Dissociation Constant Of Acids And Bases Pdf


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Acids and Bases: The Voyage of the Proton. Chemistry: Acids and Bases Transcript. Acids and bases are found all around us, and the currency of acid-base chemistry is the proton, or hydrogen ion.

Salt effect on the dissociation constant of acid-base indicators

The acid dissociation constant pK a is among the most frequently used physicochemical parameters, and its determination is of interest to a wide range of research fields. We present a brief introduction on the conceptual development of pK a as a physical parameter and its relationship to the concept of the pH of a solution. This is followed by a general summary of the historical development and current state of the techniques of pKa determination and an attempt to develop insight into future developments.

Fourteen methods of determining the acid dissociation constant are placed in context and are critically evaluated to make a fair comparison and to determine their applications in modern chemistry. Additionally, we have studied these techniques in light of present trends in science and technology and attempt to determine how these trends might affect future developments in the field. The centennial of the concept and of the quantitative measurement of pH was celebrated not long ago. The related concept of the acid dissociation constant pK a as a substance property is recognized as being among the most commonly used parameters in modern-day chemistry.

Both pH and pK a are essential for understanding the behavior of chemical substances in everyday life. This realization came gradually and under different names. The first notion of acids comes from ancient Greece, where people noticed that some substances tasted sour. Measurements of osmotic pressures and conductivity of solutions gave insight into the degree of dissociation. Although Henderson defined K in terms of a concentration ratio in , it was not until that Hasselbalch 7 proposed their now famous equation 1 , which remains the most commonly used equation to calculate pK a values: the Henderson-Hasselbalch equation.

In many experimental methods to determine pK a values, a certain parameter is measured as a function of pH. This results in a characteristic sigmoid curve Fig. Generally speaking, for acidic components X ranges from a bulk property of a solution of only non-dissociated acid to the situation where only dissociated acid is present.

A classic example of a sigmoidal curve created by plotting a measured quantity versus pH. The inflection point corresponds to pK a. Combining Equations 1 and 2 leads to for anions :. Since the Henderson-Hasselbalch equation only gives accurate results for dilute acids in aqueous solutions, another formula for the quantification of acid strength was developed by Hammett. However, difficulty associated with accurate determination of the parameters in this model has kept it from being as widely used as the Henderson-Hasselbalch equation.

This is commonly used to interpolate to other temperatures. An example of the pK a temperature dependency is shown in Figure 2. Various pK a values as function of temperature.

Each component has an own dependency. Values in water at infinite dilution, data from Everaerts et al. Measured pK a values also depend on the ionic strength of the solution under investigation. The ionic strength is defined as the sum of concentrations c of all ionic species, corrected for their charge number z :. Because pK a depends on the activity coefficients, the ionic strength will also influence pK a , especially at higher charge numbers z.

An example of this dependency is shown in Figure 3. Data from Cohn et al. Because the acid-base equilibrium occurs in solution, the solvent composition can also influence the pK a values Fig. Measuring pH of mixtures of organic solvents and water is in itself far from straightforward, and are beyond the scope of the present contribution. When regarding an acid dissociation reaction, three thermodynamic steps are considered: 1 the dissolution of the acid from the solvent into the gas phase, 2 the dissociation of the acid into the ions, and 3 the solution step of the ions into the solvent.

In the first and the last step, the solvent is involved. When considering the influence of the solvent, the difference between the solvation energies of the acid and the dissociated acid influences the final pK a value.

Hence the pH range of one solvent may differ from that of another solvent. Data from Sarmini and Kenndler. It must be stressed that when performing pK a measurements, all parameters mentioned will have to be kept constant in order to produce a meaningful result.

It is often overlooked that this is also the case with pH measurements. Before use, pH meters should be calibrated under the same conditions of temperature, ionic strength, and solvent. Reporting pK a values in literature also require that exact conditions of temperature, ionic strength, and solvent be stated.

If these details are omitted, it cannot be assumed they were measured in water, at room temperature, and extrapolated to infinite dilution. Many of the techniques mentioned in the subsequent section measure solutions in which not only the analyte is present, but also various other components for buffering. One has to be sure that there is no ionic or other interaction between the analyte and these other components.

Depending on the sample and matrix under investigation, the choice of technique can be a difficult one, even for the case of monovalent ions, to which this paper is limited to. For multivalent components, matters are more complicated as the pK a differences are smaller.

This is because all methods, except perhaps for nuclear magnetic resonance NMR , require curve fitting in addition to the normal calculation procedure for the respective technique. Investigating both acidic and basic pK a of amphoteric compounds also requires curve fitting in a much broader range.

For such compounds, especially peptides and proteins, the isoelectric point is often relevant for identification purposes but beyond the scope of this paper.

Figure 5 displays an overview of the first time these techniques where used for this purpose. Timeline of the first notion of the various techniques to determine pK a dissociation constant, acid strength. The simplicity and low cost of potentiometric titration has made it one of the most commonly used methods for pK a determination.

In a potentiometric titration, a known volume of reagent is added stepwise to a solution of analyte. The change in potential E upon reaction is consequently measured with the use of two electrodes, an indicator, and a reference electrode. These are often integrated in what is now commonly called a combined pH electrode. Plotting the potential versus volume subsequently gives rise to a sigmoid curve, where the inflection point gives the potential at equilibrium.

With the use of standards with known pH, this potential can be linearly converted into a pH, equaling pK a.

Increasing understanding of electrochemical processes in the late 19th century gave rise to the first potentiometer. The first description of the use of a setup to determine equilibrium constants was made by Denham in It was not long until the cumbersome hydrogen electrode was replaced by the familiar glass electrode.

Completely automated and self-adjusting pH-measuring equipment by Keeler is known from as early as Because temperature influences not only the measurement but also the pK a itself, it is of vital importance to conduct the titration at constant temperature. A good review about the various errors of the electrode itself is given by Gardiner 19 while Benett 20 discusses the various ways to determine pK a from the measured potential.

Another practical complication is the pK a -measurement of substances with a low water-solubility. One example is extrapolation of measurements in solvent mixtures.

A precise determination of the pH from the titration slope has also been difficult in the earlier period of its use. However, over the years various software programs have become available to minimize most of the previous mentioned errors.

Potentiometric titration requires a relatively large amount of sample compared with separation methods such as high performance liquid chromatography HPLC and capillary electrophoresis CE. Completely automated potentiometric pH meters for a wide range of applications and with complicated calibration software are widely available. Because of this and the simplicity and relatively low costs associated with the potentiometric pH-meter, it will probably remain in use in the foreseeable future.

This of course only holds for those analytes that are available in sufficient quantity and purity. The determination of acid dissociation constants by conductometry relies on the assumption that strong electrolytes are completely dissociated at all concentrations, while weak electrolytes only attain complete dissociation at infinite dilution.

A measurement of the conductivity of a sample yields a value that is the sum of the independent contributions of all ions present in the solution:. It can also be seen from Figure 6 , however, that this linear extrapolation does not hold for weak acids or bases. This is because for these electrolytes, the assumption that all ions are independent of their counter ions does not hold as these species are not fully dissociated.

Though the limiting conductance of a weak electrolyte cannot be directly extrapolated, their value can still be obtained quite readily. Even for weak electrolytes, the Kolhrausch law 23 of the independent migration of ions holds. This law can be stated as meaning that at infinite dilution, each ion makes a specific contribution to the conductivity regardless of the ions associated with it.

Expressing the limiting conductance of a salt in terms of ionic contributions is useful as it allows the calculation of the limiting conductance of a weak electrolyte. In order to do this, the limiting conductance of other relevant salts of a strong electrolyte is extrapolated. For a weak acid formic acid relevant plots are shown in Figure 6.

One would obtain the ionic contributions for the limiting conductance of the acid from the salt of its conjugate base and from hydrochloric acid, subtracting the value for sodium chloride to eliminate the ionic contributions of the sodium and chlorine counter ions in the reference compounds.

This yields the sum shown in Equation 8 :. Equation 9 yields the dissociation constant for a given analytical concentration c and a series of measurements would yield the dissociation constant as a function of ionic strength.

It is worth noting that this determination does not require knowledge of the pH of the solution, making it easily applicable to non-aqueous systems where such a measurement of pH would be impracticable. It also means that the method, unlike many others which express their measured quantity as a function of pH, is not constrained by the precision of the pH electrode.

This requires working with pure compounds. Foundational work that enabled the development of the conductometric method was undertaken by Friedrich Kohlrausch. Among the key ideas he developed was the use of alternating current to prevent electrolysis during conductivity measurements. His work examining a variety of electrolyte solutions led to the law of independent migration of ions that is ascribed to him.

Building on this early work, and further contributions by Ostwald and Arrhenius, the method reached maturity during the late s, and the accuracy of the method approached that of modern methods. Work on the conductometric method came to a virtual standstill after the outbreak of World War II, and further research was practically abandoned after the end of the war.

Renewed activity would not come until the s, when two new conductance equations were published. These equations enabled the study of asymmetrical electrolytes, and even mixtures of electrolytes. Since then, it has been remarked that the spectacular interest for the method that existed during its early period has waned with time, leaving the subject a somewhat unfashionable topic of research.

With the developments made over its history, however, the method has achieved a high degree of precision. In voltammetry, a changing potential is applied over a sample solution and the resulting current is measured. When the potential reaches the reduction potential of the analyte this will give rise to an increase in current, followed by a decrease due to depletion of the molecule. In the case of cyclic voltammetry, for example, this will lead to results much like those shown in Figure 7.

16.4: Acid Strength and the Acid Dissociation Constant (Ka)

The magnitude of the equilibrium constant for an ionization reaction can be used to determine the relative strengths of acids and bases. The equilibrium constant for this reaction is the base ionization constant K b , also called the base dissociation constant:. Notice the inverse relationship between the strength of the parent acid and the strength of the conjugate base. Thus the conjugate base of a strong acid is a very weak base, and the conjugate base of a very weak acid is a strong base. We can use the relative strengths of acids and bases to predict the direction of an acid—base reaction by following a single rule: an acid—base equilibrium always favors the side with the weaker acid and base, as indicated by these arrows:. Hence the ionization equilibrium lies virtually all the way to the right, as represented by a single arrow:. Similarly, in the reaction of ammonia with water, the hydroxide ion is a strong base, and ammonia is a weak base, whereas the ammonium ion is a stronger acid than water.

The acid dissociation constant K a is a quantitative measure of the strength of an acid in solution. K a is the equilibrium constant for the following dissociation reaction of an acid in aqueous solution:. The K a expression is as follows:. Acid dissociation constants are most often associated with weak acids, or acids that do not completely dissociate in solution. This is because strong acids are presumed to ionize completely in solution and therefore their K a values are exceedingly large.

Salt effect on the dissociation constant of acid-base indicators

An acid dissociation constant , K a , also known as acidity constant , or acid-ionization constant is a quantitative measure of the strength of an acid in solution. It is the equilibrium constant for a chemical reaction. The dissociation constant is defined by [note 2].

The acid dissociation constant pK a is among the most frequently used physicochemical parameters, and its determination is of interest to a wide range of research fields. We present a brief introduction on the conceptual development of pK a as a physical parameter and its relationship to the concept of the pH of a solution. This is followed by a general summary of the historical development and current state of the techniques of pKa determination and an attempt to develop insight into future developments. Fourteen methods of determining the acid dissociation constant are placed in context and are critically evaluated to make a fair comparison and to determine their applications in modern chemistry. Additionally, we have studied these techniques in light of present trends in science and technology and attempt to determine how these trends might affect future developments in the field.

Yahya Nural, H. Proton affinities of potential donor atoms of the ligands were calculated by AM1 and PM3 semiempiric methods. We found, potentiometrically, three different acid dissociation constants for 1a—f.

Dissociation extraction

The effect of the ionic environment on the dissociation constant of acid-base indicators is accounted for by using the specific interaction theory SIT of Broensted, Scatchard, and Guggenheim.

Salt effect on the dissociation constant of acid-base indicators

The acid-base equilibrium of hexamethylenetetramine hexamine was analyzed with its effective electrophoretic mobility by capillary zone electrophoresis. Although hexamine is degradable in a weakly acidic aqueous solution, and the degraded products of ammonia and formaldehyde can be formed, the effective electrophoretic mobility of hexamine was measured in the pH range between 2. The monoprotic acid-base equilibrium of hexamine was confirmed through comparisons of its electrophoretic mobility with the N -ethylquinolinium ion and with the monocationic N -ethyl derivative of hexamine, as well as a slope analysis of the dissociation equilibrium. Already have an account? Login in here.

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International Journal of Analytical Chemistry

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2 Comments

Brenda D.
30.05.2021 at 15:35 - Reply

In this research work, a potentiometric technic was used to measure the acidic dissociate constants pK a ,s for glycyl aspartic acid GLY-ASP at temperatures

Patience L.
03.06.2021 at 06:24 - Reply

Conjugate acid-base pairs. Base- dissociation constant, K. KK = K. Solutions of strong acids and bases. Neutralization. Titration.

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