Chemistry
Dissociation Constant
The dissociation constant is a measure of the extent to which a compound dissociates into its ions in a solution. It is represented by the equilibrium constant for the dissociation reaction. A higher dissociation constant indicates a greater degree of dissociation, while a lower constant indicates less dissociation.
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9 Key excerpts on "Dissociation Constant"
No longer available |Learn more- (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.
No longer available |Learn more- (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.
No longer available |Learn more- (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.
eBook - PDF- 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.
- 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.
eBook - PDF- F. John Gennari, Horacio J. Adrogue, John H. Galla, Nicolaos Maddias, F. John Gennari, Horacio J. Adrogue, John H. Galla, Nicolaos Maddias(Authors)
- 2005(Publication Date)
- CRC Press(Publisher)
The Dissociation Constant ( K 0 or p K 0 ) of a given acid at standard conditions is a measure of its strength because it connotes the ease with which H þ is released from the undissociated acid. The tenacity of H þ binding varies greatly among acid–base pairs. Strong acids, such as hydrochloric and sulfuric acid, dissociate freely when placed in solution because their conjugate bases have virtually no affinity for H þ . For strong acids, the equilibrium concentration of the undissociated acid ([HA]) is negligible and [H þ ] is high. As a result, K 0 is high and p K 0 is low. Most organic acids dissociate only partially and are less strong. For example, acetic acid only dissociates about 1% at equilibrium in solution at 25 C (i.e., 99% exists in the form CH 3 OOH and only 1% exists as CH 3 OO ). These acids have a much lower K 0 and higher p K 0 . BUFFERING When a weak acid is placed in solution, it dissociates only partially and both the undissociated acid and its conjugate base are present. Such solutions have the ability to resist a change in acidity following the addition of another acid or base. This property is termed buffering and the acid– conjugate base pairs are termed buffers. In essence, buffered solutions act like bases with regard to added acids and like acids with respect to added bases. When a strong acid is added to a solution containing a buffer, some of the added H þ combines with the base form of the buffer (increasing the acid form) rather than remaining free in solution. Thus, the increase in [H þ ] (or decrease in pH) is less than would have occurred in the absence of the 6 Gennari and Galla buffer. Similarly, addition of a strong base to a buffered solution limits the decrease in [H þ ] (or increase in pH) that would otherwise occur. Buffering can be illustrated by considering the change in [H þ ] that occurs when 10 mmol of HCl is added to 1 L of water.
eBook - ePub- 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
eBook - PDF- Gary D. Christian, Purnendu K. Dasgupta, Kevin A. Schug(Authors)
- 2013(Publication Date)
- Wiley(Publisher)
Example 6.5 illustrates this principle. Example 6.5 Assume that A and B are an ion pair, which can dissociate into A (a cation) and B (an anion). Recalculate the concentration of A in Example 6.4, assuming that the solution also contains 0.20 M B. Solution We can represent the equilibrium concentration as follows: [AB] [A] [B] Initial 0.10 0 0.20 Change (x = mmol/mL of AB dissociated) −x +x +x Equilibrium 0.10 − x x 0.20 + x ≈ 0.10 ≈ 0.20 The value of x will be smaller now than before because of the common ion effect of B, so we can certainly neglect it compared to the initial concentrations. Substituting in the equilibrium constant expression, (x)(0.20) (0.10) = 3.0 × 10 −6 x = 1.5 × 10 −6 M The concentration of A was decreased nearly 400-fold. The common ion effect can be used to make analytical reactions more favorable Adjusting the pH is a common way of shifting the equilibrium. or quantitative. The adjustment of acidity, for example, is frequently used to shift equilibria. Titrations with potassium dichromate, for example, are favored in acid solution, since protons are consumed in the reaction. Titrations with iodine, a weak oxidizing agent, are usually done in slightly alkaline solution to shift the equilibrium toward completion of the reaction, for example, in titrating arsenic(III): H 3 AsO 3 + I 2 + H 2 O H 3 AsO 4 + 2I − + 2H + 204 CHAPTER 6 GENERAL CONCEPTS OF CHEMICAL EQUILIBRIUM 6.13 SYSTEMATIC APPROACH TO EQUILIBRIUM CALCULATIONS — HOW TO SOLVE ANY EQUILIBRIUM PROBLEM Now that some familiarity has been gained with equilibrium problems, we will consider a systematic approach for calculating equilibrium concentrations that will work with all equilibria, no matter how complex. It consists of identifying the unknown concentrations involved and writing a set of simultaneous equations equal to the number of unknowns.
eBook - ePub- James F. Pankow(Author)
- 2019(Publication Date)
- CRC Press(Publisher)
5 Quantitative Acid/Base Calculations for Any Solution of Acids and Bases5.1 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 = 15.2.1 Introduction
There are four species in a solution of HA (in addition to H2 O):H +A −HA .OH−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=][OH] first chemical equilibrium equation[H+(5.1) K a=second chemical equilibrium equation[H +] [A −][ HA ](5.2) [HA] +] =[A−A T= C mass balance equation (MBE) on total A(5.3) [H +] =] +[A−] electroneutrality equation (ENE) .[OH−(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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