Chemical Equilibria ChemEquilibrium, Lecture notes of Chemistry

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A N A LY T I C A L C H E M I S T R Y C H E M 2 0 0 9. 1

Lecture-7:

Chemical Equilibria-

2021-2022 Fall Semester

What we have learned about the concept of equilibrium

 Chemical Reactions: The Rate Concept

in equilibrium

molar equilibrium constant

Reversible reaction

Attainment of equilibrium in a reversible reaction

What we have learned about the concept of equilibrium

 Le Châtelier’s Principle

 Le Châtelier’s Principle describes how the position of equilibria changes to favor the forward or backward reaction

 Equilibrium shifts to reduce change.

 When conditions are changed the reaction will do everything it can to counteract the change.

 Therefore if the temperature changes the reaction will shift to favor the side that will reduce the temp (e.g. endothermic).

 If the concentration increases on one side they are more likely to react together shifting the equilibrium so the concentration decreases.

 Increasing the pressure shifts the reaction to favor the reaction that produces a smaller number of products.

What we have learned about the concept of equilibrium

 Increasing temperature makes the reaction go in the

endothermic direction

 Decreasing temperature makes the reaction

go in the exothermic direction

 A + B ⇄ C ΔH = - 100 kJmol-^1

 If I heat up the reaction will it go to the left or to the right?

Changing temperature

What we have learned about the concept of equilibrium

 Increasing concentration sends the equilibrium towards

the opposite side

 Decreasing concentration sends the equilibrium towards

the same side

 If I increase concentration of C,

which way will this reaction go?

 2A+B ⇄ C + D

Changing Concentration

What we have learned about the concept of equilibrium

Table: Equilibria and Equilibrium Constants Important in Analytical Chemistry

Activity and Activity Coefficients

 Generally, the presence of diverse salts (not containing ions

common to the equilibrium involved) will cause an increase in dissociation of a weak electrolyte or in the solubility of a precipitate.

 Cations (+) attract anions (-), and vice versa, and so the

cations of the analyte attract anions of the diverse electrolyte and the anions of the analyte are surrounded by the cations of the diverse electrolyte.

 The attraction of the ions participating in the equilibrium of

interest by the dissolved electrolyte effectively shields them, decreasing their effective concentration and shifting the equilibrium.

Activity and Activity Coefficients

 We say that an “ion atmosphere” is formed about the cation

and anion of interest.

 As the charge on either the diverse salt or the ions of the

equilibrium reaction is increased, the diverse salt effect generally increases.

 This “effective concentration” of an ion in the presence of an

electrolyte is called the activity of the ion.

 To quantitatively describe the effects of salts on equilibrium

constants, one must use activities, not concentrations.

 In potentiometric measurements, it is activity that is measured,

not concentration.

CALCULATION OF ACTIVITY COEFFICIENTS

 In 1923, Debye and Hückel derived a theoretical expression for

calculating activity coefficients.

 The original Debye-Hückel equation is given as Equation 7 .1a

but it is of limited use as it can be used only in extremely dilute solutions:

 They later provided a more useful equation, known as the

Extended Debye - Hückel equation :

(7.1a)

(7.1b)

CALCULATION OF ACTIVITY COEFFICIENTS

» A= 0.51 and B= 0.33 are constants for water at 25 oC.

» At other temperatures, the values can be computed from

and

» where D is the dielectric constant and T is the absolute temperature;

» ai is the ion size parameter , which is the effective diameter of the hydrated ion in angstrom units, A˚. (An angstrom is 10−10^ m).

» A limitation of the Debye-Hückel equation is the accuracy to which ai can be evaluated.

(7.1b)

CALCULATION OF ACTIVITY COEFFICIENTS

 Table 7.1: The list of ion size parameters for some common ions.

CALCULATION OF ACTIVITY COEFFICIENTS

 Table 7.1 (cont.): The list of ion size parameters for some

common ions.

CALCULATION OF ACTIVITY COEFFICIENTS

Example 7.2:

» Calculate the activity coefficients for K+^ and SO 4 2−^ in a 0.020 M solution of potassium sulfate.

Solution:

» The ionic strength is 0.060, so we would use Equation 7 .1b.

» From Table 7. 1 , we find that a K+ = 3 ˚A and a( SO4) 2 − = 4.0 A˚. For K+, we can use Equation 7 .2:

CALCULATION OF ACTIVITY COEFFICIENTS

Solution:

» For SO 4 2−, use Equation 7.1b:

» This latter compares with a calculated value of 0.396 using Equation 7 .2.

» Note the decrease in the activity coefficients compared to 0. M K 2 SO 4 , especially for the SO 4 2−^ ion. from 0.713 to 0.