Síntesis de complejos ortometalados de paladio con bases de Schiff

When chemical energy is converted to an electri-cal current (a flow of electrons) in a galvanic cell, the term electrochemistry is used. Electrochemi-cal reactions are characterized by a loss of elec-trons (oxidation) at the positive pole (anode) and a simultaneous gain of electrons (reduction) at the negative pole (cathode). The galvanic cell is made up of two parts called half-cells, each containing a metal in a solution of one of its salts. These meth-ods involve the measurement of electrical signals associated with chemical systems that are within an electrochemical cell.

Electroanalytical chemistry uses electrochemis-try for analysis purposes. In the clinical laboratory, electroanalytical methods are used to measure ions, drugs, hormones, metals, and gases. Meth-ods are available for the rapid analysis of ana-lytes present in relatively high concentrations in blood and urine, such as blood electrolytes (Na⁺, K⁺, Cl, HCO₃⁻), and other analytes present in very low concentrations, such as heavy metals and drug metabolites. There are three general electrochemical techniques used in the clinical laboratory: potentiometric, voltammetric, and coulometric. This section discusses the electro-chemical methods of potentiometry, coulometry, and electrophoresis.

Potentiometry

Potentiometry measures the potential of an trode compared with the potential of another elec-trode. The method is based on the measurement of a voltage potential difference between two elec-trodes immersed in a solution under zero-current conditions. This difference in voltage between the two electrodes is usually measured on a pH or voltage meter. One electrode is called the indicator electrode; the other is the reference electrode. The reference electrode is an electrochemical half-cell that is used as a fixed reference for the cell potential measurements. One of the most common reference electrodes used for potentiometry is the silver or silver chloride electrode. The indicator electrode is the main component of potentiometric techniques.

It is important that the indicator electrode be able to respond selectively to analyte species. The most common indicator electrode used in clinical chem-istry is the ion-selective electrode (ISE).

The use of ion-selective electrodes (ISEs) is based on the measurement of a potential that develops across a selective membrane. The elec-trochemical cell response is based on an interac-tion between the membrane and the analyte being measured that alters the potential across the mem-brane. The specificity of the membrane interaction for the analyte determines the selectivity of the potential response to an analyte.

Electrodes and Ionic Concentration

Potentiometric methods of analysis involve the direct measurement of electrical potential caused by the activity of free ions. ISEs are designed to be sensitive toward individual ions. An ISE uni-versally used in the clinical laboratory is the pH electrode. Specialized probes such as ISEs can measure concentrations of ionic species other than hydrogen ions [H⁺], including fluoride, chloride, ammonia, sodium, potassium, calcium, sulfide, and nitrate ions.

An electrode is an electronic conductor in con-tact with an ionic conductor, the electrolyte. Passive (inert) electrodes act as electron donors or elec-tron acceptors; active (participating) electrodes act as ion donors or ion acceptors. The electrode reac-tion is an electrochemical process in which charge transfer takes place at the interface between the electrode and the electrolyte.

By means of its potential, an indicator elec-trode shows the activity of an ion in a solution.

The relationship between the potential and the activity is given by the Nernst equation (see later).

The potential between an electrode and a solution cannot be directly measured; a reference electrode is needed. The reference electrode should have

a known, or at least a constant, potential value under the prevailing experimental conditions.

The most essential component of a pH elec-trode is a special sensitive glass membrane that permits the passage of hydrogen ions but no other ionic species. When the electrode is immersed in a test solution containing hydrogen ions, the exter-nal ions diffuse through the membrane until an equilibrium is reached between the external and internal concentrations. Thus there is a buildup of charge on the inside of the membrane that is pro-portional to the number of hydrogen ions in the external solution.

Because of the need for equilibrium conditions, there is very little current flow. Therefore, this potential difference between electrode and solu-tion can only be measured relative to a separate and stable reference system that is also in contact with the test solution but is unaffected by it. A sensitive, high-impedance millivolt meter or digi-tal measuring system must be used to measure this potential difference accurately.

In fact, the potential difference developed across the membrane is directly proportional to the loga-rithm of the ionic concentration in the external solution. To determine the pH of an unknown solu-tion, it is only necessary to measure the potential difference in two standard solutions of known pH, construct a straight-line calibration graph by plot-ting millivolts versus pH = (log [H⁺]), then read off the unknown pH from the measured voltage.

To measure the electrode potential developed at the ion-selective membrane, the ISE/pH electrode must be immersed in the test solution together with a separate reference system, and the two must be connected by a millivolt measuring system. At equilibrium, the electrons added or removed from the solution by the ISE membrane (depending on whether it is cation or anion sensitive) are bal-anced by an equal and opposite charge at the refer-ence interface. This causes a positive or negative deviation from the original stable reference voltage that is registered on the external measuring system.

The relationship between the ionic concentra-tion (activity) and the electrode potential is given by the Nernst equation, as follows:

E= E⁰+ (2.303RT/nF)×Log(A) where:

E = Total potential (in mV) developed between the sensing and reference electrodes

E⁰ = Constant that is characteristic of the par-ticular ISE/reference pair (sum of all the liquid junction potentials in the electro-chemical cell; see later)

2.303 = Conversion factor from natural to base-10 logarithm

R = Gas constant (8.314 joules/degree/mole)

T = Absolute temperature n = Charge on the ion (with sign) F = Faraday constant (96,500 coulombs) Log(A) = Logarithm of the activity of the

measured ion

Note that 2.303RT/nF is the slope of the line (from the straight-line plot of E versus log[A], which is the basis of ISE calibration graphs). This is an important diagnostic characteristic of the electrode; generally the slope gets lower as the electrode gets old or contaminated, and the lower the slope, the higher the errors on the sample measurements.

For practical use in measuring pH, it is not normally necessary for the operator to construct a calibration graph and interpolate the results for unknown samples. Most pH electrodes are con-nected directly to a special pH meter, which performs the calibration automatically. This determines the slope mathematically and calculates the unknown pH value for immediate display on the meter.

These basic principles are exactly the same for all ISEs, so it would appear that all can be used as easily and rapidly as the pH electrode—simply by calibrating the equipment by measuring two known solutions, then immersing the electrodes in any test solution and reading the answer directly from a meter. Some other ions can be measured in this simple way, but this is not the case for most ions.

pH Electrodes and Meters

The pH measurement was originally used by the Danish biochemist Soren Sorensen to represent the hydrogen ion concentration, expressed in equivalents per liter, of an aqueous solution: pH = log[H⁺]. In expressions of this type, enclosure of a chemical symbol within square brackets denotes that the concentration of the symbolized species is the quantity being considered.

The first commercially successful electronic pH meter was invented by Dr. Arnold Beckman in 1934. This instrument was the forerunner of mod-ern electrochemical instrumentation and became an indispensable tool in analytical chemistry.

Beginning in the 1950s, electrodes were also devel-oped for other ions, such as F, Na⁺, K⁺, and Ag⁺. The pH meter and ISE have now become indis-pensable scientific tools. Today, Beckman Coulter pH meters provide precise pH and concentration (ISE) measurements in handheld, bench-top, and high-performance meters for research, pharmaceu-tical, chemical, and environmental applications.

The term pH refers to the concentration of hydrogen ions ([H⁺], also called protons) in a solu-tion. For aqueous solutions, the scale ranges from 0 to 14, with pure water in the middle at 7. The more

acid a solution is, the lower the pH reading (0-6.9), and alkaline solutions come in at the high end of the scale (7.1-14).

Paper test strips (e.g., urine dipsticks) are good for measuring approximate pH values, but chemi-cal laboratories require more exact measurements.

A pH meter is a boxy-looking instrument attached to a glass or plastic tube called a probe. Handheld pH meters have a probe directly attached to the instrument body. The probe has a glass bulb on one end and an electrical wire on the other. The wire sends data to the instrument when the glass bulb is dipped into a sample solution.

The pH meter measures H⁺ concentration by sensing differences in the electric charges inside and outside the probe. The glass bulb is made from silica (SiO₂) that contains added metal ions. Most of the oxygen atoms in the glass are surrounded by silicon and metal atoms. However, the oxygen atoms on the inside and outside surfaces of the bulb are not completely surrounded, and they can

“grab” positively charged ions from the solution.

When the bulb is dipped into an acid solution, H⁺ ions bond with the outside surface of the glass bulb, forming electrically neutral Si–OH groups.

The Si–O– groups on the inside surface are in con-tact with a reference solution. The difference in electrical charge between the two surfaces creates an electrical potential, or voltage, and this causes an electric current to flow through the wire at the other end of the probe.

Alkaline solutions have low concentrations of H⁺ ions and higher concentrations of negative ions such as OH. The excess negative charges are balanced with positively charged metal ions such as Na⁺, and these positive ions hover close to the surface of the bulb rather than binding to the Si–

O– groups. This sets up a different sort of charge separation, and the resulting electrical signal reg-isters a high pH.

Types of Meters ANALOG METERS

The earliest type of pH meters were simple analog devices with a resolution of only 1 or 2 mV. The original meters were calibrated in mil-livolts, and the corresponding pH value was read from a calibration graph. Because pH electrodes are reasonably uniform and reproducible instru-ments, it was discovered that it is not necessary to have a unique calibration graph for each elec-trode. In this case, the meters can be calibrated directly in pH units by the manufacturer and can simply be recalibrated each time they are used (to compensate for temperature changes or slight differences in electrode response) by immersing the electrode in just one pH buffer solution and

adjusting the meter output to give the correct reading. This type of meter is simple and quick to use and is perfectly adequate for many pH mea-surements because it requires a change of more than 5 mV to change the pH value by more than 0.1 pH units.

Simple precalibrated analog meters are not appropriate for ISE measurements.

DIGITAL METERS

A major advance was made when digital meters were introduced with a resolution of 0.1 or even 0.01 mV. This enabled the analyst to measure and read the voltage with much greater accuracy and meant that the stability and repro-ducibility of the electrode response became the main limiting factors in determining the accu-racy and precision.

SELF-CALIBRATING, DIRECT-READING ION METERS

The next major advance occurred when micro-processors were introduced. These contained simple programs to calculate the slope and inter-cept from the calibration data, which were then used to calculate the sample concentration from the millivolt reading in the sample. The analyst can simply enter the concentrations of the stan-dards and measure the millivolts, then immerse the electrodes in the sample and read the sample concentration directly from the meter. These meters are often confusing to operate, with small keypads and multifunction switches, and they are not suitable for working in the nonlinear range of the electrodes, using different slopes for different parts of the calibration range, or measuring more than one ion at a time. It is often difficult for the analyst to assess the quality of the calibration or detect errors in data entry, and it is still necessary for the results to be transferred manually to a per-manent record.

Coulometry

Coulometry measures the amount of current pass-ing between two electrodes in an electrochemi-cal cell. The principle of coulometry involves the application of a constant current to generate a titrating agent; the time required to titrate a sam-ple at constant current is measured and is related to the amount of analyte in the sample. The amount of current is directly proportional to the amount of substance produced or consumed by the elec-trode. Clinical applications of coulometry include the FreeStyle Connect blood glucose monitoring system (Abbott Labs) in the point-of-care set-ting (hospitals and medical clinics) and an older application for the measurement of chloride ions in serum, plasma, urine, and other body fluids.

Electrophoresis

Electrophoresis is the migration of charged solutes or particles in an electrical field. When charged particles are made to move, differences in molecular structure can be seen because different molecules have different velocities in an electrical field. The assay using electrophoresis involves the movement of charged particles when an external electric cur-rent is produced in a liquid environment.

The electrical field is applied to the solution through oppositely charged electrodes placed in the solution. Specific ions then travel through the solution toward the electrode of the oppo-site charge. Cations (positively charged particles) move toward the negatively charged electrode (cathode), and anions (negatively charged parti-cles) move toward the positively charged electrode (anode) (Fig. 6-15).

Electrophoresis is a technique for separa-tion and purificasepara-tion of ions, proteins, and other molecules of biochemical interest. It is used

F I G U R E   6 - 1 5  Application of electrical field to solution of ions makes ions move. (From Kaplan LA, Pesce AJ: Clinical chemistry: theory, analysis, correlation, ed 5, St Louis, 2010, Mosby.)

separate serum proteins. The equipment needed for electrophoresis generally consists of a sample applicator; a solid medium (e.g., agar gel); a buffer system; an electrophoresis chamber, which houses the solid medium and the sample; electrodes and wicks; a timer; and a power supply. Additional supplies might be stains for proteins or other sub-stances being assayed and reagents used to remove the stains and to transform the solid media into a stable carrier for further densitometry studies or for preservation needs, depending on the require-ments of the laboratory.

Serum proteins including immunoglobulins are commonly separated by electrophoresis. Serum electrophoresis results in the separation of pro-teins into five fractions using cellulose acetate as a support medium (Fig. 6-16). The immunologic applications of electrophoresis include identifica-tion of monoclonal proteins in either serum or urine, immunoelectrophoresis, and various blot-ting techniques.

Immunoelectrophoresis

Immunoelectrophoresis (IEP) involves the elec-trophoresis of serum or urine followed by immu-nodiffusion. The size and position of precipitin

bands provide the same type of information regard-ing equivalence or antibody excess as the double immunodiffusion method. Proteins are differenti-ated not only by their electrophoretic mobility but by their diffusion coefficient and antibody specificity.

IEP is a combination of the techniques of elec-trophoresis and double immunodiffusion (Fig.

6-17) and consists of two phases: electrophoresis and diffusion. In the first phase, serum is placed in an appropriate medium (e.g., cellulose acetate or agarose), then electrophoresed to separate its constituents according to electrophoretic mobili-ties: albumin, alpha₁-, alpha₂-, beta-, and gam-maglobulin fractions. In the second phase after electrophoresis, the fractions are allowed to act as antigens and interact with their correspond-ing antibodies. When a favorable antigen-to-antibody ratio exists (equivalence point), the antigen-antibody complex becomes visible as pre-cipitin lines or bands. Diffusion is halted by rins-ing the plate in 0.85% saline. Unbound protein is washed from the agarose with saline, and the antigen-antibody precipitin arcs are stained with a protein-sensitive stain.

Each line represents one specific protein. Pro-teins are thus differentiated not only by their electrophoretic mobility but by their diffusion coefficient and antibody specificity. Antibody dif-fuses as a uniform band parallel to the antibody trough. If the proteins are homogeneous, the anti-gen diffuses in a circle, and the antianti-gen-antibody precipitation line resembles a segment or arc of a circle. If the antigen is heterogeneous, the antigen-antibody line assumes an elliptical shape. One arc of precipitation forms for each constituent in the antigen mixture. This technique can be used to resolve the protein of normal serum into 25 to 40 distinct precipitation bands. The exact number depends on the strength and specificity of the anti-serum used.

NORMAL APPEARANCE OF PRECIPITIN BANDS

Immunoprecipitation bands should be of normal curvature, symmetry, length, position, intensity, and distance from the antigen well and antibody trough. In normal serum, IgG, IgA, and IgM are present in sufficient concentrations of 10 mg/mL, 2 mg/mL, and 1 mg/mL, respectively, to produce precipitin lines. The normal concentra-tions of IgD and IgE are too low to be detected by IEP.

CLINICAL APPLICATIONS OF IEP

IEP is most commonly used to determine quali-tatively the elevation or deficiency of specific classes of immunoglobulins. It is a reliable and

Concentration Albumin α1–globulin α2–globulin β–globulin γ–globulins

Anode side

Distance moved

Cathode side F I G U R E   6 - 1 6  Example of effect of disease (hepatic  cirrhosis)  on  serum  protein  electrophoretic  pattern.  Upper  profile,  distribution  characteristic  of  healthy  people.  (From Kaplan LA, Pesce AJ: Clinical chemistry: theory, analysis, cor-relation, ed 5, St Louis, 2010, Mosby.)

accurate method for detecting both structural abnormalities and concentration changes in pro-teins. The most common application of IEP is in the diagnosis of a monoclonal gammopathy, a con-dition in which a single clone of plasma cells pro-duces elevated levels of a single class and type of Ig. The most important application of IEP of urine is the demonstration of Bence Jones (BJ) protein, a diagnostic sign of multiple myeloma.

Immunofixation Electrophoresis

Immunofixation electrophoresis (IFE), or simply immunofixation, has replaced IEP in the evalua-tion of monoclonal gammopathies because of its rapidity and ease of interpretation. It is a two-stage procedure using agarose gel protein electropho-resis in the first stage and immunoprecipitation in the second. The test specimen may be serum, urine, cerebrospinal fluid, or other body fluids. The primary use of IFE in clinical laboratories is for the characterization of monoclonal Igs.

Capillary Electrophoresis

In capillary electrophoresis, the classic separation techniques of zone electrophoresis, isotachopho-resis, isoelectric focusing, and gel electrophoresis are performed in small-bore (10 to 100 μm) fused-silica capillary tubes from 20 to 200 cm in length.

In capillary electrophoresis, the classic separation techniques of zone electrophoresis, isotachopho-resis, isoelectric focusing, and gel electrophoresis are performed in small-bore (10 to 100 μm) fused-silica capillary tubes from 20 to 200 cm in length.