Lead(IV) oxide

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Lead(IV) oxide
Names
IUPAC name
Lead(IV) oxide
Other names
Lead dioxide
Lead peroxide
Lead superoxide
Plattnerite (β-PbO2)
Plumbic oxide
Scrutinyite (α-PbO2)
Properties
PbO2
Molar mass 239.1988 g/mol
Appearance Dark brown or brown-black
Odor Odorless
Density 9.773 g/cm3 (α-PbO2)[1]
9.55 g/cm3 (β-PbO2)[2]
Melting point 290 °C (554 °F; 563 K) (decomposes)
Boiling point Decomposes
Insoluble
Solubility Reacts with mineral acids
Soluble in acetic acid, alkalis
Insoluble in alcohols, ketones, hydrocarbons
Vapor pressure ~0 mmHg
Thermochemistry
76.6 J·mol-1·K-1
-270 kJ/kmol
Hazards
Safety data sheet Sigma-Aldrich
Flash point Non-flammable
Related compounds
Related compounds
Lead(II) oxide
Lead(II,IV) oxide
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
Infobox references

Lead(IV) oxide also called plumbic oxide, anhydrous plumbic acid (sometimes wrongly called lead peroxide), though more commonly called lead dioxide, is a chemical compound with the formula PbO2. It is an oxide where lead is in an oxidation state of +4 and the bond type is predominantly covalent.

Properties

Chemical

Lead dioxide decomposes upon heating in air as follows:

PbO2 → Pb12O19 → Pb12O17 → Pb3O4 → PbO

Lead dioxide is an amphoteric compound with prevalent acidic properties. It dissolves in strong bases to form the hydroxyplumbate ion, [Pb(OH)6]2−:

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 hot acids, converting to the more stable Pb2+ state and liberating oxygen, though the reaction is slow:

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

Lead dioxide is well known for being a good oxidizing agent, with an example reaction listed below:[3]

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[4]

Physical

Lead dioxide is an odorless dark-brown crystalline powder, practically insoluble in water. It exists in two crystalline forms:

  • The alpha phase (α-PbO2): 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). The lead atoms are six-coordinate.
  • The beta phase (β-PbO2): has tetragonal symmetry, space group P42/mnm (No. 136), Pearson symbol tP6, lattice constants a = 0.491 nm, c = 0.3385 nm, Z = 2[6] and related to the rutile structure and can be envisaged as containing columns of octahedra sharing opposite edges and joined to other chains by corners.

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, with a resistivity as low as 10−4 Ω·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.[5]

Availability

Lead(IV) oxide is sold by many chemical suppliers.

It's encountered in nature as the minerals plattnerite (β-PbO2) and scrutinyite (α-PbO2).

Preparation

Chemical

Lead(IV) oxide can be produced chemically by several methods, such as oxidation of Pb3O4 with chlorine in alkaline medium or with dilute nitric acid:

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

Oxidizing Pb+2 salts, such as lead(II) acetate or lead(II) chloride, with calcium hypochlorite and sodium hypochlorite will also yield lead dioxide:

Pb(CH3COO)2 + Ca(ClO)2 → PbO2 + Ca(CH3COO)2 + CaCl2
PbCl2 + 2 NaClO → PbO2 + 2 NaCl + Cl2

Due to the decomposition of NaOCl to NaOH, stoichiometric amounts of PbO2 react with NaOH to form the hexahydroxoplumbate(IV) ion [Pb(OH)6]2−, soluble in water.

Electrochemical

Lead dioxide forms on pure lead metal, 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.[6] 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 deposit PbO2 onto them from lead(II) nitrate in static or flowing nitric 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.[7]

The electrochemical method has been widely used by amateur chemists to make cheap PbO2 electrodes, useful for the production of chlorates and perchlorates.

Projects

  • Make PbO2 electrodes (β-PbO2 is more desired for this purpose than the α-PbO2 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)
  • Oxidizing agent
  • Make chromates
  • Mineral collecting

Handling

Safety

Lead dioxide is extremely toxic and ingestion may be fatal.

Storage

In closed bottles.

Disposal

Should be taken to waste disposal companies. Scrap metal facilities that collect lead may accept it.

References

  1. White, W. B.; Dachille, F.; Roy, R.; Journal of the American Ceramic Society; vol. 44; (1961); p. 170 - 174
  2. Povarennikh, O. S.; Dopovidi Akad. Nauk Ukr.RSR; (1963); p. 805 - 808; C. A.; vol. 59; (1963); p. 13411,
  3. https://books.google.com/books?id=PpTi_JAx7PgC&pg=PA387
  4. https://en.wikipedia.org/wiki/Lead_dioxide
  5. https://books.google.com/books?id=_PGzaO48Rz0C&pg=PA184
  6. http://www.mindat.org/min-3237.html
  7. https://books.google.com/books?id=ArsfQZig_9AC&pg=PA573

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