Redox and Potentiometric Titrations: Unveiling Electrochemical Insights

Redox titrations are a type of volumetric analysis based on an oxidation-reduction (redox) reaction between the analyte and a standard titrant. These titrations involve the transfer of electrons and are incredibly versatile for determining the concentration of various oxidizing or reducing species. Potentiometry, an electroanalytical technique, provides a powerful and precise way to monitor these redox reactions by measuring the potential of an electrochemical cell under static (zero-current) conditions. This measured potential is then related to the concentration of species involved in the redox reaction, making it an excellent method for endpoint detection.

1. Fundamentals of Redox Reactions

1.1. Oxidation and Reduction

• Oxidation: Loss of electrons by a chemical species, resulting in an increase in its oxidation state.
• Reduction: Gain of electrons by a chemical species, resulting in a decrease in its oxidation state.
• Oxidizing Agent (Oxidant): A species that causes oxidation by accepting electrons (itself gets reduced).
• Reducing Agent (Reductant): A species that causes reduction by donating electrons (itself gets oxidized).

1.2. Electrochemical Cells and Potential

A redox reaction can be split into two half-reactions: an oxidation half-reaction and a reduction half-reaction. When these half-reactions are separated into two half-cells, an electrochemical cell is formed, and a potential difference (voltage or electromotive force, EMF) develops between the electrodes.

• Cell Potential (Ecell​): The overall potential of the electrochemical cell, which drives the electron flow. For a spontaneous reaction, Ecell​>0. Ecell​=Ereduction half-cell​−Eoxidation half-cell​ (where both E values are written as reduction potentials).

1.3. The Nernst Equation

The potential of a half-cell (electrode potential) under non-standard conditions is related to the concentrations (more accurately, activities) of the electroactive species by the Nernst Equation:

For a general half-reaction: aA+bB+ne−⇌cC+dD (written as reduction)

E=E0−nFRT​ln[A]a[B]b[C]c[D]d​

Where:

• E: Electrode potential under non-standard conditions.
• E0: Standard electrode potential (at 25∘C, 1 M concentration for solutes, 1 atm pressure for gases).
• R: Ideal gas constant (8.314 J K−1 mol−1).
• T: Absolute temperature (in Kelvin).
• n: Number of electrons involved in the half-reaction.
• F: Faraday constant (96485 C mol−1).
• ln: Natural logarithm.
• […]: Activities of the species (often approximated as concentrations, especially in dilute solutions).

At 25∘C (298.15 K), the Nernst equation can be simplified using common logarithm (log10​):

E=E0−n0.05916​log[reactants][products]​ (for reduction reactions)

This equation is fundamental to understanding how changes in analyte or titrant concentrations lead to measurable changes in potential during a redox titration.

2. Redox Titrations

Redox titrations involve the quantitative determination of an analyte by reacting it with a standard solution of an oxidizing or reducing agent. The equivalence point is reached when the analyte has completely reacted with the titrant according to the stoichiometry of the redox reaction.

2.1. Common Oxidizing Agents (Titrants)

• Potassium Permanganate (KMnO4​): A very strong oxidizing agent. It is self-indicating (deep purple MnO4−​ to colorless Mn2+ in acidic solution). However, it is not a primary standard and needs standardization. Reactions are often carried out in acidic medium.
• Potassium Dichromate (K2​Cr2​O7​): A strong oxidizing agent, but less powerful than KMnO4​. It is a primary standard (highly pure). Color change from orange (Cr2​O72−​) to green (Cr3+). Requires an external or internal indicator.
• Cerium(IV) Sulfate (Ce(SO4​)2​): A powerful oxidizing agent, often used as a primary standard. Color change from yellow (Ce4+) to colorless (Ce3+). Requires an indicator.
• Iodine (I2​): A moderately strong oxidizing agent. Used in iodimetry.

2.2. Common Reducing Agents (Titrants)

• Sodium Thiosulfate (Na2​S2​O3​): Used to titrate iodine (in iodometry). It is not a primary standard.
• Iron(II) Salts (e.g., FeSO4​): Can be used as a reducing titrant, but readily oxidized by air.
• Ascorbic Acid (Vitamin C): A mild reducing agent, can be used for some titrations.

2.3. Redox Indicators

Redox indicators are substances that change color at a specific electrode potential, indicating the equivalence point of a redox titration. Their color change is reversible and depends on their own oxidation and reduction states.

• Mechanism: The indicator itself is a redox couple (Inox​/Inred​). Its color changes as its oxidation state changes. The indicator potential (EIn0​) must lie within the steep potential jump around the equivalence point of the titration.
• Examples:
  • Starch Solution: Forms a deep blue complex with I2​. Used for titrations involving iodine (iodometry/iodimetry). It is a specific indicator, not a general redox indicator.
  • Ferroin (1,10-phenanthroline iron(II) complex): A widely used general redox indicator. It changes from light blue (Fe(phen)33+​) to red (Fe(phen)32+​). Its EIn0​ is ∼+1.14 V.
  • Diphenylamine Sulfonate: Changes from colorless to violet.
  • Potassium Permanganate (KMnO4​): Acts as a self-indicator. The endpoint is signaled by the first persistent pale pink color of excess MnO4−​.

2.4. Redox Titration Curves

A redox titration curve plots the potential of the electrochemical cell (y-axis) against the volume of titrant added (x-axis).

• Shape: Typically shows a steep sigmoid (S-shaped) curve, with a large and abrupt change in potential occurring around the equivalence point.
• Potential before Equivalence Point: The potential is dominated by the redox couple of the analyte.
• Potential at Equivalence Point: At the equivalence point, neither the analyte nor the titrant is in large excess. The potential can be calculated based on the standard potentials of both half-reactions. The inflection point of the curve is the equivalence point.
• Potential after Equivalence Point: The potential is dominated by the redox couple of the titrant.
• Factors Affecting Curve:
  • Strength of Oxidizing/Reducing Agents: Stronger agents lead to larger potential jumps.
  • Concentrations: Higher concentrations lead to larger potential jumps.
  • Stoichiometry: Reactions involving a large number of electrons (large ‘n’) will also lead to sharper changes.
  • Presence of Complexing Agents or Precipitating Agents: These can alter the activities of the ions and thus the potentials.

3. Potentiometric Measurement for Titrations

Potentiometry is a powerful electroanalytical technique used to measure the potential of an electrochemical cell under static (zero-current) conditions. This measured potential is then related to the concentration of a specific analyte in a solution. In redox titrations, it provides a very precise and objective way to determine the equivalence point, especially when a visual indicator is unavailable, provides a poor color change, or for automated systems.

3.1. Electrochemical Cell for Potentiometric Titration

A potentiometric cell for titration consists of two half-cells connected by a salt bridge:

• Indicator Electrode: The electrode whose potential depends on the concentration (or more accurately, activity) of the redox couple being measured (either the analyte or the titrant). This is typically an inert metallic electrode (e.g., platinum or gold wire) that simply provides a surface for electron transfer but does not participate in the reaction itself.
• Reference Electrode: An electrode whose potential is constant and independent of the analyte or titrant concentration. It provides a stable reference point. Common choices are the Calomel Electrode (SCE) or the Silver/Silver Chloride (Ag/AgCl) Electrode.

The cell potential (Ecell​) measured between these two electrodes is given by: Ecell​=Eindicator​−Ereference​+Eliquid junction​ The liquid junction potential (Elj​) arises at the interface between the salt bridge and the sample solution due to differential ion mobilities. While ideally minimized by proper salt bridge design, it is a source of uncertainty in potentiometric measurements.

3.2. Endpoint Determination in Potentiometric Titrations

The equivalence point is found by identifying the point of maximum slope on the potentiometric titration curve. This can be done accurately by:

• First Derivative Plot: Plotting the change in potential with respect to the change in titrant volume (ΔE/ΔV) against the average titrant volume (Vavg​). The maximum of this plot corresponds precisely to the equivalence point.
• Second Derivative Plot: Plotting the second derivative (Δ2E/ΔV2) against Vavg​. The point where this plot crosses zero (goes from positive to negative) corresponds to the equivalence point.
• Advantages of Potentiometric Endpoint Detection for Redox Titrations:
  • Objective and Precise: Eliminates subjective judgment associated with visual color changes.
  • Applicable to Colored or Turbid Solutions: The endpoint can be determined even when the solution itself is colored or opaque, obscuring visual indicators.
  • Automated Systems: Easily integrated into automated titration systems.
  • Weak/Dilute Solutions: Provides clearer endpoints for reactions that give poor visual indications.

4. Reference Electrodes (Detailed Review)

Reference electrodes are crucial for providing a stable, known potential against which the indicator electrode’s potential can be measured. They are designed to maintain a constant potential despite changes in the sample solution.

4.1. Standard Hydrogen Electrode (SHE)

• Description: The primary reference electrode, defined as having a standard potential of exactly 0.000 V at all temperatures. It consists of a platinum electrode immersed in a 1 M H+ solution, over which hydrogen gas at 1 atm pressure is continuously bubbled.
• Reaction: 2H+(aq,1 M)+2e−⇌H2​(g,1 atm)
• Disadvantages: Impractical for routine use due to the need for a constant supply of pure hydrogen gas, precise pressure control, and highly pure platinum. Used as the theoretical standard.

4.2. Calomel Electrode (SCE)

• Description: A widely used secondary reference electrode. It consists of mercury in contact with mercurous chloride (Hg2​Cl2​, known as calomel) and a saturated solution of potassium chloride (KCl).
• Reaction: Hg2​Cl2​(s)+2e−⇌2Hg(l)+2Cl−(aq)
• Potential: Its potential depends on the concentration of Cl−. A saturated KCl solution (saturated calomel electrode, S.C.E.) provides a very stable potential (+0.241 V vs. SHE at 25∘C).
• Advantages: Stable, reliable, relatively easy to use.
• Disadvantages: Contains mercury (toxic), sensitive to temperature changes (potential changes), and Cl− leakage can contaminate samples.

4.3. Silver/Silver Chloride Electrode (Ag/AgCl)

• Description: Another common and robust secondary reference electrode. It consists of a silver wire coated with silver chloride (AgCl) immersed in a KCl solution of known concentration (often saturated).
• Reaction: AgCl(s)+e−⇌Ag(s)+Cl−(aq)
• Potential: The potential depends on the Cl− concentration. For a saturated KCl solution, its potential is +0.197 V vs. SHE at 25∘C.
• Advantages: Non-toxic, very stable, more robust than SCE, less temperature-dependent than SCE, widely used for field and continuous measurements.
• Disadvantages: Silver chloride can be reduced in strongly alkaline solutions or by strong reducing agents.

5. Applications of Redox and Potentiometric Titrations

These techniques are indispensable in various scientific and industrial fields:

1. Quantitative Determination of Oxidizing/Reducing Agents:
   • Iron Determination: Titration of Fe2+ with KMnO4​ or K2​Cr2​O7​.
   • Vitamin C (Ascorbic Acid) Assay: Titration with iodine.
   • Hydrogen Peroxide Determination: Titration with KMnO4​.
   • Chlorine in Water: Iodometric titration.
2. Environmental Analysis: Measuring dissolved oxygen in water, chemical oxygen demand (COD), and various pollutants.
3. Pharmaceutical Analysis: Assay of active pharmaceutical ingredients that undergo redox reactions.
4. Food Analysis: Determining antioxidants, vitamin content, and certain metal ions.
5. Clinical Analysis: Indirectly measuring substances like glucose (after an enzymatic redox reaction).
6. Industrial Quality Control: Monitoring redox processes in manufacturing, e.g., in plating baths, textile bleaching, and pulp and paper industries.
7. Potentiometric Titrations (General): Beyond redox, potentiometry is used to monitor acid-base, precipitation, and complexometric titrations, offering precise endpoint detection where visual indicators may fail.

Conclusion

Redox titrations and their potentiometric measurement represent a cornerstone of quantitative analytical chemistry. By leveraging the principles of electron transfer and electrochemical potentials, these methods offer high precision and versatility for determining a wide array of chemical species. The combination of redox chemistry with objective potentiometric detection ensures accurate and reliable analytical results across diverse scientific and industrial applications.

Redox and Potentiometric Titrations: Multiple Choice Questions

Instructions: Choose the best answer for each question. Explanations are provided after each question.

1. What is the fundamental process occurring in a redox titration? a) Acid-base neutralization b) Precipitation c) Electron transfer d) Complex formation e) Dissolution

Explanation: Redox reactions are defined by the transfer of electrons from one species to another.

2. In a redox reaction, a species that loses electrons is said to be: a) Reduced b) Oxidized c) Precipitated d) Neutralized e) Complexed

Explanation: Oxidation is the loss of electrons, resulting in an increase in oxidation state.

3. What is an oxidizing agent (oxidant)? a) A species that gains electrons and causes oxidation. b) A species that loses electrons and causes reduction. c) A species that causes reduction by gaining electrons. d) A species that causes oxidation by accepting electrons (itself getting reduced). e) A species that is always in a lower oxidation state.

Explanation: An oxidizing agent accepts electrons from another species, thus causing that species to be oxidized while itself being reduced.

4. Which equation relates electrode potential to the activity (or concentration) of electroactive species in a redox half-reaction? a) Beer-Lambert Law b) Ohm’s Law c) Nernst Equation d) Arrhenius Equation e) Henderson-Hasselbalch Equation

Explanation: The Nernst equation is central to understanding how concentration changes affect the potential of an electrode in a redox system.

5. What is the defined standard potential of the Standard Hydrogen Electrode (SHE)? a) +0.241 V b) −0.197 V c) 0.000 V d) 1.000 V e) Variable, depending on temperature

Explanation: The SHE serves as the universal zero reference point for all standard electrode potentials.

6. Which common oxidizing agent acts as its own indicator in acidic solutions, turning from deep purple to colorless? a) Potassium dichromate (K2​Cr2​O7​) b) Cerium(IV) sulfate (Ce(SO4​)2​) c) Potassium permanganate (KMnO4​) d) Iodine (I2​) e) Sodium thiosulfate (Na2​S2​O3​)

Explanation: Permanganate ion (MnO4−​) is intensely purple, while its reduction product Mn2+ is nearly colorless, making it a self-indicator.

7. A redox indicator changes color based on: a) The pH of the solution. b) Its concentration. c) The electrode potential of the solution. d) The temperature of the solution. e) The volume of titrant added.

Explanation: Redox indicators are themselves redox active species whose oxidized and reduced forms have different colors, and their color transition depends on the potential of the solution.

8. What type of indicator electrode is typically used to monitor the potential change during a potentiometric redox titration? a) Glass electrode b) Ion-selective electrode (ISE) c) Calomel electrode d) Inert metallic electrode (e.g., platinum) e) Silver/Silver Chloride electrode

Explanation: An inert electrode like platinum provides a surface for electron transfer of the redox couple but does not participate in the reaction itself, allowing it to measure the solution potential.

9. What is plotted on the y-axis of a potentiometric redox titration curve? a) Volume of titrant b) pH c) Cell potential (Ecell​) d) Concentration of analyte e) Current

Explanation: The curve plots the continuously measured potential against the volume of titrant added.

10. How is the equivalence point determined from a first derivative plot (ΔE/ΔV vs. Vavg​) of a potentiometric titration? a) The point where the plot crosses zero. b) The maximum value of the plot. c) The minimum value of the plot. d) The point where the slope is zero. e) The inflection point of the curve.

Explanation: The maximum of the first derivative corresponds to the steepest part of the S-shaped titration curve, which is the equivalence point.

11. What is the primary advantage of using potentiometric detection for redox titrations over visual indicators? a) Faster analysis time. b) Cheaper equipment. c) Applicable to colored or turbid solutions. d) Requires less sample. e) Does not require calibration.

Explanation: Potentiometric methods eliminate subjective color judgments and can be used when the solution’s color obscures visual indicators.

12. Which reference electrode is often preferred over the Saturated Calomel Electrode (SCE) due to being non-toxic and more robust, especially for field measurements? a) Standard Hydrogen Electrode (SHE) b) Silver/Silver Chloride Electrode (Ag/AgCl) c) Platinum electrode d) Glass electrode e) Lead electrode

Explanation: The Ag/AgCl electrode is a widely used and reliable alternative, being mercury-free.

13. In the Nernst equation, what does ‘n’ represent? a) Concentration of reactants. b) Number of moles. c) Number of electrons involved in the half-reaction. d) Electrode potential. e) Temperature.

Explanation: ‘n’ is the stoichiometric number of electrons transferred in the balanced half-reaction, which dictates the slope in the Nernst equation.

14. If a chemical species gains electrons, it is said to be: a) Oxidized b) Reduced c) Neutralized d) Precipitated e) Complexed

Explanation: Reduction is the gain of electrons, leading to a decrease in oxidation state.

15. What type of agent is Fe2+ when it reacts to form Fe3+? a) Oxidizing agent b) Reducing agent c) Complexing agent d) Precipitating agent e) Indicator

Explanation: Fe2+ loses an electron to become Fe3+, thus it is acting as a reducing agent (it gets oxidized).

16. Which of the following is a disadvantage of using Potassium Permanganate (KMnO4​) as a titrant? a) It is not a strong enough oxidizing agent. b) It is not a primary standard. c) It requires an external indicator. d) It is colorless. e) It only works in basic solutions.

Explanation: KMnO4​ is not stable enough or pure enough to be a primary standard, so its solution concentration must be determined by standardization against a known primary standard (e.g., sodium oxalate).

17. What is the color change of Ferroin indicator from its reduced to its oxidized form? a) Red to blue b) Blue to red c) Colorless to pink d) Yellow to orange e) Violet to colorless

Explanation: Ferroin, in its Fe(phen)32+​ (reduced) form, is red. In its Fe(phen)33+​ (oxidized) form, it is light blue.

18. At the equivalence point of a symmetric redox titration (e.g., Ce4+ vs. Fe2+), the potential is typically: a) Equal to the E0 of the analyte. b) Equal to the E0 of the titrant. c) An average of the two standard potentials involved. d) Zero. e) Infinite.

Explanation: For reactions where the number of electrons exchanged is equal, the equivalence point potential is often the average of the two standard potentials. More generally, it depends on the stoichiometry.

19. Why is a high-impedance voltmeter required for potentiometric measurements? a) To increase the current flow. b) To make the measurement faster. c) To minimize current drawn from the cell, thus preventing disturbance of the equilibrium potential. d) To generate a potential. e) To heat the sample.

Explanation: Drawing current would change the concentrations at the electrode surface and thus alter the equilibrium potential being measured. High impedance prevents this.

20. What is the purpose of a salt bridge in an electrochemical cell setup for potentiometry? a) To provide a surface for electron transfer. b) To generate light. c) To complete the electrical circuit and maintain charge neutrality. d) To separate the two electrodes physically. e) To amplify the signal.

Explanation: The salt bridge allows ions to flow between the two half-cells, maintaining electrical neutrality and completing the circuit without allowing the solutions to mix.

21. In a potentiometric titration curve, what does the second derivative plot (Δ2E/ΔV2 vs. Vavg​) help determine? a) The initial potential. b) The potential after the equivalence point. c) The point where the slope of the first derivative is maximum. d) The point where the second derivative crosses zero (the inflection point of the first derivative). e) The concentration of the titrant.

Explanation: The zero-crossing point of the second derivative corresponds precisely to the maximum slope of the first derivative, indicating the equivalence point.

22. Which of the following is a common reducing agent used as a titrant in iodometric titrations? a) Potassium permanganate b) Potassium dichromate c) Sodium thiosulfate d) Cerium(IV) sulfate e) Iodine

Explanation: Sodium thiosulfate (Na2​S2​O3​) is widely used to titrate iodine, reducing I2​ to I−.

23. What type of titration would typically use starch solution as an indicator? a) Acid-base titration b) Complexometric titration c) Precipitation titration d) Titrations involving iodine (iodometry/iodimetry) e) Red-ox titration (general)

Explanation: Starch forms a distinct blue complex with iodine, making it a specific and highly sensitive indicator for the presence of I2​.

24. The term “oxidant” is synonymous with: a) Reducing agent b) Electron donor c) Oxidizing agent d) Precipitate e) Ligand

Explanation: An oxidant is an oxidizing agent, meaning it causes oxidation by taking electrons.

25. If the potential of a redox reaction measured with a platinum electrode is very high and positive, it suggests the presence of a: a) Strong reducing agent b) Strong oxidizing agent c) Non-redox active species d) Weak acid e) Weak base

Explanation: A highly positive potential indicates a strong tendency for reduction to occur, characteristic of a strong oxidizing environment or the presence of a strong oxidizing agent.

26. Why is Potassium Dichromate (K2​Cr2​O7​) often preferred over Potassium Permanganate (KMnO4​) in some titrations? a) It is a stronger oxidizing agent. b) It acts as a self-indicator. c) It is a primary standard. d) It is colorless. e) It only works in basic solutions.

Explanation: Potassium dichromate is highly pure, stable, and non-hygroscopic, making it suitable for direct preparation of standard solutions (primary standard).

27. What is the process of losing electrons and increasing oxidation state called? a) Reduction b) Precipitation c) Neutralization d) Oxidation e) Complexation

Explanation: Oxidation is defined as the loss of electrons.

28. Which application involves determining the amount of Fe2+ by titrating with KMnO4​ or K2​Cr2​O7​? a) Water hardness determination b) Chloride analysis c) Iron determination d) pH measurement e) Complexometric titration

Explanation: The determination of iron, often in its Fe2+ form, by oxidation to Fe3+ is a very common redox titration.

29. The region of a redox titration curve before the equivalence point is dominated by the potential of the: a) Titrant’s redox couple. b) Indicator’s redox couple. c) Analyte’s redox couple. d) Reference electrode. e) Liquid junction.

Explanation: Before the equivalence point, the solution contains mostly the unreacted analyte and its oxidized/reduced form, so the potential is governed by the analyte’s redox system.

30. What is the fundamental difference between an indicator electrode and a reference electrode in potentiometry? a) Only indicator electrodes use membranes. b) Indicator electrode potential changes with analyte/titrant, reference electrode potential is constant. c) Reference electrodes are always larger. d) Indicator electrodes are always made of platinum. e) Reference electrodes generate current, indicator electrodes do not.

Explanation: The key distinction is their function: one responds to changes in the solution, the other provides a stable baseline.

31. What are the units for the Faraday constant (F) in the Nernst Equation? a) Joules per Kelvin per mole (J K−1 mol−1) b) Volts (V) c) Coulombs per mole (C mol−1) d) Liters per mole (L mol−1) e) Grams per mole (g mol−1)

Explanation: The Faraday constant represents the charge of one mole of electrons.

32. Which of the following would NOT typically be determined using a redox titration? a) Concentration of Vitamin C (Ascorbic Acid) b) Dissolved oxygen in water c) Hardness of water d) Hydrogen peroxide concentration e) Chlorine content in bleach

Explanation: Water hardness is typically determined by complexometric titration (EDTA), not redox titration.

33. What is the main reason for the sharp potential change around the equivalence point in a redox titration curve? a) Change in pH. b) Sudden change in the ratio of oxidized and reduced forms of the analyte/titrant. c) Formation of a precipitate. d) Depletion of the indicator. e) Increase in temperature.

Explanation: Around the equivalence point, there is a rapid shift from the dominance of the analyte’s redox couple to the dominance of the titrant’s redox couple, leading to a dramatic potential change.

34. In the simplified Nernst equation at 25∘C, the constant 0.05916 results from the combination of: a) R, T, and F divided by n. b) R, T, and F and conversion from natural to common logarithm. c) Only R and T. d) Only F. e) n and T.

Explanation: It’s derived from RT/F multiplied by ln(10) to convert from natural to common logarithm, at a specific temperature (298.15 K).

35. If an oxidizing agent is very strong, the E0 for its reduction half-reaction will typically be: a) Very negative. b) Close to zero. c) Very positive. d) Highly variable. e) Equal to its concentration.

Explanation: A very positive standard reduction potential indicates a strong tendency to gain electrons, meaning it is a strong oxidizing agent.

36. A plot of the potential (Ecell​) versus the volume of titrant is used for potentiometric titrations. This method is considered a type of: a) Direct potentiometry b) Indirect potentiometry c) Titrimetry d) Coulometry e) Amperometry

Explanation: Potentiometric titration falls under the broader category of titrimetry, where the volume of titrant is used to determine the analyte. It also uses potentiometric measurement.

37. Which of the following is an example of an inert metallic indicator electrode? a) Ag/AgCl electrode b) Glass electrode c) Calomel electrode d) Platinum electrode e) Lead electrode

Explanation: Platinum electrodes are commonly used as inert indicators because they provide a surface for electron transfer without reacting themselves.

38. What is the consequence of having a significant liquid junction potential in a potentiometric measurement? a) The measured potential will be higher than the true potential. b) The measured potential will be lower than the true potential. c) It introduces uncertainty and affects the accuracy of the potential measurement. d) It speeds up the electrode response time. e) It makes the electrode more selective.

Explanation: The liquid junction potential adds an unknown and often variable component to the measured cell potential, reducing the accuracy of the potential reading.

39. In a redox titration of Fe2+ with Ce4+, what happens to the concentration of Fe2+ as Ce4+ is added before the equivalence point? a) It increases. b) It remains constant. c) It decreases. d) It becomes zero. e) It oscillates.

Explanation: As Ce4+ is added, it reacts with and consumes Fe2+, causing the concentration of Fe2+ to decrease.

40. What clinical application sometimes uses redox reactions, indirectly measured by potentiometry, for substance determination? a) Blood pressure monitoring. b) Body temperature measurement. c) Glucose determination (e.g., in enzymatic biosensors). d) Bone density measurement. e) Heart rate monitoring.

Explanation: Enzymatic glucose sensors, for example, involve a redox reaction catalyzed by an enzyme, and the resulting change in electron flow or potential can be measured potentiometrically to determine glucose concentration.