Sodium persulfate is the inorganic compound with the formula Na2S2O8. It is the sodium salt of H2S2O8, an oxidizing agent, it is a white solid. It is non-hygroscopic and has good shelf-life; the salt is prepared by the electrolytic oxidation of sodium hydrogen sulfate: 2 NaHSO4 → Na2S2O8 + H2Oxidation is conducted at a platinum anode. In this way about 165,000 tons were produced in 2005; the standard redox potential of sodium persulfate into hydrogen sulfate is 2.1 V, higher than that of hydrogen peroxide but lower than ozone. The sulfate radical formed in situ has a standard electrode potential of 2.7 V. However, there are a few drawbacks in utilizing platinum anodes to produce the salts. Thus, boron-doped diamond electrodes have been proposed as alternatives to the conventional platinum electrodes, it is used as a radical initiator for emulsion polymerization reactions for styrene based polymers such as Acrylonitrile butadiene styrene. Applicable for accelerated curing of low formaldehyde adhesives.
It is a bleach, both standalone and as a detergent component. It is a replacement for ammonium persulfate in etching mixtures for zinc and printed circuit boards, is used for pickling of copper and some other metals, it is used as a soil conditioner and for soil and groundwater remediation and in manufacture of dyestuffs, modification of starch, bleach activator, desizing agent for oxidative desizing, etc. Sodium persulfate is a specialized oxidizing agent in chemistry, classically in the Elbs persulfate oxidation and the Boyland–Sims oxidation reactions, it is used in radical reactions. The salt is an oxidizer and forms combustable mixtures with organic materials such as paper
A chemical oscillator is a complex mixture of reacting chemical compounds in which the concentration of one or more components exhibits periodic changes, They are a class of reactions that serve as an example of non-equilibrium thermodynamics with far-from-equilibrium behavior. The reactions are theoretically important in that they show that chemical reactions do not have to be dominated by equilibrium thermodynamic behavior. In cases where one of the reagents has a visible color, periodic color changes can be observed. Examples of oscillating reactions are the Belousov-Zhabotinsky reaction, the Briggs-Rauscher reaction, the Bray-Liebhafsky reaction and the chlorine dioxide – iodine – malonic acid reaction; the earliest scientific evidence that such reactions can oscillate was met with extreme scepticism. In 1828, G. T. Fechner published a report of oscillations in a chemical system, he described an electrochemical cell. In 1899, W. Ostwald observed that the rate of chromium dissolution in acid periodically increased and decreased.
Both of these systems were heterogeneous and it was believed and through much of the last century, that homogeneous oscillating systems were nonexistent. While theoretical discussions date back to around 1910, the systematic study of oscillating chemical reactions and of the broader field of non-linear chemical dynamics did not become well established until the mid-1970s. Chemical systems cannot oscillate about a position of final equilibrium because such an oscillation would violate the second law of thermodynamics. For a thermodynamic system, not at equilibrium, this law requires that the system approach equilibrium and not recede from it. For a closed system at constant temperature and pressure, the thermodynamic requirement is that the Gibbs free energy must decrease continuously and not oscillate; however it is possible that the concentrations of some reaction intermediates oscillate, that the rate of formation of products oscillates. Theoretical models of oscillating reactions have been studied by chemists and mathematicians.
In an oscillating system the energy-releasing reaction can follow at least two different pathways, the reaction periodically switches from one pathway to another. One of these pathways produces a specific intermediate; the concentration of this intermediate triggers the switching of pathways. When the concentration of the intermediate is low, the reaction follows the producing pathway, leading to a high concentration of intermediate; when the concentration of the intermediate is high, the reaction switches to the consuming pathway. Different theoretical models for this type of reaction have been created, including the Lotka-Volterra model, the Brusselator and the Oregonator; the latter was designed to simulate the Belousov-Zhabotinsky reaction. A Belousov–Zhabotinsky reaction is one of several oscillating chemical systems, whose common element is the inclusion of bromine and an acid. An essential aspect of the BZ reaction is its so-called "excitability" — under the influence of stimuli, patterns develop in what would otherwise be a quiescent medium.
Some clock reactions such as the Briggs–Rauscher reactions and the BZ using the chemical ruthenium bipyridyl as catalyst can be excited into self-organising activity through the influence of light. Boris Belousov first noted, sometime in the 1950s, that in a mix of potassium bromate, cerium sulfate, propanedioic acid and citric acid in dilute sulfuric acid, the ratio of concentration of the cerium and cerium ions oscillated, causing the colour of the solution to oscillate between a yellow solution and a colorless solution; this is due to the cerium ions being reduced by propanedioic acid to cerium ions, which are oxidized back to cerium ions by bromate ions. The Briggs–Rauscher oscillating reaction is one of a small number of known oscillating chemical reactions, it is well suited for demonstration purposes because of its visually striking color changes: the freshly prepared colorless solution turns an amber color changing to a dark blue. This fades to colorless and the process repeats, about ten times in the most popular formulation.
The Bray–Liebhafsky reaction is a chemical clock first described by W. C. Bray in 1921 with the oxidation of iodine to iodate: 5 H2O2 + I2 → 2 IO3− + 2 H+ + 4 H2Oand the reduction of iodate back to iodine: 5 H2O2 + 2 IO3− + 2 H+ → I2 + 5 O2 + 6 H2O There are other classes of chemical oscillators: Catalytic oscillator Mercury beating heart Video of BZ reaction History of oscillating reactions
Redox is a chemical reaction in which the oxidation states of atoms are changed. Any such reaction involves both a reduction process and a complementary oxidation process, two key concepts involved with electron transfer processes. Redox reactions include all chemical reactions; the chemical species from which the electron is stripped is said to have been oxidized, while the chemical species to which the electron is added is said to have been reduced. It can be explained in simple terms: Oxidation is the loss of electrons or an increase in oxidation state by a molecule, atom, or ion. Reduction is a decrease in oxidation state by a molecule, atom, or ion; as an example, during the combustion of wood, oxygen from the air is reduced, gaining electrons from carbon, oxidized. Although oxidation reactions are associated with the formation of oxides from oxygen molecules, oxygen is not included in such reactions, as other chemical species can serve the same function; the reaction can occur slowly, as with the formation of rust, or more in the case of fire.
There are simple redox processes, such as the oxidation of carbon to yield carbon dioxide or the reduction of carbon by hydrogen to yield methane, more complex processes such as the oxidation of glucose in the human body. "Redox" is a portmanteau of the words "reduction" and "oxidation". The word oxidation implied reaction with oxygen to form an oxide, since dioxygen was the first recognized oxidizing agent; the term was expanded to encompass oxygen-like substances that accomplished parallel chemical reactions. The meaning was generalized to include all processes involving loss of electrons; the word reduction referred to the loss in weight upon heating a metallic ore such as a metal oxide to extract the metal. In other words, ore was "reduced" to metal. Antoine Lavoisier showed. Scientists realized that the metal atom gains electrons in this process; the meaning of reduction became generalized to include all processes involving a gain of electrons. Though "reduction" seems counter-intuitive when speaking of the gain of electrons, it might help to think of reduction as the loss of oxygen, its historical meaning.
Since electrons are negatively charged, it is helpful to think of this as reduction in electrical charge. The electrochemist John Bockris has used the words electronation and deelectronation to describe reduction and oxidation processes when they occur at electrodes; these words are analogous to protonation and deprotonation, but they have not been adopted by chemists worldwide. The term "hydrogenation" could be used instead of reduction, since hydrogen is the reducing agent in a large number of reactions in organic chemistry and biochemistry. But, unlike oxidation, generalized beyond its root element, hydrogenation has maintained its specific connection to reactions that add hydrogen to another substance; the word "redox" was first used in 1928. The processes of oxidation and reduction occur and cannot happen independently of one another, similar to the acid–base reaction; the oxidation alone and the reduction alone are each called a half-reaction, because two half-reactions always occur together to form a whole reaction.
When writing half-reactions, the gained or lost electrons are included explicitly in order that the half-reaction be balanced with respect to electric charge. Though sufficient for many purposes, these general descriptions are not correct. Although oxidation and reduction properly refer to a change in oxidation state — the actual transfer of electrons may never occur; the oxidation state of an atom is the fictitious charge that an atom would have if all bonds between atoms of different elements were 100% ionic. Thus, oxidation is best defined as an increase in oxidation state, reduction as a decrease in oxidation state. In practice, the transfer of electrons will always cause a change in oxidation state, but there are many reactions that are classed as "redox" though no electron transfer occurs. In redox processes, the reductant transfers electrons to the oxidant. Thus, in the reaction, the reductant or reducing agent loses electrons and is oxidized, the oxidant or oxidizing agent gains electrons and is reduced.
The pair of an oxidizing and reducing agent that are involved in a particular reaction is called a redox pair. A redox couple is a reducing species and its corresponding oxidizing form, e.g. Fe2+/Fe3+ Substances that have the ability to oxidize other substances are said to be oxidative or oxidizing and are known as oxidizing agents, oxidants, or oxidizers; that is, the oxidant removes electrons from another substance, is thus itself reduced. And, because it "accepts" electrons, the oxidizing agent is called an electron acceptor. Oxygen is the quintessential oxidizer. Oxidants are chemical substances with elements in high oxidation states, or else electronegative elements that can gain extra electrons by oxidizing another substance. Substances that have the ability to reduce other substances are said to be reductive or reducing and are known as
Sulfuric acid known as vitriol, is a mineral acid composed of the elements sulfur and hydrogen, with molecular formula H2SO4. It is a colorless and syrupy liquid, soluble in water, in a reaction, exothermic, its corrosiveness can be ascribed to its strong acidic nature, and, if at a high concentration, its dehydrating and oxidizing properties. It is hygroscopic absorbing water vapor from the air. Upon contact, sulfuric acid can cause severe chemical burns and secondary thermal burns. Sulfuric acid is a important commodity chemical, a nation's sulfuric acid production is a good indicator of its industrial strength, it is produced with different methods, such as contact process, wet sulfuric acid process, lead chamber process and some other methods. Sulfuric acid is a key substance in the chemical industry, it is most used in fertilizer manufacture, but is important in mineral processing, oil refining, wastewater processing, chemical synthesis. It has a wide range of end applications including in domestic acidic drain cleaners, as an electrolyte in lead-acid batteries, in various cleaning agents.
Although nearly 100% sulfuric acid can be made, the subsequent loss of SO3 at the boiling point brings the concentration to 98.3% acid. The 98.3% grade is more stable in storage, is the usual form of what is described as "concentrated sulfuric acid". Other concentrations are used for different purposes; some common concentrations are: "Chamber acid" and "tower acid" were the two concentrations of sulfuric acid produced by the lead chamber process, chamber acid being the acid produced in the lead chamber itself and tower acid being the acid recovered from the bottom of the Glover tower. They are now obsolete as commercial concentrations of sulfuric acid, although they may be prepared in the laboratory from concentrated sulfuric acid if needed. In particular, "10M" sulfuric acid is prepared by adding 98% sulfuric acid to an equal volume of water, with good stirring: the temperature of the mixture can rise to 80 °C or higher. Sulfuric acid reacts with its anhydride, SO3, to form H2S2O7, called pyrosulfuric acid, fuming sulfuric acid, Disulfuric acid or oleum or, less Nordhausen acid.
Concentrations of oleum are either expressed in terms of % SO3 or as % H2SO4. Pure H2S2O7 is a solid with melting point of 36 °C. Pure sulfuric acid has a vapor pressure of <0.001 mmHg at 25 °C and 1 mmHg at 145.8 °C, 98% sulfuric acid has a <1 mmHg vapor pressure at 40 °C. Pure sulfuric acid is a viscous clear liquid, like oil, this explains the old name of the acid. Commercial sulfuric acid is sold in several different purity grades. Technical grade H2SO4 is impure and colored, but is suitable for making fertilizer. Pure grades, such as United States Pharmacopeia grade, are used for making pharmaceuticals and dyestuffs. Analytical grades are available. Nine hydrates are known, but four of them were confirmed to be tetrahydrate and octahydrate. Anhydrous H2SO4 is a polar liquid, having a dielectric constant of around 100, it has a high electrical conductivity, caused by dissociation through protonating itself, a process known as autoprotolysis. 2 H2SO4 ⇌ H3SO+4 + HSO−4The equilibrium constant for the autoprotolysis is Kap = = 2.7×10−4The comparable equilibrium constant for water, Kw is 10−14, a factor of 1010 smaller.
In spite of the viscosity of the acid, the effective conductivities of the H3SO+4 and HSO−4 ions are high due to an intramolecular proton-switch mechanism, making sulfuric acid a good conductor of electricity. It is an excellent solvent for many reactions; because the hydration reaction of sulfuric acid is exothermic, dilution should always be performed by adding the acid to the water rather than the water to the acid. Because the reaction is in an equilibrium that favors the rapid protonation of water, addition of acid to the water ensures that the acid is the limiting reagent; this reaction is best thought of as the formation of hydronium ions: H2SO4 + H2O → H3O+ + HSO−4 Ka1 = 2.4×106 HSO−4 + H2O → H3O+ + SO2−4 Ka2 = 1.0×10−2 HSO−4 is the bisulfate anion and SO2−4 is the sulfate anion. Ka1 and Ka2 are the acid dissociation constants; because the hydration of sulfuric acid is thermodynamically favorable and the affinity of it for water is sufficiently strong, sulfuric acid is an excellent dehydrating agent.
Concentrated sulfuric acid has a powerful dehydrating property, removing water from other chemical compounds including sugar and other carbohydrates and producing carbon and steam. In the laboratory, this is demonstrated by mixing table sugar into sulfuric acid; the sugar changes from white to dark brown and to black as carbon is formed. A rigid column of black, porous carbon will emerge as well; the carbon will smell of caramel due to the heat generated. C 12 H 22 O 11 ⏞ sucrose → H 2 SO 4 12 C + 11 H 2
Ammonium persulfate is the inorganic compound with the formula 2S2O8. It is a colourless salt, soluble in water, much more so than the related potassium salt, it is a strong oxidizing agent, used in polymer chemistry, as an etchant, as a cleaning and bleaching agent. The dissolution of the salt in water is an endothermic process. Ammonium persulfate is prepared by electrolysis of a cold concentrated solution of either ammonium sulfate or ammonium bisulfate in sulfuric acid at a high current density; the method was first described by Hugh Marshall. As an oxidizing agent and a source of radicals, APS finds many commercial applications. Salts of sulfate are used as radical initiators in the polymerization of certain alkenes. Commercially important polymers prepared using persulfates include styrene-butadiene rubber and polytetrafluoroethylene. In solution, the dianion dissociates to give radicals: 2− ⇌ 2 •−The sulfate radical adds to the alkene to give a sulfate ester radical, it is used along with tetramethylethylenediamine to catalyze the polymerization of acrylamide in making a polyacrylamide gel, hence being important for SDS-PAGE and western blot.
Illustrative of its powerful oxidizing properties, it is used to etch copper on printed circuit boards as an alternative to ferric chloride solution. This property was discovered many years ago. In 1908, John William Turrentine used a dilute ammonium persulfate solution to etch copper. Turrentine weighed copper spirals before placing the copper spirals into the ammonium persulfate solution for an hour. After an hour, the spirals were weighed again and the amount of copper dissolved by ammonium persulfate was recorded; this experiment was extended to other metals such as nickel and iron, all of which yielded similar results. The oxidation equation is thus: 1⁄2 S2O2−8 + e− → SO2−4. Ammonium persulfate is a standard ingredient in hair bleach. Persulfates are used as oxidants in organic chemistry. For example in the Minisci reaction Airborne dust may be irritating to eye, throat and skin upon contact. Exposure to high levels of dust may cause difficulty in breathing, it has been noted. Furthermore, it has been suggested that exposure to ammonium persulfate can cause asthmatic effects in hair dressers and receptionists working in the hair dressing industry.
These asthmatic effects are proposed to be caused by the oxidation of cysteine residues, as well as methionine residues. International Chemical Safety Card 0632
Sodium bisulfite is a chemical compound with the chemical formula NaHSO3. Sodium bisulfite is a food additive with E number E222; this salt of bisulfite can be prepared by bubbling sulfur dioxide in a solution of sodium carbonate in water. Sodium bisulfite in contact with chlorine bleach will generate heat and form sodium bisulfate and sodium chloride. Sodium bisulfite can be prepared by bubbling excess sulfur dioxide through a solution of suitable base, such as sodium hydroxide or sodium bicarbonate. SO2 + NaOH → NaHSO3 SO2 + NaHCO3 → NaHSO3 + CO2 Sodium bisulfite is a weakly acidic species with a pKa of 6.97. Its conjugate base is the sulfite ion, SO32−: HSO3− ↔ SO32− + H+The theoretical protonated species is sulfurous acid, it forms a bisulfite adduct with aldehyde groups and with certain cyclic ketones to give a sulfonic acid. This reaction is useful for purification procedures. Contaminated aldehydes in a solution precipitate as the bisulfite adduct which can be isolated by filtration; the reverse reaction takes place in presence of a base such as sodium bicarbonate or sodium hydroxide and the bisulfite is liberated as sulfur dioxide.
Examples of such procedures are described for benzaldehyde, 2-tetralone, the ethyl ester of pyruvic acid and glyoxal. In the ring-expansion reaction of cyclohexanone with diazald, the bisulfite reaction is reported to be able to differentiate between the primary reaction product cycloheptanone and the main contaminant cyclooctanone; the other main use of sodium bisulfite is as a mild reducing agent in organic synthesis in particular in purification procedures. It can efficiently remove traces or excess amounts of chlorine, iodine, hypochlorite salts, osmate esters, chromium trioxide and potassium permanganate. A third use of sodium bisulfite is as a decoloration agent in purification procedures because it can reduce coloured oxidizing agents, conjugated alkenes and carbonyl compounds. Sodium bisulfite is the key ingredient in the Bucherer reaction. In this reaction an aromatic hydroxyl group is replaced by an aromatic amine group and vice versa because it is a reversible reaction; the first step in this reaction is an addition reaction of sodium bisulfite to an aromatic double bond.
The Bucherer carbazole synthesis is a related organic reaction that uses sodium bisulfite as a reagent. While the related compound, sodium metabisulfite, is used in all commercial wines to prevent oxidation and preserve flavor, sodium bisulfite is sold by some home winemaking suppliers for the same purpose. In fruit canning, sodium bisulfite is used to kill microbes. In the case of wine making, sodium bisulfite releases sulfur dioxide gas when added to water or products containing water; the sulfur dioxide kills yeasts and bacteria in the grape juice before fermentation. When the sulfur dioxide levels have subsided, fresh yeast is added for fermentation, it is added to bottled wine to prevent the formation of vinegar if bacteria are present, to protect the color and flavor of the wine from oxidation, which causes browning and other chemical changes. The sulfur dioxide reacts with oxidation by-products and prevents them from causing further deterioration. Sodium bisulfite is added to leafy green vegetables in salad bars and elsewhere, to preserve apparent freshness, under names like LeafGreen.
The concentration is sometimes high enough to cause allergic reactions. On July 8, 1986, sodium bisulfite was banned from use by the FDA on fresh fruits and vegetables in the United States following the deaths of 13 people and many illnesses among asthmatics. Sodium bisulfite is used in the analysis of the methylation status of cytosines in DNA. In this technique, sodium bisulfite deaminates cytosine into uracil, but does not affect 5-methylcytosine, a methylated form of cytosine with a methyl group attached to carbon 5; when the bisulfite-treated DNA is amplified via polymerase chain reaction, the uracil is amplified as thymine and the methylated cytosines are amplified as cytosine. DNA sequencing techniques are used to read the sequence of the bisulfite-treated DNA; those cytosines that are read as cytosines after sequencing represent methylated cytosines, while those that are read as thymines represent unmethylated cytosines in the genomic DNA. Sodium bisulfite is a common reducing agent in the chemical industry, as it reacts with dissolved oxygen: 2 NaHSO3 + O2 → 2 NaHSO4It is added to large piping systems to prevent oxidative corrosion.
In biochemical engineering applications, it is helpful to maintain anaerobic conditions within a reactor. Sodium bisulfite should not be confused with sodium bisulfate, used as a pH lowering chemical for swimming pools. In drinking water treatment, sodium bisulfite is added after super chlorination, to reduce the residual chlorine before discharging to the service reservoir. In wastewater treatment, sodium bisulfite is added following disinfection with chlorine prior to discharging the effluent to the receiving water. Residual chlorine can have a negative impact on aquatic life. In steam boilers, sodium bisulfite has been a reliable oxygen scavenger in boiler feedwater for 60 years; this compound is characterized as having fast reaction times, low use-cost, years of proven performance, availability. Sodium bisulfite when used in steam boilers has USDA approvals. Sodium metabisul
Hans Heinrich Landolt
Hans Heinrich Landolt was the Swiss chemist who discovered iodine clock reaction. He is one of the founders of Landolt-Börnstein database. Landolt was born in Zurich and at the age of nineteen entered the university there to study chemistry and physics, he attended the lectures of Carl Jacob Löwig and published his first work on stibmethyl in Schriften der Naturforschenden Gesellschaft. He was appointed assistant to Lowig and followed him in 1853 to Breslau; the same year he obtained the degree of Doctor of Philosophy with a thesis "Ueber die Arsenäthyle", a notable contribution to the law of chemical valence. After the defense, he went to Berlin to attend lectures of Eilhard Mitscherlich, Johannes Muller and Dubois. Facilities for experimental research in chemistry were non-existent in Berlin at the time, therefore Landolt left for Heidelberg for a newly founded institute of Robert Bunsen. After devoting himself for a short time to the electrolytic production of calcium and lithium, Landolt started an investigation of the gases produced in the Bunsen burner, constructed in the winter of 1854–55.
In 1856 Landolt returned to Breslau, where he was soon afterwards joined by Lothar Meyer and Friedrich Konrad Beilstein. In the same year he became a lecturer in chemistry on the strength of his monograph on "Chemische Vorgange in der Flamme der Leuchtgase". In 1857, he was called to Bonn where he studied the effect of the atomic composition of liquids containing carbon and oxygen on the transmission of light; the results were published in 1862–1864 and were a continuation of the previous researches of John Hall Gladstone. In his life he elaborated the work of Hertz and demonstrated that light waves are differentiated from electric waves by the wavelength, in 1892 he extended his early work to measurements of the molecular refractivity of organic substances for radiowaves. At Bonn, in 1859, Landolt married the daughter of Swiss parents settled in Bonn. In 1869, he was appointed to the head of the newly founded technical college at Aachen, where a chemical institute was built according to his plans.
His work there was concerned with the relations between physical properties and chemical constitution. In particular, he studied optical rotation by various chemicals. In 1880, he was called by the Prussian Ministry of Agriculture to the newly founded Agricultural College in Berlin, where he remained until 1891. There he constructed new laboratories and collaborated with Richard Börnstein in the compilation of the "Physikalisch-chemischen Tabellen", their third edition was published in 1905 with the assistance of Wilhelm Meyerhoffer and a generous financial support by the Berlin Academy of Sciences. In 1882 Landolt became a member of the Berlin Academy. Around that time he made remarkable investigations into the kinetics of the iodine clock reaction between iodic acid and sulfurous acid. From 1891 till his retirement in 1905, he served as director of the second chemical institute of the Berlin University. There he worked on three major problems: relation between the melting point and molecular weight, effect of crystallinity on the optical rotation and change in weight during chemical reactions.
The negative result for the last experiments was regarded as an accurate experimental confirmation of the conservation laws of mass and energy. Landolt was known for his humor, friendliness and cigar, he was fit and worked as usual until the week before his death, when he had a sudden failure of heart and kidney. He was buried, in accordance with his desire, at Bonn where he spent most memorable years of his life; this article incorporates text from Obituary notices, by Otto N. Witt, a publication from 1911 now in the public domain in the United States. Richard Pribram. "Obituary: Hans Heinrich Landolt". Berichte der deutschen chemischen Gesellschaft. 44: 3337–3394. Doi:10.1002/cber.191104403209. Works by or about Hans Heinrich Landolt at Internet Archive