Buffer Capacity

8 Key excerpts on "Buffer Capacity"

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  eBook - ePub

  Analytical Chemistry

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

    (Publisher)

  + and B are equal to the stoichiometric concentrations and, second, their values are not influenced by the ionization of B.

  Buffer Capacity

  The ability of a buffer to consume acid or base does not continue indefinitely. Each prepared buffer is capable of resisting only a certain amount of added acid or base. A measure of this quantity is referred to as Buffer Capacity and can be defined as the number of moles of strong base or strong acid required to cause a unit increase or decrease in pH in 1 liter of a buffered solution.

  There are two experimental techniques for producing a high Buffer Capacity, which is usually desirable. First, a large concentration of the buffer components can be used. Or, second, the concentration of the weak acid (base) that is present should equal the concentration of its salt. Hence, maximum capacity is obtained when the dissociation constant for the weak acid (K w /K b for a weak base) is equal to the hydronium ion concentration that is desired. Usually, a combination of these two techniques is utilized.

  Example 8-14

  Calculate the change in pH if 1.000×10−3 mole of HCl were added to the buffer prepared in Example 8-9 . Assume that the volume does not change.

  An approximate answer is obtained if it is assumed that the HCl stoichiometrically converts C2 H3 O2 − to HC2 H3 O2 . This is reasonable because the amount of acid that is added, even though it is a strong acid, is very small in comparison to the concentration of the buffer components. Therefore,

  Substitution into Eq. (8-35) yields

  The addition of the acid to the buffer solution in Example 8-9

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  College Chemistry

  • Steven Boone, Drew H. Wolfe(Authors)
  • 2011(Publication Date)
  • Collins Reference

    (Publisher)

  CHAPTER 17

  Aqueous Equilibria: Buffers, Titrations, Solubility, and Complex Ion Equilibria

  I n this chapter, our discussion of acids and bases continues by considering buffer solutions, those that resist changes in the pH, and titrations, volumetric procedures in which unknown acid or base solutions are neutralized with known base or acid solutions. Additionally, solubility and complex ion equilibria will be discussed. A solubility equilibrium establishes between undissolved solutes and their dissolved ions. Complex ion equilibria establish between coordination complexes and their ions.

  17.1 BUFFER SOLUTIONS

  Solutions that contain both an acid and its conjugate base, or a base and its conjugate acid in sufficient quantity are called buffer solutions. The Nature of Buffer Solutions Buffer solutions maintain a nearly constant pH despite the addition of small amounts of either acids or bases. In other words, buffer solutions resist changes in pH. Components of a Buffer Solution

  If an acid is added to a buffer solution, it must have a basic component to neutralize the acid, and if a base is added to a buffer solution, it must have an acidic component to neutralize the base. Hence, buffer solutions are prepared by mixing either a weak acid and its conjugate base, or a weak base and its conjugate acid.

  Buffering Action: Adding Acid to a Buffer Solution

  To understand how buffer solutions maintain a constant pH, let’s consider a buffer prepared by mixing equimolar amounts of the weak acid acetic acid, HC2 H3 O2 , and the sodium salt of its conjugate base, sodium acetate, NaC2 H3 O2 . The net ionic equation for the equilibrium that establishes in this buffer solution is as follows.

  The Na+ ion from the sodium acetate is present as a spectator ion. If a small amount of a strong acid such as HC1 is added to this buffer solution, the resulting H+ ions shift the equilibrium to the left. The net result of adding the HC1(aq) is to decrease the acetate ion concentration and increase the acetic acid concentration. Almost all the added H+

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  Environmental Process Analysis

  Principles and Modeling

  • Henry V. Mott(Author)
  • 2013(Publication Date)
  • Wiley

    (Publisher)

  Bases present in an aqueous solution comprise the capacity of that solution to resist depression of the pH upon addition of proton donating substances (acids). Acids present in an aqueous solution comprise the capacity of that solution to resist elevation of the pH upon addition of proton accepting substances (bases). Here we must discern between alkalinity, acidity, and buffering capacity.

  Alkalinity (also called acid neutralizing capacity, [ANC]) is a measured value based on a standard laboratory test involving titration of an aqueous sample using a strong acid from its initial pH to a standard end point pH value.

  Acidity (also called base neutralizing capacity, [BNC]) is a measured value based on a standard laboratory test involving titration of an aqueous solution using a strong base from its initial pH to a standard end point pH value.

  TABLE 6.2 Abundance (Ionization) Fractions for Mono- Through Tetraprotic Acids

  Monoprotic acidsa , b

  Diprotic acidsa , b

  Triprotic acidsa , b

  Tetraprotic acidsa , b

  a α 0 refers to the fully protonated acid; α 1 – α 4 refer to species resulting from donation of 1–4 protons, respectively, from the fully protonated acid.

  b n is the residual charge on the fully protonated acid, and is generally equal to the number of R–NH3 + groups present within the fully protonated acid.

  Buffering capacity is a more general or qualitative term that describes the capacity of a solution in general to resist changes in pH from additions of acids or bases. Some authors use the term buffer intensity as the specific buffering capacity of an aqueous solution at a specific value of pH.

  Let us first focus upon alkalinity [ANC] as a standard laboratory test. This test was devised in most part to characterize the carbonate system species present in natural waters. The acid buffering capacity of the vast majority of natural waters is due solely to the presence of carbonate system species in the solution.

  Briefly, the alkalinity laboratory test involves the measurement of the initial pH of a water sample, and titration of a known volume of that water sample to an end point pH in the range of ~4.3 using a strong acid. Historically, methyl orange has been used as an indicator in the total alkalinity titration, turning from orange to colorless at pH ~4.3. Phenolphthalein has been used in titrations of samples with initial pH above 8.3, as solutions containing phenolphthalein turn from pink to colorless at pH ~8.3. The quantity and normality of the acid added to the solution to the pH 8.3 (phenolphthalein) end point and pH 4.3 (methyl orange) end point are recorded. Most often the results are converted from eqacid /Lsolution (or meq/L) into mg/L as CaCO 3

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  Buffer Solutions

  • 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 values

  We 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

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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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  Aquaculture Engineering

  • Odd-Ivar Lekang(Author)
  • 2019(Publication Date)
  • Wiley-Blackwell

    (Publisher)

  Below this point the carbonate system will only be represented by carbonic acid. A definition of total alkalinity can also be the requirement in the number of equivalents per litre of a strong acid to titrate the water sample down to a pH of 4.5. There can still be alkalinity down to these low pH's. Alkalinity can also be expressed as P alkalinity, meaning alkalinity above phenolphthalein indicator with endpoint pH around 8.2–8.4 but this is used in aquaculture. 5.3.3 A buffer A buffer may be added to increase the alkalinity in water when it is low. A good buffer keeps the pH constant at the required pH. For aquaculture purpose, buffers that keep pH stable at around 7 are good buffers. What is happening is that the buffer builds up alkalinity and the alkalinity is typically used/consumed in a narrow pH interval for a good buffer. Outside this interval, the pH will drop fast. The buffering capacity is strongest when the pH is close to the p K (equilibrium constant) for the media, acid buffer system. As said the typical natural buffer in water is made of the carbonate system. Here the carbonic acid–bicarbonate equilibrium has a p K value around 6.3 and the bicarbonate–carbonate equilibrium has a p K value of 10.3. This means that the titration curve flattens out in those areas (Fig. 5.4)

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  Aquatic Chemistry Concepts, Second Edition

  • James F. Pankow(Author)
  • 2019(Publication Date)
  • CRC Press

    (Publisher)

  8 Buffer Intensity β

  8.1 Introduction

  If any amount of strong base is added to any aqueous solution, the pH will go up. Conversely, if any amount of strong acid is added, the pH will go down. Chemists are often interested in knowing how pH-sensitive a system is to incremental addition of strong base or acid; the parameter that quantifies this concept is the buffer intensity β . In the environment, many aquatic organisms depend on the pH of the water to be in some particular range, and deviations from that range cause stress, or in extreme cases, mass die-offs. In the human body, a great many important biochemical reactions are very pH sensitive, so buffering is again extremely important. For humans, nothing makes this point more clearly than noting that human blood needs to be in a very narrow pH range so that pHdependent reactions can proceed at needed rates. For arterial blood, the normal range is 7.35 to 7.45; for venous blood, 7.32 to 7.42. Having recently passed through the lungs, arterial blood is slightly more alkaline because of loss of some metabolically generated CO2 to exhaled air. “Acidosis” is the condition when blood pH is too low, and “alkalosis” is the condition when blood pH is too high. (Acidosis in humans often leads to tachypnea/hyperventilation as a means to off-gas CO2 and raise blood pH.) Most natural water organisms benefit when their aquatic environment is characterized by an adequately large β so as to prevent significant swings in pH due to acid/base additions or losses.

  If adding a relatively large amount of strong base or acid causes only a small pH change, then β is large. If adding a small amount of strong base or acid causes a large pH change, then β is small. With (C B  − C A

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  Chemistry

  Concepts and Problems, A Self-Teaching Guide

  • Richard Post, Chad Snyder, Clifford C. Houk(Authors)
  • 2020(Publication Date)
  • Jossey-Bass

    (Publisher)

  + concentration), creating a greater than normal acidity. Electrolyte balance inside and outside the cells of our body is also pH dependent. To maintain the proper balance requires a carefully regulated system wherein the pH of the system remains virtually constant within very specific limits. We now discuss this system that is so prevalent in our bodies and that is extensively used in commercial processes to maintain a constant pH.

  BUFFER SOLUTIONS

  When chemists wish to keep the pH of a solution fairly constant even if some small amount of strong acid or base is added, they will use a buffer solution. A buffer solution involves a chemical equilibrium between either a weak acid and its salt or a weak base and its salt, and shows the common ion effect.

  A typical buffer solution is one made up of acetic acid (

  HC2 H3 O2

  ), which dissociates to a small degree into H+ and ions, and sodium acetate, a salt of acetic acid that dissociates completely into Na+ and ions. Which ion is common to acetic acid and sodium acetate? __________

  Answer: , the acetate ion

  A buffer solution can consist of a weak acid and its salt or a weak base and its salt, depending upon the desired pH of the buffer solution. A buffer solution with a pH in the acidic range (1 – 7) can be made from a solution of a weak acid and its salt. A buffer solution with a pH in the basic range (7–14) can be made from a solution of a weak base and its salt. HC2 H3 O2 and its salt NaC2 H3 O2 are useful for making a buffer solution with a pH in the _________ range.

  Answer: acidic (HC2 H3 O2 is a weak acid.)

  The key to understanding the action of a buffer solution is to remember that a weak acid (or weak base) is only dissociated to a very small degree. Most of the HC2 H3 O2 is still in molecular form when in aqueous solution. The salt, in contrast, is completely dissociated. All of the NaC2 H3 O2 becomes Na+ and

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