Iodine and Iodine-Containing Compounds



Iodine and Iodine-Containing Compounds


Paul L. Bigliardi

Maren Eggers

Marc Cataldo

Ram Kapil

Manjunath Shet

Michael K. Pugsley



Iodine, a nonmetallic, essential element discovered in 1812 by the French scientist Bernard Courtois, was named by Joseph Louis Gay-Lussac in 1814 after the Greek word meaning “violet,” which is the color of iodine vapor. It is not found in elemental form in nature but occurs sparingly in the form of iodides primarily in seawater, other brackish waters, in sea deposits, in Chilean saltpeter, and nitrate-bearing earth. Besides the single stable isotope, which is iodine 127 (127I), there are more than 30 artificial isotopes with half-lives that range between 0.2 second and 1.57 × 107 years. Some of these isotopes have resulted from the dangerous fallout following nuclear accidents or attacks, whereas others are used in nuclear medicine (mainly iodine 131 [131I] [8.04 d] and iodine 123 [123I] [13.2 h]).

The first use of iodine in medical practice was as a remedy for bronchocele (a segment of a bronchus that is closed because of congestion).1 Soon afterward, Lugol2 treated scrofuloderma (tuberculous lesions of the skin) with an iodine/iodide solution bearing his name that is still in use today (known as strong iodine solution).3 Iodine was officially recognized by the United States Pharmacopeia in 1830, specifically as tinctura iodini (tincture of iodine). Clinicians and microbiologists described a great quantity of experimental data and numerous clinical applications, which can be found in numerous reviews.4,5,6,7,8,9,10

Despite the successes that have been achieved with iodine, it was ascertained early that it also possesses properties, such as unpleasant odor, staining the skin, or reactivity with iron and other metals, that could make it unsuitable for some practical applications.11 Furthermore, iodine solutions are not stable (under certain circumstances); it is well known that they can irritate tissue and may be a poison at certain doses or concentrations. The adverse side effects associated with acute iodine toxicity include vomiting, diarrhea, metabolic acidosis, seizure, stupor, delirium, and collapse.12 High levels of iodine are also linked to disruption of thyroid hormone metabolism, the thyroid-pituitary axis, and the compensatory mechanisms that exist to protect such metabolism against low or high levels of iodine intake12 and associated with pain when administered as an iodine-alcohol solution on open wounds. Additionally, some risk of allergic reactions in the past 100 years led to the development of many iodine-based preparations that were designed to reduce these unwanted side effects without a significant loss of antimicrobial efficiency. This problem can be approached by using the appropriate concentration for different applications and developing formulations allowing controlled and regular release of active iodine for optimal antimicrobial activity. The iodophors (eg, povidone-iodine) were the first such compounds largely to achieve this goal. More recently, graphene-iodine nanocomposites have been developed, which can inhibit bacterial growth and have been found to be potentially biocompatible with human cells with very low cytotoxic effects for human cell cultures, although they have not yet been proven effective in clinical practice.13


CHEMISTRY

Iodine is the halogen with the highest atomic weight (126.9 amu) of the common halogens. It is the weakest oxidizing agent of the group unlike its iodide anion, I–, which is the strongest reducing agent of the halogens and forms grayish-black metallic scales that melt at 113.5°C to a black, mobile liquid. Iodine boils at 184.4°C (at atmospheric pressure) to produce the characteristic violetcolored vapor. Despite the high boiling point, it has an appreciable vapor pressure at room temperature (22°C) and sublimes before it melts if it is not heated too fast and with too high degree of heat.

Elemental iodine is only slightly soluble in water (ie, 1 g dissolves in 3450 mL at 20°C), forming a brown solution. Its solubility in water is increased with the addition of alkali iodides by which triiodide and higher polyiodides
are formed (see equations [3], [5], and [6] in the following text). In polar organic solvents (such as alcohols, ketones, carbonic acids), iodine forms a brown solution whose color is explained by the formation of an electrostatic attraction between iodine and the solvent molecule that provides a stabilizing force for the molecular complex. This complex is known as an electron-donor-acceptor complex or charge transfer complex. In nonpolar solvents (such as carbon tetrachloride [CCl4], benzene, hydrocarbons), iodine dissolves to a violet color that is explained by the presence of iodine in the free state (I2), as in the gas phase.


Properties of Disinfecting Iodine Solutions

Iodine-based disinfectants can be divided into three main groups according to the solvent and substances interfering (by complexing) with the iodine species: (1) pure aqueous solutions, (2) alcoholic solutions, and (3) iodophoric preparations or combinations of two or three components. Each of them exhibits intrinsic differences in their chemical, antimicrobial, and toxic/irritant properties.

A reliable understanding of the processes occurring at disinfection, which includes not only killing of microorganisms but also interactions with the material or host tissue (eg, innate surfaces, living tissue, body fluids), is essential based on knowledge of the microbial mixture and the environment, the formulation, and the elimination of the active ingredient over time.


Aqueous Solution

For the iodine-water system, nine different equilibria (equations [1], [2], [3], [4], [5], [6], [7], [8] and [9]) are specified14 that produce at least 10 iodine species: I–, I2, I3–, I5–, I62-, HOI, OI–, HI2O–, I2O2-, H2OI+, and IO3–.










As noted earlier, this is a system of appreciable complexity, with several associated equilibria governed mainly by H+ and I– ions, which implies that pH and additional iodide influence equilibrium concentrations. Another important feature is the reaction rate; whereas the reactions in equations (1), (2), (3), (4), (5), (6), (7) and (8) are thought to occur instantaneously, disproportionation to iodate (equation [9]) proceeds comparatively slowly, with a rate highly influenced by pH and additional iodide, as can easily be deduced from the rate law15:


where [ ] means equilibrium concentration of the bracketed species.

Because the reactions in equations (1), (2), (3), (4), (5), (6), (7), (8) and (9) are well studied, with focus on the fate of radioiodine species that emerge in the course of nuclear accidents,14 a calculation may represent the easiest way to approach equilibrium concentrations. It is difficult to determine all iodine species experimentally because for some species, limited analytical methods are available.16 In natural water samples, concentrations of I– and IO3– have been recently determined with a two-column high-performance liquid chromatography system; I– was detected with amperometry, and IO3– was detected with spectrophotometry.17

Several investigations into the equilibrium concentrations of aqueous iodine solutions conducted differ mainly regarding the equilibria considered and the regulating parameters, pH, and additional iodide. One study investigated all the immediately established equilibria (equations [1], [2], [3], [4], [5], [6], [7] and [8]) and both regulating parameters.18 It dealt with fresh iodine solutions not altered by disproportionation (iodate formation) and provided results about the equilibrium concentrations of the species I–, I2, I3–, I5–, I62-, HOI, OI–, HI2O–, IO2–, and H2OI–. The results for selected variations of total iodine and iodide, Lugol solution and its dilutions, and the rates of iodate formation (Figures 16.1, 16.2, 16.3 and 16.4) are the basis for most of the following conclusions:



  • Additional iodide and pH have a marked influence on the individual equilibrium concentrations, and consequently, conditions can be indicated in which the number of species of importance is substantially reduced. In the most common case, only I–, I2, and I3– play a role for iodine in the presence of additional iodide at pH 6 or less (Figure 16.2).


  • In such a system, HOI and all species derived from it (OI–, HI2O–, I2O2-, H2OI+) and the higher polyiodides can be neglected without any noticeable loss of precision. In other words, in this case, only the triiodide equilibrium (equation [3]) is relevant and it is not influenced by pH. This has two consequences: (1) the distribution of the three species is the same at pH 6 or less, and (2) a sufficiently precise evaluation
    can be based solely on the determination of [I2] (eg, potentiometrically19 or by dialysis20) and [I–] (iodide electrode), whereas triiodide is calculated from both; however, there are also methods that measure these species in a single operation.16,21






    FIGURE 16.1 Calculated equilibrium concentrations in aqueous iodine solutions without additional iodide. A, c[I2] = 10-3 mol/L. B, c[I2] = 10-4 mol/L. C, c[I2] = 10-5 mol/L. D, c[I2] = 10-6 mol/L.


  • Exceptions for the aforementioned quoted restriction to I–, I2, and I3– in presence of additional iodide are systems with very high iodide and iodine concentrations in which the equilibria in equations (5) and (6) are also important but are independent of pH as is the case with the equilibrium in equation (3). For example, in high-level Lugol solution, the species I5– and I62- make up 8.2% of the oxidation capacity and should not be neglected (Figure 16.3A). On the other hand, in the absence of additional iodide, at pH 8 to 9 and at high dilution (c[I2] ≤10-5 mol/L), HOI accounts for over 90% of the oxidation capacity (Figure 16.1C). Absence of iodide is also obligatory for the presence of the iodine cation H2OI+,22 but because this is associated with an extremely acid milieu, this has little relevance in practice.


  • The problem of stability (ie, the rate of iodate formation) arising at pH above 7 can be reduced to the equilibrium concentration of HOI, which manifests in the simple rate law:

    d [IO3–] / dt = 0.25 [HOI]3 / [H+]

    which allows for an estimate of stability at weak alkaline conditions.18 Figure 16.4 shows the initial rates of iodate formation (as a measure of stability) for iodine solutions (10-6 to 10-3 mol/L) without additional iodide and for a 0.001 mol/L iodine solution in the presence of additional iodide (10-4 to 10-1 mol/L).

The poor solubility of elemental iodine in water (338.3 ppm, 25°C, pH 5) can be increased by the addition of iodide, as first demonstrated by Lugol.2 Lugol solution is
a high-concentration iodine formulation with 5% (vol/vol) iodine (0.197 mol/L) and 10% (vol/vol) potassium iodide (KI) (0.6024 mol/L) and the following equilibrium concentrations: 6.129e-4 mol/L (155.6 ppm) free molecular iodine, 0.406 mol/L (51 650 ppm) iodide, 0.1803 mol/L (68 640 ppm) triiodide, 9.95 × 10-4 mol/L (631 ppm) pentaiodide, and 7.03 × 10-3 mol/L (5350 ppm) hexaiodide. Lugol solution, with a threefold molar excess of iodide, is completely soluble at any dilution. This applies to all iodine/iodide ratios at least down to a twofold excess of iodide. Less iodide results in a gap of solubility of free molecular iodine.






FIGURE 16.2 Calculated equilibrium concentrations in aqueous 10-3 mol/L iodine solutions in presence of additional iodide. A, No additional iodide. B, c[I–] = 10-3 mol/L. C, c[I–] = 10-2 mol/L. D, c[I–] = 10-1 mol/L.


Alcoholic Solution

Iodine equilibrates with alcohols by undergoing “outer” and “inner” complexes that finally result in the formation of triiodide, a reaction that takes approximately 24 hours23



Therefore, as in the aqueous system, we have several oxidizing iodine species: I2, ROH · I2, ROHI+, I3–, and ROH · I3–. However, calculations concerning their distribution, even if possible, are of no use in a bactericidal context in a solvent that is itself a strong disinfectant.


Solutions of Iodophors

Iodophors are polymeric organic molecules (alcohols, amides, sugars) capable of complexing iodine species, resulting in reduced equilibrium concentrations of the species compared with those of pure aqueous solutions with the same total iodine and total iodide concentrations.

Because iodophoric preparations always contain appreciable iodide, the relevant species that must be considered
are restricted to I–, I2, and I3–, for which the following (simplified) complexing reactions can be written:






FIGURE 16.3 Calculated equilibrium concentrations in Lugol solution and its dilutions. A, Undiluted. B, 1:100. C, 1:10 000.




where R = structural regions of the iodophor molecule that can form complexes by steric or electronic effects.

To the extent that the chemistry of aqueous disinfectant solutions containing iodophors is understood, both electronic and steric effects are thought to be responsible for these interactions.24 Thus, taking as an analogy known interactions with oxygen compounds of low molecular weight, such as amides, esters, ketones, and ether,25,26 it can be assumed that between molecular iodine and iodophor, molecules without exception contain such functional oxygen-containing groups (eg, povidone contains a carbonyl oxygen of the amide function in the pyrrolidinone ring), donor-acceptor complexes are formed (see also equation [11]), and iodine plays the part of the acceptor:


The iodophors, especially at high concentrations, are able to surround the iodine species in the manner of clathrates and withdraw it from equilibrium (equations [11], [12] and [13]) because of the spatial arrangement of the dissolved polymer molecules that exists with near regions of helix-like structure.20 This interaction must be important for the iodide ion and particularly for the large-mass triiodide ion, which cannot form a donor-acceptor complex because of their negative charge.

However, because no quantitative data (mass law constants) are available, an exact calculation for iodophoric preparations is not feasible. Nevertheless, qualitative investigations of the interactions with the iodophoric molecule polyvinylpyrrolidone (PVP) reveal that KI– is much less than KI2– and KI3–. With regard to the normal
conditions of use (ie, presence of appreciable iodide and pH <7), this has the following consequences18:






FIGURE 16.4 Initial rates of iodate formation expressed as percentage loss of initial oxidation capacity. A, c[I2] = 10-3, 10-4, 10-5, 10-6 mol/L; no additional iodide. B, c[I2] = 10-3 mol/L; c[I–] = 0, 10-4, 10-3, 10-2, 10-1 mol/L. From these curves, it can be determined that, for example, in a 10-4 mol/L iodine solution at pH 8, the initial rate of loss of oxidation capacity is approximately 10% of the initial concentration per minute.



  • The HOI and the species derived from it (OI–, HI2O–, I2O2-, and H2OI+) can be neglected.


  • Because the reactions in equations (11), (12) and (13) reduce the equilibrium concentrations of I2, I3–, and, to a certain degree, I–, the species I5– and I62- can be ignored as well.

Therefore, in iodophoric preparations, only the triiodide equilibrium (equation [3]) and the interactions of the iodophoric molecules with I2, I3–, and I– are important, all of which are independent of pH. Because HOI is virtually absent, stability problems concern only interactions with oxidizable components but not disproportionation to iodate.


Influence of Temperature

Although not usually considered, temperature should not be overlooked. In a study dealing with 10 different povidone-iodine preparations, the results for the relative alteration of free iodine with temperature fitted to an exponential function of the form


which is valid from 10°C to 40°C.27 Following this equation, [I2] increases approximately 5.4% and 100% if the temperature rises 1.0°C and 13.1°C, respectively. This increase of [I2] must be considered in the application of povidone-iodine preparations as disinfectants or antiseptics on living tissues. Because of their higher temperature (30°C-36°C), povidone-iodine preparations used on living tissues exhibit a significantly higher [I2] than they do at room temperature (Δt = 10°C-16°C; Δ% [I2] = 70%-130%).

In a clinical study examining the application of iodophoric preparations, no difference was found in the germicidal effectiveness of povidone-iodine (10% vol/vol) against Staphylococcus aureus, Enterococci, Escherichia coli, or group B Streptococcus that was warmed to 32°C versus povidone-iodine that was used at 25°C during application.28


Atypical Behavior of Iodophors at Dilution

If 10% povidone-iodine is diluted, the concentration of free molecular iodine unexpectedly increases and passes through a maximum in the 0.1% solution. As can be seen in Figure 16.5, the concentration of free iodine in a 10% povidone-iodine solution is approximately 2.0 mg and 8 × 10-6 mol/L and rises in a 1:100 dilution nearly 10-fold. On further dilution, after passing the maximum ([I2] ≈ 10-4 mol/L), the free iodine behaves increasingly “normally”—that is, it decreases—and below 0.01%, the povidone-iodine solution can be regarded as a simple aqueous solution of iodine.

Because [I2] depends not only on the concentration of povidone-iodine but also on total iodine (in general, 1%) and total iodide (iodine-iodide ratio; see Pinter et al30), and the presence of iodine-complexing pharmaceutical additives, it undergoes considerable variation. Figure 16.5 shows the typical course of [I2] of a pure aqueous povidone-iodine at dilution. The ordinate of the maximum (and to a lesser degree, its abscissa) therefore is not a constant for different iodophors and preparations containing
iodophors. Figure 16.3 also shows the behavior of Lugol solution on dilution, which explains the drastic reduction of free iodine caused by the complexing properties of the povidone molecules.






FIGURE 16.5 Total available (dashed lines) and free molecular (solid lines) iodine in aqueous povidone-iodine (determined potentiometrically) (Gottardi19) and in Lugol solution (calculated after Gottardi29).

This atypical behavior has been confirmed in a study that compared three distinct solutions of povidone-iodine, ranging in povidone-iodine concentrations from 7.5% (wt/vol) to 10% (wt/vol).31 The maximum free iodine concentration after dilution ranged from 31 to 51 mg/L. All three solutions exhibited a similar trend in that the amount of free iodine increased with increasing dilution factor, up to a dilution factor of 1 and then began decreasing for the subsequent dilution factors of 10 and 100.


Individual Reactivity of Iodine Species

Iodide (I–), as a nonoxidizing species, has no antiseptic activity. This is true for iodate (IO3–) as well, which acts as an oxidant only at acidic pH, as HIO3 (pH <4).

Free iodine (I2) is the only species with a proved correlation between equilibrium concentration and bactericidal activity. Its solvated forms, I2 · H2O or I2 · ROH, are thought to be the effective antimicrobial agents in aqueous and alcoholic solution. The term free serves to distinguish this I2 from complex-bound I2, as discussed later for iodophoric preparations, and refers to the solvated forms I2 · H2O and I2 · ROH.

Triiodide (I3–) probably has no antiseptic activity, which was deduced from the negative effect that increasing the iodide concentration has on inactivation of poliovirus.32 On the other hand, triiodide represents the reservoir oxidation capacity in noniodophoric preparations (Lugol solution). It is the main species responsible for staining of tissue.33

The HOI is thought to contribute to bactericidal action, which is a plausible analogy with Cl2/H2O solutions, where HOCl is the true active species (see chapter 15). Chang34 claimed differing behaviors of I2 and HOI against certain microbes, reports on the bactericidal properties of HOI, and on attempts to establish a difference between I2 and HOI. However, these observations should be treated cautiously.18 It is possible to manipulate an aqueous iodine solution to exhibit greater than 90% of its oxidation capacity as HOI (pH ≈ 8.5, no additional iodide); such systems in general have no practical importance, mainly because of stability problems. On the other hand, given a similar reactivity of I2 and HOI in highly diluted systems (c[I2] ≤10-5 mol) without additional iodide (as is usual in drinking water disinfection) and in the pH range of 3 to 9, a more or less constant bactericidal activity of iodine in aqueous solution can generally be expected.35

Iodine cation (H2OI+) is thought to be a very potent iodinating agent. Many publications identify this species as being responsible for disinfection, which can be traced back to a comprehensive and often-cited study dealing with the halogens in disinfection.32 However, exact calculations18 show that the iodine cation has some importance, if any, only under very acidic conditions (pH <1) and in the total absence of additional iodide, where it amounts to only approximately 0.3% of the total iodine at high dilutions (Figure 16.2D). However, under conditions used in practice (ie, in the presence of iodide to regulate the concentration of free molecular iodine and improve stability), the iodine cation is virtually absent and therefore likely of no significance. For example, a solution with c[I2] = 0.001 and c[I–] = 0.01 mol/L generates [I2] = 1.31 × 10-4 mol/L or 33.3 ppm at pH <8 (Figure 16.3C). The concentration of the iodine cation, however, is [H2OI+] = 2.15 × 10-13 mol/L (not shown in Figure 16.3C), which is approximately nine orders of magnitude less than [I2]. Even if a higher reactivity is attributed to the iodine cation, which is thought to play an important role in certain substitutions, this is not likely able to explain any real contribution to a disinfecting process.


Virtual Impossibility of Discriminating Antimicrobial Activities of I2 and HOI

The contribution of free molecular iodine, I2, and HOI to the disinfection processes and differences in their bactericidal power has been previously discussed.32 A solution containing predominantly I2 requires a pH of less than 5 and absence of iodide, whereas a solution containing
predominantly HOI needs a pH of approximately 8.4. A comparison of the killing effect of I2 and HOI presupposes that the susceptibility of bacteria for interaction with these iodine species is the same in both pH ranges.


Stability of Iodine-Based Disinfectants

Iodine and other disinfectants based on halogens in the oxidation states 0 or +1, if they are not present as pure substances (ie, without a solvent), can gradually lose their antimicrobial properties (eg, during storage). This is due to (1) substitutions of covalent hydrogen (eg, O-H, N-H, C-H, as a result of reactions with solvent molecules and other formulation additives); (2) additions to olefinic double bonds; and (3) the disproportionation of hypohalous acid to halate in aqueous preparations (equation [9]), which has little to no antimicrobial properties. Although substitutions, which in the case of iodine are thought to be fewer than with chlorine and bromine, and additions can be avoided by an appropriate composition, the equilibria in equations (1), (2), (3), (4), (5), (6), (7) and (8) are established in any case if water is present, and iodate formation (equation [9]) can begin.

Based on calculated equilibrium concentrations, reaction times, and initial rates of iodate formation, the following conclusions have been drawn concerning the stability of iodine-containing disinfecting solutions14,18,35:



  • Below pH 6, a decrease in disinfection efficacy because of the formation of iodate can be excluded.


  • Above pH 7, the formation of iodate, the extent of which largely depends on the pH value as well as on the iodide concentration, must be taken into consideration. Raising the pH value lowers the stability (iodate formation increases), whereas raising the iodide concentration improves the stability (iodate formation is reduced).


  • Because of the stabilizing effect of the iodide ion, provided its concentration is sufficiently high, the opposing effect of the pH value can be overcompensated. As a result, iodine-based preparations can also exhibit sufficient stability for practice in the weak alkaline range (eg, Lugol solution, at pH <9).


  • In highly diluted iodine solutions (<10-5 mol/L, or 2.54 mg/L), which are used to disinfect potable water or swimming pool water, only a slow iodate formation can be expected even in absence of additional iodide and pH 8 or lower. In accordance with this, in iodine-based disinfection plants, no significant amounts of iodate have been detected.36


  • Atemnkeng et al31 studied the dilution of three commercial povidone-iodine formulations (Braunol®, standardized Betadine®, and unstandardized iso-Betadine®). The pH values were found to increase as a function of the dilution ranging from 5.55 to 5.80 for unstandardized iso-Betadine®, whereas being higher (5.65-6.20) for Braunol® and standardized Betadine® solutions. The values remain constant with increase in dilution factors. The observed pH for all three solutions is advantageous because their values approximate the pH of the skin at any dilution. The pH >5 in formulations is also important to control Dushman reaction, which could raise free iodine concentrations to an intolerable level for skin application.31


The Absence of N-iodo Compounds in Aqueous Solution

The well-known equilibria of halogen with N-H compounds is shown below:



These lead to numerous N-chloro compounds that play an important role in chlorine-based disinfection practice; this, however, is not the case with iodine. Although N-iodo compounds can be synthesized in a nonaqueous system,37,38 in contact with water, they immediately hydrolyze:


This assertion complies with the pH range relevant for disinfection (pH 4-8) and not with alkaline conditions at 0°C, where N-iodo alkylamines can be synthesized even in aqueous solutions.39

The HOI immediately begins to disproportionate according to equation (9); iodide and protons develop, and a comproportionation or synproportionation chemical reaction in which each reactant, containing the same element but with different oxidation numbers, forms a product with the same oxidation number; and the reverse of reaction in equation (1) also takes place, forming molecular iodine. Within minutes, the reaction settles, and the ratio of resulting products, I2 and IO3–, complies with the stoichiometry set forth in equation (16)40:


Because the equilibria of equations (15) and (16) are lying far on the right side, practically no N-iodo compound is present and the reverse reactions of equations (14a) and (14b) do not take place. N-iodo compounds, therefore, are virtually without any relevance for disinfection.


Reaction With Proteins: Antimicrobial Action and Consumption Effects

Halogens can react with both living and dead microorganisms and with dissolved proteins. In contrast to chlorine, where oxidizing (and antimicrobial) N-chloro
compounds still emerge, reactions with compounds of iodine (consumption effects) are associated only with a loss of oxidation capacity (equations [16], [17] and [18]) because N-iodo compounds are not formed (see earlier).

Interactions with thiol groups in proteins (S-H compounds) can result in formation of different compounds whereby oxidation to disulfides (equation [16a]) can occur on the one hand and to sulfur-oxygen acids (ie, sulfenic, sulfinic, and sulfonic acids; equation [16b]) on the other hand. Because these reactions occur with similar speed, the proportions of the diverse products are mainly governed by the mode of mixing. Equation (17) describes substitution at activated aromatic compounds (eg, the amino acids tyrosine, histidine, and the nucleosides cytosine and uracil), whereas equation (18) refers to the addition of I2 to the olefinic function of unsaturated fatty acids






PREPARATIONS CONTAINING OR RELEASING FREE IODINE


Solutions of Iodine and Iodide

To this group belongs a great variety of preparations containing elemental iodine and potassium (or sodium) iodide in water, ethyl alcohol, and glycerol, or in mixtures of these solvents. They rank with the oldest disinfectants and have survived more than 150 years owing to their efficacy, economy, and stability. The following are the official preparations according to the United States Pharmacopeia and the National Formulary3: (1) iodine topical solution, an aqueous solution containing 2.0% iodine and 2.4% sodium iodide; (2) strong iodine solution (Lugol solution), an aqueous solution containing 5% iodine and 10% KI; (3) iodine tincture, containing 2.0% iodine and 2.4% sodium iodide in aqueous ethanol (1:1); and (4) strong iodine tincture, containing 7% iodine and 5% KI in 95% ethanol. Because all these preparations contain large amounts of iodide (0.16-0.6 mol/L), only the triiodide equilibrium (equation [3]) becomes important. As a result, these solutions virtually contain only molecular iodine, iodide, and triiodide and are therefore very stable because there is no HOI present (see earlier). Because of their high content of free molecular iodine (eg, Lugol solution [I2] = 155.6 ppm), they are powerful disinfectants with the disadvantage of staining and a toxic potential that should not be underestimated.


Preparations Containing Organic Complexing Agents

In addition to preparations with complexing agents of low molecular weight, such as tetraglycine hydroperiodide,10 the inclusion compound iodine-maltosylcyclodextrin41 or quaternized chitosan with iodine,42 this group includes the important iodophors, a term indicating in general the combination of iodine with a carrier (as these complexing agents usually are called) of high molecular weight. In aqueous solution, iodophors form comparable iodine species as pure aqueous iodine solutions (see earlier overview). However, the polymer carriers, because of their complexing properties, partly reduce the equilibrium concentrations of the iodine species and give the iodophor preparations properties that make them superior in some respects to solutions containing only iodine and iodide.


Iodophors

An iodophor is a complex of iodine with a carrier that has at least three functions: (1) to increase the solubility of iodine, (2) to provide a sustained-release reservoir of the halogen, and (3) to reduce the equilibrium concentration of free molecular iodine. The carriers are usually neutral polymers, such as PVP, nonylphenoxy polyethoxyethanol, polyether glycols, polyvinyl alcohols, polyacrylic acid, polyamides, polyoxyalkylenes, and polysaccharides.

In the solid state, iodophors form deep brown to black, crystalline powders that usually do not smell of iodine, indicating a tight bonding with the carrier molecules. The solubility of iodophors in water is good but depends on the chain length of the polymeric molecules. In the case of povidone-iodine, the solubility can vary between 5% (type 90/04, average molecular weight near 1 000 000) and more than 20% (type 17/12, average molecular weight near 10 000). The best known iodophor is povidone-iodine, a compound of l-vinyl-2-pyrrolidone polymer with iodine, which according to United States Pharmacopeia and National Formulary3 contains no less than 9.0% and no more than 12.0% of available iodine when calculated on the dried basis. On the basis of spectroscopic investigations,43 it was found that povidone-iodine (in the solid state) is an adduct not with molecular iodine (I2) but with hydrotriiodic acid (HI3), where the proton is fixed by a short hydrogen bond between two carbonyl groups of two pyrrolidone rings and the triiodide anion is bound ionically to this cation (Figure 16.6).







FIGURE 16.6 Structure of solid povidone-iodine. Reprinted from Schenck et al.43 Copyright © 1979 Wiley-Liss, Inc., A Wiley Company. With permission.

A completely different situation occurs in solution, where this structure no longer exists and equilibria between I2, I–, I3–, and the polymeric organic molecules are established (equations [3] and [11], [12] and [13]). The high concentration of carrier molecules (approximately 90 g/L) results in the content of free molecular iodine being greatly reduced in such preparations (10% aqueous solution of povidone-iodine: c[I2] ≈ c[I–] ≈ 0.04 mol/L, [I2] ≈ 1 × 10-5 mol/L or 2.54 ppm) compared with pure aqueous solutions with the same total iodine and total iodide content (aqueous iodine solution: c[I2] ≈ c[I–] ≈ 0.04 mol/L, pH 5: [I2] = 5.77 × 10-3 mol/L or 1466 ppm). Note that this is a hypothetical value because the solubility of molecular iodine is 334 ppm (at 25°C), and an aqueous solution of this composition contains unquantifiable, undissolved iodine. The high content of free iodide (which varies between 10-3 and 10-1 mol/L, according to the preparation) also means that HOI can be disregarded and only I2 is responsible for disinfection (see earlier).


Forms of Application

According to the USP 41-NF 36,3 the following application forms of povidone-iodine are approved: povidone-iodine topical solution, povidone-iodine cleansing solution, povidone-iodine ointment, and povidone-iodine topical spray solution. Regarding available iodine, all must contain no less than 85% and no more than 120% of the labeled amount. In general, povidone-iodine preparations contain 1% to 10% povidone-iodine, which is equivalent to 0.1% to 1.0% available iodine; the cleansing solutions contain one or more surface-active agents.


Iodine Use on the Efficacy of Povidone-iodine Preparations

Because iodophoric preparations are mainly used as antiseptics, the influence of iodine-consuming body fluids is a very important feature regarding antimicrobial potency and rate. Mainly in presence of blood, which is characterized by numerous SH group functions, the reservoir of available iodine is substantially diminished and, because of the formed iodide, the triiodide equilibrium (equation [3]) is shifted to the right. Both effects decrease the proportion of free molecular iodine (see earlier).

On the other hand, when povidone-iodine preparations are contaminated with liquid substrates, the dilution effect (see earlier) causes an increase in the equilibrium concentration of free molecular iodine (see Figure 16.5). The extent to which this effect compensates for the other two depends on the content of reducing substances. Thus, with whole blood, a large decrease in the concentration of free molecular iodine occurs, whereas in the presence of plasma or exudates, the concentration remains unchanged if the ratio of blood to povidone-iodine is not too high.44

Quantitative analysis of iodine (and other oxidizing substances) with blood has poor reproducibility due to the multiple different reactions of iodine with various SH groups (equations [16a] and [16b]). In practice, no substantial decrease in the bactericidal efficacy of 10% povidone-iodine preparations is likely with body fluids having a composition similar to plasma (volume substrate/volume 10% povidone-iodine ≤0.6). However, contamination by 25% or more of whole blood can significantly limit the observed antimicrobial activity.


Iodophoric Preparations and “Active Agents”

In the search for the ideal disinfectant, particularly antiseptics including the combination of immediate microbial kill with complete lack of unwanted side effects on host tissue, much work was done by microbiologists comparing various preparations and formulations containing different antimicrobials (eg, chlorine, iodine, aldehydes, peroxides, chlorhexidine, and quaternary ammonium compounds). The results of such studies usually are presented in such terms as, for example, “0.25% chlorhexidine and 0.025% benzalkonium chloride was more (or less) effective than 10% povidone-iodine.” Such descriptions imply that povidone-iodine is an active agent; however, this is not the case. The basic requirements to designate a substance as an active agent are (1) a defined molecule and (2) a positive correlation between the concentration of these defined molecules and antimicrobial activity. Despite its extensive use, both criteria are not fulfilled when using povidone-iodine preparations and may lead to misleading results.

Regarding the nature of povidone-iodine and its disinfecting properties, the following points have been made45:

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May 9, 2021 | Posted by in MICROBIOLOGY | Comments Off on Iodine and Iodine-Containing Compounds

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