Chemistry
Buffers Preparation
Buffers preparation involves creating a solution that resists changes in pH when an acid or base is added. This is typically achieved by mixing a weak acid and its conjugate base, or a weak base and its conjugate acid, in specific proportions. The resulting buffer solution is important in various chemical and biological processes, as it helps maintain a stable pH environment.
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eBook - ePub- Jeffrey Gaffney, Nancy Marley(Authors)
- 2017(Publication Date)
- Elsevier(Publisher)
buffer ) is a solution which resists changes in pH when small quantities of a strong acid or a strong base are added to it. So, buffer solutions are used as a means of keeping the pH at a nearly constant value. They are used widely in both biochemical and industrial applications. In biochemistry, buffers are necessary for assays used to study the activity of enzymes. Since enzyme activity varies with pH, the pH must remain constant during the assay to get accurate results. Buffer solutions are also used in medicines that require a constant pH to maintain their activity. The textile industry uses buffer solutions to keep the pH within narrow limits during fabric dyeing and processing. Many laundry detergents also use buffers to prevent their natural ingredients from breaking down.Buffer solutions usually consist of a weak acid and its conjugate base in relatively equal concentrations. Practically, this is achieved by mixing a soluble compound that contains the conjugate base with an aqueous solution of the acid. For example, an acetate buffer is made by the addition of sodium acetate to an aqueous solution of acetic acid. Buffer solutions achieve their resistance to pH change because of the excess amount of the conjugate base and the equilibrium between the weak acid (HA) and its conjugate base (A− ). When a small amount of a strong acid is added to a buffer solution, the excess conjugate base present in the buffer consumes the added hydronium ion from the strong acid converting it into water and the weak acid of the conjugate base.A −aq +H 3O +aq →H 2O l + HA aqThis results in a decrease in the amount of the excess conjugate base and an increase in the amount of the unionized weak acid in the solution. So, the pH of the buffer solution remains relatively stable and may decrease by only a very small amount when a small amount of the strong acid is added to it. Similarly, when a small amount of a strong base is added to a buffer solution, the added hydroxide ions are consumed by the weak acid forming water and the conjugate base of the acid.OH −aq + HA aq →H 2O l +A −aqThe result is that the amount of the weak acid decreases and the amount of the conjugate base increases. This consumption of hydroxide prevents the pH of the solution from rising significantly, which would occur if the buffer system was not present.
eBook - ePub- Professor Rob Beynon, J Easterby(Authors)
- 2004(Publication Date)
- Taylor & Francis(Publisher)
1 Basic concepts 1. What are buffers?In 1900, Fembach and Hubert, studying the enzyme amylase, noted that a partially neutralized solution of phosphoric acid acted as a ‘protection against abrupt changes in acidity or alkalinity: the phosphates behave as a sort of buffer’. This definition is unchanged today, although we now have a better understanding of the phenomenon, and the reasons why it is important to make use of it.• Definition of buffers• Why understand them?• How to approach the maths• Why pH 7.0 is ‘neutral’Let us assume for the moment that you are familiar with the pH scale to express acidity or alkalinity of a solution, and that you know that in aqueous systems pH 7.0 is neutral, solutions with pH values below 7 are acidic, and solutions with pH values greater than 7 are alkaline.It is easy to demonstrate the phenomenon of buffering (Figure 1.1 ). We could take 1 ml of a 1 M solution of a strong acid, such as hydrochloric acid, and add it to a litre of pure water. The pH would drop (to a close approximation) from 7 to 3, a decrease of 4 pH units. If instead of pure water we had used one litre of a solution of 0.1 M sodium phosphate at the same initial pH of 7, the pH would only drop by 0.02 pH units to 6.98. Thus, the phosphate solution ‘buffers’ pH changes. A buffered solution resists changes in pH when the solution is exposed to acids or alkalis that would otherwise cause dramatic changes in pH. No buffer can resist these changes completely; at best they can minimize the effects.Figure 1.1 The concept of pH bufferingAdd a few drops of strong acid to pure water or a solution of sodium phosphate, both at the same pH, and note the pH shift in both instances. If you are unsure about the derivation and use of the pH scale, don’t worry—it will be discussed in some detail shortly.
eBook - ePub- John Kenkel(Author)
- 2020(Publication Date)
- CRC Press(Publisher)
0. 500 mL of the 100 ppm solution would be measured into a 500 mL flask and diluted to volume.3.7 BUFFER SOLUTIONS
Buffer solutions, or solutions which resist changes in pH even when a strong acid or base is added, are almost always composed of a weak acid or weak base and the salt of this weak acid or base. The reason for the resistance to pH change is that the weak acid or weak base ionization equilibrium shifts in these solutions such that the H+ or OH− added are consumed, thus resulting in no net pH change.Although commercially prepared buffer solutions are available, these are most often utilized solely for pH meter calibration and not for adjusting or maintaining a chemical reaction system at a given pH. It is not surprising, therefore, that the analyst often needs to prepare his/her own solutions for this purpose. It then becomes a question of what proportions of the acid, or base, and its salt should be mixed to give the desired pH.The answer is in the expression for the ionization constant, Ka or Kb , where the ratio of the salt concentration to the acid concentration is found. In the case of a weak acid,HA ⇌H +A −(3.33)
and in the case of a weak base,K a=[H +][A −][ HA ](3.34) B+H 2O ⇌BH ++OH −(3.35) K b=[BH +][OH −][ B ](3.36) Knowing the value of Ka or Kb , for a given weak acid or base and knowing the desired pH value, one can calculate the ratio of salt concentration to acid (or base) concentration that will produce the given pH. Rearranging Equation 3.34, for example, would show the method for calculating this ratio in the case of a weak acid and its salt.=K a[H +][A −][ HA ](3.37) This is one form of the Henderson-Hasselbalch equation for dealing with buffer solutions. It says that one would simply divide the Ka by the [H+ ] to obtain the required ratio. It should be stressed that since the Ka , or Kb , enters into the calculation, how weak the acid or base is dictates what is a workable pH range for that acid or base. Table 3.3
eBook - ePubChemistry
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 ionA 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
eBook - ePub- Juan Moreno, Rafael Peinado(Authors)
- 2012(Publication Date)
- Academic Press(Publisher)
The pH of wine changes little when a moderate quantity of a strong acid or a strong base is added. This ability of wine to resist large changes in pH is an indication of its buffering capacity, a phenomenon common to all solutions that contain a weak acid and its salt.Buffering capacity is explained by an equilibrium that is established between the molecular form of an acid and its conjugate base. The addition of a strong acid to a wine introduces H+ ions that react with the A− anion, leading to a reduction in its concentration and an increase in the molecular form (HA) of the acid. When the equilibrium between the molecular and dissociated forms of the acid is re-established in the wine, a new concentration of H+ is obtained that is derived solely from this dissociation and that is responsible for the new pH value. Similarly, the addition of a strong base to a wine leads to a neutralization reaction with the molecular form of the acid, which reduces the concentration of HA and increases the concentration of A− . The re-establishment of the acid-base equilibrium generates a new concentration of H+ and therefore a new pH value in the wine.2 Buffering Capacity of Weak Acid Solutions
The titration curve of a weak acid with a strong base contains a region in which the pH rises sharply during the addition of the first few milliliters of the base. This is followed by a period during which the pH remains relatively constant despite continued addition of the base. A similar situation occurs in the titration curve for a weak base with a strong acid, except that in this case there is an initial sharp reduction in pH followed by a period of relative stability.In each case, the slow change in pH upon addition of a base or acid is indicative of the buffering capacity of the solution. Hence, a buffer solution is any solution that undergoes only limited changes in pH when H+ or OH−
eBook - ePub- Peter Kam, Ian Power(Authors)
- 2015(Publication Date)
- CRC Press(Publisher)
K) of the substance.P H SYSTEMH+ 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 = log101[H +]pH = log10[H +]H +=10− pHpH = p K + log base/acidTable 8.1 Relationship between pH and hydrogen ion concentrationpH 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+ ].BUFFERSA 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 (HYDROGEN ION BALANCEHCO 3 −). The interaction between this buffer pair forms the basis of the measurement of acid–base 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+
eBook - ePub- Gillian Cockerill, Stephen Reed(Authors)
- 2011(Publication Date)
- Wiley(Publisher)
Note : [base] and [conj. acid] must be in same units)We can predict from the Henderson-Hasselbalch equation that buffers are most efficient in ‘mopping-up’ free protons when there are approximately equal concentrations of base and acid present, i.e. acid:base ratio of 1:1 and pH = pKa . Indeed, in a test-tube, buffer efficiency is maximal when the actual solution pH = pKa ± 1 pH unit. This rule does not always hold true for physiological systems, however, because unlike the situation in a test-tube, homeostatic regulatory mechanisms can ‘top-up’ the concentration of the individual conjugate pairs, perhaps in response to a pH challenge to ensure efficiency of the buffer system.
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SECTION 1.vii
Self-assessment exercise 1.2
1. Complete the following table by calculating values indicated by ?? Arrange the acids as a list with the strongest at the top and weakest at the bottom.Acid Ka pKa Acetic 1.83 × 10−5 ?? Carbonic 7.9 × 10−7 ?? Acetoacetic 2.6 × 10−4 ?? Fumaric Ka1 = 9.3 × 10−4 Ka2 = 2.88 × 10−5 ?? ?? Phosphoric Ka2 = 1.6 × 10−7 ?? Succinic Ka1 = 6.6 × 10−5 Ka2 = 2.75 × 10−6 ?? ?? Pyruvic ?? 2.50 Aspartic side chain COOH ?? pKa2 = 3.9 Histidine ?? pKa3 = 6.0 2. How does knowledge of the ionic behaviour of amino acids help us understand the buffer action of proteins?3. Write balanced chemical equations to show the dissociation of pyruvic acid (CH3 .CO.COOH) and of phosphoric acid (H3 PO4 ).4. State, giving reasons, whether (i) histidine (imidazole side chain pKa = 6.0), and (ii) the side chain amino group of arginine (pKa = 12.6) would be protonated or deprotonated at typical cytosolic pH of 6.85.5. What proportions (ratio) of lactic acid and sodium lactate would be required to prepare a buffer with a final pH of 3.5?6. Use your answer to question 5 above and state how many grams of lactic acid and sodium lactate would be needed to prepare 2.5 litres of the buffer (pH = 3.5).7. Refer to the diagram of citric acid shown in Section 1.v above. Draw a chemical equation to show the successive dissociation of the three ionisable protons. (Note : The COOH attached to the central carbon has the highest pKa
eBook - ePub- Arnost Kotyk, Jan Slavik(Authors)
- 2020(Publication Date)
- CRC Press(Publisher)
A and can be expressed (for low degrees of dissociation) as
(cf. Equation 12).pH = PK A+ log (c/A −c HA)(44) On adding a strong base to the solution, to a final concentration CB , a salt is formed at practically the same concentration so that CB′ ≅ c A− . This necessarily leads to a decrease of the concentration of undissociated acid, such that cHA ≅ CT − CB' . This then turns into the so-called Henderson-Hasselbalch equationpH = pK A+ log [c B′/ (c T−c Bʹ) ](45) Analogously, starting with a weak base (cT = cB + cBH+) we may writepH = pK BH+ log(c B/c)BH +(46) Addition of a strong acid to a final concentration cA' leads then to salt formation such that cA ' ≅ cBH+ and, necessarily cB ≅ cT − cA' . The pH is then defined bypH = pK BH+ log[(c T−c Aʹ) /c Aʹ](47) It is easy to show that the “titration” of, say, a weak acid with a strong base, proceeds through a “buffer zone” where pH changes very little in spite of adding large amounts of the base (Figure 2 ). The figure indicates that the buffer zone centers about a point where pH = PKAʹ , this being also the point where cBʹ = 1/2 cA .The efficiency of a buffer solution is usually expressed by its buffering power β which is the expression for a change of pH brought about by adding a certain amount of strong acid or base. Thus,FIGURE 2. Titration of a weak acid A with a strong base B′ The pH observed at the different base-to-acid ratios is defined by the expressions on the right. The zone of high buffering power lies here at about 0.2 to 0.8.β = dc Bʹ/ dpH = 2 .3c Bʹ(1 −c Bʹ/c A) for an acid buffer(48a) β = dc Aʹ/ dpH = 2 .3c Aʹ(c Aʹ/c B− 1 ) for an alkaline buffer(48b) The buffering power thus depends on cBʹ , or cAʹ and its maximum is reached (when the first derivative of Equation 48a or 48b is equal to zero) at cBʹ = 1/2 cA or cAʹ = 1/2 cB (Figure 3 ).FIGURE 3. The buffering power β (in mmol dm−3 ) of a weak acid buffer with pKA
