Titrimetry Methods
Redox Titrations
Redox reactions are the most diverse of the four main classes of inorganic aqueous reactions
(acid-base, pptn, complexation and redox). In principle, then, redox titrations can be used to analyze for any oxidizing or reducing agent. However, many redox reactions are either too slow or have inconsistent stiochiometry. The stability of titrant and analyte solutions can also be a problem. Nevertheless, a wide variety of analytes can be conveniently determined by redox titrations.
General Considerations
Consider a generic redox half-reaction (charges omitted for clarity):
ox + ne– l red
A chemical (i.e., ox in this equation) that pulls electrons from another substance is an oxidizing
agent, while a chemical (red) that forces another substance to accept electrons is a reducing agent. Together, ox/red form a redox couple; redox couples are analogous to acid/base conjugate pairs. And just like acid-base reactions, the “conjugate” of a strong oxidizing agent is a weak reducing agent. The strength of oxidizing/reducing agents can be deduced by the standard reduction potential: a very positive standard potential indicates a strong oxidizing agent, while a low positive or a negative potential is characteristic of a strong reducing agent. The strength of an oxidizing or reducing agent is very often dependent on pH. There is a general rule of thumb: acidic conditions tend to make oxidizing agents more powerful and render reducing agents less reactive. Some few redox reagents are relatively insensitive to pH, which can be an advantage.
Most redox reagents are stable (if they are stable at all!) only within a certain pH range.
Sample treatment is often necessary to adjust the oxidation state of the analyte. The analyte is either pre-reduced or pre-oxidized. For pre-reduction of the analyte, many metals (many of which are strong reducing agents) can be used. It is common to use a reductor, which is a column of granulated metal through which the sample solution is poured. Two common reductors are: the Jones Reductor, which uses amalgamated zinc (ZnHg) granules, and the Walden Reductor, which uses silver granules (chloride is added to the sample, usually as HCl). The Walden Reductor is more selective (i.e., a less powerful reducing agent) than the Jones Reductor.
Pre-oxidation is not as common as pre-reduction, since the analyte is usually desired in a reduced form for titration with an oxidizing agent. However, when pre-oxidation is necessary, sodium bismuthate, NaBiO3, ammonium peroxydisulfate, (NH4)2S2O8, or hydrogen peroxide may be used.
Common Titrants
Reducing Agents
• reducing agents are not stable in air (undergo air oxidation) and so are not often used. Here are a few titrants
• the two most common reducing titrants are ferrous ammonium sulfate (FAS) and sodium thiosulfate. Procedures using these titrants are capable of determining the concentrations analytes that are (at least) moderately strong oxidizing agents. Ferrous Ammonium Sulfate (FAS or Mohr’s salt), (NH4)2Fe(SO4)2
• the ferrous ion is a fairly weak reducing agent:
Fe2+ t Fe3+ + e– E°= 0.771V
The use of ferrous ion as a titrant is limited to the analysis of moderately strong oxidizing agents; it is used for the direct titration of a few metals such as U(VI), Mo(VI) and V(IV). Probably the most important use of FAS is in back-titrations of dichromate and other reasonably strong oxidants.
• solutions of FAS are most stable under acidic conditions (in 0.5M H2SO4); still, the solution is
stable only for about a day. Standardization is with potassium dichromate, K2Cr2O7.
Sodium Thiosulfate, Na2S2O3
• thiosulfate is a moderately strong reducing agent:
2S2O32– t S4O62– + 2e– E°= 0.09 V
• thiosulfate is actually not suitable for the direct analysis of most oxidizing agents, since reactions with thiosulfate tend to produce also produce sulfite and sulfate. However, it is widely used in back-titrations of iodine that is produced by the reactions of oxidizing agents with iodide, another reducing agent (this procedure is called iodometry).
• thiosulfate solutions are standardized with iodine which has been prepared by acidifying primary standard potassium iodate in the presence of a slight excess of potassium iodide:
acidic solution IO3– + 5I– + 6H+ t 3I2(aq) + 3H2O
The titration reaction between iodine and thiosulfate is fairly straightforward:
I2 + 2S2O32– t 2I– + S4O62–
• alkaline solutions of sodium thiosulfate are fairly stable Oxidizing Agents
• used for the analysis of reducing agents. Pre-reduction of analyte is common; analyte is often
unstable in reduced form, and care must be taken in sample handling Potassium Permanganate, KMnO4
• used since the mid-1800’s - one of the earliest titrimetric agents
• a strong oxidant
MnO4– + 4H+ + 3e– t MnO2(s) + 2H2O E°= 1.692 V
• standardized with sodium oxalate, Na2C2O4. Not a very stable titrant unless precautions are taken; should be standardized fairly often.
• can be used for the analysis of many reducing agents, weak or strong. Examples are given in table 16-3 in Harris: e.g., Br–, H2O2, NO2–, Fe2+, As3+, Sb3+, Mo3+, W3+, U4+, Ti3+
Ceric Sulfate, Ce(SO4)2
• another strong oxidant, just about as strong as permanganate
Ce4+ + e– t Ce3+ E°= 1.44 V (in H2SO4)
• standardized with Na2C2O4. Alternately, primary standard (NH4)2Ce(NO3)6 can be used (expensive!). Titrant is very stable in acid solutions; it ppts in alkaline solutions.
• almost anything that can be done with potassium permanganate can be done more conveniently with ceric sulfate.
Potassium Dichromate, K2Cr2O7
• historically important; like permanganate, used since mid-1800’s
• a moderately strong oxidizing agent; oxidizing ability depends strongly on pH, decreasing rapidly as solution becomes more neutral
Cr2O72– + 14H+ + 6e– t 2Cr3+ + 7H2O E°= 1.36 V
• available in sufficient purity to be its own primary standard; in fact, it is the most common reagent used to standardize reducing titrants. If necessary, dichromate solutions can be standardized with Na2C2O4.
• most common applications: analysis of iron content of ores and COD of wastewaters. Advantage for iron ore analysis: no problem with HCl solutions, unlike permanganate (which oxidizes chloride to chlorine). Back-titrations involving FAS are also common: FAS may be added in excess for the analysis of oxidizing agents (back-titration with dichromate) or FAS may be used for to analyze excess dichromate (as in COD measurements).
Titrations involving Iodine, I2
General Applicability
• an important class of techniques: can be used to analyze moderately strong oxidants or reductants. Advantage of moderate strength as a redox reagent: better selectivitity. Permanganate and ceric oxidize almost everything present.
• the standard reduction of iodine is
I2(aq) + 2e– t 2I– E°= 0.621 V
• iodine is a moderate oxidizing agent; iodide is a moderate reducing agent. There are two classes of titrations involving iodine:
1. Iodimetry, which is based on the direct reaction between the analyte and iodine. Since iodine is an oxidizing agent, iodimetry is used for the analysis of reductants.
2. Iodometry, which is based on the reaction between the analyte and an unmeasured excess of
iodide to produce iodine, which is measured by titration with thiosulfate. The amount of iodine produced by this reaction is stoichiometrically related to the amount of analyte originally present in the solution. Since iodide is an reductant, iodometry is used for the analysis of oxidants.
• applications of iodimetry and iodometry are extensive; see Harris table 16-4 for more details.
Remember: iodimetry (analyte reacts with iodine) is for the analysis of reducing agents, while
iodometry (production of iodine by reaction of analyte with iodide, followed by back-titration with thiosulfate) is for the analysis of oxidizing agents. Titrations involving iodine are more selective than those involving more powerful redox reagents.
Iodine Aquatic Chemistry
Iodine crystals are only sparingly soluble in water, so the standard potential listed earlier for iodine gives a misleading impression of the strength of iodine as an oxidizing agent. Usually iodine is prepared by dissolution in a solution of concentrated potassium iodide, due to the formation of the triiodide ion:
I2(aq) + I– l I3– K = 710
This reaction allows iodine to dissolve. However, the actual concentration of I2(aq) remains low; thus, the oxidizing power still does not approach that of 1M I2(aq). Due to the presence of triiodide, the following reaction is often used to represent iodine oxidation during a titration:
I3–(aq) + 2e– t 3I–(aq) E°= 0.545 V
The oxidizing ability of iodine solutions is not very dependent on pH; however, in alkaline solutions (pH > 8), iodine disproportionates to iodate and iodide:
alkalin solution 3I2 + 3H2O l IO3– + 5I– + 6H+
Note that this reaction is quite reversible: upon acidification, the reaction shifts to the left as iodate reacts with iodide to form iodine.
Solution Preparation of Titranons
• iodine titrant solutions are usually prepared by dissolving solid iodine in potassium iodide solutions. The solution may be standardized with primary standard sodium oxalate or with sodium thiosulfate that has been previously standardized.
• sodium thiosulfate is the reducing agent that is universally used for the back-titration of iodine
produced in iodometry. This titration reaction is stoichiometric and fairly rapid:
I2 + 2S2O32– t 2I– + S4O62–
• sodium thiosulfate titrant is prepared simply by dissolving the salt in water and storing under
slightly basic conditions. It is standardized with potassium iodate that has been acidified in excess iodide.
Iodine/triiodide solutions are unstable for a variety of reasons. First of all, aqueous iodine exerts a significant vapor pressure. Also, under acidic conditions iodide is slowly air-oxidized to produce
iodine. Finally, under alkaline conditions, iodine will disproportionate to produce iodide and iodate, as mentioned previosusly.
Thus, iodine solutions are generally most stable at neutral pH values. Iodine
titrant solutions must be standardized fairly frequently.
Endpoint Detection for Redox Titrations
• it is probably worthwhile to mention that starch is an excellent chemical indicator for titrations
involving iodine
• potentiometric detection with an inert indicator electrode (e.g., Pt) is a general method for following redox titrations
• amperometric detection can also be used in many cases
• many redox titrants are colored (e.g., permanganate or iodine) and so photometric detection can also be used to follow the course of the titration
Applications of Redox Titrations
Example Applications
• DO by Winkler (iodometric) titration
• COD by dichromate back-titration (using FAS)
• analysis of iron in ores by dichromate titration
• analysis of residual chlorine by iodometric titration
• analysis of ascorbic acid (vitamin C), hydrogen peroxide, bleach, ...
Summary of Applications of Redox Reactions
• analytes that are oxidizing agents are most conveniently analyzed by addition of excess reducing agent and then back-titrating. There are two common ways of doing this: (i) addition of a measured excess of ferrous ammonium sulfate and back-titrating the unreacted excess with dichromate titrant;
(ii) addition of an unmeasured excess of potassium iodide and using thiosulfate to back-titrate the iodine produced by reaction of iodide with analyte.
• analytes that are reducing agents may be analyzed by a variety of oxidizing agents: potassium
permanganate and ceric sulfate are strong oxidants, potassium dichromate is a moderately strong oxidizing titrant (especially suitable for the analysis of ferrous iron, or back-titrations with FAS) and iodine is a milder, more selective oxidizing agent that may be used for the direct analysis of a number of reducing agents (iodimetry), as well as the indirect analysis of reducing agents (back-titration with thiosulfate in iodometry; see above)
• pre-treatment of the analyte with an oxidizing agent or a reducing agent is often needed in redox titrations
Literature
Arthur I. Vogel,"A Text-Book of Quantitative Inorganic Analysis", 2 nd edition.
HA Laitinen, WE Harris, "Chemical Analysis", 2 nd edition.
Redox Titrations
Redox reactions are the most diverse of the four main classes of inorganic aqueous reactions
(acid-base, pptn, complexation and redox). In principle, then, redox titrations can be used to analyze for any oxidizing or reducing agent. However, many redox reactions are either too slow or have inconsistent stiochiometry. The stability of titrant and analyte solutions can also be a problem. Nevertheless, a wide variety of analytes can be conveniently determined by redox titrations.
General Considerations
Consider a generic redox half-reaction (charges omitted for clarity):
ox + ne– l red
A chemical (i.e., ox in this equation) that pulls electrons from another substance is an oxidizing
agent, while a chemical (red) that forces another substance to accept electrons is a reducing agent. Together, ox/red form a redox couple; redox couples are analogous to acid/base conjugate pairs. And just like acid-base reactions, the “conjugate” of a strong oxidizing agent is a weak reducing agent. The strength of oxidizing/reducing agents can be deduced by the standard reduction potential: a very positive standard potential indicates a strong oxidizing agent, while a low positive or a negative potential is characteristic of a strong reducing agent. The strength of an oxidizing or reducing agent is very often dependent on pH. There is a general rule of thumb: acidic conditions tend to make oxidizing agents more powerful and render reducing agents less reactive. Some few redox reagents are relatively insensitive to pH, which can be an advantage.
Most redox reagents are stable (if they are stable at all!) only within a certain pH range.
Sample treatment is often necessary to adjust the oxidation state of the analyte. The analyte is either pre-reduced or pre-oxidized. For pre-reduction of the analyte, many metals (many of which are strong reducing agents) can be used. It is common to use a reductor, which is a column of granulated metal through which the sample solution is poured. Two common reductors are: the Jones Reductor, which uses amalgamated zinc (ZnHg) granules, and the Walden Reductor, which uses silver granules (chloride is added to the sample, usually as HCl). The Walden Reductor is more selective (i.e., a less powerful reducing agent) than the Jones Reductor.
Pre-oxidation is not as common as pre-reduction, since the analyte is usually desired in a reduced form for titration with an oxidizing agent. However, when pre-oxidation is necessary, sodium bismuthate, NaBiO3, ammonium peroxydisulfate, (NH4)2S2O8, or hydrogen peroxide may be used.
Common Titrants
Reducing Agents
• reducing agents are not stable in air (undergo air oxidation) and so are not often used. Here are a few titrants
• the two most common reducing titrants are ferrous ammonium sulfate (FAS) and sodium thiosulfate. Procedures using these titrants are capable of determining the concentrations analytes that are (at least) moderately strong oxidizing agents. Ferrous Ammonium Sulfate (FAS or Mohr’s salt), (NH4)2Fe(SO4)2
• the ferrous ion is a fairly weak reducing agent:
Fe2+ t Fe3+ + e– E°= 0.771V
The use of ferrous ion as a titrant is limited to the analysis of moderately strong oxidizing agents; it is used for the direct titration of a few metals such as U(VI), Mo(VI) and V(IV). Probably the most important use of FAS is in back-titrations of dichromate and other reasonably strong oxidants.
• solutions of FAS are most stable under acidic conditions (in 0.5M H2SO4); still, the solution is
stable only for about a day. Standardization is with potassium dichromate, K2Cr2O7.
Sodium Thiosulfate, Na2S2O3
• thiosulfate is a moderately strong reducing agent:
2S2O32– t S4O62– + 2e– E°= 0.09 V
• thiosulfate is actually not suitable for the direct analysis of most oxidizing agents, since reactions with thiosulfate tend to produce also produce sulfite and sulfate. However, it is widely used in back-titrations of iodine that is produced by the reactions of oxidizing agents with iodide, another reducing agent (this procedure is called iodometry).
• thiosulfate solutions are standardized with iodine which has been prepared by acidifying primary standard potassium iodate in the presence of a slight excess of potassium iodide:
acidic solution IO3– + 5I– + 6H+ t 3I2(aq) + 3H2O
The titration reaction between iodine and thiosulfate is fairly straightforward:
I2 + 2S2O32– t 2I– + S4O62–
• alkaline solutions of sodium thiosulfate are fairly stable Oxidizing Agents
• used for the analysis of reducing agents. Pre-reduction of analyte is common; analyte is often
unstable in reduced form, and care must be taken in sample handling Potassium Permanganate, KMnO4
• used since the mid-1800’s - one of the earliest titrimetric agents
• a strong oxidant
MnO4– + 4H+ + 3e– t MnO2(s) + 2H2O E°= 1.692 V
• standardized with sodium oxalate, Na2C2O4. Not a very stable titrant unless precautions are taken; should be standardized fairly often.
• can be used for the analysis of many reducing agents, weak or strong. Examples are given in table 16-3 in Harris: e.g., Br–, H2O2, NO2–, Fe2+, As3+, Sb3+, Mo3+, W3+, U4+, Ti3+
Ceric Sulfate, Ce(SO4)2
• another strong oxidant, just about as strong as permanganate
Ce4+ + e– t Ce3+ E°= 1.44 V (in H2SO4)
• standardized with Na2C2O4. Alternately, primary standard (NH4)2Ce(NO3)6 can be used (expensive!). Titrant is very stable in acid solutions; it ppts in alkaline solutions.
• almost anything that can be done with potassium permanganate can be done more conveniently with ceric sulfate.
Potassium Dichromate, K2Cr2O7
• historically important; like permanganate, used since mid-1800’s
• a moderately strong oxidizing agent; oxidizing ability depends strongly on pH, decreasing rapidly as solution becomes more neutral
Cr2O72– + 14H+ + 6e– t 2Cr3+ + 7H2O E°= 1.36 V
• available in sufficient purity to be its own primary standard; in fact, it is the most common reagent used to standardize reducing titrants. If necessary, dichromate solutions can be standardized with Na2C2O4.
• most common applications: analysis of iron content of ores and COD of wastewaters. Advantage for iron ore analysis: no problem with HCl solutions, unlike permanganate (which oxidizes chloride to chlorine). Back-titrations involving FAS are also common: FAS may be added in excess for the analysis of oxidizing agents (back-titration with dichromate) or FAS may be used for to analyze excess dichromate (as in COD measurements).
Titrations involving Iodine, I2
General Applicability
• an important class of techniques: can be used to analyze moderately strong oxidants or reductants. Advantage of moderate strength as a redox reagent: better selectivitity. Permanganate and ceric oxidize almost everything present.
• the standard reduction of iodine is
I2(aq) + 2e– t 2I– E°= 0.621 V
• iodine is a moderate oxidizing agent; iodide is a moderate reducing agent. There are two classes of titrations involving iodine:
1. Iodimetry, which is based on the direct reaction between the analyte and iodine. Since iodine is an oxidizing agent, iodimetry is used for the analysis of reductants.
2. Iodometry, which is based on the reaction between the analyte and an unmeasured excess of
iodide to produce iodine, which is measured by titration with thiosulfate. The amount of iodine produced by this reaction is stoichiometrically related to the amount of analyte originally present in the solution. Since iodide is an reductant, iodometry is used for the analysis of oxidants.
• applications of iodimetry and iodometry are extensive; see Harris table 16-4 for more details.
Remember: iodimetry (analyte reacts with iodine) is for the analysis of reducing agents, while
iodometry (production of iodine by reaction of analyte with iodide, followed by back-titration with thiosulfate) is for the analysis of oxidizing agents. Titrations involving iodine are more selective than those involving more powerful redox reagents.
Iodine Aquatic Chemistry
Iodine crystals are only sparingly soluble in water, so the standard potential listed earlier for iodine gives a misleading impression of the strength of iodine as an oxidizing agent. Usually iodine is prepared by dissolution in a solution of concentrated potassium iodide, due to the formation of the triiodide ion:
I2(aq) + I– l I3– K = 710
This reaction allows iodine to dissolve. However, the actual concentration of I2(aq) remains low; thus, the oxidizing power still does not approach that of 1M I2(aq). Due to the presence of triiodide, the following reaction is often used to represent iodine oxidation during a titration:
I3–(aq) + 2e– t 3I–(aq) E°= 0.545 V
The oxidizing ability of iodine solutions is not very dependent on pH; however, in alkaline solutions (pH > 8), iodine disproportionates to iodate and iodide:
alkalin solution 3I2 + 3H2O l IO3– + 5I– + 6H+
Note that this reaction is quite reversible: upon acidification, the reaction shifts to the left as iodate reacts with iodide to form iodine.
Solution Preparation of Titranons
• iodine titrant solutions are usually prepared by dissolving solid iodine in potassium iodide solutions. The solution may be standardized with primary standard sodium oxalate or with sodium thiosulfate that has been previously standardized.
• sodium thiosulfate is the reducing agent that is universally used for the back-titration of iodine
produced in iodometry. This titration reaction is stoichiometric and fairly rapid:
I2 + 2S2O32– t 2I– + S4O62–
• sodium thiosulfate titrant is prepared simply by dissolving the salt in water and storing under
slightly basic conditions. It is standardized with potassium iodate that has been acidified in excess iodide.
Iodine/triiodide solutions are unstable for a variety of reasons. First of all, aqueous iodine exerts a significant vapor pressure. Also, under acidic conditions iodide is slowly air-oxidized to produce
iodine. Finally, under alkaline conditions, iodine will disproportionate to produce iodide and iodate, as mentioned previosusly.
Thus, iodine solutions are generally most stable at neutral pH values. Iodine
titrant solutions must be standardized fairly frequently.
Endpoint Detection for Redox Titrations
• it is probably worthwhile to mention that starch is an excellent chemical indicator for titrations
involving iodine
• potentiometric detection with an inert indicator electrode (e.g., Pt) is a general method for following redox titrations
• amperometric detection can also be used in many cases
• many redox titrants are colored (e.g., permanganate or iodine) and so photometric detection can also be used to follow the course of the titration
Applications of Redox Titrations
Example Applications
• DO by Winkler (iodometric) titration
• COD by dichromate back-titration (using FAS)
• analysis of iron in ores by dichromate titration
• analysis of residual chlorine by iodometric titration
• analysis of ascorbic acid (vitamin C), hydrogen peroxide, bleach, ...
Summary of Applications of Redox Reactions
• analytes that are oxidizing agents are most conveniently analyzed by addition of excess reducing agent and then back-titrating. There are two common ways of doing this: (i) addition of a measured excess of ferrous ammonium sulfate and back-titrating the unreacted excess with dichromate titrant;
(ii) addition of an unmeasured excess of potassium iodide and using thiosulfate to back-titrate the iodine produced by reaction of iodide with analyte.
• analytes that are reducing agents may be analyzed by a variety of oxidizing agents: potassium
permanganate and ceric sulfate are strong oxidants, potassium dichromate is a moderately strong oxidizing titrant (especially suitable for the analysis of ferrous iron, or back-titrations with FAS) and iodine is a milder, more selective oxidizing agent that may be used for the direct analysis of a number of reducing agents (iodimetry), as well as the indirect analysis of reducing agents (back-titration with thiosulfate in iodometry; see above)
• pre-treatment of the analyte with an oxidizing agent or a reducing agent is often needed in redox titrations
Literature
Arthur I. Vogel,"A Text-Book of Quantitative Inorganic Analysis", 2 nd edition.
HA Laitinen, WE Harris, "Chemical Analysis", 2 nd edition.
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