or weeks), they cannot be accounted for in the buffer capacity term and have to be included as separate source (dCldl) terms in Eq. (8). Such source terms are likely to be highly soil specific and difficult to measure (11). Many rhizosphere modelers have chosen to ignore them altogether, either by dealing with soils in which they are of limited importance or by growing plants for relatively short periods of time, where their contribution is small. Where such terms have been included, it is common to find first-order kinetic equations being used to describe the rate of interconversion (12). 

FIG. 14 A model for the uptake of weakly basic compounds into lipid bilayer membrane (inside acidic) in response to the pH difference. For compounds with appropriate pki values, a neutral outside pH results in a mixture of both the protonated form AH (membrane impermeable) and unprotonated form A (membrane permeable) of the compound. The unprotonated form diffuse across the membrane until the inside and outside concentrations are equal. Inside the membrane an acidic interior results in protonation of the neutral unprotonated form, thereby driving continued uptake of the compound. Depending on the quantity of the outside weak base and the buffering capacity of the inside compartment, essentially complete uptake can usually be accomplished. The ratio between inside and outside concentrations of the weakly basic compound at equilibrum should equal the residual pH gradient. 

Electrophoretic separations occur in electrolytes. The type, composition, pH, concentration, viscosity, and temperature of the electrolytes are all crucial parameters for separation optimization. The composition of the electrolyte determines its conductivity, buffer capacity, and ion mobility and also affects the physical nature of a fused silica surface. The general requirements for good electrolytes are listed in Table 1. Due to the complex effects of the type, concentration, and pH of the separation media buffer, conditions should be optimized for each particular separation problem. 

Parenteral products should be formulated to possess sufficient buffer capacity to maintain proper product pH. Factors that influence pH include product degradation, container and stopper effects, diffusion of gases through the closure, and the effect of gases in the product or in the headspace. However, the buffer capacity of a formulation must be readily overcome by the biological fluids thus, the concentration and ratios of buffer ingredients must be carefully selected. 

Thus, the buffering capacity depends on the composition of the buffer, i.e. on the concentration of the salt a or b. The maximum value found by differentiation of Eq. (1.4.27) with respect to a corresponds, for an acidic buffer, to b = s. 

At equilibrium, the concentration of H+ will remain constant. When a strong acid (represented by H+ or HA) is introduced into solution, the concentration of H+ is increased. The buffer compensates by reacting with the excess H ions, moving the direction of the above reaction to the left. By combining with bicarbonate and carbonate ions to form the nonionic carbonic acid, equilibrium is reestablished at a pH nearly the same as that existing before. The buffer capacity in this case is determined by the total concentration of carbonate and bicarbonate ions. When no more carbonate or bicarbonate ions are available to combine with excess H+ ions, the buffer capacity has been exceeded and pH will change dramatically upon addition of further acid. 

If there is no buffering capacity [k - O and C(HA) - O] and if the diffusion away is very slow, the concentration of H+ ions at the interface will grow until the second term in Eq. (36) becomes equal to the first one. Substituting such a condition in Eq. (35), one can see that the rate of growth of the oxide becomes zero, i.e., the oxide attains a constant thickness. (In fact, some hydrogen ions will always escape by diffusion and, hence, complete equality of the two terms in Eq. (35) can never be attained so that some growth will have to continue.)... 

Calbindin Provides storage and buffering capacity for the regulation of Ca2+ concentrations (357)... 

The tailings comprise 5-10 wt. % pyrrhotite a highly reactive sulfide mineral that releases protons and Fe3+ into adjacent pore waters on oxidation. Further, the concentration of carbonate minerals in the tailings is low providing little buffering capacity above pH 5. Therefore, the tailings continued to acidify until they reach the pH of AI(OH)3 (pH 4-4.5) and Fe(OH)3 (pH 2.5-3.5) buffering. 

These equations allow us to calculate the pH or pOH of the buffer solution knowing Kof the weak acid or base and the concentrations of the conjugate weak acid and its conjugate base. Also, if the desired pH is known, along with K, the ratio of base to acid can be calculated. The more concentrated these species are, the more acid or base can be neutralized and the less the change in buffer pH. This is a measure of the buffer capacity, the ability to resist a change in pH. 

The buffer capacity is a quantitative measure of the ability of a buffer to resist a change in pH. The more concentrated the acid-base components of the buffer, the higher its buffer capacity. 

Table IV gives minimum steam requirement (infinite stages) at several different solution capacities. The factor attribu-able to equilibrium nonlinearity increases as more SO2 is absorbed, because the buffer capacity is consumed to a greater extent. Any capacity for SO2 absorption can be achieved by varying Na concentration (pH) in the solution. At low pH ([Na] = 1.5 M) the solution capacity for SO2 absorption is small, but the nonlinearity factor is also small (1.05). Solution capacity can be increased by operating at higher pH ([Na] = 2.5 M), but nonlinearity is more severe (1.32).

These four buffer solutions have the same initial pH but different concentrations (shown by the numbers beside or on the bars). The pH increases with the addition of a certain amount of strong base. The more concentrated the buffer solution is (that is, the higher its buffer capacity), the smaller the change in pH is. 

---