Potassium nitrate
 KNO3 recrystallized from aqueous solution.
| |
| Names | |
|---|---|
| IUPAC name
Potassium nitrate
| |
| Other names
Chinese salt
Chinese snow Indian saltpeter Nitrate of potash Saltpeter Saltpetre | |
| Properties | |
| KNO3 | |
| Molar mass | 101.1032 g/mol |
| Appearance | White solid |
| Odor | Odorless |
| Density | 2.1 g/cm3 (20 °C) |
| Melting point | 334 °C (633 °F; 607 K) |
| Boiling point | 400 °C (752 °F; 673 K) (decomposes) |
| 13.3 g/100 ml (0 °C) 31.6 g/100 ml (20 °C) 246 g/100 ml (100 °C) | |
| Solubility | Soluble in ammonia, glycerol Slightly soluble in ethanol, ethylenediamine Insoluble in diethyl ether, isopropanol, toluene |
| Solubility in ethanol | 0.6 g/100 ml (20 °C) |
| Solubility in hydrazine | 14 g/100 ml (20 °C) |
| Solubility in methanol | 3 g/100 ml (20 °C) |
| Vapor pressure | ~0 mmHg |
| Thermochemistry | |
| Std enthalpy of
formation (ΔfH |
-494.0 kJ/mol |
| Hazards | |
| Safety data sheet | Sigma-Aldrich |
| Flash point | Non-flammable |
| Lethal dose or concentration (LD, LC): | |
| LD50 (Median dose)
|
1,901 mg/kg (rabbit, oral) 3,750 mg/kg (rat, oral) |
| Related compounds | |
| Related compounds
|
Lithium nitrate Sodium nitrate Rubidium nitrate Calcium nitrate |
| Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa). | |
| Infobox references | |
Potassium nitrate, also known as saltpeter or saltpetre or Indian saltpeter, is an important chemical compound in both industry and chemistry. Known for centuries for its role in early pyrotechnics, the nitrate ion makes KNO3 a widely used oxidizer with a variety of uses inside and outside pyrotechnics.
Contents
Properties
Chemical
Potassium nitrate is a good source of both potassium and nitrate ions.
When heated to to temperatures between 550 and 790 °C, under an oxygen atmosphere, it loses oxygen and converts into potassium nitrite:
- 2 KNO3 → 2 KNO2 + O2
Molten potassium nitrate will react with carbon dioxide and water vapors to form potassium carbonate and hydroxide.[1]
Potassium nitrate will react with hydrochloric acid to release nitric acid, that will give off nitrogen dioxide fumes:
- 2 KNO3 + 2 HCl + H2O → 2 KCl + HNO3 + NO2 + H2O
By adding an excess of HCl, this reaction can be used to generate "poor man's aqua regia".
A substantial amount of energetic materials are based on the nitrate ion, and it is arguably the most important ion when it comes to explosives, such as black powder.
- 10 KNO3 + 3 S + 8 C → 2 K2CO3 + 3 K2SO4 + 6 CO2 + 5 N2
A less known reaction is the synthesis of potassium cyanide, by reacting a mixture of potassium nitrate and charcoal in a cast iron bowl, in an inert atmosphere to prevent combustion or oxidation to potassium cyanate:
- 2 KNO3 + 7 C → KCN + KCNO + 5 CO[2]
If you attempt to try this reaction, AVOID ADDING ACID TO THE RESULTED SLAG AS IT WILL GIVE OFF HYDROGEN CYANIDE GAS WHICH CAN BE DEADLY (see the Sciencemadness thread below).
Presence of iron may also lead to formation of Prussian blue.
A mixture of sucrose (most sugars work) and potassium nitrate known as rock candy can be ignited, producing pink flames and large amounts of white smoke, useful for improvised smoke bombs. The solid end product of this reaction is potassium carbonate, plus some of the leftover reactants.
- KNO3 + C12H22O11 → K2CO3 + CO2 + N2
The ignition also releases copious amounts of caramelized sugar particles, giving the smoke a sweet fragrance.
Physical
Potassium nitrate is white solid with a melting point of 334 °C. Its solubility curve makes recrystallization easy, being only somewhat soluble in freezing water (13.3 g/100 ml at 0 °C), but very soluble in boiling water (246 g/100 ml at 100 °C). It is not hygroscopic, absorbing about 0.03% water in 80% relative humidity over a period of 50 days.
Potassium nitrate is soluble in anhydrous basic solvents, like ammonia and hydrazine, as well as glycerol, but only slightly soluble in ethanol and methanol. Pure KNO3 is odorless and has a slight bitter taste.
Availability
In many American hardware stores, potassium nitrate is sold in a reasonably pure form as a stump remover, which can be purified by recrystallization from water. Care must be taken if it is being bought this way, as some stump removers mostly consist of other agents, such as sodium metabisulfite, which can be a rude awakening to one treating it as potassium nitrate.
It can also be bought as fertilizer, either pure or mixed with traces of other substances. If it's present as prills, recrystallization might be required as the granules are difficult to grind properly in a mortar. Some types of prills tend to be covered with an anti-caking agent, such as calcium stearate, and filtration before recrystallization is required for purification. It's easy to determine if the saltpeter source has anticaking agent, as when the solution gets more concentrated solution, it will darken.
Potassium nitrate occurs naturally as the mineral niter (or nitre).
Preparation
Potassium nitrate can be prepared by neutralizing KOH with nitric acid or adding ammonium nitrate to it.
- KOH + HNO3 → KNO3 + H2O
- KOH + NH4NO3 → KNO3 + NH3 + H2O
It can also be made by a double displacement reaction between ammonium nitrate and potassium sulfate or potassium chloride, and then recrystallizing the KNO3 out at low temperatures. This reaction can also be done using potassium sulfate and sodium nitrate, but instead of crystallizing potassium nitrate, sodium sulfate is crystallized out along with some of the potassium nitrate at low temperatures, leaving most of the desired product in solution.
A reaction between a small amount of ice or water, potassium hydroxide, and ammonium nitrate easily proceeds to completion very exothermically, driving ammonia gas out of the mixture and producing potassium nitrate as an end product. This can also be done with potassium carbonate, but the carbonate route requires heating the two in solution.
Another method is a double displacement from barium nitrate with potassium carbonate or potassium sulfate. The resulting barium carbonate or sulfate can be filtered off, leaving potassium nitrate in solution.
The nitrate ion itself is very difficult to make from other OTC chemicals, and despite the increasing difficulty in obtaining this salt, buying it in pure or impure from is the only real viable option.
An old method of making saltpeter involves separating nitrates from manure, and adding potash. The resulting solution is recrystallized repeatedly until fairly pure saltpeter is obtained. This method requires large amounts of manure, but can be done fairly easy if one has a farm or a big garden.
Projects
- Making nitric acid
- Potassium nitrite synthesis
- Nitrocellulose synthesis
- Black powder
- Candy rocket
- Various flash powders
- Dissolving gold when mixed with concentrated HCl (poor man's aqua regia)
- Electrolyte in a salt bridge
- Potassium cyanide synthesis
- Grow large crystals
Handling
Safety
Potassium nitrate is virtually non-toxic, and is even approved for use in food as a preservative (E252). Soldiers of the early modern era even used the saltpeter-based black powder as a table salt substitute to sprinkle on their rations.
While generally quite stable, it is an oxidizing agent, so mixtures with reducing agents like sucrose, red phosphorus, sulfur, magnesium or aluminium should not be ground or pressed together, and the risk of spontaneous ignition is ever present. Sulfur is especially dangerous, as acidic traces can greatly increase the sensitivity of energetic mixtures. As it is not hygroscopic, like its sodium cousin, it poses a greater fire hazard.[3]
Storage
Potassium nitrate should be stored in closed bottles, away from any flammable materials as well as organic and acidic vapors. It is not hygroscopic, so there's no need to seal it.
Disposal
Potassium nitrate is a good fertilizer, and unless it's contaminated with heavy metals or other toxicants, it can be disposed pretty much anywhere.
Gallery
Crystals of potassium nitrate, dried
Potassium nitrate, fresh out of solution
References
- ↑ Proceedings of the Symposia on Corrosion in Batteries and Fuel Cells and Corrosion in Solar Energy Systems, Chris J. Johnson, Steven L. Pohlman, 1984, p. 407-408
- ↑ http://www.google.com/patents/US579988
- ↑ Tdep - experience working with and setting fire to said chemical.
