Acid-Base Reactions and Buffers
Acid-base reactions involve the transfer of protons from an acid to a base. Buffers are solutions that resist changes in pH when an acid or base is added. They are composed of a weak acid and its conjugate base, or a weak base and its conjugate acid, and help maintain the stability of pH in various chemical systems.
7 Key excerpts on "Acid-Base Reactions and Buffers"
eBook - ePub- Professor Rob Beynon, J Easterby(Authors)
- 2004(Publication Date)
- Taylor & Francis(Publisher)
...Under physiological conditions, the species on both sides of the equation can co-exist in substantial amounts—compare this with a strong acid such as HCl, which is virtually completely ionized to H + and Cl –. There are other more rigorous definitions of weak acids and bases which were alluded to in 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 (phosphate 1–) 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 changes A 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 3 COOH ⇌ CH 3 COO − + 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 CH 3 COO – ‘mops up’ the added protons (Figure 3.1). Electrical neutrality is preserved because every H + that reacts with a CH 3 COO – anion to form the neutral CH 3 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...
eBook - ePub- Jeffrey Gaffney, Nancy Marley(Authors)
- 2017(Publication Date)
- Elsevier(Publisher)
...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 3 O + aq → H 2 O l + HA aq This 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 2 O l + A − aq The result is that the amount of the weak acid decreases and the amount of the conjugate base increases...
eBook - ePubChemistry
Concepts and Problems, A Self-Teaching Guide
- Richard Post, Chad Snyder, Clifford C. Houk(Authors)
- 2020(Publication Date)
- Jossey-Bass(Publisher)
...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 (HC 2 H 3 O 2), 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. HC 2 H 3 O 2 and its salt NaC 2 H 3 O 2 are useful for making a buffer solution with a pH in the _________ range. Answer: acidic (HC 2 H 3 O 2 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 HC 2 H 3 O 2 is still in molecular form when in aqueous solution. The salt, in contrast, is completely dissociated...
eBook - ePubFood Chemistry
A Laboratory Manual
- Dennis D. Miller, C. K. Yeung(Authors)
- 2022(Publication Date)
- Wiley(Publisher)
...1 Acids, Bases, and Buffers 1.1 Learning Outcomes After completing this exercise, students will be able to: Explain the roles of acids and bases in food products. Measure the pH of selected food products. Prepare and evaluate a buffer system. Measure the buffering capacity of a common beverage. 1.2 Introduction Many food components may be classified as acids or bases due to their capacity to donate or accept protons (hydrogen ions). These components perform numerous important functions including flavor enhancement, control of microbial growth, inhibition of browning, alteration of texture, prevention of lipid oxidation, and pH control. Acids and bases are key metabolites in living plant and animal organisms, for example as intermediates in the TCA cycle, and are mostly retained when the plant is harvested or the animal is slaughtered so they are naturally present in foods. They may also be added during processing or synthesized during fermentation to produce desired characteristics in the final food product. The concentration and relative proportion of acids and bases determine the pH of a food, an extremely important characteristic. pH can affect the flavor, color, texture, stability, and behavior in food processing situations. For example, commercial sterilization of acid foods (pH less than 4.6) [1] can be achieved under milder processing conditions than in foods with a higher pH. 1.2.1 Acids Acids serve a variety of functions in foods including flavor enhancement, control of microbial growth, protein coagulation, emulsification, control of browning, buffering action, and metal chelation (to control lipid oxidation). All acids have a sour taste but different acids produce distinctively different sour flavors...
eBook - ePub- Peter Kam, Ian Power(Authors)
- 2015(Publication Date)
- CRC Press(Publisher)
...The ability of a substance to donate or accept a proton (i.e., to act as an acid or a base) depends on the concentration of H + ions in solution (pH of the solution) and the degree of dissociation (p 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 (H 2 CO 3) 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 + ion turnover in a normal adult is approximately 150 moles...
eBook - ePub- Gillian Cockerill, Stephen Reed(Authors)
- 2011(Publication Date)
- Wiley(Publisher)
...When neutral amino acids such as alanine are combined together in proteins, the carboxylic acid and amino groups of adjacent molecules are used in peptide bond formation, so only the N-terminal and C-terminal residues show the acid-base behaviour illustrated above. However, a few amino-acid residues such as glutamate, aspartate (both dicarboxylic), histidine, arginine and lysine (which have additional proton-accepting basic groups) are able to donate or accept protons even when found within the peptide chain of a protein. Histidine is particularly important because it is the only amino acid that has a proton which is able to dissociate in the neutral pH range. The ability to donate or accept protons makes proteins very important buffers, as we shall see in Part 3. SECTION 1.vi Buffers and the Henderson-Hasselbalch equation The normal processes of metabolism produce large quantities of acid each day, which if left unchecked would very soon cause irreparable damage to cells and the organism. Efficient means of dealing with the acid insult, including chemical buffering, are therefore of the utmost importance for survival. A buffer resists changes in pH of a solution and is composed of a weak acid (shown here as HB) and a salt (represented as NaB) of that weak acid. The salt provides anions (B −) to react with, and therefore neutralise, any addition of protons. The weak acid is present to furnish (by dissociation) more anions as they are ‘used up’ by added protons. A weak acid obeys the laws of chemical equilibria, so when some anions are ‘removed’ (as a result of the neutralisation), the position of the equilibrium of the weak acid will change to liberate more anions...
eBook - ePub- Arnost Kotyk, Jan Slavik(Authors)
- 2020(Publication Date)
- CRC Press(Publisher)
...IONIZATION OF ACIDS AND BASES Although water is by far the most abundant component of all living systems, its dissociation into oxonium ions is so weak that the pH of a salt solution, either extra- or intracellular, is determined by the presence of components that readily dissociate or readily bind an oxonium ion, i.e., acids and bases, respectively. Throughout modern electrochemistry, three theories of acids and bases came into prominence. The first theory, that of Arrhenius 3 dates back to 1887 when he postulated the universal existence of dissociation of electrolytes in solution, supporting his views by conductometric measurements. He calculated the degree of dissociation α from the ratios of equivalent conductivities at a given and at an infinite dilution. Thus, α = Λ / Λ ∞ It was Arrhenius who defined acids and bases in a simple way, stating that an acid (HA) is characterized by dissociation of hydrogen ions HA ⇌ H + + A − (7a) while a base (BOH) is recognized as a substance dissociating hydroxide ions BOH ⇌ B + + OH − (7b) Although many of his deductions are still valid, particularly with respect to aqueous solutions, he could not foresee the behavior of acids and bases in nonaqueous solvents and the role of interactions between the solute and the solvent. The second theory is the one that is most relevant to our considerations of pH in aqueous solutions and is due to Brönsted 4 who replaced Arrhenius’ definition of acids and bases by stating more generally that an acid is any substance that can dissociate a proton; a base is then any substance that can bind a proton...
