Potentiometric Electrodes and Potentiometry: Unveiling Electrochemical Insights

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. It’s a non-destructive method widely applied in various fields, from environmental monitoring and clinical diagnostics to industrial process control and fundamental chemical research.

1. Fundamentals of Potentiometry

At the heart of potentiometry lies the measurement of an electromotive force (EMF), or cell potential, of an electrochemical cell. This potential develops across an interface between two phases and is sensitive to the activity (or effective concentration) of certain ions.

1.1. Electrochemical Cells

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

• Indicator Electrode: The electrode whose potential depends on the concentration of the analyte being measured.
• Reference Electrode: An electrode whose potential is constant and independent of the analyte concentration. It provides a stable reference point.

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.

1.2. The Nernst Equation

The relationship between the electrode potential and the concentration (more accurately, activity) of electroactive species is described by the Nernst Equation:

For a half-reaction: aA+bB+ne−⇌cC+dD

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, 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 for common logarithm (log10​):

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

This equation is fundamental to understanding how changes in analyte concentration lead to measurable changes in potential.

2. Reference Electrodes

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.

2.1. 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.

2.2. 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.

2.3. 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.

3. Indicator Electrodes

Indicator electrodes respond selectively to the concentration (activity) of a specific analyte ion in the sample solution. Their potential changes according to the Nernst equation.

3.1. Metallic Indicator Electrodes

These electrodes rely on electron transfer at a metal surface.

• First Kind: Pure metal electrodes immersed in a solution containing their own ions (e.g., Ag electrode in Ag+ solution).
  • Ag+(aq)+e−⇌Ag(s)
  • Limited use due to susceptibility to oxidation, interference, and not all metals form stable potentials.
• Second Kind: Metal electrode coated with a sparingly soluble salt of that metal, immersed in a solution containing the anion of the salt (e.g., Ag/AgCl electrode as an indicator for Cl−).
  • AgCl(s)+e−⇌Ag(s)+Cl−(aq)
  • Useful for determining the concentration of the anion.
• Third Kind: Respond to a third ion by a two-step equilibrium (less common).
• Redox Electrodes (Inert Metal): Inert metal electrodes (e.g., platinum, gold) used to measure the potential of a redox system. They do not participate in the reaction but provide a surface for electron transfer.
  • Example: Fe3+/Fe2+ system using a Pt electrode. Potential depends on the ratio of oxidized and reduced species.

3.2. Membrane Indicator Electrodes (Ion-Selective Electrodes – ISEs)

ISEs are the most important class of indicator electrodes due to their remarkable selectivity. They develop a potential across a selective membrane that is proportional to the logarithm of the analyte’s activity.

• Mechanism: A selective membrane separates two solutions. A potential develops across this membrane due to the difference in activity of the ion of interest on either side.
• General Response: For a singly charged cation: E=K+z0.05916​logaion​ (where z is the charge of the ion).
• Properties:
  • Selectivity: The ability to preferentially respond to one ion over others. Expressed by the selectivity coefficient (KA,Bpot​), which quantifies the electrode’s response to an interfering ion B relative to the analyte A. A smaller KA,Bpot​ indicates better selectivity.
  • Linear Range: The range of concentrations over which the electrode exhibits a Nernstian (linear) response.
  • Detection Limit: The lowest concentration at which the electrode provides a reliable signal.
  • Response Time: How quickly the electrode reaches a stable potential after a change in analyte concentration.

Types of ISEs:

1. Glass Electrode (for pH measurement):
   • Membrane: Thin, hydrated glass membrane (usually made of a special lithium-barium silicate glass).
   • Mechanism: The potential develops across the outer surface of the glass membrane due to selective exchange of H+ ions between the solution and silanol groups (Si-O−) in the hydrated gel layer. The inner surface has a fixed H+ reference solution.
   • Integrated Design: Most modern pH electrodes are “combination electrodes,” integrating both the glass indicator electrode and the Ag/AgCl reference electrode into a single probe for convenience.
   • Alkaline Error: At very high pH (low H+ concentrations) and high concentrations of alkali metal ions (e.g., Na+), the glass membrane can respond to these metal ions as well, leading to a measured pH that is lower than the true pH.
   • Acid Error: At very low pH, the electrode response can deviate from Nernstian behavior, often reading higher than the true pH.
2. Liquid-Membrane Electrodes:
   • Membrane: A porous, hydrophobic polymer film (e.g., PVC) impregnated with a water-insoluble liquid ion exchanger (a neutral carrier or charged organic complexing agent).
   • Mechanism: The ion exchanger selectively binds the target ion, facilitating its transport across the membrane, generating a potential.
   • Applications: Ca2+, K+, NO3−​, ClO4−​, etc. The Ca2+ electrode uses an organophosphate as an ion exchanger. The K+ electrode uses valinomycin as a neutral carrier (highly selective).
3. Solid-State Electrodes:
   • Membrane: A single crystal or a polycrystalline pellet of an inorganic salt.
   • Mechanism: Conduction occurs by movement of either cations or anions through crystal lattice defects or vacancies.
   • Applications: F− electrode (uses a LaF3​ crystal doped with EuF2​ for conductivity). Also Cl−, Br−, I−, S2−, Ag+, Cu2+ electrodes.
4. Gas-Sensing Electrodes:
   • Membrane: A gas-permeable membrane that separates the sample solution from a thin layer of an internal electrolyte solution containing a conventional pH electrode and a reference electrode.
   • Mechanism: Dissolved gas from the sample diffuses across the gas-permeable membrane into the internal electrolyte, changing its pH. The pH change is then measured by the internal pH electrode.
   • Applications: CO2​ (measures H+ from CO2​+H2​O⇌H2​CO3​⇌H++HCO3−​), NH3​, SO2​, NOx​.

4. Potentiometric Titrations

Potentiometry is an excellent technique for determining the equivalence point of a titration, especially when a visual indicator is unavailable or provides a poor color change.

• Procedure: An indicator electrode (responsive to the analyte or titrant) and a reference electrode are immersed in the analyte solution. The potential of the cell is measured as a function of the volume of titrant added.
• Titration Curve: A plot of cell potential (Ecell​) vs. titrant volume. The curve typically shows a steep sigmoid shape, with the most significant change in potential occurring around the equivalence point.
• Endpoint Determination: The equivalence point is found by identifying the point of maximum slope on the titration curve. This can be done by:
  • First Derivative Plot: Plotting ΔE/ΔV vs. Vavg​. The maximum of this plot corresponds to the equivalence point.
  • Second Derivative Plot: Plotting Δ2E/ΔV2 vs. Vavg​. The point where this plot crosses zero corresponds to the maximum slope in the first derivative, hence the equivalence point.
• Advantages:
  • Applicable to colored or turbid solutions where visual indicators fail.
  • Higher accuracy and precision than visual methods.
  • Suitable for weak acids/bases or very dilute solutions where color changes are indistinct.
  • Can be automated.

5. Instrumentation for Potentiometry

A basic potentiometric setup consists of:

• Potentiometer/pH Meter: A high-impedance voltmeter capable of measuring cell potential without drawing significant current (which would disturb the equilibrium). pH meters are specialized potentiometers for pH measurement.
• Indicator Electrode: Chosen based on the analyte.
• Reference Electrode: Provides a stable potential.
• Stirrer: To ensure homogeneous mixing of the solution.

6. Applications of Potentiometry

Potentiometry is a versatile analytical tool with broad applications:

1. pH Measurement: The most common application, using a glass electrode to measure the activity of hydrogen ions in virtually any aqueous sample.
2. Ion Concentration Determination (Direct Potentiometry): Using various ISEs to directly measure the concentration of specific ions (e.g., F− in water, K+ in blood serum, Ca2+ in dairy products). Requires careful calibration.
3. Potentiometric Titrations: As described above, for determining equivalence points in acid-base, precipitation, redox, and complexometric titrations.
4. Environmental Monitoring: Measuring pH of rain and natural waters, fluoride in drinking water, ammonia in wastewater, nitrate in agricultural runoff.
5. Clinical Analysis: Measuring electrolytes (Na+,K+,Ca2+,Cl−) in blood, urine, and other biological fluids. Blood gas analysis (pCO2​) also uses gas-sensing electrodes.
6. Industrial Process Control: Monitoring and controlling pH in chemical reactions, fermentation processes, and wastewater treatment plants.
7. Food Analysis: Measuring acidity in fruit juices, salt content in processed foods, calcium in dairy.
8. Pharmaceutical Analysis: Quality control of drugs, dissolution testing, endpoint determination in titrations.

Conclusion

Potentiometry stands as a fundamental and indispensable electroanalytical technique. Its ability to provide rapid, accurate, and selective measurements of ion activities and concentrations makes it invaluable across countless scientific and industrial applications. From the ubiquitous pH meter to sophisticated ion-selective electrodes, potentiometric methods continue to offer unique insights into the electrochemical world of solutions.

Potentiometric Electrodes and Potentiometry: Multiple Choice Questions

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

1. What is the primary measurement made in potentiometry? a) Current b) Resistance c) Cell potential (EMF) d) Conductance e) Absorbance

Explanation: Potentiometry measures the potential difference (voltage) between two electrodes in an electrochemical cell under zero-current conditions.

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

Explanation: The Nernst equation is fundamental to potentiometry, describing the quantitative relationship between potential and concentration.

3. What is the role of the “reference electrode” in a potentiometric cell? a) Its potential varies with analyte concentration. b) It conducts current through the solution. c) It provides a stable, known potential for comparison. d) It is directly responsive to the analyte. e) It introduces interference.

Explanation: A reference electrode must maintain a constant and stable potential to serve as a reliable benchmark for measuring the indicator electrode’s changing potential.

4. 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 is the universal reference point for electrode potentials, defined as 0.000 V under standard conditions.

5. Which of the following is a common secondary reference electrode that contains mercury? a) Standard Hydrogen Electrode (SHE) b) Silver/Silver Chloride Electrode (Ag/AgCl) c) Calomel Electrode (SCE) d) Glass electrode e) Platinum electrode

Explanation: The Calomel Electrode (SCE) uses a mercury-mercurous chloride paste in a KCl solution.

6. Which class of indicator electrodes is known for its remarkable selectivity to specific ions? a) Metallic indicator electrodes of the first kind b) Metallic indicator electrodes of the second kind c) Redox electrodes d) Membrane indicator electrodes (Ion-Selective Electrodes – ISEs) e) Glass electrodes (specifically for non-H+ ions)

Explanation: Ion-Selective Electrodes (ISEs) are designed with selective membranes that respond primarily to a single type of ion.

7. The potential of a glass electrode (for pH measurement) primarily develops across which component? a) The inner Ag/AgCl wire. b) The salt bridge. c) The thin, hydrated glass membrane. d) The internal buffer solution. e) The ceramic frit.

Explanation: The potential is generated by the selective exchange of hydrogen ions at the surface of the hydrated glass membrane.

8. What is the “alkaline error” observed with glass electrodes? a) The electrode reads too high at low pH. b) The electrode reads too low at high pH and high alkali metal ion concentration. c) The electrode becomes insensitive at high pH. d) The electrode breaks in alkaline solutions. e) The potential becomes unstable in alkaline solutions.

Explanation: At very high pH, the glass membrane can start responding to alkali metal ions (like Na+), causing the measured pH to be lower than the true pH.

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

Explanation: A potentiometric titration curve typically plots the measured cell 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 first derivative plot shows a sharp peak (maximum) at the point of greatest potential change, which corresponds to the equivalence point.

11. What is the main advantage of potentiometric titrations over visual titrations? 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 titrations are advantageous when visual indicators are obscured by sample color or turbidity.

12. The selectivity coefficient (KA,Bpot​) for an ISE quantifies its response to: a) The analyte A. b) An interfering ion B relative to the analyte A. c) The overall cell potential. d) The pH of the solution. e) The temperature.

Explanation: The selectivity coefficient indicates how much an interfering ion (B) affects the electrode’s measurement of the primary ion (A). A smaller value means better selectivity.

13. A metallic indicator electrode of the “first kind” consists of: a) An inert metal in a redox system. b) A metal coated with its sparingly soluble salt. c) A pure metal immersed in a solution containing its own ions. d) A glass membrane. e) A gas-permeable membrane.

Explanation: A first-kind electrode is a simple metal wire or bar immersed in a solution containing its corresponding metal ions.

14. Which reference electrode is often preferred over 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) Mercury/Mercurous Sulfate Electrode d) Glass electrode e) Lead electrode

Explanation: The Ag/AgCl electrode is a widely used and reliable alternative to the SCE, particularly in situations where mercury is a concern.

15. What is a key application of the F− ion-selective electrode? a) Measuring chloride in blood. b) Determining calcium in water. c) Measuring fluoride in drinking water. d) Monitoring oxygen levels in gases. e) Quantifying proteins in solutions.

Explanation: The fluoride ISE, which uses a LaF3​ crystal membrane, is highly selective and widely used for fluoride analysis.

16. Which type of ISE uses a gas-permeable membrane and an internal pH electrode to measure dissolved gases? a) Glass electrode b) Liquid-membrane electrode c) Solid-state electrode d) Gas-sensing electrode e) Enzyme electrode

Explanation: Gas-sensing electrodes work by allowing a dissolved gas to diffuse across a membrane, altering the pH of an internal electrolyte, which is then measured by a pH electrode.

17. 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: Potentiometric measurements are based on equilibrium potentials. Drawing current would change the concentrations at the electrode surface and alter the potential.

18. What is the purpose of a salt bridge in an electrochemical cell? 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 ion migration to balance charges and complete the circuit without mixing the solutions of the two half-cells.

19. What term describes the lowest concentration at which an ISE provides a reliable signal? a) Linear range b) Selectivity coefficient c) Response time d) Detection limit e) Nernstian slope

Explanation: The detection limit is the minimum concentration of the analyte that can be accurately measured by the electrode.

20. Which analytical application of potentiometry is most commonly performed in laboratories and everyday settings? a) Potentiometric titrations for redox reactions. b) Determination of enzyme activity. c) pH measurement. d) Measurement of heavy metal pollutants. e) In-situ analysis of gases.

Explanation: The pH meter, using a glass electrode, is by far the most widespread application of potentiometry.

21. 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.

22. If a sample solution has a high concentration of an interfering ion, what happens to the accuracy of an ISE measurement of the analyte? a) It increases. b) It remains unchanged. c) It decreases. d) It depends only on temperature. e) It depends only on the activity coefficient.

Explanation: Interfering ions that the ISE also responds to will lead to an inaccurate measurement of the target analyte’s concentration.

23. What type of titration is illustrated by a plot of E_cell vs. volume of titrant? a) Gravimetric titration b) Spectrophotometric titration c) Potentiometric titration d) Conductometric titration e) Thermometric titration

Explanation: Potentiometric titration directly measures the cell potential as a function of titrant volume.

24. The inner reference electrode within a combination glass pH electrode is typically: a) SHE b) SCE c) Ag/AgCl d) Platinum e) Lead

Explanation: A miniature Ag/AgCl electrode is almost universally used as the inner reference electrode within a combination glass pH electrode.

25. A liquid-membrane ISE for Ca2+ often uses what type of compound as its ion exchanger? a) Inorganic salt b) Polymer c) Organophosphate d) Valinomycin e) Glass

Explanation: Organophosphates are commonly used as ion exchangers in calcium liquid-membrane electrodes due to their selective binding of Ca2+.

26. Which type of ISE is best suited for measuring the activity of K+ ions with high selectivity, often in biological samples? a) Glass electrode b) F− solid-state electrode c) Liquid-membrane K+ electrode with valinomycin d) Gas-sensing CO2​ electrode e) Metallic indicator electrode

Explanation: Valinomycin is a cyclic ether that exhibits exceptional selectivity for K+ over Na+, making it ideal for potassium ISEs.

27. What is a “linear range” for an ISE? a) The maximum concentration it can measure. b) The range of concentrations over which the electrode exhibits a Nernstian (linear) response. c) The minimum concentration it can measure. d) The range of pH values it can operate in. e) The range of temperatures it can operate in.

Explanation: Within the linear range, the plot of potential vs. logarithm of concentration is a straight line with the expected Nernstian slope.

28. If the Nernstian slope for a monovalent cation at 25∘C is +0.05916 V/decade, what would be the slope for a divalent cation? a) +0.05916 V/decade b) +0.11832 V/decade c) +0.02958 V/decade d) −0.05916 V/decade e) −0.02958 V/decade

Explanation: For a divalent cation, z=2, so the slope is 0.05916/2=0.02958 V/decade.

29. What is the primary disadvantage of the Standard Hydrogen Electrode (SHE) for routine laboratory use? a) It is too large. b) It is expensive. c) It is impractical and difficult to set up and maintain. d) It is not accurate. e) It interferes with most samples.

Explanation: The requirement for pure hydrogen gas at constant pressure and a precisely 1 M H+ solution makes SHE very cumbersome for everyday use.

30. Which term describes the potential that arises at the interface between the salt bridge and the sample solution? a) Standard potential b) Liquid junction potential c) Reference potential d) Indicator potential e) Redox potential

Explanation: The liquid junction potential is generated by the unequal diffusion rates of ions across the interface.

31. How does temperature affect the Nernst equation? a) It only affects the E0 term. b) It affects the slope term (RT/nF). c) It has no effect. d) It only affects the concentration terms. e) It only affects the number of electrons ‘n’.

Explanation: Temperature (T) is directly present in the RT/nF term, meaning the slope of the Nernstian response changes with temperature.

32. What is a typical application of potentiometry in industrial process control? a) Measuring mass of products. b) Controlling furnace temperature. c) Monitoring and controlling pH in chemical reactors. d) Analyzing the elemental composition of solids. e) Determining molecular structure.

Explanation: Maintaining specific pH levels is crucial for many industrial chemical reactions and processes, making pH measurement a vital control parameter.

33. An Ag/AgCl electrode is used as an indicator electrode for Cl−. This makes it what kind of metallic electrode? a) First kind b) Second kind c) Third kind d) Redox e) Inert

Explanation: A second-kind electrode is a metal coated with its sparingly soluble salt, responsive to the anion of that salt. The Ag/AgCl electrode responds to Cl−.

34. 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 second derivative crosses zero at the point of maximum slope in the first derivative plot, which precisely indicates the equivalence point.

35. What is the fundamental difference between an indicator electrode and a reference electrode? a) Only indicator electrodes use membranes. b) Indicator electrode potential changes with analyte, 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 the responsiveness to the analyte: indicator electrodes show a potential change, while reference electrodes provide a stable baseline.

36. A CO2​ gas-sensing electrode works by measuring a pH change in an internal solution caused by the diffusion of CO2​. What is the reaction that causes this pH change? a) CO2​+H2​O⇌H2​CO3​⇌H++HCO3−​ b) CO2​+OH−⇌HCO3−​ c) CO2​⇌C+O2​ d) CO2​+e−⇌CO2−​ e) CO2​ reacting with the electrode material.

Explanation: Dissolved CO2​ reacts with water to form carbonic acid, which then dissociates, releasing H+ ions and thus changing the pH measured by the internal pH electrode.

37. If an ISE for a cation is placed in a solution, and the measured potential is more positive than expected, what might this indicate about the cation’s activity in the solution? a) It is lower than expected. b) It is higher than expected. c) It is zero. d) It is independent of the potential. e) The solution is acidic.

Explanation: For a cation, a more positive potential corresponds to a higher activity (concentration) according to the Nernst equation (E=K+z0.05916​logaion​ for cations).

38. What is a significant clinical application of potentiometry? a) Measuring blood glucose levels directly. b) Quantifying proteins in urine. c) Measuring electrolyte concentrations (Na+,K+,Ca2+) in blood. d) Detecting bacterial infections. e) Analyzing DNA sequences.

Explanation: ISEs are widely used in clinical laboratories to rapidly and accurately measure key electrolyte levels in biological fluids.

39. Why is careful calibration necessary for direct potentiometric measurements of ion concentration? a) To avoid interfering with the salt bridge. b) To account for differences in cell potential between different instruments. c) To relate the measured potential to the true activity of the analyte. d) To prevent electrode breakage. e) To ensure the electrode is sterile.

Explanation: Direct potentiometry relies on comparing the measured potential to a calibration curve generated using standard solutions of known activities to accurately determine the unknown concentration.

40. What is the effect of changing temperature on the measured potential in potentiometry (assuming the Nernst equation holds)? a) The potential will always increase. b) The potential will always decrease. c) The slope of the Nernstian response will change. d) There will be no effect on the potential. e) Only the E0 value changes.

Explanation: The Nernst equation includes temperature (T) in the RT/nF term, which is the slope factor. Therefore, changes in temperature directly alter the sensitivity of the electrode’s response to concentration changes.