Point stoichiometric

A TWC catalyst must be able to partition enough CO present in the exhaust for each of these reactions and provide a surface that has preference for NO adsorption. Rhodium has a slight preference for NO adsorption rather than O2 adsorption Pt prefers O2. Rh also does not cataly2e the unwanted NH reaction as does Pt, and Rh is more sinter-resistant than Pt (6). However, the concentrations of O2 and NO have to be balanced for the preferred maximum reduction of NO and oxidation of CO. This occurs at approximately the stoichiometric point with just enough oxidants (O2 and NO ) and reductants (CO, HC, and H2). If the mixture is too rich there is not enough O2 and no matter how active the catalyst, some CO and HC is not converted. If the mixture is too lean, there is too much O2 and the NO caimot effectively compete for the catalyst sites (53—58). 

In reductions of dienes to monoenes, it is important to monitor the hydrogen consumed and to stop the reduction at the stoichiometric point few reductions will stop spontaneously. 

With a knowledge of the pH at the stoichiometric point and also of the course of the neutralisation curve, it should be an easy matter to select the appropriate indicator for the titration of any diprotic acid for which K1/K2 is at least 104. For many diprotic acids, however, the two dissociation constants are too close together and it is not possible to differentiate between the two stages. If K 2 is not less than about 10 7, all the replaceable hydrogen may be titrated, e.g. sulphuric acid (primary stage — a strong acid), oxalic acid, malonic, succinic, and tartaric acids. 

The stoichiometric point is also called the equivalence point. 

In a typical acid—base titration, the analyte is a solution of a base and the titrant is a solution of an acid or vice versa. An indicator a water-soluble dye (Section J), helps us detect the stoichiometric point, the stage at which the volume of titrant added is exactly that required by the stoichiometric relation between titrant and analyte. For example, if we titrate hydrochloric acid containing a few drops of the indicator phenolphthalein, the solution is initially colorless. After the stoichiometric point, when excess base is present, the solution in the flask is basic and the indicator is pink. The indicator color change is sudden, so it is easy to detect the stoichiometric point (Fig. L.3). Toolbox L.2 shows how to interpret a titration the procedure is summarized in diagram (3), where A is the solute in the titrant and B is the solute in the analyte. 

Suppose that 25.00 mL of a solution of oxalic acid, H2C204 (4), which has two acidic protons, is titrated with 0.100 m NaOH(aq) and that the stoichiometric point is reached when 38.0 mL of the solution of base is added. Find the molarity of the oxalic acid solution. 

SOLUTION Because each acid molecule provides two protons, if the concentrations of acid and base were the same, the titration would require a volume of base equal to twice the volume of acid. The volume of base added to reach the stoichiometric point is less than twice the volume of acid so we can expect that the acid is less concentrated than the base. Proceed as in Toolbox L.2. 

Self-Test L.3A A student used a sample of hydrochloric acid that was known to contain 0.72 g of hydrogen chloride in 500.0 mL of solution to titrate 25.0 ml. of a solution of calcium hydroxide. The stoichiometric point was reached when 15.1 mL of acid had been added. What was the molarity of the calcium hydroxide solution ... 

The stoichiometric point is reached when all the Fe2+ has reacted and is detected when the purple color of the permanganate ion persists. A sample of ore of mass 0.202 g was dissolved in hydrochloric acid, and the resulting solution needed 16.7 mL of 0.0108 M KMn04(aq) to reach the stoichiometric point, (a) What mass of iron(ll) ions is present (b) What is the mass percentage of iron in the ore sample ... 

A sample of industrial waste was analyzed for the presence of arsenic(III) oxide by titration with 0.0100 M KMn04. It took 28.15 mL of the titrant to reach the stoichiometric point. How many grams of arsenic(III) oxide did the sample contain ... 

L.10 A sample of oxalic acid, H,C,04 (with two acidic protons), of volume 25.17 mL was titrated to the stoichiometric point with 25.67 mL of 0.327 M NaOH(aq). (a) What is the molarity of the oxalic acid (b) Determine the mass of oxalic acid in the solution. 

L.ll A sample of barium hydroxide of mass 9.670 g was dissolved and diluted to the mark in a 250.0-mL volumetric flask. It was found that 11.56 mL of this solution was needed to reach the stoichiometric point in a titration of 25.0 ml. of a nitric acid solution, (a) Calculate the molarity of the HN03 solution. 

L.18 A tablet of vitamin C was analyzed to determine whether it did in fact contain, as the manufacturer claimed, 1.0 g of the vitamin. One tablet was dissolved in water to form 100.00 mL of solution, and 10.0-mL of that solution was titrated with iodine (as potassium triiodide). It required 10.1 mL of 0.0521 M I3-(aq) to reach the stoichiometric point in the titration. Given that 1 mol L,- reacts with 1 mol vitamin C in the reaction, is the manufacturer s claim correct The molar mass of vitamin C is 176 g-mol-1. 

Suppose that 200. mL of hydrogen chloride at 690. Torr and 20.°C is dissolved in 100. mL of water. The solution was titrated to the stoichiometric point with 15.7 mL of a sodium hydroxide solution. What is the molar concentration of the NaOH in solution ... 

As we saw in Section L, titration involves the addition of a solution, called the titrant, from a buret to a flask containing the sample, called the analyte. For example, if an environmental chemist is monitoring acid mine drainage and needs to know the concentration of acid in the water, a sample of the effluent from the mine would be the analyte and a solution of base of known concentration would be the titrant. At the stoichiometric point, the amount of OH " (or 11,0 ) added as titrant is equal to the amount of H30+ (or OH-) initially present in the analyte. The success of the technique depends on our ability to detect this point. We use the techniques in this chapter to identify the roles of different species in determining the pH and to select the appropriate indicator for a titration. 

A plot of the pH of the analyte solution against the volume of titrant added during a titration is called a pH curve. The shape of the pH curve in Fig. 11.4 is typical of titrations in which a strong acid is added to a strong base. Initially, the pH falls slowly. Then, at the stoichiometric point, there is a sudden decrease in pH through 7. At this point, an indicator changes color or an automatic titrator responds electronically to the sudden change in pH. Titrations typically end at this point. However, if we were to continue the titration, we would find that the pH... 

Experimentally, we know that the pH changes abruptly close to the stoichiometric point. Suppose we reach the stoichiometric point in the titration described in Example 1 1.4 and then add an additional 1.00 mL of HCl(aq). To find out by how much the pH changes, we work through the steps in Toolbox 11.1 as in Example 11.4, except that now the acid is in excess. We find that, after the addition, the pH has fallen to 2.1 (point D in Fig. 11.4). This point is well below the pH (of 7) at the stoichiometric point, although only 1 mL more acid has been added. 

In many titrations, one solution—either the analyte or the titrant—contains a weak acid or base and the other solution contains a strong base or acid. For example, if we want to know the concentration of formic acid, the weak acid found in ant venom (1), we can titrate it with sodium hydroxide, a strong base. Alternatively, to find the concentration of ammonia, a weak base, in a soil sample, titrate it with hydrochloric acid, a strong acid. Weak acids are not normally titrated with weak bases, because the stoichiometric point is too difficult to locate. 

FIGURE 11.5 The variation of pH during a typical titration of a strong acid (the analyte) with a strong base (the titrant). The stoichiometric point (S) occurs at pH - 7. 

Figures 11.6 and 11.7 show the different pH curves that arc found experimentally for these two types of titrations. Notice that the stoichiometric point does not occur at pH = 7. Moreover, although the pH changes reasonably sharply near the stoichiometric point, it does not change as abruptly as it does in a strong acid-strong base titration.

The pH at the stoichiometric point depends on the type of salt produced in the neutralization reaction. At the stoichiometric point of the titration of formic acid, HCOOH, with sodium hydroxide,... 

STRATEGY Decide whether the salt present at the stoichiometric point provides an ion that acts as a weak base or as a weak acid. If the former, expect pH > 7 if the latter, expect pH < 7. To calculate the pH at the stoichiometric point, proceed as in Example 10.10 or 10.11, noting that the amount of salt at the stoichiometric point is equal to the initial amount of acid and the volume is the total volume of the combined analyte and titrant solutions. The Kb of a weak base is related to the / a of its conjugate acid by Ka X Kb = fCw Ka is listed in Table 10.1. Assume that the autoprotolysis of water has no significant effect on the pH, and then check that assumption. 

SOLUTION The salt present at the stoichiometric point, sodium formate, provides basic formate ions, and so we expect pH > 7. From Table 10.1, Ka = 1.8 X 10 4 for formic acid therefore, Kb = KJKa = 5.6 X 1CT11. 

Now that we know the composition of the solution at the stoichiometric point, we can calculate the pH of the solution as described in Toolbox 10.2. The equilibrium to consider is... 

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