CHAPTER 12 ELECTRODE POTENTIALS AND THEIR APPLICATIONS TO XIDATION/REDUCTION TITRATIONS Introduction to Analytical Chemistry

CHAPTER 12ELECTRODE POTENTIALS AND

THEIR APPLICATIONS TO XIDATION/REDUCTION

TITRATIONS

Introduction toIntroduction toAnalytical ChemistryAnalytical Chemistry

Copyright © 2011 Cengage Learning3-2

Example 12-1

Calculate the thermodynamic potential of the following cell and the free energy change associated with the cell reaction.

(12-2)

(12-3)


Example 12-1


Example 12-4

Calculate the cell potential for

Note that this cell does not require two compartments (nor a salt bridge) because molecular H2 has little tendency to react directly with the low concentration of Ag in the electrolyte solution.


Example 12-4


Example 12-4

The negative sign indicates that the cell reaction as considered,

is nonspontaneous.


12B Calculating Redox Equilibrium Constants

At chemical equilibrium, we may write

(12-5)

(12-4)

(12-6)



Rearrangement of Equation 12-7 gives

(12-7)

(12-8)



At 25°C

(12-9)


Example 12-5

Calculate the equilibrium constant for the reaction shown in Equation 12-4. Substituting numerical values into Equation 12-8 yields


12C-1 Electrode Potentials during Redox Titrations

Let us now consider the redox titration of iron(II) with a standard solution of cerium(IV).

This reaction is rapid and reversible so that the system is at equilibrium at all times throughout the titration.



If a redox indicator has been added to this solution, the ratio of the concentrations of its oxidized and reduced forms must adjust so that the electrode potential for the indicator, EIn, is also equal to the system potential.

Because , data for a titration curve can be obtained by applying the Nernst equation for either the cerium(IV) half-reaction or the iron(III) half-reaction.



Equivalence-Point PotentialsAt the equivalence point in the titration of iron(II) with

cerium(IV), the potential of the system Eeq is controlled by both half reactions:



The definition of equivalence point requires that

(12-10)



(12-11)


Example 12-8

Obtain an expression for the equivalence-point potential in the titration of 0.0500 M U⁴⁺ with 0.1000 M Ce⁴⁺. Assume that both solutions are 1.0 M in H₂SO₄.


Example 12-8

To combine the log terms, we must multiply the first equation by 2 to give


Example 12-8

At equivalence


Example 12-8

The equivalence-point potential for this titration is pH-dependent.

12-19


12C-2 The Titration Curve

Let us first consider the titration of 50.00 mL of 0.0500 M Fe²⁺ with 0.1000 M Ce⁴⁺ in a medium that is 1.0 M in H₂SO₄ at all times.



Initial Potential we lack sufficient information to calculate an initial

potential.Potential after the Addition of 5.00 mL of Cerium(IV)



Redox reactions used in titrimetry are sufficiently complete

Ce⁴⁺ is minuscule with respect to the other species present in the solution.



Equivalence-Point Potential Substitution of the two formal potentials into Equation

12-11 yields



Potential after the Addition of 25.10 mL of Cerium(IV)

− the iron(II) concentration is negligible


Figure 12-3

Figure 12-3 Titration curves for 0.1000 M Ce4 titration. A: Titration of 50.00 mL of 0.05000 M Fe2. B: Titration of 50.00 mL of 0.02500 M U4.


12C-3 Effect of Variables on Redox Titration Curves

Reactant Concentration titration curves for oxidation/reduction reactions are usually

independent of analyte and reagent concentrations.Completeness of the Reaction

The change in the equivalence-point region of an oxidation/reduction titration becomes larger as the reaction becomes more complete.


Figure 12-6

Figure 12-6 Effect of titrant electrode potential on reaction completeness. The standard electrode potential for the analyte is 0.200 V; starting with curve A, standard electrode potentials for the titrant are 1.20, 1.00, 0.80, 0.60, and 0.40, respectively. Both analyte and titrant undergo a oneelectron change.


12D-1 General Redox Indicators

General oxidation/reduction indicators are substances that change color upon being oxidized or reduced.

(12-12)



A color change is seen when

changes to



The potential change required to produce the full color change of a typical general indicator

a typical general indicator exhibits a detectable color change when a titrant causes the system potential to shift from to or about (0.118/n) V.



Starch/Iodine Solutions A starch solution containing a little triiodide or iodide ion

can also function as a true redox indicator.


12D-2 Specific Indicators

The best-known specific indicator is starch, which forms a dark blue complex with triiodide ion as discusssed above. This complex signals the end point in titrations in which iodine is either produced or consumed.


12E Potentiometric End Points

End points for many oxidiation/reduction titrations are readily observed by making the solution of the analyte part of the cell:


12F Auxiliary Oxidizing And Reducing Reagents

The analyte in an oxidation /reduction titration must be in a single oxidation state at the outset.

When an iron-containing sample is dissolved usually contains a mixture of iron(II) and iron(III) ions.


12F Auxiliary Oxidizing And Reducing Reagents

We must first treat the sample solution with an auxiliary reducing agent to convert all the iron to iron(II).

To be useful as a preoxidant or a prereductant, a reagent must react quantitatively with the analyte. In addition, any reagent excess must be readily removable because the excess reagent usually interferes by reacting with the standard solution.


12F-1 Auxiliary Reducing Reagents

A number of metals are good reducing agents and have been used for the prereduction of analytes. Included among these are zinc, aluminum, cadmium, lead, nickel, copper, and silver.


12F-2 Auxiliary Oxidizing Reagents

Sodium Bismuthate Sodium bismuthate is a powerful oxidizing agent; it is

capable, for example, of converting manganese(II) quantitatively to permanganate ion.

The half-reaction for the reduction of sodium bismuthate can be written as

12-37


12F-2 Auxiliary Oxidizing Reagents

Ammonium PeroxydisulfateSodium Peroxide and Hydrogen Peroxide


12G-1 Iron(II) Solutions

Numerous oxidizing agents are conveniently determined by treatment of the analyte solution with a measured excess of standard iron(II) followed by immediate titration of the excess with a standard solution of potassium dichromate or cerium(IV)


12G-2 Sodium Thiosulfate

The scheme used to determine oxidizing agents involves adding an unmeasured excess of potassium iodide to a slightly acidic solution of the analyte. Reduction of the analyte produces a stoichiometrically equivalent amount of iodine. The liberated iodine is then titrated with a standard solution of sodium thiosulfate, Na₂S₂O₃.


12G-2 Sodium Thiosulfate

(12-13)


12H-1 The Strong Oxidants: Potassium Permanganate and

Cerium(IV)

The formal potential shown for the reduction of cerium(IV) is for solutions that are 1 M in sulfuric acid. In 1 M perchloric acid and 1 M nitric acid, the potentials are 1.70 and 1.61 V, respectively. Solutions of cerium(IV) in the latter two acids are not very stable.

The half-reaction shown for permanganate ion occurs only in solutions that are 0.1 M or greater in strong acid.


12H-5 Determining Water with the Karl Fischer Reagent

Determination of water in various types of solids and organic liquids.

(12-15)

(12-14)


Detecting the End PointEnd points are obtained by electroanalytical

measurements.

12H-5 Determining Water with the Karl Fischer Reagent


THE END

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CHAPTER 12 ELECTRODE POTENTIALS AND THEIR APPLICATIONS TO XIDATION/REDUCTION TITRATIONS Introduction to Analytical Chemistry