Bleaches

When chlorine gas is bubbled through a cylinder of tomato juice, the chlorine/tomato juice mixture turns almost completely white within five minutes. This spectacular change is a result of the chemical action of chlorine, acting as an oxidizing bleaching agent, on the pigments in tomato juice. When old newspaper clippings, discolored through aging and exposure to light, are treated with 1 percent aqueous sodium borohydride solution, the paper is dramatically whitened within twenty minutes. In this instance, the paper has been restored to its original white color by the action of sodium borohydride acting as a reducing bleaching agent.

A bleaching agent is a substance that can whiten or decolorize other substances. Colored substances generally contain groups of atoms, called chromophores , that can absorb visible light having specific, characteristic wavelengths, and reflect or transmit the part of light that is not absorbed. For example, if a chromophore absorbs blue light, it will reflect light of the complementary color, and the chromophore-containing substance will appear yellow. Bleaching agents essentially destroy chromophores (thereby removing the color), via the oxidation or reduction of these absorbing groups. Thus, bleaches can be classified as either oxidizing agents or reducing agents .

Some of the use of bleaching agents are:

• The bleaching of textiles and fabrics
• The bleaching of wood pulp
• The removal of stains
• Commercial and household laundering and cleaning
• As ingredients in scouring cleansers and dishwashing products
• The bleaching of hair

Oxidizing Bleaches

A large number of oxidizing bleaches were reviewed by Jules A. Szilard in Bleaching Agents and Techniques (1973). The oxidizing bleaches (and bleaching agents) in common use today are: chlorine, chlorine dioxide, alkaline hypochlorites, hydrogen peroxide, peroxygen compounds, and sunlight and artificial light.

Chlorine (Cl ₂). The discovery of chlorine by the Swedish chemist Carl Wilhelm Scheele in 1774 marked the beginning of the modern era of bleaching. According to Sidney M. Edelstein in a 1948 journal article titled "The Role of Chemistry in the Development of Dyeing and Bleaching," French chemist Claude-Louis Berthollet was the first to use chlorine to bleach cotton and linen fabrics.

Chlorine has been used to bleach wood pulp. Many pulp mills employing the Kraft pulping process prepare sodium hydroxide (needed to digest wood chips) on-site via the electrolysis of brine , a concentrated aqueous solution of sodium chloride.

2NaCl + 2H₂O → 2NaOH + H₂ + Cl₂          (1)

Chlorine is a side product. Subsequent chlorine bleaching of the brown pulp gives a product that can be used for the manufacture of writing and printing paper. Unfortunately, organic compounds in the pulp are both oxidized and chlorinated, yielding small quantities of organochlorine compounds, including

dioxins. In fact, the most abundant dioxin produced by the pulp and bleaching process, 2,3,7,8-tetrachlorodibenzo-p -dioxin (2,3,7,8-TCDD), has been found to be both a carcinogen and a deadly toxin . Thus, chlorine as a bleaching agent is being replaced by the safer bleaching agents chlorine dioxide and hydrogen peroxide. In fact, the trend in the pulp and paper industries is toward totally chlorine free (TCF) bleaching. Chlorine is now used in the bleaching industry mainly to prepare hypochlorite solutions and dry bleaches such as calcium hypochlorite.

Chlorine Dioxide (ClO ₂). Chlorine dioxide has been used as a bleaching agent both in its gaseous phase and in aqueous solution. Because of its explosive nature, chlorine dioxide in the gaseous phase is often diluted with nitrogen or carbon dioxide. If stored or shipped, chlorine dioxide is passed through cold water and kept under refrigeration.

Chlorine dioxide is prepared industrially via the reduction of sodium chlorate by sulfur dioxide in aqueous solution.

2NaClO₃ + SO₂ + H₂SO₄ → 2ClO₂ + 2NaHSO₄          (2)

A relatively safe method for the preparation of ClO₂ involves the reaction between sodium chlorite (NaClO₂) and formaldehyde (H₂CO).

H₂CO + H⁺ + ClO₂⁻ → HOCl + HCOOH          (3)

As reaction 3 proceeds, the pH of the solution drops (due to the production of formic acid [HCOOH]). The increased acidity of the solution promotes the formation of ClO₂, shown in equation 4.

HCOOH + HOCl + 2ClO₂⁻ → 2ClO₂ + Cl⁻ + H₂O + HCOO⁻          (4)

In acidic solution, chlorine dioxide behaves as an oxidizing agent. The complete reduction of ClO₂ is shown in equation 5.

ClO₂ + 4H⁺ + 5 e⁻ → Cl⁻ + 2H₂O          (5)

The individual steps of this overall reduction reaction produce HClO₂, HOCl, and Cl₂, which all behave as oxidizing agents. An acidic medium is required, as ClO₂ disproportionates in alkaline solution, as shown in equation 6.

2ClO₂ + 2OH⁻ → ClO₃⁻ + ClO₂⁻ + H₂O          (6)

Chlorine dioxide is mainly used for pulp bleaching.

Hypochlorites (OCl ⁻). Hypochlorite bleach solutions are made from NaOCl and, to a lesser extent, Ca(OCl)₂. Hypochlorites are used in laundering, as disinfectants, in the bleaching of pulp and textiles, and in the removal of ink from recycled paper. Commercial bleaching solutions are obtained by passing chlorine gas through cold, dilute, aqueous sodium hydroxide, as shown in equation 7.

Cl₂ + 2OH⁻ → OCl⁻ + Cl⁻ + H₂O          (7)

Alternatively, the hypochlorite ion can be generated by the hydrolysis of organic nitrogen-chlorine compounds. Some of the more important nitrogen-chlorine compounds used in this way are the chlorinated isocyanurates. These find use in cleansing and dishwasher products.

To be an effective bleach, the hypochlorite solution should be kept alkaline (pH > 9.0), in order to suppress the hydrolysis of OCl⁻ (see equation 8) and prevent the formation of unstable HOCl.

OCl⁻ + H₂O → HOCl + OH⁻          (8)

In acidic solutions, HOCl forms and decomposes.

3HOCl → HClO₃ + 2HCl          (9)

HOCl will also react with HCl, one of the decomposition products.

HOCl + HCl → H₂O + Cl₂          (10)

Hypochlorite bleaching solutions must not contain heavy metal cations, as these cations (like light or heat) promote the decomposition of HOCl, as shown in equation 11.

2HOCl → 2HCl + O₂          (11)

The active ingredients in hypochlorite bleaches vary with pH. At pH < 2, Cl₂ is the main component in solution; at pH 4 to 6, HOCl is the dominant species; at pH > 9, OCl⁻ is the only component present. It is the hypochlorite ion in basic solution that is the active ingredient in household bleach, which is typically about 5 to 6 percent NaOCl. The OCl⁻ ion oxidizes chromophores in colored materials, and is itself reduced to chloride and hydroxide ions.

OCl⁻ + H₂O + 2 e⁻ → Cl⁻ + 2OH⁻          (12)

The whitening process effected by commercial hypochlorite bleach is often enhanced by the use of optical brighteners, compounds that absorb incident ultraviolet light and emit visible light, making the fabric appear brighter and whiter.

Hydrogen Peroxide (H ₂O ₂) . Hydrogen peroxide can be prepared by the reaction of barium peroxide and sulfuric acid (see equation 13). As barium sulfate precipitates out, hydrogen peroxide is easily separated.

BaO₂ + H₂SO₄ → BaSO₄ + H₂O₂          (13)

Hydrogen peroxide, as a bleaching agent used in the pulp and paper industry, has the advantage that it is nonpolluting. Because of the instability of pure hydrogen peroxide, aqueous solutions are employed in bleaching. At room temperature, hydrogen peroxide very slowly decomposes to water and oxygen.

2H₂O₂ → H₂O + O₂          (14)

However, the presence of transition metal cations (particularly Fe³⁺, Mn²⁺, and Cu²⁺) and other catalysts dramatically accelerates this reaction. As a result, aqueous hydrogen peroxide must be stabilized with complexing agents that sequester transition metal cations.

The active bleaching species in hydrogen peroxide is the perhydroxyl anion , OOH⁻, formed through the ionization of H₂O₂.

H₂O₂ + H₂O → H₃O⁺ + OOH⁻          (15)

The acid ionization constant of hydrogen peroxide is very low (K_a = 2 × 10⁻¹²) with the result that solutions of H₂O₂ must be made alkaline in order

to raise the concentration of OOH⁻. In the absence of an alkaline medium, hydrogen peroxide is no longer effective as a bleaching agent. For example, the bleaching stage of hair dyeing often employs hydrogen peroxide (5–6%), but also ammonia to provide an alkaline medium.

At the same time the pH must not rise above 11, as at this point the decomposition of OOH⁻ begins to occur.

2OOH⁻ → O₂ + 2OH⁻          (16)

Peroxygen Compounds. A number of solid peroxygen compounds that release hydrogen peroxide when dissolved in water exist. These include sodium perborate (NaBO₃ z 4H₂O or NaBO₂ z H₂O₂ z 3H₂O) and sodium carbonate peroxyhydrate (2Na₂CO₃ z 3H₂O₂). The structure of sodium perborate contains the peroxoanion B₂(O₂)₂(OH)₄²⁻, which contains two O–O linkages that join two tetrahedral BO₂(OH)²⁻ groups. These peroxygen compounds are used in detergents, denture cleaners, and tooth powders.

Bleaching with Light. Bleaching that involves either natural sunlight or artificial light has been used to remove stains from paper artifacts and to treat textiles. The material to be bleached is first immersed in an alkaline solution of either calcium or magnesium bicarbonate, and then protected from ultraviolet radiation by covering it with Plexiglas, Lexan, or Mylar. Exposure to light is then allowed to take place for two to four hours, for natural sunlight, and two to twelve hours, for artificial light.

Reducing Bleaches

Reducing agents used in bleaching include sulfites, bisulfites, dithionites, and sodium borohydride, all of which are used in pulp and textile bleaching.

Sulfites (SO ₃²⁻ ) and Bisulfites (HSO ₃⁻ ). The oxidation state of sulfur in both SO₃²⁻ and HSO₃⁻ is +4, and oxidation to +6 occurs readily, with the formation of SO₄²⁻ and HSO₄⁻, respectively, making sulfites and bisulfites good reducing agents.

Dithionites (S ₂O ₄²⁻ ) . Both sodium and zinc dithionite have found use in the bleaching of mechanical pulps and textiles. The preparation of the dithionite ion is accomplished via the reduction of the bisulfite ion and sulfur dioxide with Zn dust.

2HSO₃⁻ + SO₂ + Zn → Zn²⁺ + S₂O₄²⁻ + SO₃²⁻ + H₂O          (17)

The dithionite ion, S₂O₄²⁻, which has sulfur in the +3 oxidation state, behaves as a strong reducing agent in alkaline solution.

S₂O₄²⁻ + 4OH⁻ → 2SO₃²⁻ 2H₂O + 2 e⁻          (18)

As the pH is lowered, the reducing power of the dithionite ion drops off, as predicted by LeChatelier's principle.

Dithionites are useful in removing rust stains, and neutral citrate solutions of Na₂S₂O₄ were used to remove iron corrosion products from objects recovered from the Titanic.

Sodium Borohydride (NaBH ₄). Sodium borohydride has been used mainly in the industrial bleaching of mechanical pulps. The BH₄⁻ ion is a strong reducing agent in alkaline solution.

BH₄⁻ + 10OH⁻ → BO₃³⁻ + 8 e⁻ + 7H₂O          (19)

One problem with using sodium borohydride is that the BH₄⁻ ion slowly decomposes in aqueous solution.

BH₄⁻ + 4H₂O → B(OH)₄⁻ + 4H₂          (20)

As an alternative method, BH₄⁻ salts may be dissolved in either CH₃OH or the less toxic C₂H₅OH. The decomposition of the BH₄⁻ ion in alcohols occurs at a much slower rate:

BH₄⁻ + 4ROH → B(OR)₄⁻ + 4H₂ (R = CH₃, C₂H₅)          (21)

Conclusion

A bleaching agent can whiten or decolorize a substance by reacting with the chromophores that are responsible for the color of the substance. Depending on the nature of the chromophores, the bleaching agent will either be an oxidizing or reducing agent. That is, the chromophore is either oxidized or reduced to produce a colorless or whitened substance. Bleaching agents and their commercial uses are summarized in Table 1.

see also Berthollet, Claude-Louis; Chlorine; Detergents; Scheele, Carl.

Henry A. Carter

Bibliography

Allen, R. L. M. (1971). ColorChemistry. New York: Appleton-Century-Crofts.

Carter, Henry A. (1995). "A Simple Recipe for Whitening Old Newspaper Clippings." Journal of Chemical Education 72(7):651.

Carter, Henry A. (1996). "The Chemistry of Paper Preservation. Part 2: The Yellowing of Paper and Conservation Bleaching." Journal of Chemical Education 73(11): 1,068–1,073.

Cotton, F. Albert; Wilkinson, Geoffrey; Murillo, Carlos A., et al. (1999). Advanced Inorganic Chemistry, 6th edition. New York: Wiley-Interscience.

Duffield, P. A. (1986). Review of Bleaching. West Yorkshire, UK: Dyeing, Printing, and Bleaching Textile Technology Group.

Edelstein, Sidney M. (1948). "The Role of Chemistry in the Development of Dyeing and Bleaching." Journal of Chemical Education 25(3):144–149.

Emsley, John (1998). Molecules at an Exhibition. New York: Oxford University Press.

Freemantle, Michael (1994). "Chemical Techniques Help Conserve Artifacts Raised from Titanic Wreck." Chemical and Engineering News 72(42):49–52.

Holst, Gustaf (1954). "The Chemistry of Bleaching and Oxidizing Agents." Chemical Reviews 1954:169–194.

Jones, Mark M.; Johnston, David O.; Netterville, John T., et al. (1987). Chemistry and Society, 5th edition. New York: Saunders.

Kraft, F. (1969). "Bleaching of Wood Pulps." In The Pulping of Wood, 2nd edition, ed. R. G. MacDonald, and J. N. Franklin. New York: McGraw-Hill.

Lee, John D. (1996). Concise Inorganic Chemistry, 5th edition. New York: Chapman & Hall.

Nemetz, Thomas M., and Ball, David W. (1993). "Bleaching with Chlorine: Another Tomato Juice Demonstration." Journal of Chemical Education 70(2):154–155.

Swaddle, Thomas W. (1990). Applied Inorganic Chemistry. Calgary: University of Calgary Press.

Szilard, Jules A. (1973). Bleaching Agents and Techniques. Park Ridge, NJ: Noyes Data Corporation.

Waring, David R., and Hallas, Geoffrey, eds. (1990). The Chemistry and Application of Dyes. New York: Plenum Press.

Internet Resources

Linak, Erik; Leder, Andy; and Takei, Naoka. CEH Report: Hypochlorite Bleaches. Available from <http://ceh.sric.sri.com/Public/Reports/>.