| Silver nitrate | |
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Other names
Nitric acid silver(1+) salt |
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| Identifiers | |
| CAS number | 7761-88-8 |
| PubChem | 24470 |
| ChemSpider | 22878 |
| UNII | 95IT3W8JZE |
| ChEBI | CHEBI:32130 |
| ChEMBL | CHEMBL177367 |
| Jmol-3D images | Image 1 |
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| Properties | |
| Molecular formula | AgNO3 |
| Molar mass | 169.87 g mol−1 |
| Appearance | white solid |
| Density | 4.35 g cm−3 |
| Melting point |
212 °C, 485 K, 414 °F |
| Boiling point |
444 °C, 717 K, 831 °F (decomp.) |
| Solubility in water | 1.22 kg/L (0 °C) 2.16 kg/L (20 °C) 4.40 kg/L (60 °C) 7.33 kg/L (100 °C) |
| Solubility | soluble in ethanol and acetone |
| Hazards | |
| EU classification | C N |
| R-phrases | R8,R34, R50/53 |
| S-phrases | (S1/2), S26, S45, S60, S61 |
| NFPA 704 |
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| (verify) (what is: /?) Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa) |
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| Infobox references | |
Silver nitrate is an inorganic compound with chemical formula AgNO3. This compound is a versatile precursor to many other silver compounds, such as those used in photography. It is far less sensitive to light than the halides. It was once called lunar caustic because silver was called luna by the ancient alchemists, because they believed that silver was associated with the moon.[1]
In solid silver nitrate, the silver ions are three-coordinated in a trigonal planar arrangement.[2]
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Albertus Magnus, in the 13th century, documented the ability of nitric acid to separate gold and silver by dissolving the silver.[3] Magnus noted that the resulting solution of silver nitrate could blacken skin. Its common name at the time was nitric acid silver.
Silver nitrate can be prepared by reacting silver, such as a silver bullion or silver foil, with nitric acid:
This is performed under a fume hood because of toxic nitrogen oxide(s) evolved during the reaction.[4]
A typical reaction with silver nitrate is to suspend a rod of copper in a solution of silver nitrate and leave it for a few hours. The silver nitrate reacts with copper to form hairlike crystals of silver metal and a blue solution of copper nitrate:
Silver nitrate also decomposes when heated:
Most metal nitrates thermally decompose to the respective oxides, but silver oxide decomposes at a lower temperature than silver nitrate, so the decomposition of silver nitrate yields elemental silver instead.
Silver nitrate is the least expensive salt of silver; it offers several other advantages as well. It is non-hygroscopic, in contrast to silver fluoroborate and silver perchlorate. It is relatively stable to light. Finally, it dissolves in numerous solvents, including water. The nitrate can be easily replaced by other ligands, rendering AgNO3 versatile. Treatment with solutions of halide ions gives a precipitate of AgX (X = Cl, Br, I). When making photographic film, silver nitrate is treated with halide salts of sodium or potassium to form insoluble silver halide in situ in photographic gelatin, which is then applied to strips of tri-acetate or polyester. Similarly, silver nitrate is used to prepare some silver-based explosives, such as the fulminate, azide, or acetylide, through a precipitation reaction.
Treatment of silver nitrate with base gives dark grey silver oxide:[5]
Standard solution molarity = 0.05006 M, depends on solubility and concentration however.
The silver cation, Ag+, reacts quickly with halide sources to produce the insoluble silver halide, which is a cream precipitate if Br- is used, a white precipitate if Cl- is used and a yellow precipitate if I- is used. This reaction is commonly used in inorganic chemistry to abstract halides:
where X− = Cl−, Br−, or I−.
Other silver salts with non-coordinating anions, namely silver tetrafluoroborate and silver hexafluorophosphate are used for more demanding applications.
Similarly, this reaction is used in analytical chemistry to confirm the presence of chloride, bromide, or iodide ions can be tested by adding silver nitrate solution. Samples are typically acidified with dilute nitric acid to remove interfering ions, e.g. carbonate ions and sulfide ions. This step avoids confusion of silver sulfide or silver carbonate precipitates with that of silver halides. The color of precipitate varies with the halide: white (silver chloride), pale yellow/cream (silver bromide), yellow (silver iodide). AgBr and especially AgI photo-decompose to the metal, as evidence by a grayish color on exposed samples.
Silver nitrate is used in many ways in organic synthesis, e.g. for deprotection and oxidations. Ag+ binds alkenes reversibly, and silver nitrate has been used to separate mixtures of alkenes by selective absorption. The resulting adduct can be decomposed with ammonia to release the free alkene.[6]
In histology, silver nitrate is used for silver staining, for demonstrating reticular fibers, proteins and nucleic acids. For this reason it is also used to demonstrate proteins in PAGE gels. It is also used as a stain in scanning electron microscopy..
Silver salts have antiseptic properties. Until the development and widespread adoption of antibiotics, dilute solutions of AgNO3 used to be dropped into newborn babies' eyes at birth to prevent contraction of gonorrhea from the mother. Eye infections and blindness of newborns was reduced by this method; incorrect dosage, however, could cause blindness in extreme cases. This protection was first used by Credé in 1881.[8][9][10] Fused silver nitrate, shaped into sticks, was traditionally called "lunar caustic". It is used as a cauterizing agent, for example to remove granulation tissue around a stoma. Dentists sometimes use silver nitrate infused swabs to heal oral ulcers. Silver nitrate is also used by some podiatrists to kill cells located in the nail bed. Silver nitrate is also used to cauterize superficial blood vessels in the nose to help prevent nose bleeds.
The Canadian physician C. A. Douglas Ringrose researched the use of silver nitrate for sterilization procedures on women. A specialist in obstetrics and gynaecology, Ringrose believed that the corrosive properties of silver nitrate could be used to block and corrode the fallopian tubes, in a process that he called "office tubal sterilization".[11] The technique was ineffective; in fact at least two women underwent abortions. Ringrose was sued for malpractice, although these suits were unsuccessful.[12]
Much research has been done in evaluating the ability of the silver ion at inactivating E. coli, a microorganism commonly used as an indicator for fecal contamination and as a surrogate for pathogens in drinking water treatment. Concentrations of silver nitrate evaluated in inactivation experiments range from 10–200 micrograms per liter as Ag+. The antimicrobial properties of silver was first observed thousands of years ago when silver containers were used to store water for preservation. Its disinfection ability has been scientifically studied for over a century.
Silver's antimicrobial activity saw many applications prior to the discovery of pharmaceutical antibiotics, when it fell into near disuse. Its association with argyria made consumers wary and led them to turn away from it when given an alternative. Since that time, as antibiotic-resistant microorganisms have emerged, interest in using the silver ion for anti-microbial purposes has resumed.[13]
Before a disinfectant can be effectively used as a water disinfectant, its inactivation kinetics must be established. Kinetics generally depend on both the dosage of disinfectant and the time of application. It is important to understand the kinetics so that the minimum dosage of disinfectant can be applied for the minimum amount of time while still effectively inactivating any pathogens in the water. Because there are many microorganisms present in water, the inactivation kinetics of each one cannot be studied extensively. Therefore, indicator organisms generally more resistant to inactivation than others are used to estimate the kinetics of microorganisms as a whole. Escherichia coli, also referred to as E. coli, is a commonly used indicator organism.
It is well documented that the silver ion is effective in the inactivation of E. coli.[14][15][16][17][18][19][20][21][22][23][24][25] However, there are many inconsistencies in the literature regarding the kinetics of the inactivation of E. coli by the silver ion. With inconsistent data, it is impossible to tell what the true inactivation kinetics are, and therefore impossible to implement any sort of large-scale water treatment.
The inconsistencies may be due to several factors. First, the kinetics may depend on the source of the silver ion being used. In recent years, research has focused largely on electrolytically generated silver ions or colloidal silver. Most studies in which the inactivation kinetics of E. coli by silver nitrate were explored extensively date back several decades. Even within this smaller group of studies, vast inconsistencies exist, likely due to inaccurate analytical methods for measuring the concentration of silver in solution.[26] Monitoring the decay of the silver ion in solution is imperative as silver tends to both adsorb readily to organic matter in the water and to be light reactive.[27] Furthermore, silver tends to adsorb to glassware, which can lead not only to a decrease in the silver concentration within a given experiment but also to a release of the silver in subsequent experiments unless measures further than general glassware washing are taken for the removal of silver from the glassware surface.[24] Therefore studies must both minimize the external factors effecting the concentration and to measure the changes in concentration that take place throughout the experiment.
Despite the inconsistencies in the literature regarding the kinetics of the inactivation of E. coli by silver nitrate, important information can still be taken from the work. A study by Wuhrmann and Zobrist investigated the effect of various parameters upon the kinetics. First, they studied the effect of several ions in the water, including calcium, phosphates and chloride, all of which were found to decrease the bactericidal effect of silver.[23] These effects are important to consider when designing an experiment. Because of the effect of phosphates, it is undesirable to use phosphate buffer to run experiments, as this creates a phosphate concentration much higher than that found in natural waters and will falsely slow the inactivation kinetics. Furthermore, it is important to avoid touching any glassware with bare hands, as chloride from sweat may contaminate the glassware, again slowing inactivation. Chambers, Proctor and Kabler established the importance of using an effective neutralizer solution made of a combination of sodium thioglycolate and sodium thiosulfate, rather than sodium thiosulfate alone, which though it is effective in neutralizing other disinfectants does not sufficiently stop the bactericidal action of silver nitrate.[24] Both tested the effect of pH on the kinetics, finding that a higher pH increased the bactericidal action.[24][27] Wuhrmann and Zobrist further established that at a higher temperature, inactivation occurs faster.[23]
A further complication of the inactivation kinetics by silver is the question of which model to use. With most disinfectants, the inactivation is effectively modeled using a first-order Chick-Watson model, which states that a certain level of disinfection will occur at a certain CT, or concentration *time value.[28] According to this model, the same amount of inactivation should take place when a concentration of 0.2 mg/L is applied for 10 minutes as when 0.02 mg/L is applied for 100 minutes. Wuhrmann and Zobrist found rate kinetics that followed this model for all conditions, which agrees fairly well with a study by Chambers and Proctor, while another study by Renn and Chesney found curves that did not follow this law.[26] It is therefore unclear whether this law sufficiently models inactivation by the silver ion.
Most recent papers regarding the disinfection of E. coli by silver nitrate have simply plotted the level of disinfection against time.[14][17][19][21][22] While this method of data analysis does not risk making false assumptions about first-order kinetics, it does nothing to account for the applied concentration, which is essential to any kinetics. Therefore, different curves need to be generated for each concentration that might be applied. Furthermore, it does not account for changes in concentration that might take place during the experiment, and which may vary based on many factors.
A third model which has been suggested for the inactivation kinetics by silver nitrate is that of Cs*T, or chemisorbed silver onto the cell body times time. This model suggests that the rate of inactivation depends not on the concentration in the water at a given time, but rather on the silver that has been chemisorbed by the bacteria. It is assumed, according to this model that C0 = C1 + C2 + C3, where C0 is the initial concentration, C1 is the silver still in solution, C2 is the silver lost to adsorption to glassware or other factors in the solution, and C3 is the silver chemisorbed to the bacteria. C0 is measured at the beginning of the experiment, C1 is measured throughout the experiment, and C2 is determine in a control experiment without bacteria. C3, or the Cs value, is then determined to be C0-C1-C2.[15][29] According to Hwang, et al., this model was successful in estimating inactivation of E. coli by silver nitrate.[15] Although it is possible that this model does not sufficiently account for all of the possible fates of the initial silver nitrate added to the solution, it is certainly a compelling method of data analysis. Because it is a new model, it has not been extensively studied by various researchers.
Repeated daily application of silver nitrate can induce adequate destruction of cutaneous warts, but occasionally pigmented scars may develop. In a placebo-controlled study of 70 patients, silver nitrate given over nine days resulted in clearance of all warts in 43% and improvement in warts in 26% one month after treatment compared to 11% and 14%, respectively, in the placebo group.[30]
As an oxidant, silver nitrate should be properly stored away from organic compounds. Despite being used in low concentrations to prevent gonorrhea and control nose bleeds, silver nitrate is toxic and corrosive.[31] Brief exposure to the chemical will not produce immediate or even any side effects other than the purple, brown or black skin stains; but with more exposure, side effects will become more noticeable, including burns. Long-term exposure may cause eye damage. Short contact can lead to deposition of black silver stains on the skin. Besides being very destructive of mucous membranes, it is a skin and eye irritant.
Although silver nitrate is currently not regulated in water sources by the Environmental Protection Agency, when between 1-5 g of silver have accumulated in the body, a condition called argyria can develop. Argyria is a permanent cosmetic condition in which the skin and internal organs turn a blue-gray color. The United States Environmental Protection Agency had a maximum contaminant limit for silver in water until 1990, but upon determination that argyria did not impact the function of organs affected, removed the regulation.[27] Argyria is more often associated with the consumption of colloidal silver solutions than with silver nitrate, especially at the extremely low concentrations present for the disinfection of water. However, it is still important to consider before ingesting any sort of silver-ion solution.
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