Bleaching agents are materials that lighten or whiten a substrate (for example a fabric or hard surface) through a chemical reaction. They are used for textile and paper and pulp bleaching as well as for home laundering. Bleaching usually involves either oxidative or reductive reactions that decompose stains and soils. These processes may involve the removal or change of larger molecules and color-bearing groups in the stain or soil into smaller, more soluble units which are more easily removed in the bleaching process. The most common bleaching agents generally fall into three categories: halogen bleaches and their related compounds (such as sodium hypochlorite and sodium dichloroisocyanurate), oxygen bleaches (such as hydrogen peroxide and sodium percarbonate), and reducing bleaches.
Halogen bleaching agents are the most cost-effective bleaching agents, and are mostly based on chlorine or bromine. Besides being very effective at removing stains and soils, they are also highly effective antimicrobial products – indeed, water treatment is the largest use of chlorine‑containing bleaching agents.
The most common halogen bleaches may be divided into four classes: chlorine, hypochlorites, N‑chloro compounds, and chlorine dioxide. The first three classes are categorized as “available chlorine” compounds or “chlorine bleaches”, although hypochlorites and N-chloro compounds do not actually contain chlorine gas. They are related to chlorine by the equilibria in Equations 1-4, which occur very quickly in aqueous solution:
|Cl2 (g) ⇌ Cl2 (aq)
|Cl2 (aq) + H2O ⇌ HOCl + H+ + Cl‑
|HOCl ⇌ H+ + OCl‑
|RR’NCl + H2O ⇌ HOCl + RR’NH
Solutions of halogen bleaches decompose on standing, at a rate which depends on solutions and storage conditions. Hypochlorous acid and hypochlorite anions decompose according to Equations 5 and 6:
|3 HOCl ® HClO3 + 2 HCl
|3 OCl‑ ® ClO3‑ + 2 Cl‑
Hypochlorite solutions are most stable above pH 11 where the rate of decomposition is nearly independent of pH. In this region, the rate of decomposition has a second-order dependence on the concentration of hypochlorite. Decomposition also increases with increasing ionic strength, thus concentrated solutions decompose much faster than dilute solutions. Further, because of an unusually high activation energy, the decomposition rate increases greatly with temperature. Nevertheless, solutions with less than about 7% sodium hypochlorite and a pH above 11 have acceptable long-term stability below about 30oC; if manufactured with best practices, consumers can expect hypochlorite bleach to contain at least the stated label strength by the time they purchase it, and to retain at least half of its original strength after one year.
Decomposition also occurs by reaction 7:
|2 OCl‑ ® O2 + 2 Cl‑
This reaction usually can be ignored unless it is exposed to light (where ultraviolet radiation initiates the reaction) or unless it is catalyzed by transition metal ions. Even very small amounts of transition metal ions like cobalt, nickel and copper cause rapid decomposition. They form very reactive intermediates that can decrease the stability of oxidizable compounds in the bleach solution and increase the damage to substrates (e.g., appearance of yellowing, metal oxide deposits, or in extreme cases, pin-holing of fabrics).
The major form of hypochlorite produced is sodium hypochlorite, NaOCl. It is invariably made and used as an aqueous solution. It is usually prepared by the chlorination of sodium hydroxide solutions as shown in Equation 8, although other bases like sodium carbonate can be used as well.
|Cl2 + 2 NaOH ⇌ NaOCl + NaCl + H2O
A 5‑7% sodium hypochlorite solution is sold for household purposes, of which the largest use is in laundry. Solutions of 10‑15% NaOCl are sold for swimming pool disinfection, institutional laundries, and industrial purposes such as cleaning of dairy equipment, restaurants, sanitizing industrial plant equipment and disinfecting cooling towers. Solutions of various strengths are used in household and industrial and institutional (I&I) cleaners, and disinfectants and sanitizers. A small amount is used in textile mills. Sodium hypochlorite is also made on site (to avoid rapid decomposition) for pulp bleaching.
Commercially important solid “available chlorine” bleaches (such as calcium hypochlorite and sodium dichloroisocyanurate) are usually more stable than concentrated hypochlorite solutions. They decompose very slowly in sealed containers. However, most of them decompose quickly as they absorb moisture from air or from other ingredients in a formulation.
Oxygen-based bleaches, also known as peroxygen bleaches, contain the peroxide linkage: (‑O‑O‑) in which one of the oxygen atoms is active. Hydrogen peroxide is one of the most common of these bleaching agents. It is the primary oxygen-based bleaching agent in the textile industry, and it is also used in pulp and paper and home laundry applications. In textile bleaching, hydrogen peroxide is the most common bleaching agent for protein fibers (e.g., wool, silk) and it is also used extensively for cellulosic fibers (e.g., cotton, rayon). Pure hydrogen peroxide has an active oxygen content of 47%. It is the least expensive source of active oxygen commercially available. Moreover, it is a liquid, making it convenient for many bleaching applications. It is a very weak acid and in aqueous solutions only dissociates slightly (Equation 9).
|H2O2 ⇌ H+ + HO2‑
Undissociated hydrogen peroxide is relatively stable, and for this reason most commercial products are adjusted to an acid pH. Hydrogen peroxide usually is sold in solutions containing 30‑35%, 50% or 65‑70 wt.% of the active material. More concentrated solutions (80‑85%, 90%) are available in limited quantities. Concentrated solutions of hydrogen peroxide are very hazardous and must be handled with extreme care: at concentrations above 5%, hydrogen peroxide can cause permanent eye damage. In concentrations of 15% and above, hydrogen peroxide is an aggressive oxidizer and is corrosive to many materials, including human skin.
Hydrogen peroxide bleaching is performed in alkaline solution where part of the hydrogen peroxide is converted to the perhydroxy anion (Equation 9). The perhydroxy anion HO2‑ is generally believed to be the active bleaching species and its concentration in solution increases with hydrogen peroxide concentration, alkalinity and temperature. The alkaline agents most commonly used to generate HO2‑ are caustic soda, carbonates, silicates, pyrophosphates and polyphosphates. Better bleaching is obtained at these alkaline conditions by increasing the temperature, and by adding stabilizers to prevent the uncontrolled decomposition reactions of hydrogen peroxide. Common stabilizers include colloidal stannates, silicates, pyrophosphates and polyaminocarboxylates. These stabilizers may be different from those used to stabilize commercial acidic hydrogen peroxide.
Hydrogen peroxide reacts with many compounds, such as borates, carbonates, pyrophosphates, sulfates, silicates and with a variety of organic carboxylic acids, esters and anhydrides to give peroxy compounds or peroxyhydrates. A number of these compounds are stable solids that hydrolyze readily to give hydrogen peroxide in solution, the most important being sodium perborate and sodium percarbonate. Solid oxygen bleaches, as in the case of halogen bleaches, are preferred for stability and compatibility with other sensitive ingredients. Another solid oxygen bleach, peroxymonosulfuric acid (the peroxygen product of hydrogen peroxide and sulfuric acid) is a powerful oxidizing agent. A commercially available salt, potassium peroxymonosulfate, is a white solid having a satisfactory shelf life and an active oxygen content of about 4.4%. It is a triple salt with the composition 2KHSO5·K2SO4·KHSO4.
Peracids are compounds containing the functional group -OOH derived from an organic or inorganic acid functionality. Typical structures include CH3C(O)OOH derived from acetic acid and HOS(O)2OOH (peroxymonosulfuric acid previously discussed) derived from sulfuric acid. Peracids have superior cold water bleaching capability versus hydrogen peroxide due to the greater electrophilicity of the peracid peroxygen group. The cold water bleaching performance and phosphate reductions in detergent systems accounts for their emergent utilization and vast literature of peracids in textile bleaching.
Peracids can be introduced into the bleaching system by two methods. Peracids can be manufactured separately and delivered to the bleaching bath with the other components or as an adjunct. Peracids can also be formed in situ utilizing the perhydrolysis reaction shown in Equation 10 (where L denotes a leaving group). The two main peracid precursors in general use are tetraacetylethylenediamine (TAED), which generates peracetic acid in the wash, and nonanoyloxybenzene sulfonate (NOBS), which produces pernonanoic acid when combined with hydrogen peroxide in the wash water.
|R-C(O)-L + HO2‑ ® R-C(O)OOH + L-
As bleaching agents, oxygen bleaches are much less effective than hypochlorite. However, they do have some advantages over halogen bleaching agents, such as less potential damage to textile fibers and dyes, and lack of a strong odor. Attempts have been made to increase the laundry bleaching power of hydrogen peroxide-based laundry bleaches by the addition of heavy metal catalysts. However, the effectiveness of these systems remains controversial; an early attempt to incorporate a catalyst into a laundry detergent led to fabric damage and was consequently withdrawn. Though catalysts have not been incorporated into commercial products in the U.S., they have found use in automatic dishwashing detergents in Europe.
Reducing agents generally used in bleaching include sulfur dioxide, sulfurous acid, bisulfites, sulfites, hydrosulfites (dithionites), sodium sulfoxylate formaldehyde and sodium borohydride. These materials are used mainly in industrial processes such as pulp and textile bleaching, however they have been used in a small number of consumer products. Sulfur dioxide and its derivatives have been used to bleach textiles since earliest times. Besides being an important bleaching agent in the pulp and paper industry, sulfur dioxide is also used in the manufacture of chlorine dioxide, sodium hydrosulfite, and sodium sulfite. When SO2 is dissolved in water, it yields a complex mixture given the trivial name sulfurous acid (H2SO3), which contains SO2, H3O+, S2O52‑, and HSO3‑. The composition of the mixture depends on the concentration of the sulfur dioxide in the water, the pH, and the temperature. Although sulfurous acid does not exist in the free state, it forms stable salts (the neutral sulfite, SO32‑, and the hydrogen sulfite or bisulfite HSO3‑) which are good reducing agents.
Sodium sulfite, which is used in pulp and paper bleaching, is usually produced by the reaction of sulfur dioxide with either caustic soda or soda ash (Equations 11 and 12):
|SO2 + 2NaOH ⇌ Na2SO3 + H2O
|SO2 + Na2CO3 ⇌ Na2SO3 + CO2
Free dithionous acid, H2S2O4, has never been isolated; however the salts of the acid (in particular zinc and sodium dithionite) have been prepared and are widely used as industrial reducing agents. The dithionite salts can be prepared by the reduction of sulfites, hydrosulfites, and sulfur dioxide with metallic substances such as zinc, iron, or zinc or sodium amalgams or by electrolytic reduction. The principal applications of these compounds are in bleaching of pulp and in dyeing, printing and stripping in the textile industry. An alternative form of sodium dithionite is sodium sulfoxylate formaldehyde, which is prepared by the reaction of formaldehyde with the dithionite. Its applications are like those of dithionite, except that it is less reactive and more stable thermally. When the sulfoxylate is used, a pH range of 3.2‑3.5 produces the best results. For both the dithionite salts and sulfoxylate, the higher the temperature, the greater the reducing strength. Sulfoxylate can be used at temperatures as high as 100oC.