Oxidation/reduction reaction and titration

Oxidation/reduction reaction and titration
Dr. Mohammad Khanfar

Oxidation/reduction reaction
In an oxidation/reduction reaction, electrons are transferred from one reactant to another. An example is the oxidation of iron(ll) ions by cerium(lV) ions. The reaction is described by the equation: In this reaction, an electron is transferred from Fe2+ to Ce4+ to form Ce3+ and Fe3+ ions. A substance that has a strong affinity for electrons, such as Ce4+, is called an oxidizing agent, or an oxidant. A reducing agent, or reductant, is a species, such as Fe2+, that easily donates electrons to another species.

We can split any oxidation/reduction equation into two half-reactions that show which species gains electrons and which loses them. In the oxidation—reduction methods of analysis a change in valence of the reacting products is a must which is contrary to precipitation and neutralization methods of analysis where no change in valence occur.

Reducing Agents Iron(II) Solution Sodium thiosulfate (Na2S2O3) Oxidizing Agents Strong Oxidants: Potassium Permanganate (KMnO4) and Cerium (Ce4+) Potassium Dichromate (K2Cr2O7) Weak Oxidizing Agent: Iodine or triiodide (I3-) Potassium Bromate as a source of Bromine (KBrO3) Karl Fischer Reagent for the determination of water contents: in industry, Karl Fischer titration procedure is the most widely technique for the determination of water in various types of solid and organic liquids.

Balancing Redox Equations H+, OH-, or H2O are added as needed to balance the equation For such reaction: First, we write and balance the two half-reactions involved, For MnO4-, we write: To account for the 4 oxygen atoms on the left-hand side of the equation we add 4 H2O on the right-hand side of the equation, which means that we must provide 8H+ on the left:

To balance the charge, we need to add 5 electrons to the left side of the equation. Thus, For the other half-reaction, we add one H2O to the left side of the equation to supply the needed oxygen and 2H+ on the right to balance hydrogen: Then we add two electrons to the right-hand side to balance the charge:

Before combining the two equations, we must multiply the first by 2 and the second by 5 so that the number of electrons lost will be equal to the number of electrons gained. We then add the two half-reactions to obtain: which then rearranges to the balanced equation

Redox reaction Reduction potential is a measure of how thermodynamically favorable it is for a compound to gain electrons. A high positive value for a reduction potential indicates that a compound is readily reduces and consequently is a strong oxidizing agent. i.e. it removes electrons from a substance with lower reduction potentials. A substance with a higher reduction potential will oxidize one with a lower reduction potential. The difference in potential between two substance is the reaction potential and is approximately the potential difference which would be measured if the substances comprised two halves of an electrical cell.

ECell := ECathode − Eanode
For the cell with the standard hydrogen electrode (has potential of zero by convention), one obtains: ECell = Ecathode − 0 = Eelectrode (the reduction potential, Eo )

Redox reaction Components of Potentiometric Cell
1- A reference electrode (Eref) is a half-cell having a known electrode potential that remains constant and is independent of the composition of the analyte solution. By convention, the reference electrode is always treated as the left-hand electrode in potentiometric measurements. 2- An indicator electrode (Eind) has a potential that varies in a known way with variation in the concentration of the analyte. The indicator electrode, which is immersed in the analyte solution, develops a potential (Eind) that is dependent on the activity of the analyte. 3- Salt bridge that prevents the components of the analyte solution from mixing with those of the reference electrode. Potassium chloride (KCl) is the ideal electrolyte for the salt bridge 4- Potentiometer.

Components of Potentiometric Cell A simple diagram
Potentiometer Indicator electrode made from inert material: platinum, gold or graphite Reference electrode Types: Ag/AgCl or Hg/Hg2Cl2 Salt bridge Analyte solution

Setup for the redox titration of Fe2+ with Ce4+.

Redox titration The titration curve for Fe2+ (analyte) against
Where a reference electrode has a reduction potential > 0, then the predicted reading of the potential for a redox pair is obtained by subtracting the reduction potential for the reference electrode, e.g. for an Ag/AgCl reference electrode = V is subtracted. The titration curve for Fe2+ (analyte) against Ce4+ (titrant)

Redox titration In carrying out redox titrations, standard Ag/AgCl or Hg/Hg2Cl2 electrodes are used as reference in conjunction with an inert redox electrode (indicator electrod), e.g. platinum, which takes its potential from the particular redox pair in the solution in which it is immersed.

Redox titration Effect of Variables on Redox Titration Curves
Reactant Concentration Esystem for an oxidation/reduction titration is usually independent of dilution. Consequently, titration curves for oxidation/reduction reactions are usually independent of analyte and reagent concentrations. This characteristic is in distinct contrast to that observed in the other types of titration curves

Redox titration Effect of Variables on Redox Titration Curves
Completeness of the Reaction Described by equilibrium constant (Keq) The change in potential in the equivalence-point region of an oxidation/reduction titration becomes larger as the reaction becomes more complete (i.e. larger keq). The greater the difference in reduction potential, the larger the reaction potential, the more reaction complete (larger Keq)

The reaction potential
between analyte and titrant

Redox titration Oxidation/reduction indicators
Two types of chemical indicators are used to obtain end points for oxidation/reduction titrations: General redox indicators and Specific indicators.

Redox titration General Redox Indicators
General oxidation/reduction indicators are substances that change color on being oxidized or reduced. In contrast to specific indicators, the color changes of true redox indicators are largely independent of the chemical nature of the analyte and titrant and depend instead on the changes in the electrode potential of the system that occur as the titration progresses. The half-reaction responsible for color change in a typical general oxidation reduction indicator can be written as

Redox titration Iron II Complexes of Orthophenanthrolines
A class of organic compounds known as 1,10-phenanthrolines, or orthophenanthrolines, form stable complexes with iron(II) and certain other ions. The parent compound has a pair of nitrogen atoms located in such positions that each can form a covalent bond with the iron (II) ion.

Redox titration Three orthophenanthroline molecules combine with each iron ion to yield a complex. The complexed iron in the ferroin undergoes a reversible oxidation/reduction that can be written: In contrast to many indicators, the oxidized form of ferroin is remarkably inert toward strong oxidizing agents. Oxidized form Reduced form

Redox titration Starch/Iodine Solutions
Starch, which form a blue complex with triiodide ion, is a widely used specific indicator in redox reactions involving iodine as oxidant or iodide ion as a reductant. A starch solution containing a little triiodide or iodide ion can also function as a true redox indicator. In the presence of excess oxidizing agent, the concentration ratio of iodine to iodide is high, giving a blue color to the solution. With excess reducing agent, on the other hand, iodide ion predominates, and the blue color is absent. Thus, the indicator system changes from colorless to blue in the titration of many reducing agents with various oxidizing agents. This color change is independent of the chemical composition of the reactants

Redox titration Specific Indicators
The best-known specific indicator is starch, which forms a dark blue complex with triiodide ion. This complex signals the end point in titrations in which iodine is either produced or consumed. Another specific indicator is potassium thiocyanate, which may be used, for example, in the titration of iron (III) (analyte) with solutions of titanium(IlI) sulfate (titrant). The end point involves the disappearance of the red color of the iron(III)/thiocyanate complex as a result of the marked decrease in the iron(IIl) concentration at the equivalence point.

Applications of Oxidation/Reduction Titrations
Auxiliary oxidizing and reducing reagents During the steps that precede the titration, such as dissolving the sample and separating interferences, can convert the analyte to a mixture of oxidation states. The analyte in an oxidation/reduction titration must be in a single oxidation state at the outset before starting the titration analysis Therefore, an auxiliary oxidizing or reducing reagent is used to generate a single oxidizing state of the analyte For example, when a sample containing iron is dissolved. the resulting solution usually contains a mixture of iron(II) and iron(III) ion. If we choose to use a standard oxidant to determine iron, we must first treat the sample solution with an auxiliary reducing agent to convert all of the iron to iron(Il). If we plan to titrate with a standard reductant however. pretreatment with an auxiliary oxidizing reagent is needed.

To be useful as a preoxidant or a prereductant, a reagent must react quantitatively with the analyte. In addition, any reagent excess must be easily removable because the excess reagent usually interferes with the titration by reacting with the standard solution. 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 Sticks or coils of the metal can be immersed directly in the analyte solution. After reduction is complete, the solid is removed manually and rinsed with water. The analyte solution must be filtered to remove granular or powdered forms of the metal. An alternative to filtration is the use of a reductor

Jones reductor As auxiliary reducing apparatus Granular metal
Perforated disc To vacuum

Auxiliary Oxidizing Reagents Sodium Bismuthate Oxidations are performed by suspending the bismuthate in the analyte solution and boiling for a brief period. The unused reagent is then removed by filtration. Ammonium Peroxydisulfate (NH4)2S208 is also a powerful oxidizing agent. In acidic solution, it converts chromium(lII) to dichromate, cerium(III) to cerium(lV), and manganese(II) to permanganate. Sodium Peroxide and Hydrogen Peroxide

Applying standard reducing agents in titrimetric analysis Standard solutions of most reductants tend to react with atmospheric oxygen, therefore, they are seldom used in direct titration. Instead, they used in back titration The most common standard reducing agents are iron (II) and sodium thiosulfate. Iron(lI) Solutions Air oxidation of iron(II) takes place rapidly in neutral solutions but is inhibited in the presence of acids, with the most stable preparations being about 0.5 M in H2S04 , such solutions are stable for no more than one day.

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) Sodium Thiosulfate Thiosulfate ion (S2032-) is a moderately strong reducing agent that has been widely used to determine oxidizing agents by an indirect procedure that involves iodine as an intermediate.

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, Na2S203. An example of is the determination of sodium hypochlorite in bleaches: Sodium thiosulfate is also used in the back titration of excess iodine (will be described later) (Adding unmeasured excess of KI) 2

Detecting End Points in Iodine/Thiosulfate Titrations 1) I2 has a distinct color, thus, provided that the analyte solution is colorless, the disappearance of the iodine color can serve as the indicator in titrations with sodium thiosulfate. 2) Titrations involving iodine are performed with a suspension of starch as an indicator. Aqueous starch suspensions decompose within a few days by bacterial action Starch undergoes decomposition in solutions with high I2 concentrations. In titrations of excess I2 with Na2S2O3, addition of the indicator must be deferred until most of the I2 has been reduced.

Standardizing Thiosulfate Solutions Potassium iodate (KIO3)is an excellent primary standard for thiosulfate solutions. In this application, weighed amounts of KIO3 are dissolved in water containing an excess of potassium iodide. When this mixture is acidified with a strong acid, the reaction occurs instantaneously. The liberated iodine is then titrated with the thiosulfate solution. The stoichiometry of the reaction is: I mol IO3- = 3 mol I2 = 6 mol S2032-

Example. A solution of sodium thiosulfate was standardized by dissolving g KIO3, ( g/mol) in water, adding a large excess of KI, and acidifying with HCl. The liberated iodine required mL of the thiosulfate solution to decolorize the blue starch/iodine complex. Calculate the molarity of the Na2S203.

Applying standard oxidizing agents The five most used oxidizing agents are: Potassium Permanganate (KMnO4), Cerium (Ce4+), Potassium Dichromate (K2Cr2O7), Iodine or tri-iodide (I3-), Potassium Bromate as a source of Bromine (KBrO3) The Strong Oxidants-Potassium Permanganate and Cerium(lV) The oxidizing strengths of permanganate and cerium(lV) solutions are comparable. Solutions of cerium(lV) in sulfuric acid are stable indefinitely, whereas permanganate solutions decompose slowly and thus require occasional restandardization.

Furtherrnore, cerium(IV) solutions in sulfuric acid do not oxidize chloride ion and can be used to titrate hydrochloric acid solutions of analytes; in contrast, permanganate ion cannot be used with hydrochloric acid solution unless special precautions are taken to prevent the slow oxidation of chloride ion that leads to overconsurnption of the standard reagent. A further advantage of cerium(IV) is that a primary-standard grade salt of the reagent is available, thus making possible the direct preparation of standard solutions. Despite these advantages of cerium solutions over permanganate solutions the latter are more widely used. One reason is the color of permanganate solutions, which is intense enough to serve as an indicator in titrations. A second reason for the popularity of permanganate solutions is their modest cost.

Detecting the End Points of Potassium permanganate 1) KMnO4 has intense purple color, which is sufficient to serve as an indicator for most titrations. The permanganate end point is not permanent because excess permanganate ions react slowly with the relatively large concentration of manganese(II) ion present at the end point, according to the reaction 2) If the solution is very dilute, diphenylamine sulfonic acid or the I,I0-phenanthroline complex of iron(II) provides a sharper end point. Keq = 1047

Detecting the End Points of Cerium (IV) Solutions of ceriurn(lV) are yellow-orange, but the color is not intense enough to act as an indicator in titrations. The most widely used indicator is the iron(II) complex of 1,10-phenanthroline or one of its substituted derivatives

Standarization of KMnO4. Sodium oxalate is a widely used primary standard. In acidic solutions, the oxalate ion is converted to the oxalic acid.

Analysis of phenols, Acetomenaphthone as an example First, acetamenaphthone (I) undergoes hydrolysis in acidic medium to yield the corresponding phenol and secondly, this phenol is oxidised quantitatively with ammonium ceric sulphate to give the resulting 1, 4-dione derivative (II).

Potassium Dichromate Dichromate titrations are generally carried out in solutions that are about I M in hydrochloric or sulfuric acid Potassium dichromate solutions are indefinitely stable, primary-standard reagent is available commercially and at a modest cost. The disadvantages of potassium dichromate compared with cerium (IV) and permanganate ion are its lower Eo and the slowness of its reaction with certain reducing agents. The orange color of a dichromate solution is not intense enough for use in end point detection. Diphenylamine sulfonic acid is an excellent indicator for titrations with this reagent.

Example. A 5.00-mL sample of ethanol antiseptic was diluted to L in a volumetric flask. The ethanol (C2H50H) in a mL aliquot of the diluted solution was distilled into mL of M K2Cr207 and oxidized to acetic acid with heating. The reaction is After cooling, mL of M Fe2+ was pipetted into the flask. The excess Fe2+ was then titrated with 7.46 mL of the standard K2Cr207 to a diphenylamine sulfonic acid end point. Calculate the percent (w/v) C2H5OH (46.07 g/mol) in the antiseptic.

Iodine Iodine is a weak oxidizing agent used primarily for the determination of strong reductants. However, this low potential is advantageous because it imparts a degree of selectivity that makes possible the determination of strong reducing agents in the presence of weak ones. Iodine is not very soluble in water (0.001 M). To obtain solutions having analytically useful concentrations of the element. iodine is ordinarily dissolved in moderately concentrated solutions of potassium iodide. In this medium, iodine is reasonably soluble as a consequence of the reaction

Iodine solutions lack stability for the following reasons: 1) Volatility of the solute. 2) Iodine slowly attacks most organic materials, Consequently, cork or rubber stoppers are never used 3) Air oxidation of iodide ion also causes changes in the molarity of an iodine Must be restandardized regularly using sodium thiosulfate

Iodometric analysis of ascorbic acid The determination depends upon the quantitative oxidaion of ascorbic acid to dehydroascorbic acid with iodine acid solution

Iodine-absorbing substances in penicillines A major stability problem in penicillines is the hydrolysis of the lactam ring Penicillin with an open lactam ring is inactive. When the lactam ring is open it will react with iodine. 1 mole of the ring-open form will react with 8 equivalents of iodine, the intact lactam ring will not react. Excess iodine is added to penicillin sample and then back titrated with Na2S2O3.

Potassium Bromate (KBrO3) as a Source of Bromine Like iodine, bromine (Br2) can act as an oxidizing agent and can react with various organic compounds (by addition and substitution reaction). However, bromine solutions lacks stability. Instead, an unmeasured excess of potassium bromide (KBr) is added to an acidic solution of the analyte. On introduction of a measured volume of standard potassium bromate ((KBrO3), a stoichiometric quantity of bromine is produced.

The primary use of standard potassium bromate is the determination of organic compounds that react with bromine. Few of these reactions are rapid enough to make direct titration feasible. Instead, a measured excess of standard bromate is added to the solution that contains the sample plus an excess of potassium bromide. After acidification, the mixture is allowed to stand in a glass-stoppered vessel until the bromine/analyte reaction is judged complete. To determine the excess bromine, an excess of potassium iodide is introduced so that the following reaction occurs:

Bromine is incorporated into an organic molecule either by substitution or by addition. Substitution Reactions Halogen substitution involves the replacement of hydrogen in an aromatic ring with a halogen. Substitution methods have been successfully applied to the determination of aromatic compounds that contain strong ortho-para-directing groups, particularly amines and phenols.

Example, analysis of sulfanilamide A g sample of an antibiotic powder was dissolved in HCI and the solution diluted to mL. A mL aliquot was transferred to a flask, followed by mL of M KBrO 3.An excess of KBr was added to form Br2, and the flask was stoppered. After 10 min, during which time the Br2 brominated the sulfanilamide, an excess of KI was added. The liberated iodine titrated with mL of M sodium thiosulfate. The reactions are:

Calculate the percent sulfanilamide (NH2C6H4S02NH, 172.21 g/mol)
in the powder.

Addition Reactions Addition reactions involve the opening of an olefinic double bond. For example, I mol of ethylene reacts with I mol of bromine in the reaction

Determining Water with the Karl Fischer Reagent One of the most widely used analytical methods for the determination of water in various types of solids and organic liquids. This important titrimetric method is based on an oxidation/reduction that is relatively specific for water. The Karl Fischer reaction is based on the oxidation of sulfur dioxide by iodine. The Karl Fischer reagent consists of a mixture of anhydrous methanol, an anhydrous base (pyridine or imidazol), iodine and sulfur dioxide (SO2).

The classical reaction has been shown to occur in two steps. In the first step, I2 and S02 react in the presence of pyridine and water to form pyridiniurn sulfite (C5H5N . SO3) and pyridinium iodide (C5H5N . HI) where I2, SO2, and SO3 are complexed by the pyridine. The second step is the consumption of water by pyridinium sulfite: Note that the stoichiometry is one mole of I2 per mole of H2O present (1:1)

Detecting the End Point 1) An end point in a Karl Fischer titration can be observed visually based on the brown color of the excess reagent. 2) More commonly, end points are obtained from electroanalytical measurements. 3) The end point can be determined spectrophotometrically by measuring the absorbance of the titration solution at 525nm (I2 absorbance)

Precautions for Karl fisher 1) Atmospheric oxygen and humidity 2) If water is released slowly from the crystals (hydrates), back titration is employed (using excess reagent then back titrate the remaining unreacted reagent with H2O/methanol. 3) Most classes of organic compounds do not interfere in the titration. So water can be measured in the presence of mineral acids, alcohol, phenols, ethers, amines, amides, halides, sulfides. Interfering chemicals, ascorbic acids, Thiols, carbonyl compounds, Carboxylic acids.

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