Lead dioxide

Lead dioxide
Lead dioxide
CAS number 1309-60-0 YesY
UN number 1872
RTECS number OGO700000
Molecular formula PbO2
Molar mass 239.2 g/mol
Appearance black powder
Density 9.38 g/cm3
Melting point

290 °C decomp.

Solubility in water insoluble
MSDS External MSDS
EU Index 082-001-00-6
EU classification Repr. Cat. 1/3
Harmful (Xn)
Dangerous for the environment (N)
R-phrases R61, R20/22, R33, R62, R50/53
S-phrases S53, S45, S60, S61
NFPA 704
NFPA 704.svg
Flash point Non-flammable
Related compounds
Other cations Carbon dioxide
Silicon dioxide
Germanium dioxide
Tin dioxide
Related lead oxides Lead(II) oxide
Lead(II,IV) oxide
Related compounds Thallium(III) oxide
Bismuth(III) oxide
 YesY (verify) (what is: YesY/N?)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Lead dioxide, PbO2, also plumbic oxide is an oxide of lead in oxidation state +4. It is an odorless dark-brown crystalline powder which is nearly insoluble in water. It exists in two crystalline forms. The alpha phase has orthorhombic symmetry; it has been first synthesized in 1941 and identified in nature as a rare mineral scrutinyite in 1988. On the contrary, the more prevailing, tetragonal beta phase was first identified as the mineral plattnerite around 1845 and later produced synthetically. Lead dioxide is a strong oxidizing agent which is used in the manufacture of matches, pyrotechnics, dyes and other chemicals. It also has several important applications in electrochemistry, in particular as a component of lead acid batteries.




Crystal structure of α-PbO2
Crystal structure of β-PbO2

Lead dioxide is an odorless dark-brown crystalline powder which is nearly insoluble in water.[1] It has two major polymorphs, alpha and beta, which occur naturally as rare minerals scrutinyite and plattnerite, respectively. Whereas the beta form was known already in 1845[2], α-PbO2 was first synthesized in 1941 and identified as a mineral only in 1988. The alpha form has orthorhombic symmetry, space group Pbcn (No. 60), Pearson symbol oP12, lattice constants a = 0.497 nm, b = 0.596 nm, c = 0.544 nm, Z = 4 (four formula units per unit cell).[3] The symmetry of the beta form is tetragonal, space group P42/mnm (No. 136), Pearson symbol tP6, lattice constants a = 0.491 nm, c = 0.3385 nm, Z = 2.[4]

Lead dioxide decomposes upon heating in air as follows:

PbO2 → Pb12O19 → Pb12O17 → Pb3O4 → PbO

The stoichiometry of the end product can be controlled by changing the temperature – for example, in the above reaction, the first step occurs at 290 °C, second at 350 °C, third at 375 °C and fourth at 600 °C. In addition, Pb2O3 can be obtained by decomposing PbO2 at 580–620 °C under oxygen pressure of 1.4 kbar. Therefore, thermal decomposition of lead dioxide is a common industrial way of producing various lead oxides.[5]


Lead dioxide is an amphoteric compound with prevalent acidic properties. It dissolves in strong bases to form the hydroxyplumbate ion, Pb(OH)62−:[1]

PbO2 + 2 NaOH + 2 H2O → Na2[Pb(OH)6]

It also reacts with basic oxides in the melt yielding orthoplumbates M4[PbO4].

Because of the instability of its Pb4+ cation, lead dioxide reacts with warm acids, converting to the more stable Pb2+ state and liberating oxygen:[5]

2 PbO2 + 2 H2SO4 → 2 PbSO4 + H2O + O2
2 PbO2 + 4 HNO3 → 2 Pb(NO3)2 + H2O + O2
PbO2 + 4 HCl → PbCl2 + 2 H2O + Cl2

Lead dioxide is well known for being a good oxidizing agent with example reaction listed below:[6]

2 MnSO4 + 5 PbO2 + 6 HNO3 → 2 HMnO4 + 2 PbSO4 + 3 Pb(NO3)2 + 2 H2O
2 Cr(OH)3 + 10 KOH + 3 PbO2 → 2 K2CrO4 + 3 K2PbO2 + 8 H2O


Although the formula of lead dioxide is nominally given as PbO2, the actual oxygen to lead ratio varies between 1.90 and 1.98 depending on the preparation method. Deficiency of oxygen (or excess of lead) results in the characteristic metallic conductivity of lead dioxide, which can be as low as 10–4 Ohm·cm and which is exploited in various electrochemical applications. Like metals, lead dioxide has a characteristic electrode potential, and in electrolytes it can be polarized both anodically and cathodically. Lead dioxide electrodes have a dual action, that is both the lead and oxygen ions take part in the electrochemical reactions.[7]


Lead dioxide is produced commercially by several methods, which include oxidation of Pb3O4 in alkaline slurry in a chlorine atmosphere,[5] reaction of lead(II) acetate with calcium chloride, or reacting Pb3O4 with dilute nitric acid:[1]

Pb3O4 + 4 HNO3 → PbO2 + 2 Pb(NO3)2 + 2 H2O

An alternative synthesis method is electrochemical: lead dioxide forms on pure lead, in dilute sulfuric acid, when polarized anodically at electrode potential about +1.5 V at room temperature. This procedure is used for large-scale industrial production of PbO2 anodes. Lead and copper electrodes are immersed in sulfuric acid flowing at a rate of 5–10 L/min. The electrodeposition is carried out galvanostatically, by applying a current of about 100 A/m2 for about 30 minutes. The drawback of the lead electrode is its softness, especially compared to the hard and brittle PbO2 which has a Mohs hardness of 5.5.[8] This mismatch in mechanical properties results in peeling of the coating. Therefore, an alternative method is to use harder substrates, such as titanium, niobium, tantalum or graphite and electrodeposit PbO2 on them from lead(II) nitrate in static or flowing sulfuric acid. The substrate is usually sand-blasted before the deposition to remove surface oxide and contamination and to increase the surface roughness and adhesion of the coating.[9]


Lead dioxide is used in the production of matches, pyrotechnics, dyes and the curing of sulfide polymers. It is also used in the construction of high-voltage lightning arresters.[5]

Lead dioxide is used as anode material in electrochemistry. Beta-PbO2 is more attractive for this purpose than the alpha form because it has relatively low resistivity, good corrosion resistance even in low-pH medium, and a high overvoltage for the evolution of oxygen in sulfuric acid and nitric acid based electrolytes. Lead dioxide can also withstand chlorine evolution in hydrochloric acid. Lead dioxide anodes are inexpensive and were once used instead of conventional platinum and graphite electrodes for regenerating potassium dichromate. They were also applied as oxygen anodes for electroplating copper and zinc in sulfate baths. In organic synthesis, lead dioxide anodes were applied for the production of glyoxylic acid from oxalic acid in a sulfuric acid electrolyte.[9]

The most important use of lead dioxide is as the cathode of lead acid batteries. Its utility arises from the anomalous metallic conductivity of PbO2. The lead acid battery stores and releases energy by shifting the equilibrium (a comproportionation) between metallic lead, lead dioxide, and lead(II) salts in sulfuric acid.

Pb + PbO2 + 2 HSO4 + 2 H+ → 2 PbSO4 + 2 H2O, E = +2.05 V


Being a strong oxidant, lead dioxide is a poison when ingested. The associated symptoms include abdominal pain and spasms, nausea, vomiting and headache. Acute poisoning can lead to muscle weakness, metallic taste, loss of appetite, insomnia, dizziness, with shock, coma and death in extreme cases. The poisoning also results in high lead levels in blood and urine. Contact with skin or eyes results in local irritation and pain.[10]


  1. ^ a b c Mary Eagleson (1994). Concise encyclopedia chemistry. Walter de Gruyter. p. 590. ISBN 3110114518. http://books.google.com/?id=Owuv-c9L_IMC&pg=PA590. 
  2. ^ Haidinger W (1845) Zweite Klasse: Geogenide. II. Ordnung. Baryte VII. Bleibaryt. Plattnerit., p. 500 in Handbuch der Bestimmenden Mineralogie Bei Braumüller and Seidel Wien pp. 499-506 (in German)
  3. ^ J. E. Taggard, Jr. et al. (1988). "Scrutinyite, natural occurrence of α-PbO2 from Bingham, New Mexico, U.S.A., and Mapimi, Mexico". Canadian Mineralogist 26: 905. http://rruff.info/uploads/CM26_905.pdf. 
  4. ^ Harada, H.; Sasa, Y.; Uda, M. (1981). "Crystal data for β-PbO2". Journal of Applied Crystallography 14 (2): 141. doi:10.1107/S0021889881008959. 
  5. ^ a b c d Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Oxford: Butterworth-Heinemann. p. 386. ISBN 0080379419. 
  6. ^ Anil Kumar De (2007). A Text Book of Inorganic Chemistry. New Age International. p. 387. ISBN 8122413846. http://books.google.com/?id=PpTi_JAx7PgC&pg=PA387. 
  7. ^ M. Barak (1980). Electrochemical power sources: primary and secondary batteries. IET. pp. 184 ff.. ISBN 0906048265. http://books.google.com/?id=_PGzaO48Rz0C&pg=PA184. 
  8. ^ Plattnerite at Mindat
  9. ^ a b François Cardarelli (2008). Materials Handbook: A Concise Desktop Reference. Springer. p. 573. ISBN 1846286689. http://books.google.com/?id=ArsfQZig_9AC&pg=PA573. 
  10. ^ Lead dioxide MSDS

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