Concentration from acid-base titration

Consider the redox reaction shown in equation (4.1). From a potentiometric titration, it is found that 12.5 cm of 06 + (0.01 moi dm ) will completely oxidize 25.0 cm of Fe + soiution. What is the concentration of the ferrous iron (Hint - remember the equation, Ci / =CtV2, from acid-base titrations.)... 

Most data are from acid-base titrations at various salt concentrations, leading to a common intersection point. Data obteilned for only two salt concentrations are, as a rule, avoided. This is also the policy for curves where the sample conditions and/or the reversibility of the titrations are insufllclently controlled or which are suspect for other reasons. Isoelectric points or points of zero charge obtained by other methods are sometimes recorded (in italics) if titration data are not available, or if such points are interesting for other reasons. When the temperature Is known = -2.303 J T(pH° - pOH°)/2, see (3.8.16), is also... 

To make comparisons with the above salt effects, however, accurate values of An from acid-base titrations and correction of concentrations to activities should be considered. At this time, however, several different graphical representations of relative efficiencies are possible. These include comparison of the relative effectiveness of changes in chemical potential, Ap, to drive T, from just above to below the operating temperature, comparison of the relative ApAn areas determined from acid-base titration curves, and comparison of the significance of different degrees of positive cooperativity, that is, the impact of changes in the Hill coefficient. 

In both models, the concentration of surface species was calculated from acid-base titration data to determine the speciation of the solution-solid interface. From a knowledge of the concentration of solution and interfacial species, the mass action... 

Adsorption of anions or neutral molecules on oxides is controlled by the same parameters (pH, concentration, temperature) as those involved in complexatiqn of cations in solution. The pH of the suspension imposes both the level of protonation of the anion and the surface charge of the oxide. As in the case of cation complexation in solution (see Chapter 5), there is an optimum pH range for anion adsorption. Adsorption is characterized by complexation equilibrium constants, determined using an approach similar to the calculation of complexation constants in solution from acid-base titration of suspensions [48-51]. 

FIGURE 11.10 The stoichiometric point of an acid base titration may be detected by the color change of an indicator. Here we see the colors of solutions containing a few drops of phenolphthalein at (from left to right) pH of 7.0, 8.5, 9.4 (its end point), 9.8, and 12.0. At the end point, the concentrations of the conjugate acid and base forms of the indicator are equal... 

In the practice of potentiometric titration there are two aspects to be dealt with first the shape of the titration curve, i.e., its qualitative aspect, and second the titration end-point, i.e., its quantitative aspect. In relation to these aspects, an answer should also be given to the questions of analogy and/or mutual differences between the potentiometric curves of the acid-base, precipitation, complex-formation and redox reactions during titration. Excellent guidance is given by the Nernst equation, while the acid-base titration may serve as a basic model. Further, for convenience we start from the following fairly approximate assumptions (1) as titrations usually take place in dilute (0.1 M) solutions we use ion concentrations in the Nernst equation, etc., instead of ion activities and (2) during titration the volume of the reaction solution is considered to remain constant. 

Other methods employ a microplate format followed by fast HPLC. Some researchers approach the determination from a different perspective. For example, an alternative method for ionizable substances is the pSol determination based on an acid-base titration.25 26 Kinetic solubility determinations involve determining the concentration of the compound in the buffer of interest when an induced precipitate first appears. 

Common chemical titrations include acid-base, oxidation-reduction, precipitation, and complexometric analysis. The basic concepts underlying all titration are illustrated by classic acid-base titrations. A known amount of acid is placed in a flask and an indicator added. The indicator is a compound whose color depends on the pH of its environment. A solution of base of precisely known concentration (referred to as the titrant) is then added to the acid until all of the acid has just been reacted, causing the pH of the solution to increase and the color of the indicator to change. The volume of the base required to get to this point in the titration is known as the end point of the titration. The concentration of the acid present in the original solution can be calculated from the volume of base needed to reach the end point and the known concentration of the base. 

In an acid-base titration, you carefully measure the volumes of acid and base that react. Then, knowing the concentration of either the acid or the base, and the stoichiometric relationship between them, you calculate the concentration of the other reactant. The equivalence point in the titration occurs when just enough acid and base have been mixed for a complete reaction to occur, with no excess of either reactant. As you learned in Chapter 8, you can find the equivalence point from a graph that shows pH versus volume of one solution added to the other solution. To determine the equivalence point experimentally, you need to measure the pH. Because pH meters are expensive, and the glass electrodes are fragile, titrations are often performed using an acid-base indicator. 

In a typical acid-base titration (Section 3.10), a solution containing a known concentration of base (or acid) is added slowly from a buret to a second solution containing an unknown concentration of acid (or base). The progress of the titration is monitored, either by using a pH meter (Figure 16.6a) or by observing the color of a suitable acid-base indicator. With a pH meter, you can record data to produce a pH titration curve, a plot of the pH of the solution as a function of the volume of added titrant (Figure 16.6b). 

Quantitation of Napelline.—The pharmacological activity of napelline (35)38 has prompted the development of a method for the quantitation of this alkaloid in raw plant materials.39 This method consists of the extraction of the total alkaloids, chromatographic separation, and a micro-scale acid-base titration of the napelline eluate in a non-aqueous medium. The total alkaloids were exhaustively extracted from a sodium carbonate suspension with chloroform. This extract was concentrated, dissolved in acetone, and chromatographed on silica-gel plates, with a standard reference of napelline as a marker. The appropriate bands were quantitatively removed and extracted, and the extracts were concentrated to dryness. These residues were dissolved in glacial acetic acid and titrated with 0.01N perchloric acid. The standard deviation of this method was 3.39 x 10 3. No limits of detection are reported. 

Any titration involves the progressive change of the activities (or concentrations) of the titrated and titrating species and, in principle, can be done potentiometrically. However, for an accurate determination it is necessary that there is a fairly rapid variation in equilibrium potential in the region of the equivalence point. Useful applications are redox, complexation, precipitation, acid-base titrations, etc. From the titration curve it is possible to calculate values of the formal potentials of the titrated and titrating species, as explained below. 

When the progression of an acid-base titration is graphed as a function of pH vs the volume of acid or base added, the curve will appear as shown below. If we recall, from general chemistry coursework, that the steepest point on the curve represents the equivalence point of the titration (the point where the amount of acid and base are equal), we can locate the point on the curve that represents the midpoint of the titration. This point is found at half the concentration of base added to acid (or acid added to base) to reach the equivalence point. Once we have done this, we recall the Henderson-Hesselbach equation (Fig. 2.8)—specifically, the term dealing with the concentrations of the ionic and the neutral species. Realizing that at the midpoint of the titration, these concentrations are equal, the logarithmic term in the Henderson-Hesselbach equation reduces to log(l), which is equal to zero. Therefore, the equation reduces to pA a = pH at the midpoint of the titration. 

An acid-base titration is a procedure that is used where a base of known concentration is added to an acid of unknown concentration (or vice-versa) in order to determine the concentration of the unknown. In addition, it is possible to determine the Ka of the acid being titrated (or Kb of the base) as well as an appropriate indicator. Acid-base titrations are often the topic of AP test questions and are frequently used in the laboratory questions. You should know about titrations from a conceptual level, be able to perform calculations for titrations, and know how to properly perform one in the laboratory. We ll begin with the conceptual explanation of titrations. 

The amphoteric nature of wool was demonstrated in the early studies of Speakman and Hirst (1933), Elod (1933), and in particular by the complete acid-base titration curve obtained by Speakman and Stott (1934). Even earlier attempts had been made to determine the isoelectric point of wool by the methods indicated in Table XXIII. Some variation in the isoelectric point is to be expected because the pH at which the net charge, including bound ions, is zero depends on the nature and concentrations of ions in the aqueous environment. For example, Sookne and Harris (1939) have shown that the early electrophoretic value of Harris (1932) was affected by the absorption of phthalate ions from the buffer solutions. With acetate buffer they obtained values of 4.2 and 4.5 for powdered wool and cortical cells, respectively. The isoelectric points listed in Table XXIII are... 

The composition of the rain—an average inorganic composition is given in Figures 3.1 and 5.1—reflects the acid-base titration that occurs in the atmosphere. Total concentrations (the sum of cations or anions) typically vary from 20 to 500 /4eq liter" and pH from 3.5 to 6. 

A 1 to 1.5 A/ solution of methylamine in benzene is prepared by leading methylamine (from a lecture bottle) through 100 mL of dry benzene during about 30 min. The actual concentration of the amine is determined by acid-base titration, using dilute hydrochloric acid and bromocresol green as indicator. 

It is important to know the dissociation constant of an indicator in order to use it properly in acid-base titrations. Spectrophotometry can be used to measure the concentration of these intensely colored species in acidic versus basic solutions, and from these data the equilibrium between the acidic and basic forms can be calculated. In one such study on the indicator wj-nitrophenol, a 6.36 X 10 M solution was examined by spectrophotometry at 390 nm and 25°C in the following experiments. In highly acidic solution, where essentially all the indicator was in the form HIn, the absorbance was 0.142. In highly basic solution, where essentially all of the indicator was in the form In , the absorbance was 0.943. In a further series of experiments, the pH was adjusted using a buffer solution of ionic strength I, and absorbance was measured at each pH value. The following results were obtained ... 

Demonstration of an acid-base titration that uses the indicator phenolphthalein. An alkaline solution is in the flask, with a few drops of the indicator added to give a purple color. Acid is slowly added from the burette above. When the acid has completely neutralized the alkali in the solution, the indicator turns colorless. This technique is used to gather data for many investigations, such as determining the concentration of the acid or alkali. (Courtesy of Jerry Mason/Science Photo Library)... 

Figure 3.4 shows that the absolute value of ionic strength, except at the CIP. Thus addition of inert electrolyte (crystalline or as concentrated solution) to suspension at CIP does not result in any pH change. At any other initial pH the addition of inert electrolyte induces a shift of pH toward the CIP. This means that the pH increases on addition of inert electrolyte at pH < CIP and decreases at pH > CIP. A series of suspensions is prepared using the same amounts of solid and of dilute solution of inert electrolyte, the suspensions are adjusted to different pH, and allowed to equilibrate. The normal precautions as by classical acid-base titration (thermostating, protection from COi) have to be respected. Then constant amount of inert electrolyte is added to all samples, and the change in pH is recorded, and the results (ApH) are plotted as a function of pH. The intersection of the obtained curve... 

Acetic acid is a weak acid. It is available at different concentrations. Highly concentrated acetic acid at 98% and above is called glacial acetic acid because its freezing point range is between 13.3 °C (98%) and 16.7 °C (100%). Glacial acetic acid is flammable. The concentration of acetic acid can easily be determined using acid-base titration with phenolphthalein as an indicator. The water used should be free from CO, prepared by boiling before use. 

Thus the change of the concentration of free base of one order of magnitude will produce an increase of potential of 59 millivolts at 25°, i.e., the increase of base concentration from 1CH to 10 mol per liter will produce the same effect on the potential as the change from 10-3 to lO2 mol per liter. A plot of E against the volume of added base will be very closely of the form given in Fig. 2 (o). A more adequate discussion of the theory of the acid-base titration, including the ionization equilibrium of water, is given later in this chapter. 

Finally we have seen in section 4-11 how acid-base titrations can be used in practice, even without any preliminary separations or sample clean-up, and what trade-offs are made in such analyses. This example illustrates a rather radical departure from the traditional emphasis on titrations as methods of high precision. As illustrated in Table 4.11-1, even when precise concentrations of well-defined chemical species cannot be derived from such complex mixtures, they nonetheless can be made to yield very useful quantitative information. 

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