Pyrotechnic colorant

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The red lithium flame leads to lithium's use in flares and pyrotechnics
Copper compounds glow green or blue-green in a flame.
Calcium compounds glow orange in a flame.
Sodium compounds glow yellow in a flame.

A pyrotechnic colorant is a chemical compound which causes a flame to burn with a particular color. These are used to create the colors in pyrotechnic compositions like fireworks and colored fires. The color-producing species are usually created from other chemicals during the reaction. Metal salts are commonly used; elemental metals are used rarely (e.g. copper for blue flames).

The color of the flame is dependent on the metal cation; the anion of the salt has very little direct influence. The anions however influence the flame temperature, both by increasing it (e.g. nitrates, chlorates) and decreasing it (e.g. carbonates, oxalates), indirectly influencing the flame brightness and brilliancy. For temperature-decreasing additives, the limit of colorant may be about 10–20 wt.% of the composition.[1]

Some common examples are:

Color Compound name Chemical formula Notes
Red Strontium nitrate Sr(NO3)2 Common. Used with chlorine donors. Excellent red, especially with metal fuels. Used in many compositions including road flares.
Red Strontium carbonate SrCO3 Common. Produces good red. Slows burning of compositions, decomposes yielding carbon dioxide. Fire retardant in gunpowders. Inexpensive, non-hygroscopic, neutralizes acids. Superior over strontium oxalate in absence of magnesium.
Red Strontium oxalate SrC2O4 Decomposes yielding carbon dioxide and carbon monoxide. In presence of magnesium fuel, carbon monoxide reduces particles of magnesium oxide, yielding gaseous magnesium and eliminating the black body radiation of the MgO particles, resulting in clearer color.
Red Strontium sulfate SrSO4 Common. High-temperature oxidizer. Used in strobe mixtures and some metal-based red compositions.
Red Strontium chloride SrCl2 Common. Produces bright red flame.
Orange Calcium carbonate CaCO3 Produces orange flame. Yields carbon dioxide on decomposition. Often used in toy fireworks as a substitute for strontium.
Orange Calcium chloride CaCl2
Orange Calcium sulfate CaSO4 High-temperature oxidizer. Excellent orange source in strobe compositions.
Orange Hydrated calcium sulfate CaSO4(H2O)x*
Gold/Yellow Charcoal powder C
Gold/Yellow Iron powder with oxygen based carbon OC12 Fe+C
Yellow Sodium bicarbonate NaHCO3 Compatible with potassium chlorate. Less burning rate decrease than sodium carbonate. Incompatible with magnesium and aluminium, reacts evolving hydrogen gas.
Yellow Sodium carbonate Na2CO3 Hygroscopic. Significantly decreases burning rate, decomposes evolving carbon dioxide. Strongly alkaline. Very effective colorant, can be used in small amounts. Corrodes magnesium and aluminium, incompatible with them.
Yellow Sodium chloride NaCl Loses hygroscopicity on heating. Corrodes metals.
Yellow Sodium oxalate Na2C2O4 Non-hygroscopic. Slightly reacts with magnesium, no reaction with aluminium.
Yellow Sodium nitrate NaNO3 Also acts as oxidizer. Bright flame, used for illumination.
Yellow Cryolite Na3AlF6 One of the few sodium salts that is nonhygroscopic and insoluble in water.
Green Barium chloride BaCl2
Green Barium chlorate Ba(ClO3)2 Classic exhibition green with shellac fuel. Sensitive to shock and friction. Oxidizer.
Green Barium carbonate BaCO3 Pretty color when ammonium perchlorate is used as oxidizer.
Green Barium nitrate Ba(NO3)2 Not too strong effect. With chlorine donors yields green color, without chlorine burns white. In green compositions usually used with perchlorates.
Green Barium oxalate BaC2O4
Blue Copper(I) chloride CuCl Richest blue flame. Almost insoluble in water.
Blue Copper(I) oxide Cu2O Lowest cost blue colorant.
Blue Copper(II) oxide CuO Used with chlorine donors. Excellent in composite stars.
Blue Copper carbonate CuCO3 Best when used with ammonium perchlorate.
Blue Basic copper carbonate CuCO3·Cu(OH)2, 2 CuCO3·Cu(OH)2 Occurs naturally as malachite and azurite. Good with ammonium perchlorate and for high-temperature flames with presence of hydrogen chloride. Not easily airborne, less poisonous than Paris Green.
Blue Copper oxychloride 3CuO·CuCl2 Good blue colorant with suitable chlorine donor.
Blue Paris Green Cu(CH3COO)2.3Cu(AsO2)2 Copper acetoarsenite, Emerald Green. Toxic. With potassium perchlorate produces the best blue colors. Non-hygroscopic. Fine powder readily becomes airborne; toxic inhalation hazard. Used in majority of Japanese blue compositions as it gives very pretty color.
Blue Copper arsenite CuHAsO3 Almost non-hygroscopic. Almost as good colorant as copper acetoarsenite. Toxic. Can be used with chlorate oxidizers.
Blue Copper sulfate CuSO4·5 H2O Can be used with nitrates and perchlorates. Acidic, incompatible with chlorates. With red phosphorus in presence of moisture liberates heat, may spontaneously ignite. Less expensive than copper acetoarsenite. Anhydrous copper sulfate is hygroscopic, can be used as a desiccant. With ammonium perchlorate produces almost as pretty blue color as achievable with copper acetoarsenite.
Blue Copper metal Cu Rarely used, other compounds are easier to work with. Yields pretty blue color in ammonium perchlorate based compositions; but reacts with ammonium perchlorate and liberates ammonia in presence of moisture. The composition must be kept dry.
Purple Combination of red and blue compounds Sr+Cu
Purple Rubidium compounds Rb rarely used
Silver/White Aluminium powder Al
Silver/White Magnesium powder Mg
Silver/White Titanium powder Ti
Silver/White Antimony (III) sulfide Sb2S3
Infrared Caesium nitrate CsNO3 two powerful spectral lines at 852.113 nm and 894.347 nm
Infrared Rubidium nitrate RbNO3

The * indicates that the compound will burn orange where x=0,2,3,5.

Radiating species[edit]

Despite the wide numbers of metal ion donors, they serve to form only a few atomic and molecular species that are useful as light emitters.[2]

In many cases, chlorine donors have to be added in order to achieve sufficiently deep colors, as the desired emitting molecules have to be generated.

Some color emitters are of atomic nature (e.g. lithium, sodium). Presence of chlorine, and the reaction to monochlorides, may actually impair their color purity or intensity.

At high temperatures, the atoms will ionize. The emission spectra of ions are different than of neutral atoms; the ions may emit in undesired spectral ranges. For example, Ba+ emits in blue wavelengths. Ionization can be suppressed by addition of an easier-to-ionize metal with weak visible emission of its own, e.g. potassium; the potassium atoms then act as electron donors, neutralizing the barium ions.[3]

The color blue is notoriously difficult to produce in fireworks, as the copper compounds need to be heated at a specific temperature for the optimal shade of blue to be produced. Thus, a deep, rich blue is usually viewed as the mark of an experienced fireworks maker.

Care should be taken to avoid formation of solid particles in the flame zone, whether metal oxides or carbon; incandescent solid particles emit black body radiation that causes "washing out" of the colors. Addition of aluminium raises the flame temperature but also leads to formation of solid incandescent particles of aluminium oxide and molten aluminium. Magnesium has less such effect and is therefore more suitable for colored flames; it is more volatile than aluminium and more likely to be present as vapors than as particulates. Formation of solid particles of magnesium oxide can further be inhibited by presence of carbon monoxide, either by negative oxygen balance of the composition in presence of organic fuels, or by addition of the colorant in the form of an oxalate, which decomposes to carbon dioxide and carbon monoxide; the carbon monoxide reacts with the magnesium oxide particles to gaseous magnesium and gaseous carbon dioxide.

Color Emitter Wavelengths Notes
Yellow Sodium (D-line) 589 nm Very strong, overpowers other colors, avoid contamination
Orange CaCl (molecular bands) most intense: 591–599 nm and 603–608 nm, and others
Red SrCl (molecular bands) a: 617–623 nm
b: 627–635 nm
c: 640–646 nm
The SrCl species tends to be oxidized to less desirable SrO; strontium-containing compositions are therefore usually formulated to be oxygen-deficient.[3]
Red SrOH(?) (molecular bands) 600–613 nm
Red Li (atomic spectral lines)
Green BaCl (molecular bands) a: 511–515 nm
b: 524–528 nm
d: 530–533 nm
Lines of BaOH and BaO are also present, emitting in yellow and yellowish-green (487, 512, 740, 828, and 867 nm for BaOH, 549, 564, 604 and 649 for BaO). The BaOH lines are much stronger than the BaO lines. In absence of chlorine, the BaCl lines are not present and only the BaOH and BaO lines are visible.


The BaCl species tends to be oxidized to less desirable BaO; barium-containing compositions are therefore usually formulated to be oxygen-deficient.
Presence of Ba+ is undesired, as it emits in a blue region at 455.4 nm. Potassium may be added to suppress barium ionization, as it ionizes easier and acts as an electron donor for the barium ions.[3]

Blue CuCl (molecular bands) several intense bands between 403–456 nm, less intense at 460–530 nm Low dissociation energy of copper compounds causes presence of free copper atoms in the flame, weakly emitting in green (lines between 325–522 nm). In presence of chlorine, CuCl is formed, emitting strongly in blue. At higher temperatures CuCl dissociates and lines of atomic copper are present in the spectrum; CuO and CuOH are also formed, emitting molecular bands at green-yellow (535–555 nm) for CuOH and at orange-red (580–655 nm) for CuOH. Adequate control of temperature is therefore required for blue-burning compositions.
Infrared Carbon particles black body radiation For good broadband infrared output, compositions producing lots of heat and carbon particles are required. The burning temperature should be lower than of visible-illuminating compounds. The intensity of the emitted radiation depends on the burn rate. Temperature can be increased by addition of magnesium. A magnesium/Teflon/Viton composition is common for missile decoy flares.[4]
Infrared CO2 (molecular bands) mostly 4300 nm Produced by carbon-containing fuels.
Infrared Cs (atomic spectral lines) two powerful spectral lines at 852.113 nm and 894.347 nm Used in infrared illumination compositions. Metal is avoided in the compositions to prevent formation of bright, visible-radiating particles.[5]
Infrared Rb (atomic spectral lines) spectral lines in near-infrared Used in infrared illumination compositions, less commonly than cesium.

References[edit]

  1. ^ B. J. Kosanke et al. Pyrotechnic Chemistry. Volume 4 of Pyrotechnic reference series, Journal of Pyrotechnics, 2004 ISBN 1889526150, p. 30
  2. ^ "The Physics of Coloured Fireworks". Cc.oulu.fi. Retrieved 2010-03-23. 
  3. ^ a b c Michael S. Russell The chemistry of fireworks, Royal Society of Chemistry, 2009 ISBN 0-85404-127-3, p. 85
  4. ^ Jai Prakash Agrawal High Energy Materials: Propellants, Explosives and Pyrotechnics, Wiley-VCH, 2010 ISBN 3-527-32610-3, p. 349
  5. ^ B. J. Kosanke et al. Pyrotechnic Chemistry, Journal of Pyrotechnics, 2004 ISBN 1889526150, p. 58