Acid Dissociation Constant

11 Key excerpts on "Acid Dissociation Constant"

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  Chemistry of Carboxylic Acid

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

    (Publisher)

  ________________________ WORLD TECHNOLOGIES ________________________ Chapter- 11 Acid Dissociation Constant Acetic acid, a weak acid, donates a proton (hydrogen ion, highlighted in green) to water in an equilibrium reaction to give the acetate ion and the hydronium ion. Red: oxygen, black: carbon, white: hydrogen. 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 known as dissociation in the context of acid-base reactions. The equilibrium can be written symbolically as: HA A − + H + , where HA is a generic acid that dissociates by splitting into A − , known as the conjugate base of the acid, and the hydrogen ion or proton, H + , which, in the case of aqueous solutions, exists as a solvated hydronium ion. In the example shown in the figure, HA represents acetic acid, and A − the acetate ion. The chemical species HA, A − and H + are said to be in equilibrium when their concentrations do not change with the passing of time. The dissociation constant is usually written as a quotient of the equilibrium concentrations (in mol/L), denoted by [HA], [A − ] and [H + ]: Due to the many orders of magnitude spanned by K a values, a logarithmic measure of the Acid Dissociation Constant is more commonly used in practice. p K a , which is equal to −log 10 K a , may also be (incorrectly) referred to as an Acid Dissociation Constant: ________________________ WORLD TECHNOLOGIES ________________________ The larger the value of p K a , the smaller the extent of dissociation. A weak acid has a p K a value in the approximate range −2 to 12 in water. Acids with a p K a value of less than about −2 are said to be strong acids; a strong acid is almost completely dissociated in aqueous solution, to the extent that the concentration of the undissociated acid becomes undetectable.

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  Advanced Equilibrium Chemistry

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

    (Publisher)

  ________________________ WORLD TECHNOLOGIES ________________________ Chapter- 4 Acid Dissociation Constant Acetic acid, a weak acid, donates a proton (hydrogen ion, highlighted in green) to water in an equilibrium reaction to give the acetate ion and the hydronium ion. Red: oxygen, black: carbon, white: hydrogen. 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 known as dissociation in the context of acid-base reactions. The equilibrium can be written symbolically as: HA A − + H + , where HA is a generic acid that dissociates by splitting into A − , known as the conjugate base of the acid, and the hydrogen ion or proton, H + , which, in the case of aqueous solutions, exists as a solvated hydronium ion. In the example shown in the figure, HA represents acetic acid, and A − the acetate ion. The chemical species HA, A − and H + are said to be in equilibrium when their concentrations do not change with the passing of time. The dissociation constant is usually written as a quotient of the equilibrium concentrations (in mol/L), denoted by [HA], [A − ] and [H + ]: Due to the many orders of magnitude spanned by K a values, a logarithmic measure of the Acid Dissociation Constant is more commonly used in practice. p K a , which is equal to −log 10 K a , may also be (incorrectly) referred to as an Acid Dissociation Constant: ________________________ WORLD TECHNOLOGIES ________________________ The larger the value of p K a , the smaller the extent of dissociation. A weak acid has a p K a value in the approximate range −2 to 12 in water. Acids with a p K a value of less than about −2 are said to be strong acids; a strong acid is almost completely dissociated in aqueous solution, to the extent that the concentration of the undissociated acid becomes undetectable.

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  Solution and Solubility in Chemistry

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

    (Publisher)

  ________________________ WORLD TECHNOLOGIES ________________________ Chapter- 13 Acid Dissociation Constant Acetic acid, a weak acid, donates a proton (hydrogen ion, highlighted in green) to water in an equilibrium reaction to give the acetate ion and the hydronium ion. Red: oxygen, black: carbon, white: hydrogen. 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 known as dissociation in the context of acid-base reactions. The equilibrium can be written symbolically as: HA A − + H + , where HA is a generic acid that dissociates by splitting into A − , known as the conjugate base of the acid, and the hydrogen ion or proton, H + , which, in the case of aqueous solutions, exists as a solvated hydronium ion. In the example shown in the figure, HA represents acetic acid, and A − the acetate ion. The chemical species HA, A − and H + are said to be in equilibrium when their concentrations do not change with the passing of time. The dissociation constant is usually written as a quotient of the equilibrium concentrations (in mol/L), denoted by [HA], [A − ] and [H + ]: Due to the many orders of magnitude spanned by K a values, a logarithmic measure of the Acid Dissociation Constant is more commonly used in practice. p K a , which is equal to −log 10 K a , may also be (incorrectly) referred to as an Acid Dissociation Constant: ________________________ WORLD TECHNOLOGIES ________________________ The larger the value of p K a , the smaller the extent of dissociation. A weak acid has a p K a value in the approximate range −2 to 12 in water. Acids with a p K a value of less than about −2 are said to be strong acids; a strong acid is almost completely dissociated in aqueous solution, to the extent that the concentration of the undissociated acid becomes undetectable.

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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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  Physical Chemistry for Engineering and Applied Sciences

  • Frank R. Foulkes(Author)
  • 2012(Publication Date)
  • CRC Press

    (Publisher)

  CHAPTE R TWENT Y ACID AND BASE DISSOCIATION 20.1 Acid Dissociation ConstantS, K a The dissociation of an acid HA can be represented as a proton transfer reaction involving a Brøn-sted acid: HA H O (aq) 2 + H O A 3 (aq) (aq) + < + with the equilibrium constant 1 defined by K = a a a a H O A HA H O eqm 3 2 + < u u £ ¤ ² ¥ ¦ ´ . . . [1] We saw earlier that the relative activity 2 of the solvent water is given by its mole fraction, which, for most solutions, is close to unity; therefore we usually set a H O 2 = 1 , so that Eqn [1] becomes K a = a a a H O A HA 3 + < u . . . [2] where K a is called the Acid Dissociation Constant . Values of the dissociation constants for several acids and bases are listed in Table 1. We saw earlier that an acid is a substance with a tendency to transfer a proton to another molecule; therefore, it follows that a strong acid is a substance with a strong tendency to transfer a proton to another molecule. Thus, referring to Eqn [3], HA ( acid 1 ) has a tendency to transfer a proton to H 2 O ( base 2 ). HA + H 2 O H 3 O + + A – acid 1 base 2 acid 2 base 1 . . . [3] The extent of the dissociation of acid 1 also depends on the strength of the tendency of base 2 to accept the proton from acid 1 . Therefore, K a depends not only on the acid , but also on the base . For example, HCl is a strong acid (~100% dissociated) in water , but weak in many non-aqueous solvents. Therefore, if we want to measure the relative strength of different acids, we should 1 The relative activities for Acid Dissociation Constants usually are based on the molal concentration scale. 2 For a solution in which the volatile liquids obey Raoult’s Law and the vapor behaves as an ideal gas. 20-2 ACID AND BASE DISSOCIATION compare them against the same base ( base 2 ), which usually is chosen to be water.

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  Basic Analytical Chemistry

  • L. Pataki, E. Zapp, R. Belcher, D Betteridge, L Meites(Authors)
  • 2013(Publication Date)
  • Pergamon

    (Publisher)

  The greater the value of the dissociation constant, the more readily the acid donates its protons, and the stronger the acid. The dissociation constants of some monoprotic acids are given in table 8. CHEMICAL EQUILIBRIA IN SOLUTION 23 Table 8 Dissociation constant of some monoprotic acids at 25 °C Acid HB0 2 HCN HOCl CH 3 COOH HF Cl 2 CHCOOH K a 6-OxlO 10 7-2xl0 10 3-2x10-8 l-SxlO 4 6-7xl0 4 5-6xl0~ 2 The strength of a base is determined by its proton-accepting abil-ity, and thus, obviously, it is always influenced by the nature of the acid which donates the proton. In dilute aqueous solutions, this acid is water, as in the reaction C e H 5 NH 2 + H 2 0 - C e H 5 NH 3 + + OH. basej acid 2 acid! base 2 Applying the law of mass action one obtains: [ C 6 H 5 N H 3 + ] [ O H -] _ ^ o r [ C 6 H 5 N H 3 + ] [ O H -] _ Z j i . _ z [C e H 5 NH 2 ][H 2 0] ' [C e H 5 NH 2 ] * where K h is the dissociation constant of the base in water, which is constant at a given temperature, depending only on the proton-accept-ing ability of the base. The greater its proton-accepting ability, the stronger the base, and, consequently, the higher the numerical value of the dissociation constant. The dissociation constants of some bases are given in table 9. Table 9 Dissociation constants for some bases at 25°C Base C e H 5 NH 2 C e H 5 N HJJNNHJJ NH 3 C^L^¥L 2 K b 4-6xl0 10 2-3x10-» 3-OxlO 6 1·8χ10~ 5 5·6χ10 4 It can be seen from equation (28) that in simple acid-base systems there is a conjugate base for every acid, and vice versa. This implies that in a given solvent, for which the ionic product of autoprotolysis is known, it is sufficient to give the dissociation constant of only one of the components, for example for the acid, in the acid-base pair, because this also defines the strength of the conjugate base.

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  Survival Guide to General Chemistry

  • Patrick E. McMahon, Rosemary McMahon, Bohdan Khomtchouk(Authors)
  • 2019(Publication Date)
  • CRC Press

    (Publisher)

  constant. Example: H 2 SO 4(aq) → H 2 O ← H + (a q) + HSO 4 − (a q) K a = [ H + ] [ H S O 4 − ] [ H 2 S O 4 ] = 6.3 × 1 0 4 HSO 4 − (aq) → H 2 O ← H + (a q) + SO 4 − 2 (a q) K a = [ H[--=PLG. O-SEPARATOR=--]+ ] [ S O 4 − 2 ] [ H S O 4 − ] = 1.2 × 1 0 −2 The equilibrium constant for acid dissociation is often listed as a derived numerical value termed pK a (the “p” refers to “power”, signifying “exponent”). The pK a is defined as the negative logarithm of the numerical value of the K a : pK a = −log 10 (K a) Example: HCl has a K a of 1 × 10 7 ; pK a = − log 10 (1 × 10 7) = − 7 CH 3 COOH has a K a of 1.8 × 10 −5 ; pK a = − log 10 (1.8 × 10 −5) = +4.75 H 2 SO 4 has a K a of 6.3 × 10 4 ; pK a = − log 10 (6.3 × 10 4) = − 4.8 HSO 4 − has a K a of 1.2 × 10 −2 ;. pK a = −log 10 (1.2 × 10 −2) = +1.9 Base strength in aqueous solution is measured as the concentration of aqueous hydroxide ion, [OH − (aq) ]. Many strong bases are hydroxide containing ionic compounds, which have high solubility in water; thus, complete separation of dissolved ions produces high concentrations of hydroxide ion. In these cases, water acts only as the solvent and does not participate in an acid/base reaction. Metal (OH) → H 2 O solvent Metal + (aq) + OH − (aq) (variable solubility) Non-hydroxide bases undergo an acid/base reaction with the role of water as an acid; this reaction with water is the reference reaction for determining base strength. Note that (OH −) is the conjugate base of water as an acid. Non-hydroxide bases that have a base equilibrium constant (K b) with a numerical value greater than one can be considered strong bases

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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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  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

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

    (Publisher)

  Abbreviating H 3 O + as H + , the equation for the ionization of the acid can be simplified as H A m H + + A - from which the expression for K a is obtained directly. K a = 3 H + 4 3 A - 4 3 H A 4 (16.10) You should learn how to write the chemical equation for the ionization of a weak acid and be able to write the equilibrium law corresponding to its K a . For example, if we have a weak acid such as HNO 2 we can write the ionization reaction as HNO 2 m H + + NO 2 - and the corresponding equilibrium law is K a = 3 H + 4 3 NO 3 - 4 3 HNO 3 4 General equation for the ionization of a weak acid ■ Some call K a the Acid Dissociation Constant. The meat products shown here contain nitrite ion as a preservative. Andy Washnik 770 Chapter 16 | Acid–Base Equilibria in Aqueous Solutions For each of the following acids, write the equation for its ionization in water and the appro- priate expression for K a : (a) HC 2 H 3 O 2 , (b) (CH 3 ) 3 NH + , (c) H 3 PO 4 . (Hint: Determine the conjugate base for each of these acids.) For each of the following acids, write the equation for its ionization in water and the appropriate expression for K a : (a) HCHO 2 , (b) (CH 3 ) 2 NH 2 , (c) H 2 PO 4 - . For weak acids, values of K a are usually quite small and can be conveniently represented in a logarithmic form similar to pH. Thus, we can define the pK a of an acid as pK a = -log K a The strength of a weak acid is determined by its value of K a ; the larger the K a , the stronger and more fully ionized the acid. Because of the negative sign in the defining equation for pK a , the stronger the acid, the smaller is its value of pK a . The values of K a and pK a for some typical weak acids are given in Table 16.2. A more complete list is located in Appendix C. Use Table 16.2 to find all the acids that are stronger than acetic acid and weaker than formic acid.

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  Aquatic Chemistry Concepts, Second Edition

  • 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 H2 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 AT . This is because C is commonly used in treatments of this problem by others. We will use C when we wish to emphasize prior treatments that the reader may have seen, and AT

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