Chemistry
Buffer Capacity
Buffer capacity refers to the ability of a buffer solution to resist changes in pH when an acid or base is added. It is determined by the concentrations of the components of the buffer and is a measure of the buffer's effectiveness in maintaining a stable pH. A higher buffer capacity indicates that the solution can withstand larger additions of acid or base without significant changes in pH.
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11 Key excerpts on "Buffer Capacity"
- Chavan, U D(Authors)
- 2021(Publication Date)
- Daya Publishing House(Publisher)
This is due to theself-ionization of water and is independent of the presence or absence of buffering agents. The Buffer Capacity of a buffering agent is at a local maximum when p[H + ] = pK a . It falls to 33% of the maximum value at p[H + ] = pK a ± 1 and to 10% at p[H + ] = pK a ± 1.5. For this reason the useful range is approximately pK a ± 1. Buffer Capacity can be defined in many ways. You may find it defined as “maximum amount of either strong acid or strong base that can be added before This ebook is exclusively for this university only. Cannot be resold/distributed. Buffer Solution 137 a significant change in the pH will occur”. This definition - instead of explaining anything - raises a question “what is a significant change?” – sometimes even change of 1 unit doesn’t matter too much, sometimes - especially in biological systems - 0.1 unit change is a lot. Buffer Capacity can be also defined as quantity of strong acid or base that must be added to change the pH of one liter of solution by one pH unit. Such definition - although have its practical applications - gives different values of Buffer Capacity for acid addition and for base addition (unless buffer is equimolar and its pH=pK a ). This contradicts intuition - for a given buffer solution its resistance should be identical regardless of whether acid or base is added. Buffer Capacity definition that takes this intuition into account is given by b = dpH dn where n is number of equivalents of added strong base (per 1 L of the solution). Note that addition of dn moles of acid will change pH by exactly the same value but in opposite direction. We will derive formula connecting Buffer Capacity with pH, pK a and buffer concentration. To make further calculations easier let’s assume that the strong base added is monoprotic, we also assume volume of 1 which will allow us to treat concentration and number of moles interchangeably.
eBook - PDF- F. R. Whatley, M. J. Kozioł(Authors)
- 2016(Publication Date)
- Butterworth-Heinemann(Publisher)
The few estimates of intracellular buffer capacities reported for plants are also critically examined. After reviewing the known effects on these parameters of exposures to acidic and basic gaseous substances, the potentially damaging repercussions of alterations in intracellular pH and buffer capabilities are considered. It is hoped that the background and scope provided in this article will serve to challenge and stimulate researchers to design experiments to test the fundamental importance of such induced intracellular pH-related changes. Basic concepts Buffer Capacity Qualitatively, a buffer is defined as a solution which resists changes in pH despite the addition of substantial quantities of acid or base. More formally, the buffer index or 313 314 Modification of plant cell buffering capacities by gaseous air pollutants 0.07 0.06 ^ 0.05 a. 0.04 ^ 0.03 0.02 0.01 Acetic acid ( C T = 0.10 M) pK =4.75 -L Fraction titrated û L H1.4 1.2 a 1.0 ^ ω CD 0.8 ë c ο 0.6 S LL 0.4 0.2 0 1.0 2.0 3.0 4.0 5.0 6.0 7.0 8.0 9.0 10.0 11.0 12.0 pH — Figure 21.1. Demonstration of the expression in equation (1) relating the Buffer Capacity β to the slope of a Potentiometrie titration curve. To facilitate the comparison, the axes of the titration curve have been reversed since, by convention, pH is plotted versus volume of base added ( = a, the fraction titrated). C T denotes the sum of the concentrations of the protonated and unprotonated forms capacity is referred to as the 'inverse slope'. This term originates from its relation to the slope of an acid/base titration curve as illustrated in Figure 21.1. Thus ß = ^ or β = -^ -(1) μ dpH μ dpH ; with β the buffer index, d C b and dC a respectively the increments of strong base or acid added, and dpH the corresponding incremental change in pH. For a monoprotic weak acid, the expression summarized in (2) may be derived from first principles (Van Slyke, 1922; Butler, 1964; Bates, 1973).
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.
eBook - PDF- Christian Surber, Christoph Abels, Howard Maibach, Christian, Surber, Christoph, Abels, Howard, Maibach, Howard, Maibach, P. Itin, G. B. E. Jemec, P., Itin, G.B.E., Jemec(Authors)
- 2018(Publication Date)
- S. Karger(Publisher)
Buffers are solutions containing a weak or me-dium strong acid (base) and the salt of the cor-responding base (acid) in about similar concen-trations. Both parts of the buffer mixed equally (in the ratio 1: 1) results in an optimal buffering capacity. Variations of this relationship are pos-sible. However, the buffering effect will be re-duced, if the pK a ±1 value will be left. According to Henderson-Hasselbalch-Equation: pH = pKs + lg (c[A–]/c[HA]), the relation should be not less than 10: 1 or 1: 10, due to dramatically re-duced Buffer Capacity. pH Buffer in Nature There is a great variety of the pH and buffers in the environment. Acid rain as the result of indus-trial produced sulfur and nitrogen compounds leading to sulfuric and nitric acid, can cause dam-age to farm and woodlands and the growth of plants will be greatly hindered. The pH of neutral soil varies from 6.6 to 7.3. H 2 CO 3 is the most im-portant source for H + -ions and carbonate-con-taining buffers keep the pH constant for a long time [1]. Forest liming, which is the application of calcium- and magnesium-rich carbonate, is per-formed to fight acidification and to increase buf-fer capacity. pH Buffers in Humans pH within the human body varies from pH 1 to 8 depending on the organ and the function. Out-side the acceptable range of pH, enzymes lose their ability to function and proteins are dena-tured. There are several buffering agents in the body that impede any change in pH. Extracellular buffers include bicarbonate and ammonia, where-as proteins and phosphate act as intracellular buf-fers. Respiratory compensation and renal com-pensation are 2 ways by which the body can ex-crete excess molecules to maintain pH. Acid-base imbalances overcoming the buffer system can be compensated in the short term by changing the rate of ventilation, thereby expelling CO 2 and re-sulting in fewer free hydrogen ions and the pH will rise back to normal.
eBook - PDF- F. John Gennari, Horacio J. Adrogue, John H. Galla, Nicolaos Maddias, F. John Gennari, Horacio J. Adrogue, John H. Galla, Nicolaos Maddias(Authors)
- 2005(Publication Date)
- CRC Press(Publisher)
With continued addition of either a strong acid or base to a solu-tion, the corresponding weak acid or base eventually becomes saturated. In the example above, this happens when all available base (HPO 4 2 ) combines with the added H þ to form the acid, H 2 PO 4 . When this occurs, all further added H þ will remain in solution, rapidly lowering the pH until H 2 PO 4 becomes a proton acceptor (pH < 3.0). This behavior leads to the typical sigmoid titration curves in Fig. 2. Quantitation of Buffer Effect on Acidity Figure 2 depicts the titration curves of three buffers: acetic acid = acetate, uric acid = monourate, and monobasic phosphate = dibasic phosphate. The mid-point of each of the curves denotes the setting in which the concentrations of the acid and its conjugate base are equal. From Eqs. (9) and (10), one can see that when these concentrations are equal, [H þ ] ¼ K 0 or pH ¼ p K 0 . As illustrated in Fig. 2, the buffering ability or capacity is also the greatest when pH ¼ p K 0 . Acid–Base Chemistry and Buffering 7 The capacity of a specific buffer to resist a change in pH is defined as its buffer value ( b ), the first derivative of the titration curve: b ¼ d ð acid or base Þ = dpH ð 14 Þ Buffer value is conventionally expressed as mol = L or mmol = L of strong acid or base added per unit change in pH. Given the curves in Figure 2, it is apparent that b varies with pH and reaches a maximum when pH ¼ p K 0 . This unit of measure has been named the slyke, in honor of Donald D. Van Slyke (15), a pioneer in the study of acids and bases in biological fluids. In a seminal paper, he deduced that b is a function of the concentration of the weak acid in a solution (C), the K 0 (or p K 0 ) of the buffer, and the prevail-ing [H þ ] (16). He also showed that for any weak acid, the maximal value for b (i.e., when [H þ ] ¼ K 0 or pH ¼ p K 0 ) is 0.575.
eBook - PDF- Krish Moorthy(Author)
- 2007(Publication Date)
- CRC Press(Publisher)
Fundamentals of Biochemical Calculations 56 Buffers A buffer solution is one that resists a change in pH on the addition of an acid or alkali. Buffers are solutions of a weak acid and one of its salts (the conjugate base) or a weak base and one of its salts (a conjugate acid). The most useful equation for dealing with quantitative aspects of buffers is the Henderson-Hasselbalch equation: pH = pK a + log 10 [A ] [HA] -All the terms in this equation have been previously described. The [ ] strictly mean molar concen-trations, but as a ratio is involved, any chemical concentration unit will do. Students must gain a good feel for the whole equation as well as the individual terms; it can be quite tricky because ratios and negative logs are involved and also because chemical terms, such as base and salt, could mean the same thing. Please check the following: pH is the variable term, the pH of the required buffer. pK a is fixed once the buffer system is chosen (e.g., for acetate buffer, pK a = 4.76). [A ] [HA] -= [salt] [acid] = [base] [acid] = [non -protonated] [protonated] The last expression is particularly descriptive when several protonated or deprotonated groups are involved, such as with amino acid structures. The ratio [A ] [HA] -or [salt] [acid] is the most important term in the Henderson-Hasselbalch equation. It is this term that determines the pH of the buffer. It is this term that students should carefully evaluate in buffer calculations and preparations. For a start, note that log of 1 = 0, log of a number > 1 is positive, and log of a number < 1 is negative. It therefore follows that: when [salt] = [acid], pH = pK (with the acetate buffer system, the pH of the buffer would be 4.76) when [acid] > [salt], pH < pK (acetate buffer, pH would be < 4.76) when [acid] < [salt], pH > pK (acetate buffer, pH would be > 4.76). The simple rule of thumb is: the greater the [acid], the lower the pH of the buffer.
eBook - ePub- Professor Rob Beynon, J Easterby(Authors)
- 2004(Publication Date)
- Taylor & Francis(Publisher)
The buffer at pH 3.81 (Case A), with the greatest imbalance between [acid] and [base] shows the biggest pH shift. The difference is substantial—over 0.2 pH units between these two buffers. Thus, our expectations are confirmed—the more equally balanced the concentrations of the two species, the ‘better’ the buffer.Note that the buffer in Case A is actually now better placed to resist a second addition of NaOH, because the two species are now more equally matched in concentration. As a general rule of thumb, this can guide buffer selection (Chapter 5 ) because if we have a system that generates protons, we should use a buffer with a pKa slightly on the alkaline side of the pH we need. Then, as protons are generated, they shift the equilibrium to improve the balance between the buffer species, and the buffer becomes more resistant to proton additions. The converse arguments apply for a proton-consuming system, of course.Thus, we should use buffers at or near their pKa values. But intuition will tell us that the ability of a buffer system to resist changes must be influenced by the buffer concentration as well as the pKa . For example, had we used a total of 0.2 M [acetate– ] + [acetic acid] in the previous worked example, the pH changes (ΔpH) would have been approximately 0.35 (Case A) and 0.17 (CaseB)—much smaller pH changes. Thus, three factors influence our choice of buffer. First is the pH that we need to maintain. Secondly, we should know whether our system generates or consumes protons. Thirdly, we should have some idea of the concentration of protons that are generated or consumed in our system. Already, the choice of a buffer is becoming more complicated!3. How good is a buffer?—β valuesWe can define the ability of a buffer to resist pH changes in terms of ‘buffering capacity’ or β. This term measures how well the buffer works. The definition of β
- Frank R. Foulkes(Author)
- 2012(Publication Date)
- CRC Press(Publisher)
CHAPTE R TWENTY-TW O BUFFER SOLUTIONS 22.1 BUFFER SOLUTIONS In our ionic equilibrium calculations thus far, we have been solving for pH at equilibrium. The rea-son pH is so important is that a great number of chemical and biochemical processes only operate satisfactorily if the pH is held within certain narrow limits. For example, to name just a few, the pH of the medium affects the characteristics of electroplated deposits; the reactivity of enzymes; the rate of metallic corrosion; the permeability of cell membranes; the efficiency of fermentations to produce beer, wine, and alcohol; the precipitation of various substances; and the growth of micro-organisms and plants. In fact, the human body itself is full of controlled pH processes: The pH of the blood should be held between 7.30 and 7.45; if your blood pH falls below 6.8 or rises above 7.8 your body enters a state known as––“death.” The pH of blood plasma should be maintained between 7.38 and 7.41; the pH of saliva usually is about 6.8; the pH within the duodenum must be held between 6.0 and 6.5; for proper digestion the pH of the gastric juices within the stomach must be kept between 1.6 and 1.8. The body maintains these various pH ranges, as needed, by means of chemical constitu-ents that resist pH change when small amounts of acid or base are added. Solutions with such regulatory pH power are called buffer solutions . Buffer solutions contain relatively large amounts of either (a) a weak acid and its salt––this kind of buffer stabilizes pH < 7, or (b) a weak base and its salt––this kind of buffer stabilizes pH > 7.
- J. A. Beran, Mark Lassiter(Authors)
- 2022(Publication Date)
- Wiley(Publisher)
What is the pH of blood and what are the blood buffers that maintain that pH? In water habitats, buffers make existence of life possible. Potentiometric titrations are also utilized for the determination of any ion’s concentration when specific ion electrodes are used, so monitoring ions in various media is also possible. Studies of alkalinity in water of natural systems and industrial water supplies also have interestingly applications. How might your understanding the role of buffers in chemical systems and gaining the skills in using the tools for these measurements, in- form and expand your discovery in things that interest you? *At the end of most experiments, the “Next Steps” section gives valuable applications of the experi- ment and challenges the student to expand the concepts of that experiment. Next Steps to Discovery* Experiment 17 243 Experiment 17 LeChâtelier’s Principle; Buffers* 2 CrO 4 2– (aq) + 2H + (aq) Cr 2 O 7 2– (aq) + H 2 O(l) Jo A. Beran Objectives Techniques and Safety Practices The chromate ion (left) is yellow, and the dichromate ion (right) is orange. An equilibrium between the two ions is affected by changes in pH. • To study the effects of concentration and temperature changes on the position of equilibrium in a chemical system • To study the effect of strong acid and strong base addition on the pH of buffered and unbuffered systems • To observe the common-ion effect on a dynamic equilibrium The following techniques and safety practices are used in the Experimental Procedure: Introduction and Applications Most chemical reactions do not produce a 100% yield of product, not because of ex- perimental technique or design, but rather because of the chemical characteristics of the reaction. The reactants initially produce the expected products, but after a period of time the concentrations of the reactants and products stop changing.
eBook - PDF- Young, William Vining, Roberta Day, Beatrice Botch(Authors)
- 2017(Publication Date)
- Cengage Learning EMEA(Publisher)
WCN 02-300 Unit 18 Advanced Acid–Base Equilibria 573 a human body occur in aqueous solutions containing buffering agents. It is not surprising that human blood is highly buffered, for if blood is not maintained at a pH near 7.4 , death can occur. A buffer solution contains a mixture of a weak acid and a weak base, typically the conjugate base of the weak acid. For example, a buffer solution commonly used in chem-istry laboratories contains both acetic acid ( CH 3 CO 2 H , a weak acid) and sodium acetate ( NaCH 3 CO 2 , the sodium salt of the conjugate base of acetic acid). Some other examples of buffers are KH 2 PO 4 > K 2 HPO 4 ( H 2 PO 4 2 is a weak acid and HPO 4 2 2 is a weak base) and NH 4 Cl > NH 3 ( NH 4 1 is a weak acid and NH 3 is a weak base). The principle property of a buffer solution is that it experiences a relatively small change in pH when a strong acid or a strong base is added. It is the weak acid and weak base components of a buffer that make it possible for buffer solutions to absorb strong acid or strong base without a significant pH change. ● ● When a strong acid is added to a buffer, the acid reacts with the conjugate base and is completely consumed. Despite the addition of a strong acid, the pH of the buffer solu-tion decreases only slightly. Example: When H 3 O 1 is added to a CH 3 CO 2 H > NaCH 3 CO 2 buffer, it consumes some of the conjugate base, forming additional acetic acid. H 3 O 1 (aq) 1 CH 3 CO 2 2 (aq) S H 2 O( / ) 1 CH 3 CO 2 H(aq) ● ● When a strong base is added to a buffer, the base reacts with the weak acid and is com-pletely consumed. Despite the addition of a strong base, the pH of the buffer solution increases only slightly. Example: When OH 2 is added to a CH 3 CO 2 H > NaCH 3 CO 2 buffer, it consumes some of the weak acid, forming additional acetate ion. OH 2 (aq) 1 CH 3 CO 2 H(aq) S H 2 O( / ) 1 CH 3 CO 2 2 (aq) Example Problem 18.2.1 Identify buffer solutions. Identify buffer solutions from the following list. a.
eBook - PDF- Gary D. Christian, Purnendu K. Dasgupta, Kevin A. Schug(Authors)
- 2020(Publication Date)
- Wiley(Publisher)
A number, partic- ularly specific enzyme reactions that might be used for analyses (see Chapter 23), may occur in the pH range of 4 to 10 or even outside of this. The proper selection of buffers Pdf _Folio:2 47 248 CHAPTER 6 ACID–BASE EQUILIBRIA for the study of biological reactions or for use in clinical analyses can be critical in determining whether or not they influence the reaction. A buffer must have the correct pK a , near physiological pH so the ratio of [A - ]∕[HA] in the Henderson–Hasselbalch equation is not too far from unity, and it must be physiologically compatible. Phosphate Buffers One useful series of buffers are phosphate buffers. Biological systems usually con- tain some phosphate already, and phosphate buffers will not interfere in many cases. By choosing appropriate mixtures of H 3 PO 4 ∕H 2 PO 4 - , H 2 PO 4 - ∕HPO 4 2- , or HPO 4 2- ∕PO 4 3- , solutions over a wide pH range can be prepared. See G. D. Christian and W. C. Purdy, J. Electroanal. Chem., 3 (1962) 363 for the compositions of a series of phosphate buffers at a constant ionic strength of 0.2. Ionic strength is a mea- sure of the total salt content of a solution (see Chapter 5), and it frequently influences reactions, particularly in kinetic studies. Hence, these buffers could be used in cases where the ionic strength must be constant. However, the buffering capacity decreases markedly as the pH approaches the values for the single salts listed, and the single salts are not buffers at all. The best buffering capacity, obtained at the half neutralization points, is within ±1pH unit of the respective pK a values, that is, 1.96 ± 1, 7.12 ± 1, and 12.32 ± 1. Example 6.24 What weights of NaH 2 PO 4 and Na 2 HPO 4 would be required to prepare 1 L of a buffer solution of pH 7.45 that has an ionic strength of 0.100? Solution Let x = [Na 2 HPO 4 ] and y = [NaH 2 PO 4 ]. There are two unknowns, and two equations are needed.
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