Perchloric acid

Perchloric acid is a mineral acid with the formula HClO4. Usually found as an aqueous solution, this colorless compound is a stronger acid than sulfuric acid and nitric acid. It is a powerful oxidizer when hot, but aqueous solutions up to approximately 70% by weight at room temperature are generally safe, only showing strong acid features and no oxidizing properties. Perchloric acid is useful for preparing perchlorate salts, especially ammonium perchlorate, an important rocket fuel component. Perchloric acid is dangerously corrosive and readily forms potentially explosive mixtures.

Perchloric acid
Perchloric acid
Hydroxidotrioxidochlorine
Perchloric acid
Hydroxidotrioxidochlorine
Names
Systematic IUPAC name
chloric(VII) acid
Other names
Hyperchloric acid[1]
Identifiers
3D model (JSmol)
ChEBI
ChEMBL
ChemSpider
ECHA InfoCard 100.028.648
EC Number
  • 231-512-4
RTECS number
  • SC7500000
UNII
UN number 1873
Properties
HClO4
Molar mass 100.46 g/mol
Appearance colorless liquid
Odor odorless
Density 1.768 g/cm3
Melting point −17 °C (1 °F; 256 K) (azeotrope)[2]
−112 °C (anhydrous)
Boiling point 203 °C (397 °F; 476 K) (azeotrope)[3]
miscible
Acidity (pKa) −15.2 (±2.0);[4] ≈ −10
Conjugate base Perchlorate
Hazards
Safety data sheet ICSC 1006
GHS pictograms
GHS Signal word Danger
H271, H290, H302, H314, H373
P210, P280, P303+361+353, P304+340, P310, P305+351+338, P371, P380, P375
NFPA 704 (fire diamond)
Flash point Non-flammable
Related compounds
Related compounds
 Hydrochloric acid
Hypochlorous acid
Chlorous acid
Chloric acid
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references

Production

Perchloric acid is produced industrially by two routes. The traditional method exploits the high aqueous solubility of sodium perchlorate (209 g/100 mL of water at room temperature). Treatment of such solutions with hydrochloric acid gives perchloric acid, precipitating solid sodium chloride:

NaClO4 + HCl → NaCl + HClO4

The concentrated acid can be purified by distillation. The alternative route, which is more direct and avoids salts, entails anodic oxidation of aqueous chlorine at a platinum electrode.[5][6]

Laboratory preparations

Treatment of barium perchlorate with sulfuric acid precipitates barium sulfate, leaving perchloric acid. It can also be made by mixing nitric acid with ammonium perchlorate and boiling while adding hydrochloric acid. The reaction gives nitrous oxide and perchloric acid due to a concurrent reaction involving the ammonium ion and can be concentrated and purified significantly by boiling off the remaining nitric and hydrochloric acids.

Properties

Anhydrous perchloric acid is an unstable oily liquid at room temperature. It forms at least five hydrates, several of which have been characterized crystallographically. These solids consist of the perchlorate anion linked via hydrogen bonds to H2O and H3O+ centers[7] Perchloric acid forms an azeotrope with water, consisting of about 72.5% perchloric acid. This form of the acid is stable indefinitely and is commercially available. Such solutions are hygroscopic. Thus, if left open to the air, concentrated perchloric acid dilutes itself by absorbing water from the air.

Dehydration of perchloric acid gives the anhydride dichlorine heptoxide:[8]

2 HClO4 + P4O10 → Cl2O7 + "H2P4O11"

Uses

Perchloric acid is mainly produced as a precursor to ammonium perchlorate, which is used in rocket fuel. The growth in rocketry has led to increased production of perchloric acid. Several million kilograms are produced annually.[5] Perchloric acid is one of the most proven materials for etching of liquid crystal displays and critical electronics applications as well as ore extraction and has unique properties in analytical chemistry.[9] Additionally it is a useful component in etching of chrome[10]

As an acid

Perchloric acid, a superacid, is one of the strongest Brønsted–Lowry acids. That its pKa is lower than −9 is evidenced by the fact that its monohydrate contains discrete hydronium ions and can be isolated as a stable, crystalline solid, formulated as [H3O+][ClO–
4].[11] The most recent estimate of its aqueous pKa is −15.2±2.0.[4] It provides strong acidity with minimal interference because perchlorate is weakly nucleophilic (explaining the high acidity of HClO4). Other acids of noncoordinating anions, such as fluoroboric acid and hexafluorophosphoric acid are susceptible to hydrolysis, whereas perchloric acid is not. Despite hazards associated with the explosiveness of its salts, the acid is often preferred in certain syntheses.[12] For similar reasons, it is a useful eluent in ion-exchange chromatography.

It is also used for electropolishing or etching of aluminium, molybdenum, and other metals.

Safety

Given its strong oxidizing properties, perchloric acid is subject to extensive regulations.[13] It is highly reactive with metals (e.g., aluminium) and organic matter (wood, plastics). Work conducted with perchloric acid must be conducted in fume hoods with a wash-down capability to prevent accumulation of oxidisers in the ductwork.

On February 20, 1947, in Los Angeles, California, 17 people were killed and 150 injured when a bath, consisting of over 1000 litres of 75% perchloric acid and 25% acetic anhydride by volume, exploded. The O'Connor Electro-Plating plant, 25 other buildings, and 40 automobiles were obliterated, and 250 nearby homes were damaged. The bath was being used to electro-polish aluminium furniture. In addition, organic compounds were added to the overheating bath when an iron rack was replaced with one coated with cellulose acetobutyrate (Tenit-2 plastic). A few minutes later the bath exploded.[14][15]

See also

References
  1. Samuel Fomon. Medicine and the Allied Sciences. 1. p. 148.
  2. Safety data for concentrated perchloric acid, ca. 70% msds.chem.ox.ac.uk
  3. Handling of Perchloric acid ameslab.gov
  4. Trummal, A.; Lipping, L.; Kaljurand, I.; Koppel, I. A.; Leito, I. "Acidity of Strong Acids in Water and Dimethyl Sulfoxide" J. Phys. Chem. A. 2016, 120, 3663-3669. doi:10.1021/acs.jpca.6b02253.
  5. Helmut Vogt, Jan Balej, John E. Bennett, Peter Wintzer, Saeed Akbar Sheikh, Patrizio Gallone "Chlorine Oxides and Chlorine Oxygen Acids" in Ullmann's Encyclopedia of Industrial Chemistry 2002, Wiley-VCH, Weinheim. doi:10.1002/14356007.a06_483.
  6. Müler, W.; Jönck, P. (1963). "Herstellung von Perchlorsäure durch anodische Oxydation von Chlor". Chemie Ingenieur Technik. 35 (2): 78. doi:10.1002/cite.330350203.; German patent DE1031288B; US patent US2846383A.
  7. Almlöf, Jan; Lundgren, Jan O.; Olovsson, Ivar "Hydrogen Bond Studies. XLV. Crystal structure of perchloric acid 2.5 hydrate" Acta Crystallographica Section B: Structural Crystallography and Crystal Chemistry 1971, volume 27, pp. 898–904. doi:10.1107/S0567740871003236.
  8. Holleman, Arnold F.; Wiberg, Egon (2001). Inorganic chemistry. Translated by Mary Eagleson, William Brewer. San Diego: Academic Press. p. 464. ISBN 0-12-352651-5.
  9. "Perchloric Acid". GFS chemicals. Archived from the original on 2015-01-31. Retrieved 2014-01-14.
  10. "Metal Etching". Thayer School of Engineering.
  11. Kathleen Sellers; Katherine Weeks; William R. Alsop; Stephen R. Clough; Marilyn Hoyt; Barbara Pugh (2006). Perchlorate: environmental problems and solutions. CRC Press. p. 16. ISBN 0-8493-8081-2.
  12. A. T. Balaban, C. D. Nenitzescu, K. Hafner and H. Kaiser (1973). "2,4,6-Trimethylpyrilium Perchlorate". Organic Syntheses.CS1 maint: multiple names: authors list (link); Collective Volume, 5, p. 1106
  13. Perchloric Acid, 60%, GR Material Safety Data Sheet Archived 2012-03-24 at the Wayback Machine Seton Resource Center.
  14. R. C. Nester; G. F. Vander Voort (1992). Safety in the Metallographic Laboratory. ASTM Standardization News. p. 34.
  15. "CALIFORNIA: The Amazing Brew". Time.com. March 3, 1947.
Compounds containing perchlorate group
HClO4 He
LiClO4 Be(ClO4)2 B(ClO
4)−
4
B(ClO4)3		 ROClO3 N(ClO4)3
NH4ClO4
NOClO4		 O FClO4 Ne
NaClO4 Mg(ClO4)2 Al(ClO4)3 Si P S ClO−
4
ClOClO3
Cl2O7		 Ar
KClO4 Ca(ClO4)2 Sc(ClO4)3 Ti(ClO4)4 VO(ClO4)3
VO2(ClO4)		 Cr(ClO4)3 Mn(ClO4)2 Fe(ClO4)3 Co(ClO4)2,
Co(ClO4)3		 Ni(ClO4)2 Cu(ClO4)2 Zn(ClO4)2 Ga(ClO4)3 Ge As Se Br Kr
RbClO4 Sr(ClO4)2 Y(ClO4)3 Zr(ClO4)4 Nb(ClO4)5 Mo Tc Ru Rh(ClO4)3 Pd(ClO4)2 AgClO4 Cd(ClO4)2 In(ClO4)3 Sn(ClO4)4 Sb TeO(ClO4)2 I Xe
CsClO4 Ba(ClO4)2   Hf(ClO4)4 Ta(ClO4)5 W Re Os Ir Pt Au Hg2(ClO4)2,
Hg(ClO4)2		 Tl(ClO4),
Tl(ClO4)3		 Pb(ClO4)2 Bi(ClO4)3 Po At Rn
FrClO4 Ra   Rf Db Sg Bh Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og
↓
La Ce(ClO4)x Pr Nd Pm Sm(ClO4)3 Eu(ClO4)3 Gd(ClO4)3 Tb(ClO4)3 Dy(ClO4)3 Ho(ClO4)3 Er(ClO4)3 Tm(ClO4)3 Yb(ClO4)3 Lu(ClO4)3
Ac Th(ClO4)4 Pa UO2(ClO4)2 Np Pu Am Cm Bk Cf Es Fm Md No Lr
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