Lead dioxide

Lead dioxide
IUPAC name Lead(IV) oxide
Other names Plumbic oxide Plattnerite
Identifiers
CAS number	1309-60-0 Y
UN number	1872
RTECS number	OGO700000
Properties
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
Hazards
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	0 3 1 OX
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
Y (verify) (what is: Y/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.

Properties

Physical

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]

Chemical

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

Electrochemical

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]

Production

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]

Applications

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

Safety

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]

References

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

External links

• National Pollutant Inventory: Lead and Lead Compounds Fact Sheet