Silver chlorate
| Names | |
|---|---|
| IUPAC name
Silver(I) chlorate
| |
| Other names
Argentous chlorate
Silver(I) chlorate(V) | |
| Properties | |
| AgClO3 | |
| Molar mass | 191.319 g/mol |
| Appearance | White colorless solid |
| Odor | Odorless |
| Density | 4.443 g/cm3 (20 °C) |
| Melting point | 230–231 °C (446–448 °F; 503–504 K) (decomposes) |
| Boiling point | 270 °C (518 °F; 543 K) (decomposes) |
| 8.52 g/100 ml (5 °C) 10 g/100 ml (15 °C) 18.03 g/100 ml (25 °C) 23.74 g/100 ml (35 °C) 50 g/100 ml (80 °C)[1][2] | |
| Solubility | Reacts with acids and bases Sparingly soluble in alcohols |
| Vapor pressure | ~0 mmHg |
| Thermochemistry | |
| Std molar
entropy (S |
141,838 J·mol-1·K-1 |
| Std enthalpy of
formation (ΔfH |
-30,292 kJ/mol |
| Hazards | |
| Safety data sheet | Sigma-Aldrich |
| Flash point | Non-flammable |
| Related compounds | |
| Related compounds
|
Silver perchlorate |
| Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa). | |
| Infobox references | |
Silver chlorate is an inorganic compound with the formula AgClO3.
Contents
Properties
Chemical
Silver chlorate decomposes upon gentle heating to release oxygen.
- AgClO3 → AgCl + 3/2 O2
It is a powerful oxidizer, and mixtures with organic compounds, like sugars, will burn vigorously.[4]
Silver chlorate can be used to prepare other chlorate salts, via double displacement with a metal chloride.
- MCl + AgClO3 → MClO3 + AgCl ↓
Silver chlorate has been proven to be an effective oxidizing agent for some organic compounds, like crotonic acid and isocrotonic acid, which are oxidized to dZ-threo-a,P-dihydroxybutyric acid and to dl-erythro-a,P-dihydroxybutyric acid, respectively.[5]
Physical
Silver chlorate is a colorless crystalline solid, soluble in water.
Explosive
Pure silver chlorate is stable in dry and moist air. After being stored for a few days the compound becomes very sensitive and explodes on touch. On exposure to daylight, AgClO3 slowly decomposes and becomes dark and turns into an extremely explosive compound. This can also happen on the stoppers of bottles and may lead to detonation when the stopper is removed. AgClO3 decomposes to AgCl with low heat or on melting, but when quickly heated it explodes or deflagrates.[6]
Availability
Silver chlorate is difficult to find, but many chemical suppliers will often have it in their stock, at high price.
Preparation
The most common route for preparing silver chlorate, involves the double displacement reaction of silver nitrate with sodium chlorate, which yields both silver chlorate and sodium nitrate. Fractional recrystallization is used to separate the two compounds from the solution.[7][8]
- AgNO3 + NaClO3 → AgClO3 + NaNO3
The most convenient synthesis route involves the double displacement reaction between barium chlorate and silver(I) fluoride. Since barium fluoride is almost insoluble in water, the resulting silver chlorate obtained this way is very pure. Main disadvantage is that silver fluoride is expensive to make.
- Ba(ClO3)2 + 2 AgF → 2 AgClO3 + BaF2 ↓
Another direct route involves the careful addition of chloric acid to silver oxide or silver carbonate.[9][10]
- 2 HClO3 + Ag2O → 2 AgClO3 + H2O
- 2 HClO3 + Ag2CO3 → 2 AgClO3 + H2O + CO2
Reaction with silver metal will also produce silver chlorate, but silver chloride is also produced as side product.[11][12]
- 6 HClO3 + 6 Ag → 5 AgClO3 + AgCl + 3 H2O
While simpler than the first route, one major issue in using HClO3 is that it is unstable at high concentrations, meaning the resulting product will be dilute, and possibly contaminated with silver perchlorate, from perchloric acid, which results from the rapid decomposition of conc. chloric acid.
There are other more accessible routes:
Bubbling chlorine gas through an aqueous suspension of silver oxide, similar to the preparation of alkali chlorates has been described as another synthesis route.[13][14]
- Ag2O + Cl2 + H2O → AgClO3 + AgCl ↓
This works because it produces hypochlorous acid/silver hypochlorite, which rapidly breaks down into chlorate and chloride. However, the resulting AgCl will lower the yield of this route, and filtration is required to remove it from the solution.[15]
Projects
- Make other chlorates
- Primary explosive
Handling
Safety
Silver chlorate is harmful, due to its chlorate anion. It is also a sensitive explosive and must not be kept for long periods of time, as if it builds up on the bottle cap, it may detonate when one opens the bottle.
Storage
Silver chlorate should not be kept for long periods of time and must be used quickly. Ampouling may be used, but the compound must not be exposed to light.
Disposal
Addition of an aq. solution of NaOH will cause the silver to precipitate as silver oxide, while the chlorate remains in the solution. The silver should be recycled, while the chlorate can be reduced to chloride using a reducing agent.
References
- ↑ Справочник химика. - Т. 3. - М.-Л.: Химия, 1965 (Handbook of a chemist. - T. 3. - M.-L.: Chemistry, 1965)
- ↑ Ефимов А.И. и др. Свойства неорганических соединений. Справочник. - Л.: Химия, 1983 (Efimov A.I. etc. Properties of inorganic compounds. Directory. - L .: Chemistry, 1983)
- ↑ Barin I. Thermochemical Data of Pure Substances. - VCH, 1995 pp. 7
- ↑ Silver Chlorate(AgClO3) Test
- ↑ Inorganic Syntheses, Silver Chlorate, D. G. Nicholson, Cearles E. Holley Jr., George W. Watt, Joe C. Evans
- ↑ Gmelins Handbuch der Anorganischen Chemie, Silber Teil B1, Verlag Chemie GmbH, Weinheim/Bergstraße, 8th edition 1971, p. 503
- ↑ Inorganic Syntheses, Silver Chlorate, D. G. Nicholson, Cearles E. Holley Jr., George W. Watt, Joe C. Evans
- ↑ How to Make Silver Chlorate (Good Prep for Beginners)
- ↑ Hendrixson; Journal of the American Chemical Society; vol. 25; (1903); p. 639
- ↑ Foote, H. W.; Saxton, B.; Journal of the American Chemical Society; vol. 36; (1914); p. 1704 - 1708
- ↑ https://scholarworks.uni.edu/cgi/viewcontent.cgi?article=7036&context=pias
- ↑ Hendrixson; Journal of the American Chemical Society; vol. 25; (1903); p. 639
- ↑ Pierron, P.; Bulletin de la Societe Chimique de France; vol. 8; (1941); p. 664 - 670
- ↑ Gmelin Handbuch der Anorganischen Chemie; vol. Cl: SVol.B2; 3, page 320 - 322
- ↑ Balard; Annales de Chimie et de Physique; vol. 57; (1834); p. 241 - 241
