Acid Dissociation Constant
The acid dissociation constant (Ka) is a measure of the strength of an acid in a solution. It represents the equilibrium constant for the dissociation of an acid into its ions. A higher Ka value indicates a stronger acid, while a lower Ka value indicates a weaker acid.
6 Key excerpts on "Acid Dissociation Constant"
eBook - ePubPhysicochemical and Biomimetic Properties in Drug Discovery
Chromatographic Techniques for Lead Optimization
- Klara Valko(Author)
- 2013(Publication Date)
- Wiley(Publisher)
...Chapter 8 Molecular Physicochemical Properties that Influence Absorption and Distribution—Acid Dissociation Constant—pKa Definition of p K a The presence of charge on the molecules dramatically influences many of their physicochemical properties, such as lipophilicity, solubility, and permeability. The presence of charge depends on the Acid Dissociation Constant of the ionizable groups and the pH of the solution/environment. The pH is defined as the negative logarithm of the proton or, more precisely, the hydronium ion concentration in aqueous solutions. The product of the concentrations of hydronium and hydroxide ions in water is constant ; thus, the pH normally ranges from 1 to 14. The Acid Dissociation Constant, or, is defined as the pH where an ionizable group is 50% in ionized form. In other words, the Acid Dissociation Constant,, is the equilibrium constant for the reaction in which a weak acid is in equilibrium with its conjugate base in aqueous solution. For example, for acetic acid, the following equilibrium takes place: 8.1 8.2 When the acetate ion concentration is equal to the acetic acid concentration, equals the concentration. The negative logarithm of the concentration is the pH. The smaller the value of, the stronger is the acid. For basic compounds, Equation 8.3 and Equation 8.4 can be used. 8.3 8.4 Again, the negative logarithm of equals the pH of the aqueous environment, where 50% of the basic group is in a protonated charged form, while 50% is in a neutral, unionized form...
eBook - ePub- Jeffrey Gaffney, Nancy Marley(Authors)
- 2017(Publication Date)
- Elsevier(Publisher)
...The strength of a weak acid in an aqueous solution is defined by the extent that it ionizes in water. The extent that a weak acid ionizes in water is expressed quantitatively by an acid ionization constant (K a), which is a form of equilibrium constant described in detail in Chapter 7. The acid ionization constant is equal to the ratio of the molar concentrations of the ionized products (conjugate acid and conjugate base) to the molar concentration of the unionized acid at equilibrium; K a = H 3 O + A − HA w (6) Although water is a reactant in the ionization of a weak acid, the concentration of water is not used in the expression of the acid ionization constant. This is because, as the solvent, the concentration of water is in great excess compared to the concentrations of the acid and the ionization products. So, the concentration of water remains constant during the reaction. Only the concentrations of the species changed by the reaction are included in the expression for the acid equilibrium constant. The stronger the acid and the more it is ionized, the higher the concentrations of the conjugate acid and conjugate base (numerator of K a) and the smaller the concentration of the unionized acid (denominator of K a), resulting in a larger K a. The weaker the acid and the less it is ionized, the lower the concentrations of the conjugate acid and conjugate base and the higher the concentration of the unionized acid, resulting in a smaller K a...
eBook - ePub- Professor Rob Beynon, J Easterby(Authors)
- 2004(Publication Date)
- Taylor & Francis(Publisher)
...It is time to introduce the Henderson–Hasselbalch equation. As with all equilibria, we can express the dissociation of a weak acid or base in terms of an equilibrium constant. For the rest of this section, we will develop a general solution for an acid HA, but illustrate with two examples, acetic acid and Tris, two commonly used buffers. Consider once again the equation that introduces the equilibrium constant K a. K a = [ H + ] [ A − ] [ HA ] K a is a constant, but is different for each buffer. For acetic acid/acetate it is 10 –4.76 M, for Tris it is 10 –8.1 M. What is the significance of these numbers, which seem to span a wide range, even judging by these two examples? To make things clearer, we shall modify this equation, using logarithmic transformation. First we rewrite the equation in a slightly different, but equivalent form: K a = [ H + ] [ A − ] [ HA ] ◊ The corresponding equations for acetic acid and Tris buffers are: K a = [ H + ] [ CH 3 COO − ] [ CH 3 COOH ] K a = [ H + ] [ Tris ] [ TrisH + ] Take log 10 of each side of the equation, and remembering that when we have two terms multiplied together, we add their. logarithms: log 10 K a = log 10 [ H + ] + log 10 [ A − ] [ HA ] Now, rearrange the equations to a different form, by swapping the two leftmost terms over, changing their signs as we do so: − log 10 [ H + ] = − log 10 K a + log 10 [ A − ] [ HA ] The left hand term should now be familiar: ‘−log 10 [H + ]’ is of course, our definition of pH (Chapter 1). Similarly, we recognise an equivalent construction in the term ‘−log 10 (K a)’ and can refer to this as p K a —we had already met this term briefly in Chapter 2. The equation now simplifies to: pH = p K a + log 10 [ A − ] [ HA ] ◊ In words: ‘The pH of a solution of a weak acid or base is given by the p K a plus the log (base 10) of the ratio of the concentrations of base to acid.’ This is the Henderson–Hasselbalch equation...
- Kevin Reel, Derrick C. Wood, Scott A. Best(Authors)
- 2014(Publication Date)
- Research & Education Association(Publisher)
...For example, at 100°C, water exists as both a liquid and a gas in equilibrium with one another. TEST TIP If pressures are given for products and reactants in an equilibrium, be sure to write the expression for K p and not K c. EXAMPLE: The value of K p for the following reaction is 8.3 × 10 –3 at 700 K. What is the value for K c ? SOLUTION: Δn = 2 moles gaseous product – 4 moles gaseous reactants = –2 The Dissociation of Weak Acids and Bases • The water dissociation constant is shown in the following reaction: This relationship allows for computation of dissociation constants for conjugate acid–base pairs as well as concentrations of H 3 O + and OH – for use in pH and pOH calculations. • The equilibrium constants K a and K b can be used to calculate the pH of a weak acid or weak base in aqueous solution. EXAMPLE: What is the pH of a 0.10 molar solution of acetic acid, CH 3 COOH, which has a K a = 1.8 × 10 –5 ? SOLUTION: 1. Write the equation of the reaction that occurs when the weak acid or base is put in water, and its corresponding equilibrium expression. When the weak acid is placed in water, some number of moles will dissociate. For each mole of weak acid that dissociates, one mole of the proton and one mole of the weak base will be formed. 2. Construct a table to determine the amount of dissociation of the weak acid. For the AP Chemistry exam, the simplifying assumption can always be made for all types of equilibrium problems. 3. Write the equilibrium expression using the final concentrations from the preceding table. 4. Solve for x, which will be the [H 3 O + ] at equilibrium. 5. When calculating pH or % dissociation, always use the equilibrium [H 3 O + ]. TEST TIP For any problems that involve equilibrium, you must create a table as shown in step 2 of the previous problem in order to determine equilibrium concentrations...
eBook - ePub- Arnost Kotyk, Jan Slavik(Authors)
- 2020(Publication Date)
- CRC Press(Publisher)
...That is why, for instance, the dissociation constants shown here are actually apparent dissociation constants that may differ substantially from the true, thermodynamic dissociation constants based on activities. These apparent dissociation constants are often distinguished by a prime from the true ones (K ′ vs. K), but the simple notation will be used throughout. A weak monobasic acid, such as are mostly encountered in biological systems, dissociates in water according to HA + H 2 O ⇌ A − + H 3 O + (11) with the corresponding dissociation constant being defined as K A = c H 3 O + ⋅ c A − / c H 2 O = c T α 2 / (1 − α) (12) where c T = c A− + c HA and α, the degree of ionization, is equal. to c A -/ c T or c H 3 O + / c T. This This degree increases toward unity as c T approaches zero. (With strong acids α ≅ 1 even at moderate concentrations.) From the above equation c H 3 O + = c F α =1/2 (− K A + K A 2 + 4 c T K A) (13) For weak acids, at medium concentrations α ⪡ 1 so that C T = C HA and c H 3 O + ≅ c T K A (14a) or pH ≅ 1/2 (p K A − log c A) (14b) The use of this relatively simple formula is justified (i.e., the error of determination is less than 1%) only with certain limitations. The value of c H 3 O + defined by Equation 14a differs from that defined by Equation 13 in the expression 1 / 2 K A / c T as may be shown by the following consideration. It follows from Equation 12 that α 2 = K A (1 − α) / c T (15a) In the first approximation (for α very small), then, α 2 = K A / c T (15b) while in the second approximation α 2 = (K A / c T) (1 − K A / c T) (15c) For K A / c T to be less than 0.01 in the expression for α, it must be approximately less than 0.02 in the expression for α 2, so that K A /c T must be less than 4 · 10 −4. Hence, for an acid with a dissociation constant of 8.7 · 10 −4 mol dm −3, such as citric acid, a more than 2 M concentration would be required to justify the use of the simplified Equation 14a...
eBook - ePub- James F. Pankow(Author)
- 2019(Publication Date)
- CRC Press(Publisher)
...5 Quantitative Acid/Base Calculations for Any Solution of Acids and Bases 5.1 Introduction As discussed in Chapter 2, the final equilibrium position that a given aqueous system will take is determined by: (1) the T - and P -dependent value(s) of the pertinent equilibrium constant(s); (2) the mass balance and other equation(s) governing the system; and (3) how the activity coefficients γ i depend on chemical composition. The γ i will be determined by the final equilibrium composition of the system, and so will not be exactly knowable a priori for use in the calculations. However, for dilute solutions, or when the reaction medium contains relatively large amount(s) of background dissolved salt(s) not participating in the reactions, then the γ i in the final equilibrium system may be estimated accurately. 5.2 Solution of the Generic Acid HA, All γ i = 1 5.2.1 Introduction There are four species in a solution of HA (in addition to H 2 O): H + A − OH − HA. This means there are four unknowns [H + ], [A − ], [OH − ], and [HA]. We will assume all γ i = 1 so that concentrations may be used in the equilibrium K expressions rather than activities. How we can address cases when γ i ≠ 1 will be considered in Section 5.7. With four unknowns, four independent equations are required to solve the problem: K w = [H + ][OH] first chemical equilibrium equation (5.1) K a = [ H + ] [ A − ] [ HA ] second chemical equilibrium equation (5.2) [HA] + [A − ] = A T = C mass balance equation (MBE) on total A (5.3) [ H + ] = [A − ] + [OH − ] electroneutrality equation (ENE). (5.4) For our initial discussions in this chapter, the subscript “a” has been included in the K for the Acid Dissociation Constant of HA. Later, we will drop this subscript. The variable C has been introduced as synonymous with A T. This is because C is commonly used in treatments of this problem by others...
