Finding Ka

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  Advanced Equilibrium Chemistry

  • (Author)
  • 2014(Publication Date)
  • Library Press

    (Publisher)

  These calculations find application in many different areas of chemistry, biology, medicine, and geology. For example, many compounds used for medication are weak acids or bases, and a knowledge of the p K a values, together with the water–octanol partition coefficient, can be used for estimating the extent to which the compound enters the blood stream. Acid dissociation constants are also essential in aquatic chemistry and chemical oceanography, where the acidity of water plays a funda-mental role. In living organisms, acid-base homeostasis and enzyme kinetics are depen-dent on the p K a values of the many acids and bases present in the cell and in the body. In chemistry, a knowledge of p K a values is necessary for the preparation of buffer solutions and is also a prerequisite for a quantitative understanding of the interaction between acids or bases and metal ions to form complexes. Experimentally, p K a values can be determined by potentiometric (pH) titration, but for values of p K a less than about 2 or more than about 11, spectrophotometric or NMR measurements may be required due to practical difficulties with pH measurements. Definitions According to Arrhenius's original definition, an acid is a substance that dissociates in aqueous solution, releasing the hydrogen ion H + (a proton): HA A − + H + . ________________________ WORLD TECHNOLOGIES ________________________ The equilibrium constant for this dissociation reaction is known as a dissociation cons-tant. The liberated proton combines with a water molecule to give a hydronium (or oxonium) ion H 3 O + , and so Arrhenius later proposed that the dissociation should be written as an acid–base reaction: HA + H 2 O A − + H 3 O + . Brønsted and Lowry generalised this further to a proton exchange reaction: acid + base conjugate base + conjugate acid. The acid loses a proton, leaving a conjugate base; the proton is transferred to the base, creating a conjugate acid.

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  Chemistry of Carboxylic Acid

  • (Author)
  • 2014(Publication Date)
  • Research World

    (Publisher)

  These calculations find application in many different areas of chemistry, biology, medicine, and geology. For example, many compounds used for medication are weak acids or bases, and a knowledge of the p K a values, together with the water–octanol partition coefficient, can be used for estimating the extent to which the compound enters the blood stream. Acid dissociation constants are also essential in aquatic chemistry and chemical oceanography, where the acidity of water plays a fundamental role. In living organisms, acid-base homeostasis and enzyme kinetics are dependent on the p K a values of the many acids and bases present in the cell and in the body. In chemistry, a knowledge of p K a values is necessary for the preparation of buffer solutions and is also a prerequisite for a quantitative understanding of the interaction between acids or bases and metal ions to form complexes. Experimentally, p K a values can be determined by potentiometric (pH) titration, but for values of p K a less than about 2 or more than about 11, spectrophotometric or NMR measurements may be required due to practical difficulties with pH measurements. Definitions According to Arrhenius's original definition, an acid is a substance that dissociates in aqueous solution, releasing the hydrogen ion H + (a proton): HA A − + H + . ________________________ WORLD TECHNOLOGIES ________________________ The equilibrium constant for this dissociation reaction is known as a dissociation constant. The liberated proton combines with a water molecule to give a hydronium (or oxonium) ion H 3 O + , and so Arrhenius later proposed that the dissociation should be written as an acid–base reaction: HA + H 2 O A − + H 3 O + . Brønsted and Lowry generalised this further to a proton exchange reaction: acid + base conjugate base + conjugate acid. The acid loses a proton, leaving a conjugate base; the proton is transferred to the base, creating a conjugate acid.

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  Solution and Solubility in Chemistry

  • (Author)
  • 2014(Publication Date)
  • Library Press

    (Publisher)

  These calculations find application in many different areas of chemistry, biology, medicine, and geology. For example, many compounds used for medication are weak acids or bases, and a knowledge of the p K a values, together with the water–octanol partition coefficient, can be used for estimating the extent to which the compound enters the blood stream. Acid dissociation constants are also essential in aquatic chemistry and chemical oceanography, where the acidity of water plays a fun-damental role. In living organisms, acid-base homeostasis and enzyme kinetics are dependent on the p K a values of the many acids and bases present in the cell and in the body. In chemistry, a knowledge of p K a values is necessary for the preparation of buffer solutions and is also a prerequisite for a quantitative understanding of the interaction between acids or bases and metal ions to form complexes. Experimentally, p K a values can be determined by potentiometric (pH) titration, but for values of p K a less than about 2 or more than about 11, spectrophotometric or NMR measurements may be required due to practical difficulties with pH measurements. Definitions According to Arrhenius's original definition, an acid is a substance that dissociates in aqueous solution, releasing the hydrogen ion H + (a proton): HA A − + H + . ________________________ WORLD TECHNOLOGIES ________________________ The equilibrium constant for this dissociation reaction is known as a dissociation constant. The liberated proton combines with a water molecule to give a hydronium (or oxonium) ion H 3 O + , and so Arrhenius later proposed that the dissociation should be written as an acid–base reaction: HA + H 2 O A − + H 3 O + . Brønsted and Lowry generalised this further to a proton exchange reaction: acid + base conjugate base + conjugate acid. The acid loses a proton, leaving a conjugate base; the proton is transferred to the base, creating a conjugate acid.

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  Physicochemical 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 pKa

  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. The percentage of the ionized molecules depends on the proton concentration (pH) and can be calculated at any pH using the Henderson–Hasselbalch equation [1]. Equation 8.5 describes the relationship between the percentage of ionized molecules and pH for a given

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  Chemistry

  The Molecular Nature of Matter

  • Neil D. Jespersen, Alison Hyslop(Authors)
  • 2021(Publication Date)
  • Wiley

    (Publisher)

  Also provided is information from which we can obtain directly at least one of the equilibrium concentrations. Thus, we might be given the measured pH of the solution, which provides the equilibrium concentration of H + . (Equilib- rium is achieved rapidly in these solutions, so when we measure the pH, the value obtained can be used to calculate the equilibrium concentration of H + .) Alternatively, we might be given the percentage ionization of the acid or base, which we define as follows: percentage ionization = moles ionized per liter ____________________ moles available per liter × 100% (16.15) Let’s look at some examples that illustrate how to determine K a and K b from the kind of data mentioned. NOTE Laws are usually mathematical equations. Therefore the mass action law is often called the mass action equation. TOOLS Percentage ionization FIGURE 16.3 The relative strengths of conjugate acid–base pairs. The stronger the acid is, the weaker is its conjugate base. The weaker the acid is, the stronger is its conjugate base. Very strong acids are 100% ionized, and their conjugate bases do not react with water to any measurable extent. Very strong acids; 100% ionized in water C 2 H 3 O 2 – OCl – NH 3 OH – NH 2 – NO 2 – F – H 2 O Cl – NO 3 – ClO 4 – Base strength increases Very strong base; reacts 100% with water Strongest proton acceptor that can exist in water Weak bases in water Very weak bases; do not react with water to a measurable extent Acid strength increases HC 2 H 3 O 2 HOCl NH 4 + H 2 O NH 3 HNO 2 HF H 3 O + HCl HNO 3 HClO 4 Strongest proton donor that can exist in water Weak acids in water Very weak acid; does not react with water as an acid 16.4 Determining K a and K b Values 805 EXAMPLE 16.1 Calculating K a and pK a from pH Lactic acid (HC 3 H 5 O 3 ), which is present in sour milk, also gives sauerkraut its tartness. It is a monoprotic acid. In a 0.100 M solution of lactic acid, the pH is 2.44 at 25 °C.

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  Chemistry

  The Molecular Nature of Matter

  • James E. Brady, Neil D. Jespersen, Alison Hyslop(Authors)
  • 2014(Publication Date)
  • Wiley

    (Publisher)

  Another way to compare the relative strengths of acids or bases is to use the negative logarithms of K a and K b , called pK a and pK b , respectively. pK a = -log K a pK b = -log K b Stronger acids or bases have smaller pK a or pK b values, respectively. Describe how to determine acid and base ionization constants from experimental data The values of K a and K b can be obtained from initial concentra- tions of the acid or base and either the pH of the solution or the percentage ionization of the acid or base. The measured pH gives the equilibrium value for [H + ]. Determine equilibrium concentrations and pH for weak acid or base solutions Problems fall into one of three categories: (1) the only solute is a weak acid (we must use the K a expression), (2) the only solute is a weak base (we must use the K b expression), and (3) the solution contains both a weak acid and its conjugate base (we can use either K a or K b ). When the initial concentration of the acid (or base) is larger than 100 times the value of K a (or K b ), it is safe to use initial concentrations of acid or base as though they were equilibrium values in the mass action expression. When this approximation cannot be used, we can use the quadratic formula. Explain how a salt solution can be acidic or basic A solution of a salt is acidic if the cation is acidic but the anion is neutral. Metal ions with high charge densities generally are acidic, but those of Groups 1A and 2A (except Be 2+ ) are not. Cations such as NH 4 + , which are the conjugate acids of weak molecular bases, are themselves proton donors and are acidic. When the anion of a salt is the conjugate base of a weak acid, the anion is a weak base. Anions of strong acids, such as Cl - and NO 3 - , are such weak Brønsted bases that they cannot affect the pH of a solution.

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  Rapid Review of Chemistry for the Life Sciences and Engineering

  With Select Applications

  • Armen S. Casparian, Gergely Sirokman, Ann Omollo(Authors)
  • 2021(Publication Date)
  • CRC Press

    (Publisher)

  b , respectively.

  For example, acetic acid (found in vinegar) is a weak acid. Its equilibrium is represented as follows:

  CH 3

  COOH

  ( aq )

  +

  H 2

  O

  ( l )

  ⇋

  H 3

  O +

  ( aq )

  +

  CH 3

  COO −

  ( aq )

  (4.6)

  Here, H3 O+ is the active acid species, and CH3 COO− is called the conjugate base of acetic acid. Recall that the double arrows indicate an equilibrium process, meaning that this reaction does not go to completion and that all four chemical species in the reaction are present at any given time in varying concentrations.

  The equilibrium acid constant expression, Ka , is as follows:

  K a

  =

  [

  H 3

  O +

  ]

  [

  CH 3

  COO −

  ]

  [

  CH 3

  COOH

  ]

  (4.7)

  The Ka value for acetic acid at room temperature is 1.8 × 10−5 .

  The average percent ionization, depending on initial concentration of the parent acid and temperature, is about 3%–5%. Generally speaking, the smaller the value of the Ka , the weaker the acid. Table 4.1 (A) gives Ka values for a number of weak acids at 25°C. It is evident then, from comparing Ka values of acetic acid with nitrous acid, that acetic acid is weaker than nitrous acid. Phenol, in turn, is much weaker than either nitrous acid or acetic acid. Polyprotic acids have more than one hydrogen ion to dissociate and ionize and so have more than one dissociation constant, i.e., Ka 1 , Ka 2 , and Ka 3 , which get successively smaller, indicating progressive weakness in acidity. For example, oxalic acid has two Ka values, while citric acid has three.

  TABLE 4.1

  Equilibrium or Dissociation Constants for (A) Weak Acids K

  a

  (Monoprotic and Polyprotic) and (B) Weak Bases K

  b

  Monoprotic Acids	Ka
  HC2 O2 CI3	Trichloroacetic acid	(Cl3 CCO2 H)	2.2 × 10−1
  HIO3	Iodic acid	.	1.69 × 10−1
  HC2 HO2 Cl2	Dichloracetic acid	(Cl2 CHCO2 H)	5.0 × 10−2
  HC2 H2 O2 CI	Chloroacetic acid	(ClH2 CCO2 H)	1.36 × 10−3
  HNO2	Nitrous acid	7.1 × 10−4

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  Absorption and Drug Development

  Solubility, Permeability, and Charge State

  • Alex Avdeef(Author)
  • 2012(Publication Date)
  • Wiley

    (Publisher)

  a of a substance is widely useful, even in the simplest of ways. For example, urine pH, normally 5.7–5.8, can be altered sufficiently with oral doses of ammonium chloride or sodium bicarbonate to accommodate reabsorption of uncharged species for therapeutic reasons, or excretion of ionized species in drug overdose/toxicological emergencies [11]. Weak acids may be excreted in alkaline urine while weak bases may be eliminated in acidic urine, a principle that may be lifesaving with overdoses of barbiturates, amphetamines, and narcotics, for example.

  Measuring p

  Ka

  values can be challenging, since many new pharmaceutical substances of interest are very poorly soluble in aqueous solution. Potentiometry can be a reliable technique for p

  Ka

  determination [1–5], provided the solubility of the substance is at least 10−4  M over a substantial pH range. Solutions as dilute as 10−5  M can still be analyzed, but special attention must be given to electrode metrology, and potential impurities (including dissolved carbon dioxide) in solution need to be assessed reliably. If the substance is only soluble to the extent of 10−6  M and possesses an analytically useful chromophore, then spectrophotometric methods need to be applied. Cosolvent methods can improve the sensitivity further, but require extra care and methodology knowledge.

  Although the knowledge of the p

  Ka

  of a molecule is very important in a number of chemical disciplines, the focus here will be directed to pharmaceutical and biopharmaceutical applications.

  3.2 METHODS OF CHOICE FOR THE DETERMINATION OF THE p

  Ka

  The glass-membrane pH electrode and high-impedance pH meters have made the potentiometric method universally applicable for determining p

  Ka

  values [1–5, 12–18]. There are many circumstances that may warrant the determination of the p

  Ka

  by spectrophotometry (UV) [19–41], as well as by capillary electrophoresis (CE) [42–45] and, in some instances, by chromatographic [46] and nuclear magnetic resonance (NMR) techniques [47–49]. The highest-precision p

  Ka

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  Analytical Chemistry

  • Gary D. Christian, Purnendu K. Dasgupta, Kevin A. Schug(Authors)
  • 2020(Publication Date)
  • Wiley

    (Publisher)

  We need to define K a , the dissociation constant for acetic acid. In this problem, HOAc, unlike HCl, is not fully ionized and C A therefore represents the combined concentration of HOAc and OAc - . Proceeding as before 1. The species present are: H 2 O, HOAc, H + , OH - , Na + , and OAc - . 2. The equilibrium expressions are: [H + ][OH - ] = K w [H + ][OAc - ] [HOAc] = K a Pdf_Folio:278 7.5 WEAK ACID VERSUS STRONG BASE—A BIT LESS STRAIGHTFORWARD 279 The expressions of the fractions of dissociating species, as described in Chapter 6 (Section 6.11) can be manipulated to give the -values. In particular, we are inter- ested in the concentration of the acetate ion, [OAc - ]: [OAc - ] =  1 C A where  1 = K a K a + [H + ] Realize also that the after accounting for the total volume, [Cl − ] at any point repre- sents the initial amount of HCl present and similarly, after accounting for the total volume, [Na + ] at any point represents the total amount of NaOH added until that point. 3. The charge balance expression will be: ([Na + ] + [H + ]) - ([OH - ] + [OAc - ]) = 0 (7.15) 4. Accounting for dilution, we have this: [OAc - ] =  1 C A V A ∕(V A + V B ) (7.16) Putting in Equation 7.4, Equation 7.16 reduces to [OAc - ] =  1 C A ∕(1 + fC A ∕C B ) (7.17) For C A = C B as in the present case, this further simplifies to [OAc - ] =  1 C A ∕(1 + f ) (7.18) As in spreadsheet Example 7.2, Eq. 7.7 [Na + ] = fC A ∕(1 + f ) Recognizing that [OH - ] = K w ∕[H + ], we can now write our desired equation by putting Equations 7.7 and 7.18 into Equation 7.15: fC A ∕ (1 + f) + [ H + ] - K w ∕ [ H + ] - C A K a ( K a + [H + ] ) (1 + f) = 0 (7.19) Construct an Excel sheet (or modify the spreadsheet 7.2.xlsx available on the book’s website). (See on the website Section 7.5. spreadsheet for HOAc vs.NaOH.) In addi- tion to CA and KW already having been specified and the names defined, type in KA in cell A3 and 1.75E-5 in cell B3 and define this as KA.

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  Analytical Chemistry

  • Gary D. Christian, Purnendu K. Dasgupta, Kevin A. Schug(Authors)
  • 2013(Publication Date)
  • Wiley

    (Publisher)

  We need to define K a , the dissociation constant for acetic acid. In this problem, HOAc, unlike HCl, is not fully ionized and C A therefore represents the combined concentration of HOAc and OAc − . Proceeding as before 1. The species present are: H 2 O, HOAc, H + , OH − , Na + , and OAc − . 2. The equilibrium expressions are: [H + ][OH − ] = K w [H + ][OAc − ] [HOAc] = K a The expressions of the fractions of dissociating species, as described in Chapter 7 (Section 7.11) can be manipulated to give the α-values. In particular, we are interested in the concentration of the acetate ion, [OAc − ]: [OAc − ] = α 1 C A where α 1 = K a K a + [H + ] Realize also that the after accounting for the total volume, [Cl − ] at any point represents the initial amount of HCl present and similarly, after accounting for the total volume, [Na + ] at any point represents the total amount of NaOH added until that point. 3. The charge balance expression will be: ([Na + ] + [H + ]) − ([OH − ] + [OAc − ]) = 0 (8.15) 294 CHAPTER 8 ACID–BASE TITRATIONS 4. Accounting for dilution, we have this: [OAc − ] = α 1 C A V A /(V A + V B ) (8.16) Putting in Equation 8.4, Equation 8.16 reduces to [OAc − ] = α 1 C A /(1 + fC A /C B ) (8.17) For C A = C B as in the present case, this further simplifies to [OAc − ] = α 1 C A /(1 + f ) (8.18) As in spreadsheet Example 8.2, Eq. 8.7 [Na + ] = fC A /(1 + f ) Recognizing that [OH − ] = K w /[H + ], we can now write our desired equation by putting Equations 8.7 and 8.18 into Equation 8.15: fC A /(1 + f ) + [H + ] − K w /[H + ] − C A K a (K a + [H + ])(1 + f ) = 0 (8.19) Construct an Excel sheet (or modify the spreadsheet 8.2.xlsx available on the book’s website). (See on the website Section 8.5. spreadsheet for HOAc vs.NaOH.) In addition to CA and KW already having been specified and the names defined, type in KA in cell A3 and 1.75E-5 in cell B3 and define this as KA.

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