Buffer Solutions

11 Key excerpts on "Buffer Solutions"

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  eBook - PDF

  Laboratory Techniques with Reagents and Solutions

  • Chavan, U D(Authors)
  • 2021(Publication Date)
  • Daya Publishing House

    (Publisher)

  Applications Given their resistance to changes in pH, Buffer Solutions are very useful for chemical manufacturing and essential for many biochemical processes. The ideal buffer for a particular pH has a pKa equal to the pH desired, since a solution of this buffer would contain equal amounts of acid and base and be in the middle of the range of buffering capacity. Buffer Solutions are necessary to keep the correct pH for enzymes in many organisms to work. Many enzymes work only under very precise conditions; if the pH strays too far out of the margin, the enzymes slow or stop working and can denature, thus permanently disabling its catalytic activity. A buffer of carbonic acid (H 2 CO 3 ) and bicarbonate (HCO 3 − ) is present in blood plasma, to maintain a pH between 7.35 and 7.45. Industrially, Buffer Solutions are used in fermentation processes and in setting the appropriate conditions for dyeing fabrics. They are also used in chemical analyses and syntheses, and for the calibration of pH meters. Buffering Agents A buffering agent adjusts the pH of an acidic or alkaline solution and stabilizes it at that pH. Buffering agents have variable properties: some are acidic, others are basic; some are more soluble than others. They are useful for a variety of applications, including agriculture, food processing, medicine, and photography. Buffering agents and Buffer Solutions are similar in that they both regulate the pH of a solution and resist changes in pH. They function based on the same chemical principles. They may, however, be distinguished by the following differences: 1. A buffer solution maintains the pH of a system, preventing large changes in it, whereas a buffering agent modifies the pH of what it is placed into. 2. A buffering agent is the active component of a buffer solution. This ebook is exclusively for this university only. Cannot be resold/distributed.

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  Chemistry

  Principles and Reactions

  • William Masterton, Cecile Hurley(Authors)
  • 2020(Publication Date)
  • Cengage Learning EMEA

    (Publisher)

  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. 14-1 Buffers 361 A solution with these properties is called a buffer ▲ because it cushions the “shock” (i.e., the drastic change in pH) that occurs when a strong acid or strong base is added to water. To prepare a buffer, we can mix solutions of a weak acid HB and the sodium salt of that acid NaB, which consists of Na 1 and B 2 ions. This mixture can react with either a strong base HB( aq ) 1 OH 2 ( aq ) 9: B 2 ( aq ) 1 H 2 O or a strong acid B 2 ( aq ) 1 H 1 ( aq ) 9: HB( aq ) These reactions have very large equilibrium constants, as we will see in Sec-tion 14-3, and so go virtually to completion. As a result, the added H 1 or OH 2 ions are consumed and do not directly affect the pH. This is the principle of buffer action, which explains why a buffered solution is much more resistant to a change in pH than one that is unbuffered (Figure 14.1). Buffers are widely used to maintain nearly constant pH in a variety of com-mercial products and laboratory procedures (Figure 14.2). For these applications By definition, a buffer is any solu-tion that resists a change in pH. Both beakers contain distilled water. After the addition of two.LTzero.LT mL of zero.LT.three.LTzero.LT M HCl, the pH of the water has dropped significantly. Both beakers contain a pH four.LT buffer. Addition of the same volume of zero.LT.three.LTzero.LT M HCl to the buffer has barely changed the pH. a c b d Figure 14.1 The effect of a buffer upon addition of acid . Note the reading on the pH meter. © Cengage Learning/Charles D.

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  Chemistry

  The Molecular Nature of Matter

  • Neil D. Jespersen, Alison Hyslop(Authors)
  • 2021(Publication Date)
  • Wiley

    (Publisher)

  Lakes and streams with a pH less than 5 often cannot support fish life. Thus, a change in pH can produce unwanted effects, and systems that are sensitive to pH must be protected from the H + or OH − that might be formed or consumed by some reaction. Buffers are mixtures of solutes that accomplish this. Such a solution is said to be buffered or it is described as a buffer solution. NOTE If neither the cation nor the anion is able to affect the pH, the salt solution will be neutral (provided no other acidic or basic solutes are present). 16.7 Buffer Solutions 817 Composition of a Buffer A buffer contains solutes that enable it to resist large changes in pH when small amounts of either strong acid or strong base are added to it. Ordinarily, the buffer consists of two sol- utes, one a weak Brønsted acid and the other its conjugate base. If the acid is molecular, then the conjugate base is supplied by a soluble salt of the acid. For example, a common buffer system consists of acetic acid plus sodium acetate, with the salt’s acetate ion serving as the Brønsted base. In your blood, carbonic acid (H 2 CO 3 , a weak diprotic acid) and the bicarbonate ion (HCO 3 − , its conjugate base) serve as one of the buffer systems used to maintain a remark- ably constant pH in the face of the body’s production of organic acids by metabolism. Another common buffer consists of the weakly acidic cation, NH 4 + , supplied by a salt like NH 4 Cl, and its conjugate base, NH 3 . One important point about buffers is the distinction between keeping a solution at a par- ticular pH and keeping it neutral—at a pH of 7. Although it is certainly possible to buffer a solution at pH 7, buffers can be made that will work at any pH value throughout the pH scale. It is important to reiterate that a buffer only resists changes in pH; it cannot keep the pH constant. How a Buffer Works To work, a buffer must be able to neutralize either a strong acid or strong base that is added.

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  Physical Chemistry for Engineering and Applied Sciences

  • 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.

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  Analytical Chemistry

  • Clyde Frank(Author)
  • 2012(Publication Date)
  • Academic Press

    (Publisher)

  Buffered solutions are of great significance. For example, many physiological processes require a fixed pH to function properly. Often the permissable pH range is very narrow. To maintain this pH, nature has included buffers into the system, and, frequently, the components of these buffers are the same ones that the chemist uses in the laboratory. Thus, if these same processes are to be studied in the laboratory, the conditions for the buffer must be clearly understood.

  The control of pH in experimental chemistry is essential in all sorts of chemical and instrumental applications. Often the success in these applications will be determined by how carefully the pH is controlled and maintained.

  The most common type of buffer solution is made by dissolving a weak acid and a salt of the same weak acid in water or by dissolving a weak base and a salt of the same weak base in water. For example, consider the first case where ionization will occur according to the reaction

  (8-32)

  If a salt of the weak acid, NaA, is added, the H3 O+ concentration will decrease (pH increases) as predicted by the Le Chatelier principle; that is, the equilibrium is shifted to the left.

  If a strong acid or some other source of hydronium ion is introduced into the solution, association with the anion A− , reverse of reaction (8-32), takes place. Since there is a large reservoir of A− little change in pH is observed.

  If a strong base or hydroxide ion is introduced into the solution, neutralization of the hydronium ion in reaction (8-32) occurs. To replace the consumed hydronium ion more HA is ionized and since there is a reservoir of HA the pH changes only slightly.

  A buffer made from a weak base (BOH) and its salt (BX) can be described in an analogous way. In this case ionization of the weak base must be considered.

  (8-33)

  As hydronium ions are added, they are neutralized by the hydroxide ion which are replaced by more ionization of BOH. If hydroxide ions are added, they are consumed by reaction with B+

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  Fundamentals of Biochemical Calculations

  • 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.

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  Chemistry

  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.

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  Principles of Physiology for the Anaesthetist

  • Peter Kam, Ian Power(Authors)
  • 2015(Publication Date)
  • CRC Press

    (Publisher)

  K) of the substance.

  P H SYSTEM

  H+ ion concentration may be measured in two ways: directly as concentrations in nanomoles per litre or indirectly as pH. pH is defined as the negative logarithm (to the base 10) of the concentration of hydrogen ions. The pH is related to the concentration of H+ as follows:

  pH   =   log

    10

  1

  [

  H +

  ]

  pH   =   log

    10

  [

  H +

  ]

  H +

    =  

  10

  − pH

  pH   =   p K +   log   base/acid

  Table 8.1 Relationship between pH and hydrogen ion concentration

  pH	Hydrogen ion concentration (nmol/L)
  7.7	20
  7.4	40
  7.3	50
  7.1	80

  It is important to note that pH and hydrogen ion concentration [H+ ] are inversely related such that an increase in pH describes a decrease in [H+ ] (Table 8.1 ). However, the logarithmic scale is nonlinear and, therefore, a change of one pH unit reflects a 10-fold change in [H+ ] and equal changes in pH are not correlated with equal changes in [H+ ]. For example, a change of pH from 7.4 to 7.0 (40 nmol/L [H+ ] to 100 nmol/L [H+ ]) represents a change of 60 nmol/L [H+ ], although the same pH change of 0.4, but from 7.4 to 7.8 (40 nmol/L [H+ ] to 16 nmol/L [H+ ]), represents a change of only 24 nmol/L [H+ ].

  BUFFERS

  A buffer is a solution consisting of a weak acid and its conjugate base, which resists a change in pH when a stronger acid or base is added, thereby minimizing a change in pH. The most important buffer pair in extracellular fluid (ECF) is carbonic acid (H2 CO3 ) and bicarbonate (

  HCO 3 −

  ). The interaction between this buffer pair forms the basis of the measurement of acid–base balance.

  HYDROGEN ION BALANCE

  Cellular hydrogen ion turnover can be described in terms of processes that produce or consume H+ ions in the body (Table 8.2 ). The total daily H+

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  Pratt's Essential Biochemistry

  • Charlotte W. Pratt, Kathleen Cornely(Authors)
  • 2020(Publication Date)
  • Wiley

    (Publisher)

  2.3 Acid–Base Chemistry • The dissociation of water produces hydroxide ions (OH − ) and protons (H + ) whose concentration can be expressed as a pH value. The pH of a solution can be altered by adding an acid (which donates protons) or a base (which accepts protons). • The tendency for a proton to dissociate from an acid is expressed as a pK value. • The Henderson–Hasselbalch equation relates the pH of a solution of a weak acid and its conjugate base to the pK and the concentrations of the acid and base. 2.4 Tools and Techniques: Buffers • A buffered solution, which contains an acid and its conjugate base, resists changes in pH when more acid or base is added. 2.5 Clinical Connection: Acid–Base Balance in Humans • The body uses the bicarbonate buffer system to maintain a con- stant internal pH. Homeostatic adjustments are made by the lungs, where CO 2 is released, and by the kidneys, which excrete H + and ammonia. SUMMARY polarity hydrogen bond ionic interaction van der Waals radius electronegativity van der Waals interaction dipole–dipole interaction London dispersion forces dielectric constant solute solvation hydration KEY TERMS 46 CHAPTER 2 Aqueous Chemistry Brief Bioinformatics Exercises 2.1 Structure and Solubility 2.2 Amino Acids, Ionization, and pK Values BIOINFORMATICS 2.1 Water Molecules and Hydrogen Bonds 1. Each CO bond in CO 2 is polar, yet the whole molecule is nonpolar. Explain. 2. The HCH bond angle in the perfectly tetrahedral CH 4 molecule is 109º. Explain why the HO H bond angle in water is only about 104.5º. 3. Which compound has a higher boiling point, H 2 O or H 2 S? Explain. 4. Consider the following molecules and their melting points listed below. How can you account for the differences in melting points among these molecules of similar size? Molecular weight Melting (g · mol –1 ) point (°C) Water, H 2 O 18.0 0 Ammonia, NH 3 17.0 −77 Methane, CH 4 16.0 −182 5.

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  Acid-Base Disorders and Their Treatment

  • 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 ].

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  General Chemistry for Engineers

  • Jeffrey Gaffney, Nancy Marley(Authors)
  • 2017(Publication Date)
  • Elsevier

    (Publisher)

  Chapter 5

  Acids and Bases

  Abstract

  This chapter explains the differences between the Brønsted-Lowry and Lewis definitions of acids and bases and gives examples of each. Since chemical reactions involving Lewis acids and bases are covered in more detail in Chapter 10, most of the chapter is dedicated to the applications of the Brønsted-Lowry concepts. This includes the strength of acids and their conjugate bases, the behavior of acids and bases in aqueous solution, the autoionization of water, and the acid ionization constants. The “p” functions, including pH, pOH, pK a , pK b , and pK w , are discussed. The function and uses of Buffer Solutions are explained along with their design using the Henderson-Hasselbalch equation. Titration procedures are discussed and their relevance to industrial situations is stressed.

  Keywords

  Brønsted-Lowry; Lewis acid; Conjugate pairs; Ionization constant; Amphoteric; Coordinate-covalent bond; Coordination complex; Buffer solution; Henderson-Hasselbalch; Titration

  Outline

  5.1  

  Defining Acids and Bases

  5.2  

  Acids and Bases in Aqueous Solution

  5.3  

  The pH Scale

  5.4  

  Other “p” Functions

  5.5  

  Buffer Solutions

  5.6  

  The Titration

  Important Terms Study Questions

  Problems

  5.1 Defining Acids and Bases

  The first modern attempt at defining acids and bases was by a Swedish chemist named Svante Arrhenius in 1887. Arrhenius defined an acid as a material that releases hydrogen ions (H+ ) when dissolved in water. Similarly, he defined a base as a material that releases hydroxide ions (OH− ) when dissolved in water. This definition only held for ionic compounds containing hydrogen or hydroxide ions and did not apply to many acids and bases that we deal with today. Since this early definition of acids and bases was so limited, two more sophisticated and general definitions of acids and bases have since been developed, which are in wide use today. These are known as the Brønsted-Lowry definition and the Lewis definition.

  In 1923, both J.N. Brønsted of Denmark and Thomas Lowry of England, working independently, defined an acid as a species that can donate a hydrogen ion to a base. A base was defined as a species that can accept a hydrogen ion from an acid. So, a Brønsted-Lowry acid

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