Chemistry
pH Curves and Titrations
pH curves and titrations are used to determine the concentration of an unknown solution by reacting it with a solution of known concentration. The pH curve shows the change in pH as the titrant is added, and the equivalence point is where the moles of acid and base are equal. Titrations are important in determining the acidity or basicity of a solution.
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12 Key excerpts on "pH Curves and Titrations"
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.
- Henry Freiser, Monika Freiser(Authors)
- 1992(Publication Date)
- CRC Press(Publisher)
Further, precise algebraic calculations of the endpoint can be carried out readily. Both of these types are useful. The linear curve is particular ly practical when dilute solutions are titrated. Logarithmic curves are convenient when the concentration of the critical variable changes by factors of more than a thousand. The manner of locating endpoints also differs in these two types of titration curves as will be explained below. Traditionally, titration curve calculations are described in terms of equations that are valid only for parts of the titration. Equations wil l be developed here that reliably describe the entire curve. This will be done first for acid-base titration curves. In following chapters, titration curves for other reaction systems (metal complexation, redox, precipitation) wil l be developed and characterized in a similar fashion. For all, graphical and algebraic means oflocating the endpoints wil l be described, colorimetric indicators and how they function will be explained, and the application of these considerations to ( 1 ) calcu lation of titration errors, (2) buffer design and evaluation, (3) sharpness oftitrations, and finally, (4) in Chapter 1 8, the use of titration curve data to the determination of equilibrium constants wil l be presented. In acid-base titration, the appropriate concentration variable is [H + ] or, most commonly, pH. We have already developed al l the necessary equations earlier (Chapter 4). The only difference is that now, instead of writing a proton balance equation (PBE) for a single set of conditions, these apply to a whole family of points, i.e ., those involved in the entire titration. 154 Concepts &. Calculations in Analytical Chemistry Use of PSE to Derive Titration Curves 1. Strong acid -strong base. Let us add a specified volume, V A mL, of a solution of C A M HCI to a vessel and titrate this with C B M NaOH.
eBook - PDF- Gary D. Christian, Purnendu K. Dasgupta, Kevin A. Schug(Authors)
- 2013(Publication Date)
- Wiley(Publisher)
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 7. The volume changes during the titration must be taken into account when determining the concentration of the species. Table 8.1 summarizes the equations governing the different portions of the titration curve. We use f to denote the fraction of analyte, which has been titrated by titrant. In Figure 8.1, 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. So the concentration of H + gradually decreases. At 90% neutralization (f = 0.9) (90 mL NaOH), only 10% of the H + remains. Neglecting the volume change, the H + concentration at this point would be 10 −2 M, and the pH would have risen by only one pH unit. (If we correct for volume change, it will be slightly higher—see the spreadsheet below.) However, as the equivalence point The equivalence point is where the reaction is theoretically complete. is approached (the point at which a stoichiometric amount of base is added), the H + concentration is rapidly reduced until at the equivalence point (f = 1), when the neutralization is complete, a neutral solution of NaCl remains and the pH is 7.0. As we continue to add NaOH (f > 1), the OH − concentration rapidly increases from 10 −7 M at the equivalence point and levels off between 10 −2 and 10 −1 M; we then have a solution of NaOH plus NaCl. Thus, the pH remains fairly constant on either side of the equivalence point, but it changes markedly very near the equivalence point. This large change allows the determination of the completion of the reaction by measurement of either the pH or some property that changes with pH (e.g., the color of an indicator or potential of an electrode).
eBook - ePub- John Kenkel(Author)
- 2013(Publication Date)
- CRC Press(Publisher)
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:(5.1)H ++OH −→H 2OAs H+ ions are consumed in the reaction flask by the OH− added from the buret, the pH of the solution in the flask will change, since pH = −log [H+ ]. In fact, the pH should increase as the titration proceeds, since the number of H+ ions decreases due to the reaction with OH− . The lower the [H+ ], the higher is the pH. This increase in pH can be monitored with the use of a pH meter. Thus, if we were to measure the pH in this manner after each addition of NaOH and graph the pH vs. volume (mL) of NaOH added we would have a graphical display of the experiment. Figure 5.1 a shows the results of such an experiment for the case in which 0.10 N HCl is titrated with 0.10 N NaOH. The graph is called a titration curve . The sharp increase in the pH at the center of the graph occurs at the equivalence point of the titration or the point at which all the acid in the flask has been neutralized by the added base. The point at which the slope of a titration curve is a maximum (a sharp change in pH such as this, whether an increase or decrease) is called an inflection point .FIGURE 5.1
eBook - PDFFoundations of Chemistry
An Introductory Course for Science Students
- Philippa B. Cranwell, Elizabeth M. Page(Authors)
- 2021(Publication Date)
- Wiley(Publisher)
The volume of acid added is noted and used to calculate the con-centration of the base. Alternatively, the burette is filled with base which is added to the acid in the flask until just neutralised. Accurate values of pH changes occurring during an acid-base titration can be monitored using a pH meter (Figure 7.6). The results from the pH titration can be plotted, as shown in Figure 7.8, where the pH of the solution in the flask is plotted against the volume of hydrochloric acid added. If 0.1 M sodium hydrox-ide is in the flask the initial pH will be 13. Addition of HCl causes a very slight decrease in pH. Close to the neutralisation point the pH changes very rapidly and a very small increase in the volume of acid causes a sharp decrease in pH. This is called the equivalence point of the titration. It is represented by the vertical region on the plot in Figure 7.8. The pH at the equivalence point for the reaction between a strong acid and strong base is pH 7. If 0.1 M HCl is used in the burette, the volume required to completely neutralise 25 cm 3 of 0.1 M NaOH is 25 cm 3 . When a pH meter is used to monitor the reaction, the equivalence point is determined from the mid point of the vertical section on the plot of volume against pH. Conical flas k Burette Figure 7.7 Titration apparatus. 7.3 Acid-base equilibria 237 7.3.7 Indicators In practice, when carrying out a titration, the equivalence point is generally determined using an indicator . An acid-base indicator is a chemical that shows a different colour at different pH values. The colour changes of some commonly used indicators are shown in Figure 7.9. Indicators may also be derived from natural plant materials such as lichens and red cabbage, as shown in Box 7.3. Indicator pH 1 2 3 4 5 6 7 8 9 10 11 12 13 14 Universal indicator Methyl orange Methyl red Bromothymol blue Phenolphthalein change change change change Figure 7.9 Colour changes of some common indicators.
eBook - PDF- Paul Flowers, Klaus Theopold, Richard Langley, William R. Robinson(Authors)
- 2015(Publication Date)
- Openstax(Publisher)
The pH of the solutions may be calculated using familiar equilibrium techniques, or it may be qualitatively determined to be acidic, basic, or neutral depending on the relative K a and K b of the ions involved. 14.5 Polyprotic Acids An acid that contains more than one ionizable proton is a polyprotic acid. The protons of these acids ionize in steps. The differences in the acid ionization constants for the successive ionizations of the protons in a polyprotic acid usually vary by roughly five orders of magnitude. As long as the difference between the successive values of K a of the acid is greater than about a factor of 20, it is appropriate to break down the calculations of the concentrations of the ions in solution into a series of steps. 14.6 Buffers A solution containing a mixture of an acid and its conjugate base, or of a base and its conjugate acid, is called a buffer solution. Unlike in the case of an acid, base, or salt solution, the hydronium ion concentration of a buffer solution does not change greatly when a small amount of acid or base is added to the buffer solution. The base (or acid) in the buffer reacts with the added acid (or base). 14.7 Acid-Base Titrations A titration curve is a graph that relates the change in pH of an acidic or basic solution to the volume of added titrant. The characteristics of the titration curve are dependent on the specific solutions being titrated. The pH of the solution at the equivalence point may be greater than, equal to, or less than 7.00. The choice of an indicator for a given titration depends on the expected pH at the equivalence point of the titration, and the range of the color change of the indicator. Chapter 14 | Acid-Base Equilibria 825 Exercises 14.1 Brønsted-Lowry Acids and Bases 1. Write equations that show NH 3 as both a conjugate acid and a conjugate base. 2. Write equations that show H 2 PO 4 − acting both as an acid and as a base.
eBook - PDFChemistry
An Atoms First Approach
- Steven Zumdahl, Susan Zumdahl, Donald J. DeCoste, , Steven Zumdahl, Steven Zumdahl, Susan Zumdahl, Donald J. DeCoste(Authors)
- 2020(Publication Date)
- Cengage Learning EMEA(Publisher)
The pH curve is shown in Fig. 14.5. 608 CHAPTER 14 Acid–Base Equilibria Copyright 2021 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. 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. CRITICAL THINKING You have read about titrations of strong acids with strong bases, weak acids with strong bases, and weak bases with strong acids. What if you titrated a weak acid with a weak base? Sketch a pH curve and defend its shape. Label the equivalence point and discuss the pos- sibilities for the pH value at the equivalence point. 14.5 Acid–Base Indicators There are two common methods for determining the equivalence point of an acid–base titration: 1. Use a pH meter (see Fig. 13.7) to monitor the pH and then plot the titration curve. The center of the vertical region of the pH curve indicates the equivalence point (for example, see Figs. 14.1 through 14.5). 2. Use an acid–base indicator, which marks the end point of a titration by chang- ing color. Although the equivalence point of a titration, defined by the stoichiom- etry, is not necessarily the same as the end point (where the indicator changes color), careful selection of the indicator will ensure that the error is negligible. The most common acid–base indicators are complex molecules that are themselves weak acids (represented by HIn). They exhibit one color when the proton is attached to the molecule and a different color when the proton is absent. For example, phenol- phthalein, a commonly used indicator, is colorless in its HIn form and pink in its In 2 , or basic, form. The actual structures of the two forms of phenolphthalein are shown in Fig.
eBook - PDF- Douglas Skoog, Donald West, F. Holler, Stanley Crouch, Douglas Skoog(Authors)
- 2021(Publication Date)
- Cengage Learning EMEA(Publisher)
❯ Notice at all points prior to the equivalence point, H 3 O 1 and thus pH is controlled entirely by the ratio of the acid concentration to the concentration of its conjugate base. The only exception is the initial point before any titrant is added. ❯ Since the principal solute species at the equivalence point is HCN, the pH is acidic. ❯ When you titrate a weak base, use an indicator with a mostly acidic transition range. When titrating a weak acid, use an indicator with a mostly basic transition range. ❯ Figure 12-8 shows hypothetical titration curves for a series of weak bases of different strengths. The curves show that indicators with mostly acidic transition ranges must be used for weak bases. 50 60 Volume of 0.1000 M HCl, mL pH 0 10 20 30 Strong base K b = 10 –6 K b = 10 –8 K b = 10 –10 K b = 10 –4 K b = 10 –2 40 12.0 10.0 8.0 6.0 4.0 2.0 Phenolphthalein transition range Bromothymol blue transition range Bromocresol green transition range FIGURE 12-8 The effect of base strength (K b ) on titra- tion curves. Each curve represents the titration of 50.00 mL of 0.1000 M base with 0.1000 M HCl. Copyright 2022 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. 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. 12D Titration Curves for Weak Bases 305 Determining the pK Values for Amino Acids Amino acids contain both an acidic and a basic group. For example, the structure of alanine is represented in Figure 12F-1. The amine group behaves as a base, and at the same time the carboxyl group acts as an acid.
- J. A. Beran, Mark Lassiter(Authors)
- 2022(Publication Date)
- Wiley(Publisher)
The readout of the pH meter is E cell , expressed in volts, but since E cell is directly proportional to pH (equation 19.1), the meter for the readout is expressed directly in pH units. The selection of an indicator for the titration of a strong acid with a strong base is rel- atively easy in that the color change at the stoichiometric point always occurs at a pH of 7 (at 25°C). Usually, phenolphthalein can be used because its color changes at a pH close to 7. However, when a weak acid is titrated with a strong base, the stoichio- metric point is at a pH greater than 7, and a different indicator may need to be selected. 1 If the weak acid is an unknown acid, then the proper indicator cannot be selected be- cause the pH at the stoichiometric point cannot be predetermined. The color change of a selected indicator may not occur at (or even near) the pH of the stoichiometric point for the titration. To better detect a stoichiometric point for the titration of an unknown weak acid, a pH meter is more reliable. In Part A of this experiment, a titrimetric analysis is used to determine the molar con- centration of a weak acid solution. A pH meter is used to detect the stoichiometric point of the titration. An acid–base indicator will not be used. A standardized sodium hydroxide solution is used as the titrant. 2 The pH of a weak acid solution increases as the standardized NaOH solution is added. A plot of the pH of the weak acid solution as the strong base is being added, pH versus V NaOH , is called the titration curve (Figure 19.2) for the reaction. The inflection point in the sharp vertical portion of the plot (about midway on the vertical rise) is the stoichiometric point. Molar Concentration of a Weak Acid Solution Titrimetric analysis: a titration procedure that is chosen for an analysis Titration curve: a data plot of pH versus volume of titrant 264 Potentiometric Analyses Courtesy of Thermo Fisher Scientific Buffer solutions are used to calibrate pH meters.
eBook - PDFGeneral, Organic, and Biological Chemistry
An Integrated Approach
- Kenneth W. Raymond(Author)
- 2012(Publication Date)
- Wiley(Publisher)
An approach to use for determining the concentration of an unknown acid is shown here. ■ Titration is used to determine the concentration of an unknown acid or base solution. 7.9 Effect of pH on Acid and Conjugate Base Concentrations 259 7.9 EFFECT OF pH ON ACID AND CONJUGATE BASE CONCENTRATIONS Some of the compounds important to living things are acids. Whether they are found in their acid or conjugate base form depends on the pH. We have seen that K a and pK a provide information about the relative concentrations of products and reactants in acid–base reactions. For example, the K a for HF is 6.6 * 10 -4 (pK a = 3.18), so at equilibrium, the concentration of HF is greater than the product of the concentrations of F - and H 3 O + . HF + H 2 O N F - + H 3 O + K a = [F - ][H 3 O + ] [HF] = 6.6 * 10 -4 ; pK a = 3.18 We have also seen that Le Châtelier’s principle allows us to predict how changing the concentration of a reactant or product will be followed by a change in the rate of the for- ward or reverse reaction as equilibrium is reestablished. For the reaction above, increasing the concentration of H 3 O + (lowering the pH) will cause the reverse reaction to speed up, which consumes F - and produces HF. Reducing the concentration of H 3 O + (raising the pH) will cause the forward reaction to speed up, which consumes HF and produces F - . Thus, the relative amounts of HF and F - present in a solution depend on the pH. An interesting relationship exists between pH and the relative concentrations of an acid and its conjugate base: • When the pH is adjusted to have the same value as the pK a (pH = pK a ), the concentration of acid equals the concentration of its conjugate base. For the acid HF (pK a = 3.18), this means that [HF] = [F - ] when the pH is 3.18 (Table 7.6). STRATEGY Calculating the concentration of the HCl solution involves determining the number of moles of HCl present in the initial 75 mL of solution.
eBook - PDF- Saeed Farrokhpay(Author)
- 2023(Publication Date)
- Arcler Press(Publisher)
Thus, in titrations employing strong acids, totally ionized acid is retained at all titrant doses; the pH at every moment during the titration is equivalent to the residual acid’s (H + ) (Figure 6.4). On contrary to strong acid titration, whenever weak acids are titrated using a strong base, the system contains a combination of the acid as well as its conjugate base. It has the effect of cushioning the solution against sudden pH fluctuations. The ability of a solution to withstand pH variation is referred to as buffering. If a weaker acid, as well as its conjugate base, are existed in a similar medium, buffering happens. Due to buffering, the graph of pH versus titrant concentration for weaker acids is very complicated as compared to stronger acids. Although, the Henderson-Hasselbalch equation may be used to approximate this correlation (Tarakçi and Kucukoner, 2003): (14) where; (HA) is the non-ionized acid concentration; (A – ) is the conjugate base concentration; pK a is the pH when non-ionized acid and conjugate base are equivalent. The equation says maximal buffering capacity is reached when pH = pK a . It is illustrated by a graph of 0.1 Normality acetic acid titration with 0.1 Normality sodium hydroxide (Figure 6.5). Di- and triprotic acids have 2 and 3 buffering areas, correspondingly. Figure 6.6 shows a pH versus citric acid titrant curve. The Henderson- Hasselbalch equation may forecast the plateau correlating to every pK a stage in polyprotic acids. Protons and conjugated bases from different ionization states (s) hinder the transition zone among stages. So, the Henderson- Acidity and Basicity in Chemistry 170 Hasselbalch equation fails around the equivalency of 2 pK a stages. But calculating the equivalent pH is simple. The pH = (pK a 1 + pK a 2)/2. These values are listed in Table 6.4 (Ryan and Irwin, 2002). The Henderson-Hasselbalch equation needs all elements to be perfect solutions.
eBook - PDF- Vu Dang Hoang(Author)
- 2017(Publication Date)
- IntechOpen(Publisher)
2014. Available from: http://www2.iq.usp.br/ docente/gutz/Curtipot_.html [Accessed: 03 March 2017] [19] Tarap č ik P, Beinrohr E. Implementation of a Universal Algorithm for pH Calculation into Spreadsheet and its Use in Teaching in Analytical Chemistry. Slovakia: Faculty of Chem-ical and Food Technology STU; 2005 [20] Efstathiou CE. Acid-Base Titration Curves. Department of Chemistry, University of Ath-ens, Greece [Internet]. 2000. Available from: http://195.134.76.37/applets/AppletTitration/ Appl_Titration2.html [Accessed: 03 March 2017] [21] Dean GJA. Lange ’ s Handbook of Chemistry. Chapter 8. New York: McGraw-Hill; 1999. pp. 8 – 113 [22] Rodríguez-Laguna N. Contribuciones teóricas al concepto de capacidad buffer considerando especies insolubles y especies polinucleares en el Sistema [thesis]. Iztapalapa, México, DF: Universidad Autónoma Metropolitana; 2015 [23] Barriada JL, Brandariz I, Sastre de Vicente ME. Acid-base equilibria of monocarboxylic acids in various saline media: Analysis of data using Pitzer equations. Journal of Chem-ical and Engineering. 2000; 45 :1173 – 1178 [24] Turner DR, Correia-dosSantos MM, Coutinho P, Gon ç alves ML, Knox S. Rapid pK mea-surements for multibasic weak acids by gradient flow injection titration. Analytica Chimica Acta. 1992; 258 :259 – 267 [25] Saha A, Saha N, Liang-Lian J, Zhao J, Fridrich G, Sajadi SAA, Song B, Sigel H. Stability of metal ion complexes formed with methyl phosphate and hydrogen phosphate. Journal of Biological Inorganic Chemistry. 1996; 1 :231 – 238 [26] Martínez-Calatayud J, Campins-Falcó P, Micó-Albert R. Flow method for the titration of weak acids or weak bases using linear titration plots. Analyst. 1987; 112 :1063 – 1066 [27] Botello JC, Morales-Domínguez E, Domínguez JM, Gutiérrez A, Rojas-Hernández A, Ramírez-Silva MT. A new nuclear magnetic resonance algorithm to determine equilib-rium constants of the species in the B(III)-H 2 O system.
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