Chemistry
Properties of Buffers
Buffers are solutions that resist changes in pH when an acid or base is added. They consist of a weak acid and its conjugate base, or a weak base and its conjugate acid. The key properties of buffers include their ability to maintain a relatively constant pH, their capacity to neutralize added acids or bases, and their dependence on the ratio of the weak acid and its conjugate base.
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10 Key excerpts on "Properties of Buffers"
- Chavan, U D(Authors)
- 2021(Publication Date)
- Daya Publishing House(Publisher)
DEFINITION/PRINCIPLES AND ABBREVIATIONS USED The function of pH buffers is to neutralize small changes in H + and OH - concentrations, thereby maintaining pH fairly constant. The basics of pH buffering and buffer calculations and preparation are discussed below. The necessary symbols and their meanings are as follows: A weak acid and its conjugate base are represented by HA and A -, respectively; a weak base and its conjugate acid are represented by R-NH 2 and R-NH, respectively, because the most commonly used weak base buffers has NH 2 basic functional group(s). In aqueous solutions, H + is hydrated and, therefore, is written as Hp+ when necessary. Molar concentration is indicated by enclosing the substance of interest in square brackets; for example, [H + ] represents the molar concentration of H + . Buffer a solution that resists pH change is an important for many reactions e.g., enzymatic methods of analysis, etc. Buffer solutions are solutions that resist changes in pH (by resisting changes in hydronium ion and hydroxideion concentrations) upon addition of small amounts of acid or base, or upon dilution. They usually consist of a weak acid and its conjugate base, or, less commonly, a weak base and its conjugate acid. Buffer solutions are used in industry for chemical manufacturing and fermentation processes, and to set the proper conditions for dyeing fabrics. In research laboratories, buffers are used for chemical analyses, syntheses, and calibration of pH meters. In living organisms, these solutions maintain the correct pH for many enzymes to work. Blood plasma contains a buffer (of carbonic acid and bicarbonate) to maintain a pH of approximately 7.4. The main component of a buffer solution, such as a weak acid or weak base, may be used as a buffering agent. The function of a buffering agent is to drive an acidic or alkaline solution to a certain pH and maintain it at that pH.
eBook - ePub- Professor Rob Beynon, J Easterby(Authors)
- 2004(Publication Date)
- Taylor & Francis(Publisher)
Chapter 2 , but these need not concern us here.◊ Nearly all pH buffers are weak acids or bases.Notice that the weak acid can be neutral (acetic acid) or carry a positive (TrisH+ ) or negative (phosphate1– ) charge. As we develop the theory of buffers, it will become clear that these charges on the buffer species have important consequences.2. Weak acids and bases resist pH changesA buffer is able to resist changes in pH because it exists in an equilibrium between a form that has a hydrogen ion bound (conjugate acid, protonated) and a form that has lost its hydrogen ion (conjugate base, deprotonated). For the simple example of acetic acid, the equation is:CH 3COOH ⇌CH 3COO −+H +Here, the protonated form is acetic acid, with a net charge of zero, whereas the deprotonated form (acetate) has a charge of −1. The two species are in equilibrium, and this equilibrium, in common with all equilibria, can be displaced by addition of one component.Consider a solution that contains equal amounts of acetic acid and acetate ions (10 mM acetic acid, 10 mM sodium acetate, for example). If we were to add a strong acid, such as HCl, to this solution, the added H+ would displace the equilibrium to the left. Binding of H+ to CH3 COO– ‘mops up’ the added protons (Figure 3.1 ). Electrical neutrality is preserved because every H+ that reacts with a CH3 COO– anion to form the neutral CH3 COOH leaves behind a chloride (Cl– ) anion in its place. Add a strong base, such as sodium hydroxide, and the OH- ion would react with the H+ and displace the equilibrium to the right. Electrical neutrality in the solution is sustained because for every CH3 COOH that is converted to CH3 COO– , a corresponding Na+
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 - PDFGeneral, Organic, and Biological Chemistry
An Integrated Approach
- Kenneth W. Raymond(Author)
- 2012(Publication Date)
- Wiley(Publisher)
To explain, let us use a different approach to describe what happens when hydroxide ion is added to this buffer: The reaction of OH - with H 3 O + reduces the concentration of each. H 3 O + + OH - ¡ 2H 2 O This causes a slowing of the reverse reaction between CH 3 CO 2 - and H 3 O + , which allows the forward reaction to predominate until equilibrium is reestablished. As a result, lost H 3 O + will be mostly replaced and the pH will remain fairly constant. CH 3 CO 2 H + H 2 O ¡ CH 3 CO 2 - + H 3 O + Interestingly, when the two reactions directly above are added together, and when reac- tants and products appearing unchanged on each side of the reaction arrow are removed, the resulting reaction equation is identical to the one at the beginning of this paragraph. H 3 O + + OH - ¡ 2 H 2 O CH 3 CO 2 H + H 2 O ¡ CH 3 CO 2 - + H 3 O + CH 3 CO 2 H + OH - ¡ CH 3 CO 2 - + H 2 O Regardless of which approach is used to explain the effect of adding OH - to this buffer, the result is the same. There is a net conversion of CH 3 CO 2 H into CH 3 CO 2 - and the pH remains constant. Buffers are most resistant to pH changes when the pH equals the pK a of the weak acid and are effective when the pH is within one unit of the pK a (pH = pK a ; 1). The pK a of acetic acid is 4.74, so an acetic acid and acetate ion buffer functions well over the pH range 3.74 to 5.74 and is most effective at a pH of 4.74. ■ Buffers resist changes in pH. 262 CHAPTER 7 Acids, Bases, and Equilibrium We have seen that a buffer can be prepared from a weak acid and its conjugate base, and that the concentration of these two substances is connected to the pH of the solution and the pK a of the acid (Table 7.6). Taking a more mathematical approach to these concepts allows us to more accurately describe the makeup of a buffer.
eBook - PDFChemistry
Structure and Dynamics
- James N. Spencer, George M. Bodner, Lyman H. Rickard(Authors)
- 2011(Publication Date)
- Wiley(Publisher)
The choice of the conjugate acid/base pair used to prepare a buffer solu- tion can be made by comparing the desired pH with a table of values of K a for weak acids. NH 4 + K a = 5.6 * 10 - 10 NH 3 K b = 1.8 * 10 - 5 OAc - K b = 5.6 * 10 - 10 HOAc K a = 1.8 * 10 - 5 NH 4 + (aq) + H 2 O(l) uv H 3 O + (aq) + NH 3 (aq) 11.16 BUFFERS AND BUFFER CAPACITY 509 E x e r c i s e 1 1 . 1 1 Describe how to prepare a pH 9.35 buffer solution. Solution The data in either Table 11.4 or Table B.8 in the Appendix suggest that the NH 4 /NH 3 acid–base pair could be used to prepare this buffer because the pK a for the ammonium ion is 9.25. NH 4 + (aq) + H 2 O(l) uv NH 3 (aq) + H 3 O + (aq) Buffers are used extensively in the chemistry laboratory. They can be pur- chased from chemical suppliers as solutions with known concentrations and pH values. They can also be purchased as packets of a mixture of a solid conjugate acid–base pair. Dissolving the packet in water yields a buffer with a pH equal to the value stated on the package. In addition, buffers can be prepared by mixing measured amounts of an appropriate conjugate acid–base pair and then adding strong acid or base to adjust the pH to the desired value. The pH of the buffer is normally monitored with a pH meter as the strong acid or strong base is added. 11.17 Buffers in the Body Buffers are very important in living organisms for maintaining the pH of biolog- ical fluids within the very narrow ranges necessary for the biochemical reactions of life processes. Three primary buffer systems maintain the pH of blood in the 510 CHAPTER 11 / ACIDS AND BASES We can start by substituting the values of the pH of the buffer solution and the pK a value for the conjugate acid into the Henderson–Hasselbalch equation. Solving for the ratio of the concentrations of the conjugate acid–base pair gives the following result.
- 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.
eBook - PDF- Charlotte W. Pratt, Kathleen Cornely(Authors)
- 2020(Publication Date)
- Wiley(Publisher)
The broad, flat shape of the titration curve shown in Figure 2.16 indicates that the pH does not change drastically with added acid or base when the pH is near the pK. The effective buffering capacity of an acid is generally taken to be within one pH unit of its pK. For acetic acid (pK = 4.76), this would be pH 3.76–5.76. Biochemists nearly always perform experiments in buffered solutions in order to maintain a constant pH when acidic or basic substances are added or when chemical reactions produce or consume protons. Without buffering, fluctuations in pH would alter the ionization state of the molecules under study, which might then behave differently. Before biochemists appreci- ated the importance of pH, experimental results were often poorly reproducible, even within the same laboratory. A buffer solution is typically prepared from a weak acid and the salt of its conjugate base (see Sample Calculation 2.5). The two are mixed together in the appropriate ratio, according to the Henderson–Hasselbalch equation, and the final pH is adjusted if necessary by adding a small amount of concentrated HCl or NaOH. In addition to choosing a buffering compound with a pK value near the desired pH, a biochemist must consider other factors, including the compound’s solubility, stability, toxicity to cells, reactivity with other molecules, and cost. pH pK 0 0.5 H + ions dissociated 1 [CH 3 COOH] > [CH 3 COO − ] [CH 3 COOH] < [CH 3 COO − ] Midpoint [CH 3 COOH] = [CH 3 COO − ] Start point End point 1 2 3 4 5 6 7 8 Effective buffering range SAMPLE CALCULATION 2.5 PROBLEM How many mL of a 2.0-M solution of boric acid must be added to 600 mL of a solution of 10 mM sodium borate in order for the pH to be 9.45? FIGURE 2.16 Titration of acetic acid. At the start point (before base is added), the acid is present mainly in its CH 3 COOH form. As small amounts of base are added, protons dissociate until, at the midpoint of the titration (where pH = pK), [CH 3 COOH] = [CH 3 COO − ].
eBook - PDF- Gary D. Christian, Purnendu K. Dasgupta, Kevin A. Schug(Authors)
- 2013(Publication Date)
- Wiley(Publisher)
Again, the change in the ratio is small. The amount of acid or base that can be added without causing a large change in pH is governed by the buffering capacity of the solution. This is determined by the The buffering capacity increases with the concentrations of the buffering species. concentrations of HA and A − . The higher their concentrations, the more acid or base the solution can tolerate. The buffer intensity or buffer index of a solution is defined as β = dC B /dpH = −dC A /dpH (7.49) where dC B and dC A represent the number of moles per liter of strong base or acid, respectively, needed to bring about a pH change of dpH. Although the terms buffer intensity and buffer capacity are often used interchangeably, the buffer capacity is the integrated form of buffer intensity (e.g., the amount of strong acid/base needed to 4 In actuality, the pH will increase slightly because the activity coefficient of the salt has been increased by decreasing the ionic strength. The activity of an uncharged molecule (i.e., undissociated acid) is equal to its molarity (see Chapter 6), and so the ratio increases, causing a slight increase in pH. See the end of the chapter. 7.8 BUFFERS——KEEPING THE PH CONSTANT (OR NEARLY SO) 241 change the pH by a certain finite amount) and is always a positive number. The larger it is, the more resistant the solution is to pH change. For a simple monoprotic weak acid/conjugate base buffer solutions of concentration greater than 0.001 M, the buffer intensity is approximated by: β = 2.303 C HA C A − C HA + C A − (7.50) where C HA and C A − represent the analytical concentrations of the acid and its salt, See Chapter 8, Section 8.11 for a derivation of buffer intensity.
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)
1 Acid–Base Chemistry and Buffering F. John Gennari University of Vermont College of Medicine, Burlington, Vermont, U.S.A. John H. Galla Department of Medicine, University of Cincinnati College of Medicine, Cincinnati, Ohio, U.S.A. INTRODUCTION Acid–base biochemistry encompasses the physical chemistry of the constitu-ents of biological solutions that influence the dissociation of, and therefore the concentration of, hydrogen ions (H þ ) in those solutions. In the biologi-cal solutions that comprise the body fluids, these constituents include elec-trolytes that are essentially completely dissociated at the solute strength that exists in the body fluids, termed ‘‘strong ions’’ (1,2), a wide variety of weak acids and, most importantly, the volatile weak acid H 2 CO 3 (carbonic acid). Central to an understanding of acid–base homeostasis is knowledge of the chemistry of weak acids and, in particular, carbonic acid. In this chapter, we review the physical chemistry that underlies acid–base homeostasis, incorporating the concepts of Brønsted and Lowry, who defined acids as H þ donors and bases as H þ acceptors (3,4). The additional role of strong ions, which are regulated independently of the dictates of acid–base home-ostasis but influence [H þ ] is discussed at the end of this chapter. Van Slyke (5) revolutionized our ability to approach and deal with the acid–base status of biological solutions, using the concepts of Brønsted and Lowry to focus on [H þ ] through evaluation of a single weak acid, H 2 CO 3 (carbonic acid). Stewart (1,2) added the constraints of electroneutrality to the assessment of [H þ ] and separated the quantities that can be manipulated 1 external to the solution, i.e., the concentrations of strong ions, buffer content, and PCO 2 , from the quantities that are dependent on the nature of the solution, i.e., [H þ ] and [HCO 3 ].
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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