Chemistry
Polyprotic Acid Titration
Polyprotic acid titration involves the process of determining the concentration of a polyprotic acid by adding a strong base to it. Since polyprotic acids can donate multiple protons, their titration curves exhibit multiple equivalence points. Each equivalence point corresponds to the neutralization of a different proton, and the pH at each point can be used to calculate the acid's dissociation constants.
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10 Key excerpts on "Polyprotic Acid Titration"
- James N. Jensen(Author)
- 2022(Publication Date)
- Wiley(Publisher)
Unfor- tunately, the data are for the titration of a mixture of 1.0×10 –4 M of a monoprotic acid with pK a = 4 and 2.0×10 –5 M of a monoprotic acid with pK a = 3. This example illustrates that there may be many chemical compositions that can account for the general shape of a titration curve. One application in which titration curve features often are lost is in the titration of natural waters. The acid–base chemistry of natural waters is complex. Naturally occurring organics can accept and donate protons. However, the strength of the proton binding sites changes during the titration as the three-dimensional structure of the organics change. As a result, titration curves of natural waters often are fairly feature- less. The rich detail of acid–base chemistry is smeared by the large number and chang- ing nature of the organic acids and bases. To illustrate this point, the titration curve is shown in Figure 12.19 for the titration of 1.85 g/L of natural organics (fulvic acid) isolated from the Göta River (near Göteburg, Sweden; data from Plechanov et al. 1983). Note the relatively featureless titration curve. Plechanov and colleagues estimated that the carboxylic acid content of the natural material was 4.0 meq/g (or 7.4 meq/L for the sample titrated in Figure 12.19). 12.7 CHAPTER SUMMARY An acid–base titration is a stepwise addition of acids or bases to a solution. Titrations describe a number of phenomena in the environment, from the creation of the early oceans to the neutralization of acidic waste streams. Each step in a titration is a sep- arate equilibrium calculation. The equilibrium concentration of most of the species in the system change at each step because the charge balance changes as strong acid anions or strong base cations are added. Graphical representations of titrations, called titration curves, can be sketched eas- ily by calculating the equilibrium pH at a few key points, called equivalence points.
eBook - PDFChemistry
The Molecular Nature of Matter
- Neil D. Jespersen, Alison Hyslop(Authors)
- 2014(Publication Date)
- Wiley(Publisher)
Successive values of K a often differ by a fac- tor of 10 4 to 10 5 . This allows us to calculate the pH of a solu- tion of a polyprotic acid by using just the value of K a 1 . If the polyprotic acid is the only solute, the anion formed in the second step of the ionization, A 2- , has a concentration equal to K a 2 . The anions of weak polyprotic acids are bases that react with water in successive steps, the last of which has the molec- ular polyprotic acid as a product. For a diprotic acid, K b 1 = K w K a 2 and K b 2 = K w K a 1 . Usually, K b 1 W K b 2 , so virtually all the OH - produced in the solution comes from the first step. The pH of the solution can be calculated using just K b 1 and the reaction A 2- + H 2 O m H A - + OH - Draw and explain titration curves for reactions of strong or weak acids and bases A graph of the pH values versus the volume of titrant gives us a titration curve. The titration curve contains important informa- tion about the chemical system including the equivalence point and equilibrium constant data. Each calculated titration curve has four distinct calculations: (a) The starting pH of a solution that contains only an acid or base (b) The pH values for titrant volumes between the start and the equivalence points can be used to calculate the (c) The pH at the equivalence point (d) The pH values for titrant volumes past the equivalence point Acid–base indicators are weak acids in which the conjugate acid and base have different colors. The sudden change in pH at the equivalence point causes a rapid shift from one color of the indicator to the other. If matched to the equivalence point, a pH indicator will change color abruptly when a titration reaches the equivalence point. Ion product constant of water, K w (Section 16.1) K w = 3 H + 4 3 OH - 4 or K w = 3 H 3 O + 4 3 OH - 4 At 25 °C the value of K w is 1.0 Ž 10 -14 .
eBook - PDFAnalytical Chemistry
A Toolkit for Scientists and Laboratory Technicians
- Bryan M. Ham, Aihui MaHam(Authors)
- 2024(Publication Date)
- Wiley(Publisher)
In Section 10.1, we are going to take a closer look at Polyprotic Acid Titrations, such as phosphoric acid using normal solutions. 10.11 Polyprotic Acid Titration Phosphoric acid is a polyprotic acid and thus possesses two dis- tinguishable equivalence points, and thus also two pH transition ranges. Figure 10.9 depicts the titration curve denoting the mid- points, the equivalence points, and the pH transition ranges. With using a pH meter, we are able to detect and measure these two distinct transition ranges. This brings us to an important aspect of visual pH indicators as listed in Figure 10.7. The analyst, know- ing that phosphoric acid is a polyprotic acid, would need to select an indicator that would have a transition range appropriate for the measurement desired. The measurement for ppm acidity as phos- phoric acid would require an indicator with a transition range in the pH range 8–11, such as phenolphthalein. Phenolphthalein, with a very clear color change, is a popular indicator for the determination of the endpoint for weak acid titrations such as the total titration of carbonic, sulfuric, and phosphoric acids. If 0 2 4 6 8 10 12 14 pH Volume of titrant added (ml 0.1 N NaOH) 2nd Equivalence point 1st Equivalence point 0 10 20 30 40 50 FIGURE 10.8 Titration results of submitted sample for % acidity as H 3 PO 4 . 0 2 4 6 8 10 12 14 Transition range 2 pH 8 to 11 Midpoint 3 pH 12.4 pH Volume of titrant added (ml 0.1 N NaOH) 2nd Eq point 1st Eq point Transition range 1 pH 3 to 6 0 10 20 30 40 50 Midpoint 2 pH 7.2 Midpoint 1 pH 2.1 FIGURE 10.9 Phosphoric acid titration curve denoting the mid- points, the equivalence points, and the pH transition ranges.
eBook - ePub- John Kenkel(Author)
- 2013(Publication Date)
- CRC Press(Publisher)
For a titrimetric analysis to be successful, the equivalence point must be easily and accurately detected; the reaction involved must be fast; and the reaction must be quantitative. If an equivalence point cannot be detected (i.e., if there is no acceptable indicator or other detection method), then the correct volume of titrant cannot be determined. If the reaction involved is not fast, then the end point cannot be detected immediately upon adding the last fraction of a drop of titrant and there would be some doubt if the end point has been reached. If the reaction is not quantitative, meaning that if every trace of reactant in the titration flask is not consumed by the titrant at the end point, then again the correct volume of titrant cannot be determined. This latter point means that equilibrium reactions that do not go essentially to completion immediately are not acceptable reactions for this type of analysis. Thus, not all reactions are acceptable reactions.In this chapter, we investigate individual types of reactions that meet all the requirements. We will also discuss “back titrations” and “indirect” titrations in which some of the limitations that we may encounter are solved. Our discussions in Chapter 4 involved acid–base reactions. Acid– base reactions will be discussed here, but we will see that there are others that are applicable reactions.5.2 Acid–Base Titrations and Titration Curves
Various acid–base titration reactions are discussed in this section, including a number of scenarios of base in the buret and acid in the reaction flask and vice versa and also various monoprotic and polyprotic acids titrated with a strong base and various weak monobasic and polybasic bases titrated with strong acids. A monoprotic acid is an acid that has only one hydrogen ion (or proton) to donate per formula. Examples are hydrochloric acid, HCl, a strong acid, and acetic acid, HC2 H3 O2 , a weak acid. A polyprotic acid is an acid that has two or more hydrogen ions to donate per formula. Examples include sulfuric acid, H2 SO4 , a diprotic acid , and phosphoric acid, H3 PO4 , a triprotic acid.A monobasic base is one that will accept just one hydrogen ion per formula. Examples include sodium hydroxide, NaOH, a strong base, ammonium hydroxide, NH4 OH, a weak base, and sodium bicarbonate, NaHCO3 , also a weak base. A polybasic base is one that will accept two or more hydrogen ions per formula. Examples include sodium carbonate, Na2 CO3 , a dibasic base , and sodium phosphate, Na3 PO4 , a tribasic base .5.2.1 Titration of Hydrochloric Acid
A graphic picture of what happens during an acid–base titration is easily produced in the laboratory. Consider again what is happening as a titration proceeds. Consider, specifically, NaOH as the titrant and HCl as the substance titrated. In the titration flask, the following reaction occurs when titrant is added:
eBook - PDF- Douglas Skoog, Donald West, F. Holler, Stanley Crouch, Douglas Skoog(Authors)
- 2021(Publication Date)
- Cengage Learning EMEA(Publisher)
Due to electronic rights, some third party content may be suppressed from the eBook and/or eChapter(s). Editorial review has deemed that any suppressed content does not materially affect the overall learning experience. Cengage Learning reserves the right to remove additional content at any time if subsequent rights restrictions require it. 13E Titration Curves for Polyfunctional Acids 325 13E Titration Curves for Polyfunctional Acids Compounds with two or more acidic functional groups yield multiple end points in a titration if the functional groups differ sufficiently in strength as acids. The compu- tational techniques described in Chapter 12 permit construction of reasonably accu- rate theoretical titration curves for polyprotic acids if the ratio K a1 >K a2 is somewhat greater than 10 3 . If this ratio is smaller, the error becomes excessive, particularly in the region of the first equivalence point, and a more rigorous treatment of the equi- librium relationships is required. Figure 13-2 shows the titration curve for a diprotic acid H 2 A with dissociation constants of K a1 5 1.00 3 10 23 and K a2 5 1.00 3 10 27 . Because the K a1 >K a2 ratio is significantly greater than 10 3 , we can calculate this curve (except for the first equiv- alence point) using the techniques developed in Chapter 12 for simple monoprotic weak acids. Thus, to calculate the initial pH (point A), treat the system as if it con- tained a single monoprotic acid with a dissociation constant of K a1 5 1.00 3 10 23 . In region B, we have the equivalent of a simple buffer solution consisting of the weak HCO 3 2 1 H 2 O m H 3 O 1 1 CO 3 22 K a2 5 3 H 3 O 1 4 3 CO 3 22 4 3 HCO 3 2 4 5 4.69 3 10 211 Note that c NaHA >K a1 W 1 so that the denominator of Equation 13-15 can be simpli- fied. In addition, K a2 c NaHA has a value of 4.69 3 10 212 , which is substantially greater than K w .
- Henry Freiser, Monika Freiser(Authors)
- 1992(Publication Date)
- CRC Press(Publisher)
Chapter 8 Titrations I B r0nsted Acids and Bases Nature of Titrations Titrimetry is an analytical process involving adding portions of a reagent in solution, called a titrant, to react, in a specific and well-defined manner, with a component of an analyte solution, until an exact equivalent (Le., when the titration reaction is exactly complete) of the reagent has been ad de d. This is called the equivalence point. In this respect, titr imetr ic analysis differs from most other types of analysis in which an excess of reagent is adde d. Polyfunctional species will lead to more than one equivalence point in a titration. In this respect, titr imetr ic analysis differs from most other types of analysis in which we add an excess of reagent. Titrimetric methods can be classified according to the type of chemical reaction involved, such as acid-base, precipitation, metal complex formation (complexo metric), oxidation-reduction, among others. In order for a useful, reliable titration to be achieved, the titration reaction must go to completion in a relatively rapid fashion and be free of any interfering side reactions. The substance used as titrant must be readily available in pure, stable form and be reasonable in cost. Further, it must be possible to readily detect the endpoint, the experimental approximation of the equivalence point, by suitable visual or instrumental ind icators. A visual indicator is generally a highly conjugated aromatic organic compound having the characteristic reactivity of the titration type (i.e ., it may be a weak acid, metal complexing agent, reductant, etc.) whose coupled forms (HIn/In, MInlIn, OxlRed) are highly and differently colored. The visual indication of the endpoint depends on observing in a nar row concentration range of the critical variable, the color change of the indicator from one form to the other.
eBook - PDF- Gary D. Christian, Purnendu K. Dasgupta, Kevin A. Schug(Authors)
- 2020(Publication Date)
- Wiley(Publisher)
Pdf _Folio:2 67 267 268 CHAPTER 7 ACID–BASE TITRATIONS 7.1 Strong Acid versus Strong Base—The Easy Titrations An acid–base titration involves a neutralization reaction in which an acid is reacted with an equivalent amount of base. By constructing a titration curve, we can easily explain how the end points of these titrations can be detected. The end point signals the completion of the reaction. A titration curve is constructed by plotting the pH of the solution as a function of the volume of titrant added. The titrant is always a strong acid or a strong base. The analyte may be either a strong base or acid or a weak base or acid. Only a strong acid or base is used as the titrant. In the case of a strong acid versus a strong base, both the titrant and the analyte are completely ionized. An example is the titration of hydrochloric acid with sodium hydroxide: H + + Cl - + Na + + OH - → H 2 O + Na + + Cl - (7.1) The H + and OH - combine to form H 2 O, and the other ions (Na + and Cl - ) remain unchanged, so the net result of neutralization is conversion of the HCl to a neutral solution of NaCl. The titration curve for 100 mL of 0.1 M HCl titrated with 0.1 M NaOH is shown in Figure 7.1, plotted from the spreadsheet exercise setup below. The calculations of titration curves simply involve computation of the pH from the concentration of the particular species present at the various stages of the titration, using the procedures given in Chapter 6. The volume changes during the titration must be taken into account when determining the concentration of the species. Table 7.1 summarizes the equations governing the different portions of the titra- tion curve. We use f to denote the fraction of analyte, which has been titrated by titrant. In Figure 7.1, at the beginning of the titration (f = 0), we have 0.1 M HCl, so the initial pH is 1.0. As the titration proceeds (0 < f < 1), part of the H + is removed from solution as H 2 O.
eBook - PDFOscillometry and Conductometry
International Series of Monographs on Analytical Chemistry
- E. Pungor, R. Belcher, L. Gordon(Authors)
- 2016(Publication Date)
- Pergamon(Publisher)
Because of this, the minimum value of con-ductivity is shifted — in good agreement with Gilbert's theory — to lower degrees of titration. The above discussion indicates that the determination of polybasic acids cannot be carried out by measuring the mini-FIG. 92. Titration curve of tartaric acid as related to dielectric constant Medium: / . distilled water, II. 25% ethanol, III. 50% ethanol mum conductivity values; the concentration can be computed only by using the inflections corresponding to the complete titration. Owing to the fact that at the final stage of the deter-mination of a polybasic acid it is a weak acid that has to be titra-ted, it is a convenient practice to choose a titrant which gives rise to titration curves that coincide with the salt-line after the end point has been reached. Accordingly, polybasic acids are preferably titrated with weak bases. Generally, it is easy to find a suitable weak acid, because they are readily soluble in water. However, it is a more difficult problem to find an adequate weak base, since these are generally only sparingly soluble. Often, the free bases precipitate from their aqueous solutions (e.g. alkaloids), or else the preparation of adequate standard solutions is difficult, on account of the 136 OSCILLOMETRY AND CONDUCTOMETRY high vapour pressure of the reagent. As can be seen from the literature, a number of authors have tried to use ammonia so-lution [181] as a possible titrant utilizing a base of low disso-ciation constant. Unfortunately, owing to its high vapour press-ure, it is impossible to prepare standard ammonia solutions. By means of oscillometric titrations performed in open titration vessels, we have found that the titre of solutions containing excess ammonia changed by as much as 10-20 per cent within 5 to 10 min, depending upon the speed of agitation, the shape of the titration vessel, etc.
eBook - PDF- Gary D. Christian, Purnendu K. Dasgupta, Kevin A. Schug(Authors)
- 2013(Publication Date)
- Wiley(Publisher)
If we know the concentration of an acid, it is rather trivial by charge balance principles to calculate how much base we need to add to get to a certain pH. Consider the titration of V A mL of a triprotic acid of molar concentration C A being titrated by C B molar NaOH, V B mL having been added at any point. Charge balance requires that [Na + ] = [OH − ] + C A (α 1 + 2α 2 + 3α 3 ) − [H + ] If V B mL of NaOH has been added with the total volume being V A + V B , the above equation becomes V B C B /(V A + V B ) = [OH − ] − [H + ] + V A C A (α 1 + 2α 2 + 3α 3 )/(V A + V B ) Multiply both sides by (V A + V B ): V B C B = (V A + V B )([OH − ] − [H + ]) + V A C A (α 1 + 2α 2 + 3α 3 ) Separating the V A and V B multipliers on the first term on the right and transposing: V B (C B − ([OH − ] − [H + ])) = V A ([OH − ] − [H + ]) + V A C A (α 1 + 2α 2 + 3α 3 ) or V B = V A {([OH − ] − [H + ]) + C A (α 1 + 2α 2 + 3α 3 )}/(C B + [H + ] − [OH − ]) (8.26) If we know the dissociation constants and we specify pH (and hence [H + ]), we can easily calculate the alpha values and all terms on the right side of Equation 8.26 are readily calculated. Let us illustrate this with 0.10 molar solution of H 3 PO 4 (K a1 = 1.1 × 10 −2 , K a2 = 7.5 × 10 −8 , K a3 = 4.8 × 10 −13 ) with 0.17 M NaOH. Refer to the spreadsheet “Sec. 8.11 derivative titrations easy method.xlsx” in the text website for the following discussion. In cells A1:B7, we put in the values of C A , C B , V A , K a1 , K a2 , K a3 (these named as KAA, KAB, and KAC so Excel will accept these names) and K w and define these names. Titles are put in cells D8:K8, for V B , pH, [H + ], Q, α 1 , α 2 , α 3 and [OH − ]. In row 9, we enter any trial value for pH in cell E9, express cell F9 as 10 ∧ −E9 and Q, α 1 , α 2 , α 3 as their customary expressions, respectively, in cells G9:J9 in terms of the K a ’s and [H + ] (F9). Finally in cell K9, we express [OH − ] as K w /[H + ]. Only this first row of
eBook - PDF- Vu Dang Hoang(Author)
- 2017(Publication Date)
- IntechOpen(Publisher)
1. Introduction Currently, there are studies that examine the progress of an acid-base titration for one or various polydonor systems, extending sometimes this study to the theme of buffer capacity [1 – 16]. In the scientific literature, there are algorithms and simulators to construct acid-base titration curves, even considering a wide range of different mixtures of polydonor systems [17 – 20]. © 2017 The Author(s). Licensee InTech. This chapter is distributed under the terms of the Creative Commons Attribution License (http://creativecommons.org/licenses/by/3.0), which permits unrestricted use, distribution, and reproduction in any medium, provided the original work is properly cited. The buffer solutions have a certain buffering capacity that is used to maintain constant the pH of a system, having only a small uncertainty. The buffer capacity , β , has been defined as the quantity of strong acid or strong base (in the buffer solution) that gives rise to a change of one pH unit in 1 L of solution, as an intensive property of the system [15]. This involves using directly the concentration of either a strong base or an acid in the buffer solution, without considering the dilution effect, as King and Kester [2], Segurado [3], Urbansky and Schock [4], De Levie [8] did, among others. Urbansky and Schock also mentioned the use of concentration to simplify the maths. Nevertheless, the dilution effect on buffer capacity was first considered by Micha ł owski, as Asuero and Micha ł owski have established in a thorough and holistic review [6]. The buffer capacity considering the effect of dilution, β dil , is defined as the added amount of strong base or strong acid required to change in one unit the pH of an initial V o volume of the buffer solution formed by species of only one polydonor system [7]. By their definition, β is an intensive property by considering the concentration, while β dil is an extensive property to include the amount of substance.
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