Chemical equations buffers

We know the initial concentration of NH3 in the buffer solution and can use the pH to find the equilibrium [OH ]. The rest of the solution is organized around the balanced chemical equation. Our first goal is to determine the initial concentration of NH/. 

Explain why each of the following solutions is a buffer solution. Write a balanced chemical equation where appropriate. 

The buffer factor. The oceanic buffer factor (or Revelle factor), by which the concentration of CO2 in the atmosphere is determined, increases as the concentration of CO2 increases. The buffer factor is discussed above in Section 8.10.3.1.3. Here, it is sufficient to describe the chemical equation for the dissolution of CO2 in seawater. 

An en2yme-catalyzed reaction was carried out in a solution buffered with 0.03 M phosphate, pH 7.2. As a result of the reaction, 0.004 mole/liter of acid was formed, (a) What was the pH at the end of the reaction (b) What would the pH be if no buffer were present (c) Write the chemical equation showing how the phosphate buffer resisted a large change in pH. 

A reference book states that a saturated aqueous solution of potassium hydrogen tartrate is a buffer with a pH of 3.56. Write two chemical equations that show the buffer action of this solution. (Tartaric acid is a diprotic acid with the formula H2C4H40g. Potassium hydrogen tartrate is KHC4H4O5.)... 

Buffers, such as those used in so-called buffered aspirin (acetylsalicylic acid) for easier digestion, or those needed for maintaining a proper pH for enzyme function in living organisms, can be analyzed in terms of chemical equilibrium. A weakly dissociating acidic compound, for example, reaches equilibrium concentrations in solution. The following chemical equation describes the forward and reverse reactions for the hypothetical weak acid HX in the aqueous state. 

Since biological systems are dynamic, with many different processes taking place and many different substances present, buffers are necessary to prevent the kind of wide variation of pH that can inhibit proper enzyme catalysis. Thus, a proper pH aids in regulating the reaction rates associated with certain enzymes and maintaining them at levels appropriate for their particular functions. Two important biological buffers are the phosphate buffer system that regulates pH for the fluid inside cells and the carbonic acid buffer system that regulates pH for blood plasma. The chemical equations for these buffers are shown below for an aqueous solution. 

Use chemical equations to show that the buffer system in blood works like the buffer in question 1. 

Write down the chemical equations describing the two buffer regions. 

Write the chemical equation for the reaction that would occur when a base is added to a solution containing the H2P04 /HP04 buffer system. 

Based on the experimental mole ratio of reactants in part C, write a balanced chemical equation describing the complexation reaction you observed assuming all reactants used (neglecting NH20H HC1 and buffer) actually formed complex. 

Write the formulas for three combinations of weak acid and salt that would act as buffered solutions. For each of your combinations, write chemical equations showing how the components of the buffered solution would consiune added acid and base. 

Write the chemical equation for the equilibrium involving the conjugate acid and base. Write the Kl expression. Put into this the values known, and x for the unknown quantity desired. Solve roughly for X using all / values equal to 1. Use approximate equations (3-3), (3-4), or (3-5) for acid, base, or buffer mixtures. 

Cellulase and all chemicals used in this work were obtained from Sigma. Hydrolysis experiments were conducted by adding a fixed amount of 2 x 2 mm oflSce paper to flasks containing cellulase in 0.05 M acetate buffer (pH = 4.8). The flasks were placed in an incubator-shaker maintained at 50 °C and 100 rpm. A Box-Behnken design was used to assess the influence of four factors on the extent of sugar production. The four factors examined were (i) reaction time (h), (ii) enzyme to paper mass ratio (%), (iii) amount of surfactant added (Tween 80, g/L), and (iv) paper pretreatment condition (phosphoric add concentration, g/L), as shown in Table 1. Each factor is coded according to the equation... 

The equations used to calculate permeability coefficients depend on the design of the in vitro assay to measure the transport of molecules across membrane barriers. It is important to take into account factors such as pH conditions (e.g., pH gradients), buffer capacity, acceptor sink conditions (physical or chemical), any precipitate of the solute in the donor well, presence of cosolvent in the donor compartment, geometry of the compartments, stirring speeds, filter thickness, porosity, pore size, and tortuosity. 

Dissolved carbon dioxide produces carbonic acid, which ionizes to bicarbonate and carbonate ions, the reactions for which are shown in Figure 5.2 (equations 1-3). This reaction sequence is extremely important because bicarbonate is a counterion to many cations, is active in buffering the soil solution, and is involved either directly or indirectly in many soil chemical reactions. Bicarbonates are generally more soluble than carbonates, which are generally insoluble. Adding acid to carbonates or bicarbonates results in the release of carbon dioxide and the formation of the salt of the acid cation. The acid is thus neutralized. 

It should be stressed that the pH value of an actual buffer solution prepared by mixing quantities of the weak acid or base and its conjugate base or acid based on the calculated ratio will likely be different from what was calculated. The reason for this is the use of approximations in the calculations. For example, the molar concentration expressions found in Equations (5.23) to (5.30), e.g., [H+], are approximations. To be thermodynamically correct, the activity of the chemical should be used rather than the concentration. Activity is directly proportional to concentration, the activity coefficient being the proportionality constant ... 

A more comprehensive analysis of the influences on the ozone solubility was made by Sotelo et al., (1989). The Henry s Law constant H was measured in the presence of several salts, i. e. buffer solutions frequently used in ozonation experiments. Based on an ozone mass balance in a stirred tank reactor and employing the two film theory of gas absorption followed by an irreversible chemical reaction (Charpentier, 1981), equations for the Henry s Law constant as a function of temperature, pH and ionic strength, which agreed with the experimental values within 15 % were developed (Table 3-2). In this study, much care was taken to correctly analyse the ozone decomposition due to changes in the pH as well as to achieve the steady state experimental concentration at every temperature in the range considered (0°C [Pg.86]

For the sake of completeness, Figure 4-5 illustrates the more general situation of isothermal, isobaric matter transport in a multiphase system (e.g., Fe/Fe0/Fe304 / 02). A sequence of phases a, (3, y,... is bounded by two reservoirs which contain both neutral components (i) and electronic carriers (el). The boundary conditions imply that the buffered chemical potentials (u,(R)) and the electrochemical potentials (//el(R)) are predetermined in R] and Rr. Depending on the concentrations and mobilities (c/, b), c, 6 ) in the various phases v, metallic conduction, semiconduction, or ionic conduction will prevail. As long as the various phases are thermodynamically stable and no decomposition occurs, the transport equations (including the boundary conditions) are well defined and there is normally a unique solution to the transport problem. 

In this paper we have used the quantity (1 — vp0) in writing equations for sedimentation equilibrium experiments. Some workers prefer to use the density increment, 1000(dp/dc)Tfn, instead when dealing with solutions containing ionizing macromolecules. This procedure was first advocated by Vrij (44), and its advantages are discussed by Casassa and Eisenberg (39). Nichol and Ogston (13) have used the density increment in their analysis of mixed associations. The subscript p. means that all of the diffusible solutes are at constant chemical potential in the buffer... 

Because so many chemical reactions in our bodies occur in buffered environments, biochemists commonly use the Henderson-Hasselbalch equation for quick estimates of pH. In practice, the equation is used to make rapid estimates of the pH of a mixed solution intended to be used as a buffer, and then the pH is adjusted to the precise value required by adding more acid or base and monitoring the solution with a pH meter. 

Ballpark Check a common error in using the Henderson-Hasselbalch equation is to invert the [base]/[acid] ratio. It is therefore wise to check that your answer makes chemical sense. If the concentrations of the acid and its conjugate base are equal, the pH will equal the pKa. If the acid predominates, the pH will be less than the pKa, and if the conjugate base predominates, the pH will be greater than the pKa. In part (a), [acid] = [NH4 + ] is greater than [base] = [NH3], and so the calculated pH (8.77) should be less than the pKa (9.25). In part (b), the desired pH is less than the pKa, so the buffer should contain more moles of acid than base, in agreement with the solution. 

Marcus5 8 taught us that the most appropriate and useful kinetic measure of chemical reactivity is the intrinsic barrier (AG ) rather than the actual barrier (AG ), or the intrinsic rate constant (kQ) rather than the actual rate constant (k) of a reaction. These terms refer to the barrier (rate constant) in the absence of a thermodynamic driving force (AG° = 0) and can either be determined by interpolation or extrapolation of kinetic data or by applying the Marcus equation.5 8 For example, for solution phase proton transfers from a carbon acid activated by a ji-acceptor (Y) to a buffer base, Equation (1), k0 may be determined from Br A ns ted-type plots of logki or... 

If the soil suspension were instead an aqueous solution, a scale of activity values for Na+ could be defined in terms of emf data obtained for standard reference solutions of prescribed (Na+), in exactly the same way as the scale of (H) values (the operational pH scale) is defined (arbitrarily) in terms of emf data for standard buffer solutions.39,40 However, the success of this extrathermo-dynamic calibration technique depends entirely on the extent to which E, and B in the standard reference solutions are the same as E, and B in the solution of interest. For the case of a soil suspension, the presence of colloidal material may cause these two parameters to differ very much from what they would be in a reference aqueous solution. If the difference is indeed large, the value of (Na+), m, or any other ionic activity estimated with the help of standard solutions and an equation like Eq. s2.23 would be of no chemical significance. 

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