Bleaches 3406
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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 + 2H 2 O → 2NaOH + H 2 + Cl 2          (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

Table 1. Bleaching agents and their commercial uses.
Table 1. Bleaching agents and their commercial uses.

Bleaching Agent Commercial Use in Bleaching
Chlorine Bleaching pulp and paper; making hypochlorites
Chlorine dioxide Bleaching kraft paper and flour
Sodium hypochlorite Household laundering and sanitizing
Calcium hypochlorite Solid bleach used in sanitizing
Sodium dichloroisocyanurate Sanitizing and dishwashing agents
Hydrogen peroxide Bleaching textiles, fur, pulp and paper, and hair
Sodium perborate Milder bleach for laundering; dry cleaning; denture cleaning; tooth powder; replacement for phosphates in detergents
Light Bleaching paper artifacts
Sulfur dioxide Preserving grapes, wine, and apples; removal of color during refinement of sugar
Sodium sulfite; sodium bisulfite Anti-chlor (a reducing agent for removing oxidizing bleaches)
Sodium dithionite Bleaching textiles, pulp and paper; removing rust stains
Sodium borohydride Bleaching pulp and paper

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.

2NaClO 3 + SO 2 + H 2 SO 4 → 2ClO 2 + 2NaHSO 4          (2)

A relatively safe method for the preparation of ClO 2 involves the reaction between sodium chlorite (NaClO 2 ) and formaldehyde (H 2 CO).

H 2 CO + H + + ClO 2 → 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 ClO 2 , shown in equation 4.

HCOOH + HOCl + 2ClO 2 → 2ClO 2 + Cl + H 2 O + HCOO          (4)

In acidic solution, chlorine dioxide behaves as an oxidizing agent. The complete reduction of ClO 2 is shown in equation 5.

ClO 2 + 4H + + 5 e → Cl + 2H 2 O          (5)

The individual steps of this overall reduction reaction produce HClO 2 , HOCl, and Cl 2 , which all behave as oxidizing agents. An acidic medium is required, as ClO 2 disproportionates in alkaline solution, as shown in equation 6.

2ClO 2 + 2OH → ClO 3 + ClO 2 + H 2 O          (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.

Cl 2 + 2OH → OCl + Cl + H 2 O          (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 + H 2 O → HOCl + OH          (8)

In acidic solutions, HOCl forms and decomposes.

3HOCl → HClO 3 + 2HCl          (9)

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

HOCl + HCl → H 2 O + Cl 2          (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 + O 2          (11)

The active ingredients in hypochlorite bleaches vary with pH. At pH < 2, Cl 2 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 + H 2 O + 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 2 O 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.

BaO 2 + H 2 SO 4 → BaSO 4 + H 2 O 2          (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.

2H 2 O 2 → H 2 O + O 2          (14)

However, the presence of transition metal cations (particularly Fe 3+ , Mn 2+ , and Cu 2+ ) 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 H 2 O 2 .

H 2 O 2 + H 2 O → H 3 O + + OOH          (15)

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

This woman is using bleach to launder clothing. Not only can bleach remove stains, but it can act as an agent to remove color.
This woman is using bleach to launder clothing. Not only can bleach remove stains, but it can act as an agent to remove color.

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 → O 2 + 2OH          (16)

Peroxygen Compounds. A number of solid peroxygen compounds that release hydrogen peroxide when dissolved in water exist. These include sodium perborate (NaBO 3 z 4H 2 O or NaBO 2 z H 2 O 2 z 3H 2 O) and sodium carbonate peroxyhydrate (2Na 2 CO 3 z 3H 2 O 2 ). The structure of sodium perborate contains the peroxoanion B 2 (O 2 ) 2 (OH) 4 2− , which contains two O–O linkages that join two tetrahedral BO 2 (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 3 2− ) and Bisulfites (HSO 3 ). The oxidation state of sulfur in both SO 3 2− and HSO 3 is +4, and oxidation to +6 occurs readily, with the formation of SO 4 2− and HSO 4 , respectively, making sulfites and bisulfites good reducing agents.

Dithionites (S 2 O 4 2− ) . 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.

2HSO 3 + SO 2 + Zn → Zn 2+ + S 2 O 4 2− + SO 3 2− + H 2 O          (17)

The dithionite ion, S 2 O 4 2− , which has sulfur in the +3 oxidation state, behaves as a strong reducing agent in alkaline solution.

S 2 O 4 2− + 4OH → 2SO 3 2− 2H 2 O + 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 Na 2 S 2 O 4 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 BH 4 ion is a strong reducing agent in alkaline solution.

BH 4 + 10OH → BO 3 3− + 8 e + 7H 2 O          (19)

One problem with using sodium borohydride is that the BH 4 ion slowly decomposes in aqueous solution.

BH 4 + 4H 2 O → B(OH) 4 + 4H 2          (20)

As an alternative method, BH 4 salts may be dissolved in either CH 3 OH or the less toxic C 2 H 5 OH. The decomposition of the BH 4 ion in alcohols occurs at a much slower rate:

BH 4 + 4ROH → B(OR) 4 + 4H 2 (R = CH 3 , C 2 H 5 )          (21)


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


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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 .

Also read article about Bleaches from Wikipedia

User Contributions:

Sandeep Goyal
I want to ask - how sodium sulphite makes brine Cl2 free in alkaline brine ? What is the reaction?
This is an excellent article. I think it could be better if it clarified one issue. It states, "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."
It would be helpful if it mentioned that it is a more effective disinfectant at lower pH, in order that the the bleaching effect is not confused with the disinfecting efficacy.

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