Alkali Metals Sodium, Potassium, NaK, and Lithium



At room temperature sodium oxidizes rapidly in moist air, but spontaneous ignitions have not been reported except when the sodium is in a finely divided form. When heated in dry air, sodium ignites in the vicinity of its boiling point (880 degrees C or 1,616 degrees F). Sodium in normal room air and at a temperature only slightly above its melting point (98 degrees C or 208 degrees F) has been ignited by placing sodium oxide particles on its surface. This indicates the possibility of ignition at temperatures below the boiling point. Once ignited, hot sodium burns vigorously and forms dense white clouds of caustic sodium oxide fumes. During combustion, sodium generates about the same amount of heat as an equivalent weight of wood.

The principal fire hazard associated with sodium is the rapid reaction with water. It floats on water (specific gravity 0.97), reacting vigorously and melting. The hydrogen liberated by this reaction may be ignited and explode from the heat of the reaction. Sodium (like other burning, reactive metals) reacts violently with most chemicals that can oxidize, halogenated hydrocarbons, with halogens such as iodine, and with sulfuric acid.


The fire hazard properties of potassium (K) are very similar to those of sodium with the difference that potassium is usually more reactive. For example, the reaction between potassium and the halogens is more violent, and, in the case of bromine, a deflagration/detonation can occur. There is an explosive reaction with sulfuric acid. Unlike sodium, potassium forms some peroxides and superoxides during combustion. These peroxides may react violently with organic contaminants (oils, etc.). Peroxides and superoxides may also react explosively with metallic potassium (see NaK below).

NaK (Sodium-Potassium Alloys)

NaK is the term used when referring to any of several sodium-potassium alloys. The various NaK alloys differ from each other in melting point, but all are liquids or melt near room temperature. NaK alloys possess the same fire hazard properties as those of the component metals except that the reactions are more vigorous. Under pressure, NaK leaks have ignited spontaneously.

The potassium in NaK will react with atmospheric oxygen to form three different oxides, potassium oxide (K{sub 2}O), potassium peroxide (K{sub 2}O{sub 2}), and potassium superoxide (KO{sub 2}). These oxides form a crust over the NaK surface. If this crust is permeated and the superoxide (KO{sub 2}) is allowed to mix with the potassium in the NaK, a very high temperature thermite-type reaction can occur. This reaction may take several minutes to develop if the NaK is stored under an inert atmosphere or it may occur instantly if stored under atmospheric oxygen.


Lithium, like sodium and potassium, cesium, and rubidium is one of the alkali metals. Lithium undergoes many of the same reactions as sodium. For example, both sodium and lithium react with water to form hydrogen; but whereas the sodium-water reaction can generate sufficient heat to ignite the hydrogen, the far less violent lithium-water reaction does not. Lithium ignites and burns vigorously at a temperature of 180 degrees C (356 degrees F), which is near its melting point. Unlike sodium and potassium, it will burn in nitrogen. The caustic (oxide and nitride) fumes accompanying lithium combustion are more profuse and dense than those of other alkali metals burning under similar conditions. Lithium is the lightest of all metals. During combustion, it tends to melt and flow.

Storage and Handling

Because of their reactivity with water, alkali metals require special precautions to prevent contact with moisture. Drums and cases containing alkali metals should be stored in a dry, fire-resistive room or building used exclusively for alkali metal storage. Since sprinkler protection would be undesirable, no combustible materials should be stored in the same area. It is good practice to store empty as well as filled alkali metal containers in the same area, and all containers should be on skids. There should be no water or steam pipes, but sufficient heat should be maintained to prevent moisture condensation caused by atmospheric changes. Natural ventilation at a high spot in the room is desirable to vent any hydrogen that may be released by accidental contact of alkali metal with moisture.

Large quantities of alkali metal are often stored outdoors in aboveground tanks. In such installations weatherproof enclosures should cover tank manholes, and the free space within the tank should contain a nitrogen atmosphere. Argon or helium atmospheres should be substituted for nitrogen in the case of lithium.

For small-scale transfer of solid alkali metal from a storeroom to the use area, a metal container with a tight cover is recommended. Alkali metal should be removed from storage in as small quantities as practicable. When stored on work benches, it should be kept under kerosene or oil in a closed container. Alkali metal, with its great affinity for moisture, may react at the time it is sealed in a container with any atmospheric moisture. Because of the possible presence of hydrogen, containers should not be opened by hammering on the lid.

Process Hazards

Liquid alkali metal is valuable as a high temperature heat transfer medium. For example, it is used in hollow exhaust valve stems in some internal combustion engines and in the transfer of heat from one type of nuclear reactor to a steam generator. In the latter process or other large-scale use of molten alkali metal, any equipment leak may result in a fire. Where molten alkali metal is used in process equipment, steel pans should be located underneath to prevent contact with concrete floors. Contact of molten alkali metal with concrete will cause spalling of the concrete and spattering of the metal.

Processing of alkali metal is essentially remelting it to form sticks or bricks or to add as a liquid to closed transfer systems. During this handling, contact with moist air, water, halogens, halogenated hydrocarbons, and sulfuric acid must be avoided.

Extinguishing Fires in Sodium, Lithium, NaK, and Potassium

The common extinguishing agents, such as water, foam, and vaporizing liquids, should never be used because of the violent reactions upon application to alkali metals. Class D dry powders developed for metal fires, dry sand, dry sodium chloride, and dry soda ash are effective. These finely divided materials blanket the fire while the metal cools to below its ignition temperature. Alkali metal burning in an apparatus can usually be extinguished by closing all openings. Blanketing with nitrogen is also effective. In the case of lithium, argon or helium atmospheres should be used.

Zirconium and Hafnium



The combustibility of zirconium increases as the average particle size decreases, but other variables, such as moisture content, also affect its ease of ignition. In massive form, zirconium can withstand extremely high temperatures without igniting, whereas clouds of dust in which the average particle size is 3 microns have ignited at room temperature. Dust clouds of larger particle size can be readily ignited if an ignition source is present, and such explosions can occur in atmospheres of carbon dioxide or nitrogen as well as in air. Zirconium dust will ignite in carbon dioxide at approximately 621 degrees C (1,150 degrees F) and nitrogen at approximately 788 degrees C (1,450 degrees F). Tests have also indicated that layers of 3-micron-diameter dust are susceptible to spontaneous ignition. The depth of the dust layer and its moisture content are important variables for ignition. Spontaneous heating and ignition are also possibilities with scrap chips, borings, and turnings if fine dust is present. Layers of 6-micron-diameter dust have ignited when heated to 190 degrees C (374 degrees F). Combustion of zirconium dust in air is stimulated by the presence of limited amounts of water (5 to 10%). When very finely divided zirconium powder is completely immersed in water, it is difficult to ignite, but once ignited it burns more violently than in air.

Massive pieces of zirconium do not ignite spontaneously under ordinary conditions, but ignition will occur when an oxide-free surface is exposed to sufficiently high oxygen concentrations and pressure. The explanation for this reaction is the same as that cited for a similar titanium reaction. Zirconium fires (like fires involving titanium and hafnium) attain very high temperatures, but generate very little smoke.

Explosions have occurred while zirconium was being dissolved in a mixture of sulfuric acid and potassium acid sulfate. Zirconium has exploded during and following pickling in nitric acid, and also during treatment with carbon tetrachloride or other halogen-containing materials. Spontaneous explosions have occurred during handling of moist, very finely divided, contaminated zirconium scrap.


Hafnium has similar fire properties to zirconium. Hafnium burns with very little flame, but it releases large quantities of heat. Hafnium in sponge form may ignite spontaneously.

Hafnium is generally considered to be somewhat more reactive than titanium or zirconium of similar form. Damp hafnium powder reacts with water to form hydrogen gas, but at ordinary temperatures this reaction is not sufficiently vigorous to cause the hydrogen to ignite. At higher temperatures, however, ignition of the hydrogen may be expected to proceed explosively.

Storage and Handling

Special storage precautions are not required for zirconium castings because of the very high temperatures that massive pieces of the metal can withstand without igniting. Zirconium powder, on the other hand, is highly combustible; consequently, it is customarily stored and shipped in 3.78-L (1-gal) containers with at least 25% water by volume. For specific details, refer to NFPA 482, Standard for the Production, Processing, Handling, and Storage of Zirconium.

Zirconium powder storerooms should be of fire-resistive construction equipped with explosion vents. Cans should be separated from each other to minimize the possibility of a fire at one can involving others and to permit checking of the cans periodically for corrosion. One plant handling zirconium has established the procedure of disposing of cans containing powder that have been on the shelf for 6 months.

Process Hazards

In general, processing recommendations for zirconium and hafnium are the same. Handling of zirconium powder, whenever possible, should be under an inert liquid or in an inert atmosphere. If zirconium or hafnium powder is handled in air, extreme care must be used because the small static charges generated may cause ignition.

To prevent dangerous heating during machining operations, a large flow of mineral oil or water-base coolant is required. In some machining operations, the cutting surface is completely immersed. Turnings should be collected frequently and stored under water in cans. Where zirconium dust is a byproduct, dust collecting equipment which discharges into a water precipitation type of collector is a necessity.

Extinguishing Fires in Zirconium and Hafnium

Zirconium and hafnium fires can be extinguished in the same way. Fires exposing massive pieces of zirconium, for example, can be extinguished with water. Limited tests conducted by Industrial Risk Insurers have indicated that the discharge of water in spray form would have no adverse effect on burning zirconium turnings. When a sprinkler opened directly above an open drum of burning zirconium scrap, there was a brief flareup after which the fire continued to burn quietly in the drum. When a straight stream of water at a high rate of flow was discharged into the drum, water overflowed and the fire went out.

Where small quantities of zirconium powder or fines are burning, the fire can be ringed with a Class D extinguishing powder to prevent its spread, after which the fire can be allowed to burn out. Special powders developed for metal fires have been effective in extinguishing zirconium fires. When zirconium dust is present, the extinguishing agent should be applied so that a zirconium dust cloud will not form. If the fire is in an enclosed space, it can be smothered by introducing argon or helium.

Calcium and Zinc



The flammability of calcium depends considerably on the amount of moisture in the air. If ignited in moist air, it burns without flowing at a somewhat lower rate than sodium. It decomposes in water to yield calcium hydroxide and hydrogen, which may burn. Finely divided calcium will ignite spontaneously in air. It should be noted that barium and strontium are very similar to calcium in their fire properties.


Zinc does not introduce a serious fire hazard in sheets, castings, or other massive forms because of the difficulty of ignition. Once ignited, however, large pieces burn vigorously. Moist zinc dust reacts slowly with the water to form hydrogen, and, if sufficient heat is released, ignition of the dust can occur. Zinc dust clouds in air ignite at 599 degrees C (1,110 degrees F). Burning zinc generates appreciable smoke.

Storage, Processing, and Extinguishing Fires in Calcium and Zinc

The storage, handling, and processing recommendations for magnesium are generally applicable to calcium and zinc.

Metals Not Normally Combustible


The usual forms of aluminum have a sufficiently high ignition temperature so that its burning is not a factor in most fires. However, very fine chips and shavings are occasionally subject to somewhat the same type of combustion as described for magnesium. Powdered or flaked aluminum in its pure form can ignite spontaneously in air and can be explosive in air.

Iron and Steel

Iron and steel are not usually considered combustible; in a massive form (as in structural steel, cast iron parts, etc.), they do not burn in ordinary fires. Steel in the form of fine steel wool or dust may be ignited in the presence of heat from, for example, a torch, yielding a form of sparking rather than actual flaming in most instances. Fires have been reported in piles of steel turnings and other fine scrap which presumably contained some oil and were perhaps also contaminated by other materials that facilitated combustion. Spontaneous ignition of water-wetted borings and turnings in closed areas, such as ship hulls, has also been reported. Pure iron has a melting point of 1,535 degrees C (2,795 degrees F). Ordinary structural steel has a melting point of 1,430 degrees C (2,606 degrees F).


Plutonium is one of the most widely used pyrophoric materials in the DOE Complex. Some of the most serious fires occurring within the Complex are caused by the ignition of finely divided plutonium particles. Several plutonium compounds are pyrophoric. The radioactive decay of plutonium creates additional concerns such as dispersal of particles in a fire, pressurization of storage containers, and the production of hydrogen gas during decomposition of absorbed water.


Metal, Oxides, and Oxidation

Large pieces of plutonium metal react slowly with the oxygen in air at room temperature to form plutonium oxides. The rate of oxidation is dependent on a number of factors. These include (a) temperature, (b) surface area of the reacting metal, (c) oxygen concentration, (d) concentration of moisture and other vapors in the air, (e) the type and extent of alloying, and (f) the presence of a protective oxide layer on the metal surface. The rate of oxidation increases with increases in the first four factors and decreases with the last. Alloying can either increase or decrease the oxidation rate, depending on the alloying metal. Of all these factors, moisture has a large effect on the oxidation rate and is especially significant in evaluating conditions for storing plutonium metal and oxide.

Several plutonium oxides can be formed from oxidation of metal or decomposition of plutonium compounds. Oxide phases corresponding to sesquioxide (Pu{sub 2}O{sub 3}) and dioxide (PuO{sub 2}) compositions have been identified and are well characterized. Pu{sub 2}O{sub 3} is pyrophoric in air and rapidly forms plutonium dioxide while releasing heat. The dioxide is unreactive in air, but reportedly heats slowly with water vapor at elevated temperatures.


Plutonium hydride (PuH{sub x}, 2 < x < 3) forms during corrosion of plutonium metal by hydrogen from water, organic materials, and other sources. Hydride is rapidly oxidized by dry air at room temperature to produce PuO{sub 2} and H{sub 2} and reacts with nitrogen at 250 degrees C to form plutonium nitride (PuN). The quality of hydride produced depends on the rate of hydrogen formation and on the magnitude of the hydrogen- containing source. The reactivity of plutonium hydride in air depends on factors such as particle size, presence/absence of protective oxide layer, and the hydrogen:plutonium ratio, x. Finely divided hydride is pyrophoric in air at room temperature. Thus, the only safe practice is to handle and store hydride in a dry, oxygen-free atmosphere.

Carbides and Nitride

Plutonium carbides, oxycarbides, and nitride are reactive and potentially pyrophoric materials that could pose handling problems if exposed to air or oxygen-containing atmospheres. These compounds react readily with moisture to form gaseous products such as methane, acetylene, and ammonia. Because plutonium compounds of this type have been prepared at several sites and may have been "temporarily" stored under special conditions (hermetically sealed within an inert atmosphere) without first oxidizing them, caution should be exercised in opening cans that might contain such materials.

Reactions Involving Water

Water vapor accelerates the oxidation of plutonium by oxygen and reacts directly with the metal. Oxidation is about ten times higher in humid air than in dry air at room temperature. For this reason, plutonium metal has routinely been handled in a very dry atmosphere such as one with a -40 degrees C dew point. Inerting of glove boxes and enclosures for handling plutonium with nitrogen or argon is effective in reducing metal oxidation only if it also excludes water vapor. Rapid oxidation does not occur if oxygen is present at a level of 5% in nitrogen or argon. However, if 1.3% moisture (50% relative humidity) accompanies the oxygen, then rapid metal oxidation can be anticipated.

Plutonium dioxide can adsorb up to 8% of its weight as water on the surface. The quantity absorbed is a direct function of the surface area of the oxide. The principal hazard associated with absorbed water is pressurization of a sealed oxide container through any of several separate processes including evaporation of water, radiolysis to form oxygen and hydrogen, or direct reaction with the oxide to form a higher oxide and hydrogen gas.

Pressurization of oxide containers can be prevented by use of sealed containers fitted with durable, high-efficiency metal filters. Although gases can escape without release of plutonium-containing particles, air (possibly moist) is able to enter the container.


When heated to its ignition temperature, plutonium reacts at an accelerated oxidation rate, which sustains continued oxidation. The burning temperature depends on the rate of heat dissipation to the surroundings and the rate of heat generation, which is dependent on the surface area of oxidizing metal. Temperatures of plutonium fires usually exceed the melting temperature of plutonium metal (640 degrees C) which causes the material to consolidate into a molten configuration. As such, finely divided metal, turnings, and casting skulls tend to ignite readily and achieve a high initial temperature which lasts until melting occurs and the surface area is reduced.

The oxide layer that forms during burning limits the oxidation rate of plutonium. The burning process is similar to that of a charcoal briquette. The ignition temperature of plutonium metal depends on the factors that increase the oxidation rate. Finely divided plutonium metal, such as metal powder of fine machine turnings, ignites near 150 degrees C. This temperature is easily reached if a coexisting pyrophoric material such as a hydride spontaneously ignites at room temperature. Bulk or massive plutonium characterized as having a specific surface area less than 10 cm{sup 2}/g requires temperatures in excess of 400 degrees C to ignite. Many plutonium fires have occurred because samples containing finely divided metal have spontaneously ignited. Fires have not occurred with well- characterized metal existing in large pieces that have higher ignition temperatures. Thus, massive plutonium is not considered pyrophoric or capable of self-ignition.

An investigation of two instances in which kilogram-sized plutonium pieces were observed to "spontaneously ignite" in air at room temperature showed that they had been exposed to unlimited sources of hydrogen for extended periods, and that the samples were thermally insulated when ignition occurred. The amount of hydride present on these massive pieces apparently generated sufficient heat to cause ignition. These observations emphasize the need for well-characterized materials.

Storage and Handling

Plutonium should be stored as pure metal (Pu) or in its dioxide (PuO{sub 2}) form in a dry, inert or slightly oxidizing atmosphere. The formation of oxide from metal is accompanied by a large volume expansion (up to 70%) which may bulge or breach the primary container. Case studies show that mechanical wedging resulting from this expansion can even breach a second metal container, resulting in localized contamination release and possible exposure of personnel. Oxidation of the metal and rupture of the container by mechanical wedging are prevented if the storage container is hermetically sealed. Plutonium radioactively decays producing alpha particles and helium molecules. Over long-term storage, helium buildup can contribute to the pressurization of containers.

Plutonium sesquioxide and hydride should be converted to plutonium dioxide before storage. Primary and secondary containers should be hermetically sealed and contain no plastics or other materials that decompose as a result of radiation exposure.

For a more complete discussion of plutonium storage issues, refer to DOE/DP-123T, Assessment of Plutonium Storage Issues at Department of Energy Facilities, January 1994.

Extinguishing Plutonium Fires

Plutonium fires should not be approached without protective clothing and respirators unless the fire is enclosed in a glove box. The most effective agent for extinguishing plutonium fires has been found to be magnesium oxide sand. Glove boxes which contain pyrophoric forms of plutonium should also contain an amount of magnesium oxide adequate for extinguishment. The burning plutonium should be completely covered with the sand to as great a depth as possible. The magnesium oxide extinguishes the fire by providing a heat sink which cools the plutonium and by providing a barrier which limits the availability of oxygen.

Argon is a very effective extinguishing agent, providing the oxygen content in the atmosphere is maintained at 4% or less. Above 4% oxygen, flooding with argon will not extinguish a plutonium fire. This is an important point, since it is nearly impossible to reduce the oxygen content to 4% or less during argon flooding in most fume hoods. Argon may be used effectively to cool the burning plutonium prior to application of the magnesium oxide sand.

Other agents have been tested for use on plutonium fires; however, none have proven to be as effective as magnesium oxide. Typical foam or dry chemical agents are not effective extinguishing agents. Fusible salt agents have been shown to be effective on small-scale plutonium fires. However, the expansion which accompanies the oxidation of plutonium has caused the fusible salt coating to crack, allowing the plutonium to re-ignite.

Water is generally acceptable for use as an extinguishing agent for fires involving plutonium. In rare cases where criticality safety considerations preclude the introduction of moderators such as water, suitable alternative fire protection measures need to be incorporated into the facility design. Proper housekeeping which includes removal of combustibles from pyrophoric forms of plutonium is the most important aspect of fire loss minimization.

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