Electrochemistry and chemical sensors

Abstract

Background

Chemical sensors utilizing electrochemical and optical detection methods have become routine analytical tools in clinical chemistry applications, especially for measurement of critical care analytes (blood gases, electrolytes, metabolites) directly in whole blood at point of care or near patient testing. Coupling electrochemical or optical transducers together with chemical or biological recognition elements has expanded the role of chemical sensors for measurement of analytes without direct electrochemical or optical activity.

Content

This chapter reviews fundamental aspects of electrochemical and optical measurements and their applications in routine clinical practice for measurements of blood gases (pH, partial pressure of oxygen [ P O ₂ ], partial pressure of carbon dioxide [ P CO ₂ ]), electrolytes (sodium [Na ⁺ ], potassium [K ⁺ ], calcium [Ca ²⁺ ], chloride [Cl ⁻ ], magnesium [Mg ²⁺ ], lithium [Li ⁺ ]), hematocrit [HCT], and toxic metals (lead [pb ²⁺ ]) directly in whole blood. The base principles of these electrochemical and optical sensors are applied as building blocks for enzyme-based biosensors for measurement of important metabolites directly in whole blood (glucose, lactate, urea, creatinine). Recent advances in miniaturization of enabling technologies and the demand for direct whole blood measurements have transformed these sensors as ideal measurement technologies for point-of-care testing, near patient testing, or at-home monitoring. Affinity sensors utilizing antibodies, nucleic acids, and aptamers are expanding the role of biological recognition elements in biosensors. These molecules, when coupled to electrochemical, optical, and other transducers, produce biosensors for sensitive detection of biomarkers for cancer, cardiac disease, and genetic testing. Although there are several commercial applications of affinity sensors in clinical use, continued growth in research publications and patents points to future possibilities for expanding applications to other clinical markers. This chapter discusses nanotechnology to further enhance sensitivity of biosensors. The sensor technologies routinely applied for in vitro measurements of critical care analytes (blood gases, glucose) are being adapted for in vivo and minimally invasive or noninvasive wearable applications, through developing solutions to address problems such as biocompatibility and calibration stability. Commercial products for continuous glucose monitoring (CGM) and health vital sign monitoring are a few examples of wearable sensors that are increasingly available.

Introduction

Recent advances in electromechanical, microfluidic, optical, and computer technologies have enabled the wider adoption of chemical sensors utilizing electrochemical or optical transduction methodologies (and associated biosensors) in clinical analysis systems. Sensors for measurement of blood gases, electrolytes, metabolites, trace metals, and other important biomarkers are incorporated into automated, point-of-care, and laboratory clinical analyzers (see Chapters 29 , 30 , and 37 , respectively). Miniaturization of sensors and the enabling technologies together with increasing demand for analyzers that are easy to use, portable, have wireless capability, and offer low-skilled operation with minimal process steps are driving the growth of point-of-care testing. Biosensors that incorporate the same sensing principles together with bioelements (such as enzymes, antibodies, aptamers, nucleic acids, etc.) have also been successfully applied for expanding the capabilities of these devices to measure or monitor different metabolites, coagulation reactions, biomarkers, detecting drugs, or toxic chemicals through ultrasensitive immunoassays or genetic sequences in whole blood, plasma, serum, or urine samples. When integrated into chromatographic systems (see Chapter 19 ), electrochemical detectors provide a highly sensitive and selective means of detecting a variety of other analytes, such as therapeutic drugs, neurotransmitters, glutathione, and homocysteine. In addition, the development and application of optodes, which are based on some of the same selective chemistries used in electrochemical devices, have resulted in another analytical tool for measuring blood gases and electrolytes.

In this chapter, the fundamental electrochemical principles of (1) potentiometry, (2) voltammetry and/or amperometry, (3) conductance, and (4) coulometry are summarized, and their clinical applications are presented. Optodes and biosensors are also discussed. The chapter concludes with a discussion of in vivo, minimally invasive, and noninvasive sensors that are growing in clinical use for continuous glucose monitoring (CGM) or other wearable devices.

Potentiometry and ion-selective electrodes

Potentiometry is widely used clinically for the measurement of pH, P CO ₂ , and electrolytes (Na ⁺ , K ⁺ , Cl ⁻ , Ca ²⁺ , Mg ²⁺ , Li ⁺ ) in whole blood, serum, plasma, and urine, and as the basis for some biosensors for metabolites of clinical interest (urea or blood urea nitrogen [BUN]).

Basic concepts

Potentiometry is the measurement of an electrical potential difference between two electrodes (half-cells) in an electrochemical cell ( Fig. 17.1 ) when the cell current is zero (galvanic cell). Such a cell consists of two electrodes (electron and metallic conductors) that are connected by an electrolyte solution (ion conductor). An electrode , or half-cell , consists of a single metallic conductor that is in contact with an electrolyte solution. Ion conductors are composed of one or more phases that are either in direct contact with each other or separated by membranes permeable only to specific cations or anions (see Fig. 17.1 ). One of the electrolyte solutions is the unknown or test solution; this solution may be replaced by an appropriate reference solution for calibration purposes. By convention, the cell notation is shown so that the left electrode (M _L ) is the reference electrode, and the right electrode (M _R ) is the indicator (measuring) electrode (see Eq. 17.3 ).

FIGURE 17.1, Schematic of ion-selective membrane electrode–based potentiometric cell. Ag/AgCl , Silver-silver chloride; KCl , potassium chloride.

The electromotive force ( E or EMF) is defined as the maximum difference in potential between the two electrodes (right minus left) obtained when the cell current is zero. The cell potential is measured using a potentiometer , of which the common pH meter is a special type. The direct-reading potentiometer is a voltmeter that measures the potential across the cell (between the two electrodes); however, to obtain an accurate potential measurement, it is necessary that no current flow through the cell. This is accomplished by incorporating high resistance within the voltmeter (input impedance greater than 10 ¹² Ω). Modern direct-reading potentiometers are accurate and can be modified to provide direct digital displays or printouts.

Within any one conductive phase, the potential is constant as long as the current flow is zero. However, a potential difference arises between two different phases in contact with each other. The overall potential of an electrochemical cell is the sum of all potential gradients that exist between different phases of the cell. The potential of a single electrode with respect to the surrounding electrolyte and the absolute magnitude of the individual potential gradients between phases are unknown and cannot be measured. Only potential differences between two electrodes (half-cells) can be measured. Potential gradients can be classified as (1) redox potentials, (2) membrane potentials, or (3) diffusion potentials. Generally, it is possible to devise a cell in such a manner that all potential gradients except one are constant. This potential then can be related to the activity of a specific ion of interest (e.g., hydrogen [H ⁺ ] or Na ⁺ ).

Types of electrodes

Many different types of electrodes are used for potentiometric applications. They include redox, ion-selective membrane (glass and polymer), and P CO ₂ electrodes.

Redox electrodes

Redox potentials are the result of chemical equilibria involving electron transfer reactions:

Oxidized form (Ox) + ne ⁻ ⟷ Reduced form (Red)

where n = the number of electrons involved in the reaction. Any substance that accepts electrons is an oxidant (Ox), and any substance that donates electrons is a reductant (Red). The two forms, Ox and Red, represent a redox couple (conjugate redox pair). Usually, homogeneous redox processes take place only between two redox couples. In such cases, electrons are transferred from Red ₁ to an Ox ₂ . In this process, Red ₁ is oxidized to its conjugate Ox ₁ , whereas Ox ₂ is reduced to Red ₂ :

Red ₁ + Ox ₂ ⟷ Ox ₁ + Red ₂

In an electrochemical cell, electrons may be accepted from or donated to an inert metallic conductor (e.g., platinum [Pt]). A reduction process tends to charge the electrode positively (remove electrons), and an oxidation process tends to charge the electrode negatively (add electrons). By convention, a heterogeneous redox equilibrium ( Eq. 17.2) is represented by the cell:

ML | Red ₁ − Ox ₂ ⋮⋮ Ox ₂ − Red _{2 MR}

A positive potential ( E > 0) for this cell signifies that the cell reaction proceeds spontaneously from left to right; E < 0 signifies that the reaction proceeds from right to left, and E = 0 indicates that the two redox couples are at mutual equilibrium.

The electrode potential (reduction potential) for a redox couple is defined as the couple’s potential measured with respect to the standard H ₂ electrode, which is set equal to zero (see later discussion of the H ₂ electrode). This potential, by convention, is the EMF of a cell, where the standard H ₂ electrode is the reference electrode (left electrode) and the given half-cell is the indicator electrode (right electrode). The reduction potential for a given redox couple is shown by the Nernst equation:

E = E^{0} - \frac{N}{n} × \log \frac{a Red}{a Ox} = E^{0} - \frac{0.0592 V}{n} × \log \frac{a Red}{a Ox}

where E = electrode potential of the half-cell; E ⁰ = standard electrode potential when a _Red / a _Ox = 1; n = number of electrons involved in the reduction reaction; N = ( R × T × ln 10)/ F (the Nernst factor if n = 1); N = 0.0592 V if T = 298.15 K (25 °C); N = 0.0615 V if T = 310.15 K (37 °C); R = gas constant (= 8.31431 J × K ⁻¹ × mol ⁻¹ ); T = absolute temperature (in kelvins); F = Faraday constant (= 96,487 Coulombs × mol ⁻¹ ); ln 10 = natural logarithm of 10 = 2.303; a = activity; and a Red/ a Ox = product of mass action for the reduction reaction.

Redox electrodes currently in use include (1) inert metal electrodes immersed in solutions containing redox couples and (2) metal electrodes whose metal functions as a member of the redox couple.

Inert metal electrodes.

Platinum and gold (Au) are examples of inert metals used to record the redox potential of a redox couple dissolved in an electrolyte solution. The H ₂ electrode is a special redox electrode for pH measurement. It consists of a Pt or Au electrode that is electrolytically coated (platinized) with highly porous Pt (Pt black) to catalyze the electrode reaction:

H^{+} + e^{-} \leftrightarrow \frac{1}{2} H_{2}

The electrode potential is given by

E = E^{0} - N × \log \frac{{(f ​ H_{2})}^{1 / 2}}{a H^{+}}

E = E^{0} - N × [\log {(f ​ H_{2})}^{1 / 2} - \log a H^{+}]

where E ⁰ = 0 at all temperatures (by convention); f H ₂ = fugacity of H ₂ gas; a H ⁺ = activity of H ⁺ ions; and −log a H ⁺ = negative log of H ⁺ activity (p a H ⁺ or pH).

When the partial pressure of H ₂ (pH ₂ ) in the solution (and hence, f H ₂ ) is maintained constant by bubbling H ₂ through the solution, the potential is a linear function of log a H ⁺ (= −pH). In the standard H ₂ electrode , the electrolyte consists of an aqueous solution of HCl with a HCl equal to 1.000 (or c HCl = 1.2 mol/L) in equilibrium with a gas phase, and with f H ₂ equal to 1.000 (or pH ₂ = 101.3 kPa = 1 atm). The standard H ₂ electrode is also used as a reference electrode.

Metal electrodes participating in redox reactions.

The silver-silver chloride (Ag/AgCl) electrode is an example of a metal electrode that participates as a member of a redox couple. The Ag/AgCl electrode consists of an Ag wire or rod coated with AgCl _(solid) in contact with a Cl ⁻ solution of constant activity; this sets the half-cell potential. The Ag/AgCl electrode is itself considered a potentiometric electrode because its phase boundary potential is governed by an oxidation-reduction electron transfer equilibrium reaction that occurs at the surface of Ag:

AgCl _(solid) + e ⁻ ⟷ Ag° _(solid) + Cl ⁻

The Nernst equation for the reference half-cell potential of an Ag/AgCl reference electrode is written as follows:

E_{Ag / AgCl} = E_{Ag / AgCl}^{0} + \frac{R T}{n F} × ln \frac{a_{AgCl}}{a_{Ag} × a_{{Cl}^{-}}}

Because AgCl and Ag are both solids, their activities are equal to unity ( a _AgCl = a ⁰ _Ag = 1). Therefore from Eq. (17.9) , the half-cell potential is controlled by the activity of the Cl ⁻ ion in solution ( a _Cl −) contacting the electrode.

The Ag/AgCl electrode is used both as an internal reference element in potentiometric ion-selective electrodes (ISEs) and as an external reference electrode half-cell of constant potential, which is required to complete a potentiometric cell (see Fig. 17.1 ). In both cases, the Ag/AgCl electrode must be in equilibrium with a solution of constant Cl ⁻ ion activity.

The Ag/AgCl element of the external reference electrode half-cell is in contact with a high-concentration solution of a soluble Cl ⁻ salt. Saturated KCl is commonly used. A porous membrane or frit is frequently employed to separate the concentrated KCl from the sample solution. The frit serves both as a mechanical barrier to hold the concentrated electrolyte within the electrode and as a diffusional barrier to prevent proteins and other species in the sample from coming into contact with the internal Ag/AgCl element, which could poison and alter its potential. The interface between two dissimilar electrolytes (concentrated KCl/calibrator or sample) occurs within the frit and develops the liquid–liquid junction potential (Ej ), a source of error in potentiometric measurements. The difference in Ej between the calibrator and sample (residual liquid junction potential) is responsible for this error and can be minimized and usually neglected in practice if the compositions of the calibrating solutions are matched as closely as possible to the sample with respect to ionic content and ionic strength. An equitransferrant electrolyte at high concentration as the reference electrolyte further helps to minimize the residual liquid junction potential. KCl at a concentration 2 mol/L or more is preferred. Differences of approximately −2% in the measurement of sodium by ISEs have been demonstrated when the KCl concentration in the reference electrolyte is lowered from 3 to 0.5 mol/L. The magnitude of the residual liquid junction potential may also be estimated by the Henderson equation with sufficient knowledge of ionic activities, ionic charges, and ionic mobilities for each electrolyte on both sides of the junction and the temperature. Using this estimate, a correction to the overall cell potential may be applied.

The presence of erythrocytes in the sample may affect the magnitude of the residual liquid junction potential in a less predictable manner. For example, erythrocytes in blood of normal hematocrit are estimated to produce approximately 1.8 mmol/L positive error in the measurement of Na by ISEs when an open, unrestricted liquid–liquid junction is used. This bias may be minimized if a restrictive membrane or frit is used to modify the liquid–liquid junction.

The saturated calomel electrode is another example of a metal electrode that participates as a member of a redox couple. The calomel electrode consists of mercury (Hg) covered by a layer of relatively insoluble calomel (Hg ₂ Cl ₂ ) (or present as insoluble salt dispersed in the electrolyte), which is in contact with an electrolyte solution containing Cl ⁻ . The oxidation-reduction equilibrium reaction is as shown:

Hg ₂ Cl ₂ + 2 e ⁻ ⟷ 2Hg° + 2Cl ⁻

As with the Ag/AgCl electrode, the half-cell potential is controlled by the activity of Cl ⁻ ion contacting the electrode. Calomel electrodes are frequently used as reference electrodes for pH measurements using glass pH electrodes but are not commonly used in clinical instrumentation using electrochemical sensors.

Ion-selective electrodes

Membrane potentials are caused by the permeability of certain types of membranes to selected anions or cations. Such membranes are used to fabricate ISEs that selectively interact with a single ionic species. The potential produced at the membrane–sample solution interface is proportional to the logarithm of the ionic activity or concentration of the ion in question. Measurements with ISEs are simple, often rapid, nondestructive, and applicable to a wide range of concentrations.

The ion-selective membrane is the “heart” of an ISE because it controls the selectivity of the electrode. Ion-selective membranes are typically composed of glass, crystalline, or polymeric materials. The chemical composition of the membrane is designed to achieve an optimal permselectivity toward the ion of interest. In practice, other ions exhibit finite interaction with membrane sites and display some degree of interference for determination of an analyte ion. In clinical practice, if the interference exceeds an acceptable quantity, a correction is required.

The Nikolsky-Eisenman equation describes the selectivity of an ISE for the ion of interest over interfering ions:

E = E^{0} + [\frac{2.303 R T}{z_{i} F}] \log (a_{i} + \sum K_{i / j} a_{j}^{z_{i} / z_{j}})

where a _i = activity of the ion of interest; a _j = activity of the interfering ion; and K _i/j = selectivity coefficient for the primary ion over the interfering ion. Low values indicate good selectivity for the analyte i over the interfering ion j ; z _i = charge of the primary ion; and z _j = charge of the interfering ion.

All other terms are identical to those in the Nernst equation ( Eq. 17.4 ).

Various approaches may be used to determine the selectivity of an ISE for a primary ion over an interfering ion. ^, A straightforward approach is the separate solution method , in which the potential of an ISE is determined in solutions of the primary and interfering ions separately, but at equal ionic activities. The selectivity coefficient is then calculated as follows:

\log K_{i j} = \frac{E_{j} - E_{i}}{\frac{2.303 R T}{n F}} + (1 - \frac{z_{i}}{z_{j}}) \log a_{i}

An alternate approach to determine selectivity coefficient is the fixed interference method, in which the potential response of an ISE to the primary ion is determined in solutions of constant activity of the interfering ion. The potential values obtained are plotted versus the logarithm of the activity of the primary ion _ai . The intersection of the extrapolation of the linear portions of this plot indicates the value of _ai to be used to calculate K _i/j :

K_{i / j} = a_{i} / a_{j}^{z_{i} / z_{j}}

where all terms have the same definition as in Eq. (17.12) . Traditionally, the fixed interference method has been preferred because it more closely resembles a practical application of the sensor, in that primary and interfering ions are present simultaneously in solution and must compete for complexation sites in the ISE membrane.

Most ISEs used in clinical practice have sufficient selectivity and do not require correction for interfering ions. Oesch and colleagues have published required ISE selectivity coefficients for ions commonly measured in clinical chemistry over other ions found in blood. Table 17.1 shows required selectivity coefficients for the measurement of cations of interest in clinical chemistry over potentially interfering cations, assuming an acceptable maximum interference of 1% for the ion of interest.

TABLE 17.1

Required Selectivities for Cation-Selective Ion-Selective Electrodes for Whole Blood, Plasma, and Serum Measurements

Primary Ion (i)	REQUIRED SELECTIVITY COEFFICIENT (LOGK _I/J ) FOR INTERFERING CATION (J)
Primary Ion (i)	H ⁺	Li ⁺	Na ⁺	K ⁺	Mg ²⁺	Ca ²⁺
H ⁺	—	−6.5	−8.5	−7.0	−7.7	−7.7
Li ⁺ ^a	2.1	—	−4.3	−2.8	−3.5	−3.6
Na ⁺	4.4	−0.1	—	−0.6	−1.2	−1.3
K ⁺	2.8	−1.7	−3.6	—	−2.8	−2.9
Mg ²⁺	8.9	0.1	−3.9	−0.9	—	−2.4
Ca ²⁺	9.3	0.4	−3.6	−0.6	−1.9	—

Ca , Calcium; H , hydrogen; K , potassium; Li , lithium; Mg , magnesium; Na , sodium.

a Assumes a therapeutic range for Li ⁺ between 0.7 and 1.5 mmol/L.

Glass membrane and polymer membrane electrodes are two types of ISEs that are commonly used in clinical chemistry applications.

The glass electrode.

Glass membrane electrodes are used to measure pH and Na ⁺ , and as an internal transducer for P CO ₂ sensors. The H ⁺ response of thin glass membranes was first demonstrated in 1906 by Cremer. In the 1930s, practical application of this phenomenon for measurement of acidity in lemon juice was made possible by the invention of the pH meter by Beckman. Glass electrode membranes are formulated from melts of silicon and/or aluminum oxide (Al ₂ O ₃ ) mixed with oxides of alkaline earth or alkali metal cations. By varying the glass composition, electrodes with selectivity for H ⁺ , Na ⁺ , K ⁺ , Li ⁺ , rubidium (Rb ⁺ ), cesium (Cs ⁺ ), Ag ⁺ , thallium (Tl ⁺ ), and ammonium (NH ₄ ⁺ ) have been demonstrated. However, glass electrodes for H ⁺ and Na ⁺ are currently the only types with sufficient selectivity over interfering ions to allow practical application in clinical chemistry. A typical formulation for H ⁺ selective glass is 72% silicon dioxide (SiO ₂ ), 22% Na ₂ O, and 6% calcium oxide (CaO), which has a selectivity order of H ⁺ ⪢> Na ⁺ > K ⁺ . This glass membrane has sufficient selectivity for H ⁺ over Na ⁺ to allow error-free measurements of pH in the range of 7.0 to 8.0 ([H ⁺ ] = 10 ⁻⁷ to 10 ⁻⁸ mol/L) in the presence of > 0.1 mol/L Na ⁺ . Glass pH electrodes with selectivity coefficients (K _H/Na ) over Na ⁺ of 10 ⁻⁷ and better have been realized. By altering the formulation of the glass membrane slightly to 71% SiO ₂ , 11% Na ₂ O, and 18% Al ₂ O ₃ , its selectivity order becomes H ⁺ > Na ⁺ > K ⁺ . Thus the preference of the glass membrane for H ⁺ over Na ⁺ is greatly reduced, resulting in a practical sensor for Na ⁺ at pH values typically found in blood.

Polymer membrane electrodes.

Polymer membrane ISEs are used for monitoring pH and for measuring electrolytes, including K ⁺ , Na ⁺ , Cl ⁻ , Ca ²⁺ , Li ⁺ , Mg ²⁺ , and carbonate (CO ₃ ²⁻ ) (for total CO ₂ measurements). They are the predominant class of potentiometric electrodes used in modern clinical analysis instruments.

Mechanisms of response of these ISEs fall into three categories: (1) charged, dissociated ion exchanger; (2) charged associated carrier; and (3) neutral ion carrier (ionophore). ^, An early charged-associated, ion-exchanger type ISE for Ca ²⁺ was developed and commercialized for clinical application in the 1960s based on the Ca ²⁺ -selective, ion-exchange/complexation properties of 2-ethylhexyl phosphoric acid dissolved in dioctyl phenyl phosphonate. A porous membrane was impregnated with this cocktail and mounted at the end of an electrode body. This type of sensor was referred to as the “liquid membrane” ISE. Later, a method was devised in which these ingredients could be cast into a plasticized poly(vinyl chloride) (PVC) membrane that was more rugged and convenient to use. This same approach is still used today to formulate PVC-based ISEs for clinical use.

A major breakthrough in the development and routine application of PVC-type ISEs was the discovery by Simon and colleagues that the neutral antibiotic valinomycin could be incorporated into organic liquid membranes (and later plasticized PVC membranes), resulting in a sensor with high selectivity for K ⁺ over Na ⁺ (K _K/Na = 2.5 × 10 ⁻⁴ ). The K ⁺ ISE based on valinomycin was the first example of a neutral carrier ISE and is used extensively today for the routine measurement of K ⁺ in blood. Fig. 17.2 shows the response of the valinomycin-based K ⁺ ISE in the presence of physiologic concentrations of Na ⁺ , Ca ²⁺ , and Mg ²⁺ . The wide linear range and excellent selectivity of this ISE over three orders of magnitude makes it suitable for the measurement of K ⁺ in blood and urine. The K ⁺ range in blood is only a small portion of the electrode linear range and is spanned by a total ΔEMF of approximately 9 mV. Interference from other cations, which is seen as deviation from linearity, is not apparent at K ⁺ activities more than 10 ⁻⁴ mol/L. Other, less selective polymer-based ISEs (e.g., for the measurement of Mg ²⁺ and Li ⁺ ) are subject to interference from Ca ²⁺ and Na ⁺ and Na ⁺ , respectively, requiring simultaneous determination and correction for the presence of significant concentrations of these interfering ions. ^,

FIGURE 17.2, Typical electromotive force (EMF) response of potassium (K1) selective membrane electrode to changes in activity of K1 in the sample solution. Bracketed interval represents the normal reference interval of K + concentration in blood. Ag/AgCl , Silver-silver chloride.

Studies regarding the relationship between molecular structure and ionic selectivity have resulted in the development of polymer-based ISEs using a number of naturally occurring and synthetic ionophores, with sufficient selectivity for application in clinical analysis. The chemical structures of several of these neutral ionophores are illustrated in Fig. 17.3 .

FIGURE 17.3, Structures of common ionophores used to fabricate polymer membrane–type ion-selective electrodes for clinical analysis.

Dissociated anion exchanger-based electrodes using lipophilic quaternary ammonium salts as active membrane components are still used commercially for the determination of Cl ⁻ in whole blood, serum, and plasma, despite some limitations. Selectivity for this type of ISE is controlled by extraction of the ion into the organic membrane phase and is a function of the lipophilic character of the ion (because, unlike the carriers described earlier, no direct binding interaction occurs between the exchanger site and the anion in the membrane phase). Thus the selectivity order for a Cl ⁻ ISE based on an anion exchanger is fixed as lipophilic anion R ⁻ > perchlorate (ClO ₄ ⁻ ) > iodide (I ⁻ ) > nitrate (NO ₃ ⁻ ) > bromide (Br ⁻ ) > Cl ⁻ > fluoride (F ⁻ ), where R ⁻ represents anions with greater lipophilic character than ClO ₄ ⁻ . Application of the Cl ⁻ ion-exchange electrode is therefore limited to samples without significant concentrations of anions more lipophilic than Cl ⁻ . For example, blood samples containing salicylate or thiocyanate will produce positive interference for the measurement of Cl ⁻ . Repeated exposure of the electrode to the anticoagulant heparin will lead to loss of electrode sensitivity toward Cl ⁻ because of the extraction of negatively charged heparin into the membrane. This extraction process has been used successfully to devise a method to detect heparin concentrations in blood by potentiometry and to develop a simple potentiometric technique to screen for the presence of toxic, high-charge density polyanion contaminants (e.g., oversulfated chondroitin sulfate) in biomedical-grade heparin preparations.

High selectivity for the CO ₃ ²⁻ anion can be achieved using a neutral carrier ionophore that possesses trifluoroacetophenone groups doped within a polymeric membrane. ^, Such ionophores form negatively charged adducts with CO ₃ ²⁻ anions, and the resulting electrodes have proven useful in commercial instruments for determination of total CO ₂ in serum and/or plasma, after dilution of the blood to a pH value in the range of 8.5 to 9.0, where a significant fraction of total CO ₂ will exist as CO ₃ ²⁻ anions.

A typical formulation of a PVC membrane ISE used in clinical instrumentation consists of the following in % by weight (wt%):

1 to 3 wt% ionophore;
≈ 64 wt% plasticizer;
≈ 30 wt% PVC; and <1 wt% additives.

The plasticizer is crucial in controlling the polarity of the membrane, and thus along with the ionophore, plays a pivotal role in determining the selectivity of the membrane toward the ion of interest. A large lipophilic anion (e.g., tetraphenylborate derivative) is often included as an additive for preparation of cation-selective ISE membranes. This anion serves as a counter-anion for the cation of interest as it is extracted into the membrane phase, forming a positively charged complex with the neutral ionophore. However, it is the ratio of the bound-to-unbound ionophore sites at the membrane surface that determines the magnitude of the phase boundary potential generated by the ISE membrane. Thus the selective response to the activity of the ion of interest is an interfacial property of the given ISE membrane.

Studies have demonstrated that the ultimate detection limits of polymer membrane–type ISEs are controlled in part by the leakage of analyte ions from the internal solution to the outer surface of the membrane and into the sample phase in close contact with the membrane. Hence, lower limits of detection can be achieved by decreasing the concentration of the primary analyte ion within the internal solution of the electrode. Furthermore, this leakage of analyte ions, coupled with an ion exchange process at the membrane sample interface when the selectivity of the membrane over other ions is assessed, can often yield a measured potentiometric selectivity coefficient that underestimates the true selectivity of the membrane. To determine “unbiased” selectivity coefficients by the separate solution method, the membrane should not be exposed to the analyte ion for extended periods of time, and the concentration of the analyte ion in the internal solution should be low. To avoid leakage of primary ions from the inner solution of conventional polymer membrane ISEs, new, more stable designs for solid-state ion sensors have been suggested, in which the ionophore-doped, polymer ion–sensing membrane (based on a more water-repellent poly[methylmethacrylate]/poly[decylmethacrylate copolymer]) is coated onto a conductive poly(3-octylethiophene 2,5-diyl) (POT) polymer layer on the surface of an underlying Au electrode.

An interesting application of Na ⁺ selective polymer (or glass) membrane electrodes is seen in the determination of whole blood hematrocrit. Because intracellular Na ⁺ concentrations are much lower than those in the plasma phase, the change in Na ⁺ concentration (dilution) measured potentiometrically before and after erythrocyte lysis can be used to assess the hematocrit of the blood sample. This approach can be coupled with simultaneous measurement of changes in K ⁺ ion concentration as determined with a valinomycin-based polymer membrane ISE to quantify the concentration of K ⁺ ions within red blood cells.

Recently, a reversible polymer membrane ISE for measurement of polyions was described. The sensor consists of a highly lipophilic electrolyte (tetradodecylammonium-dinonylnaphthalenesulfonate), which is free of intrinsic ion-exchange properties, added to a plasticized PVC membrane at high concentration (approximately 10 wt%). Application of a cathodic current pulse across the membrane results in reversible extraction and potentiometric response to the polycation protamine. Protamine is a polypeptide administered to neutralize heparin activity. A response to heparin in whole blood was demonstrated via protamine titration. Later, it was found that by changing the lipophilic cation in the membrane to tridodecylmethylammonium, application of an anodic current pulse resulted in a direct potentiometric response to heparin. Testing of other polyanions in addition to heparin led to the conclusion that the magnitude of the potentiometric response is a function of charge density of the polyanion.

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