Hydrogen Compounds

First Edition, 2012

ISBN 978-81-323-3467-5

© All rights reserved.

Published by: Research World 4735/22 Prakashdeep Bldg, Ansari Road, Darya Ganj, Delhi - 110002 Email: [email protected] 

Table of Contents Chapter 1 - Hydrogen Iodide Chapter 2 - Hydrogen Chloride Chapter 3 - Hypochlorous Acid Chapter 4 - Hydrogen Cyanide Chapter 5 - Chromic Acid and Perchloric Acid Chapter 6 - Carbonic Acid and Hydrogen Fluoride Chapter 7 - Hydrogen Isocyanide and Isocyanic acid Chapter 8 - Nitric Acid Chapter 9 - Hydrogen Peroxide Chapter 10 - Properties of Water Chapter 11 - Phosphorous Acid and Sulfamic Acid Chapter 12 - Phosphoric Acid Chapter 13 - Hydrogen Sulfide Chapter 14 - Sulfuric Acid

Chapter-1

Hydrogen Iodide

Hydrogen iodide

IUPAC name Hydrogen iodide Identifiers CAS number 10034-85-2 PubChem 24841 ChemSpider 23224 UNII 694C0EFT9Q EC number 233-109-9 KEGG C05590 MeSH Hydroiodic+acid ChEBI CHEBI:43451 ChEMBL CHEMBL1233550 RTECS number MW3760000 Gmelin Reference 814 Properties Molecular formula HI Molar mass 127.91 g mol−1 Exact mass 127.912293452 g mol-1 Appearance Colorless gas Density 2.85 g cm-3 (at −47 °C) Melting point -51 °C, 222 K, -60 °F Boiling point -34 °C, 239 K, -29 °F –9.5 Acidity (pKa) Basicity (pKb) 23.5 1.466 Refractive index (nD) Structure Dipole moment 0.38 D

Thermochemistry Std enthalpy of formation ΔfHo298 Specific heat capacity, C MSDS EU Index

0.0016199 kJ mol-1 0.2283 J/(g·K) Hazards External MSDS 053-002-00-9

EU classification C

R-phrases S-phrases

R35 (S1/2), S9, S26, S36/37/39, S45 Related compounds Astatane

Hydrogen bromide Hydrogen chloride Related compounds Hydrogen fluoride Indium hydride Rubidium hydride Stibine Supplementary data page Structure and n, εr, etc. properties Thermodynamic Phase behaviour data Solid, liquid, gas Spectral data UV, IR, NMR, MS

Hydrogen iodide (HI) is a diatomic molecule. Aqueous solutions of HI are known as iohydroic acid or hydroiodic acid, a strong acid. Gas and aqueous solution are interconvertible. HI is used in organic and inorganic synthesis as one of the primary sources of iodine and as a reducing agent.

Properties of hydrogen iodide HI is a colorless gas that reacts with oxygen to give water and iodine. With moist air, HI gives a mist (or fumes) of hydroiodic acid. It is exceptionally soluble in water, giving hydriodic acid. One liter of water will dissolve 425 liters of HI gas, the final solution having only four water molecules per molecule of HI.

Hydroiodic acid Hydroiodic acid is a solution of pure HI in water. Commercial hydroiodic acid usually contains 57% HI by mass. The solution forms an azeotrope boiling at 127 °C with 57% HI, 43% water. Hydroiodic acid is one of the strongest of all common acids due to the high stability of its corresponding conjugate base. The iodide ion is the largest of all common halides which results in the negative charge being dispersed over a larger space.

By contrast, a chloride ion is significantly smaller, meaning its negative charge is more concentrated, leading to a stronger interaction between the proton and the chloride ion. This weaker H+---I– interaction in HI facilitates dissociation of the proton from the anion, and is the reason HI is the strongest acid of the hydrohalides(except for hydroastatic acid [theoretically]). HI(g) + H2O(l) → H3O(aq)+ + I–(aq) Ka ≈ 1010 HBr(g) + H2O(l) → H3O(aq)+ + Br–(aq) Ka ≈ 109 HCl(g) + H2O(l) → H3O(aq)+ + Cl–(aq) Ka ≈ 108

Preparation The industrial preparation of HI involves the reaction of I2 with hydrazine, which also yields nitrogen gas. 2 I2 + N2H4 → 4 HI + N2 When performed in water, the HI must be distilled. HI can also be distilled from a solution of NaI or other alkali iodide in concentrated hypophosphorous acid (note that sulfuric acid will not work for acidifying iodides as it will oxidize the iodide to elemental iodine). Another way HI may be prepared is by bubbling hydrogen sulfide steam through an aqueous solution of iodine, forming hydroiodic acid (which is distilled) and elemental sulfur (this is filtered). H2S +I2 → 2 HI + S Additionally HI can be prepared by simply combining H2 and I2. This method is usually employed to generate high purity samples. H2 + I2 → 2 HI For many years, this reaction was considered to involve a simple bimolecular reaction between molecules of H2 and I2. However, when a mixture of the gases is irradiated with the wavelength of light equal to the dissociation energy of I2, about 578 nm, the rate increases significantly. This supports a mechanism whereby I2 first dissociates into 2 iodine atoms, which each attach themselves to a side of an H2 molecule and break the H—H bond: H2 + I2 + 578 nm radiation → H2 + 2 I → I – - – H – - – H – - – I → 2 HI In the laboratory, another method involves hydrolysis of PI3, the iodine equivalent of PBr3. In this method, I2 reacts with phosphorus to create phosphorus triiodide, which then reacts with water to form HI and phosphorous acid.

3 I2 + 2 P + 6 H2O → 2 PI3 + 6 H2O → 6 HI + 2 H3PO3

Key reactions and applications 

HI will undergo oxidation if left open to air according to the following pathway: 4 HI + O2 → 2H2O + 2 I2 HI + I2 → HI3

HI3 is dark brown in color, which makes aged solutions of HI often appear dark brown. 

Like HBr and HCl, HI add to alkenes HI + H2C=CH2 → H3CCH2I

HI is also used in organic chemistry to convert primary alcohols into alkyl halides. This reaction is an SN2 substitution, in which the iodide ion replaces the "activated" hydroxyl group (water). HI is preferred over other hydrogen halides in polar protic solvents because the iodide ion is a much better nucleophile than bromide or chloride, so the reaction can take place at a reasonable rate without much heating. The large iodide anion is less solvated and more reactive in polar protic solvents and thus causes the reaction to proceed faster because of stronger partial bonds in the transition state. This reaction also occurs for secondary and tertiary alcohols, but substitution occurs via the SN1 pathway.

HI (or HBr) can also be used to cleave ethers into alkyl iodides and alcohols, in a reaction similar to the substitution of alcohols. This type of cleavage is siginficant because it can be used to convert a chemically stable and inert ether into more reactive species. In this example diethyl ether is cleaved into ethanol and iodoethane. The reaction is regioselective, as iodide tends to attack the less sterically hindered ether carbon.

Hydroiodic acid is subject to the same Markovnikov and anti-Markovnikov guidelines as HCl and HBr. 

HI reduces certain α-substituted ketones and alcohols replacing the α substituent with a hydrogen atom.

Illicit use of hydroiodic acid

Lab using the HI/P method Hydriodic acid is currently listed as a Federal DEA List I Chemical. Owing to its usefulness as a reducing agent, reduction with HI and red phosphorus has become the most popular method to produce methamphetamine in the United States. Clandestine chemists react pseudoephedrine (recovered from decongestant pills) with hydroiodic acid and red phosphorus under heat, HI reacts with pseudoephedrine to form iodoephedrine, an intermediate which is reduced primarily to methamphetamine. Because of its listed status and closely monitored sales, clandestine chemists now use red phosphorus and iodine to generate hydroiodic acid in situ.

Use in salt industry Hydroiodic acid can be used to synthesize sodium iodide or potassium iodide for increasing iodine content of salt.

Chapter-2

Hydrogen Chloride

Hydrogen chloride

IUPAC name Hydrogen chloride Chlorane Other names Chlorohydric acid Hydrochloride Hydrochloric acid Hydrochloric acid gas Identifiers CAS number 7647-01-0 PubChem 313 ChemSpider 307 UNII QTT17582CB EC number 231-595-7 RTECS number MW4025000 Properties Molecular formula HCl Molar mass 36.46 g/mol Appearance Colorless gas, hygroscopic. Density 1.477 g/L, gas (25 °C)

Melting point Boiling point Solubility in water Acidity (pKa) Molecular shape Dipole moment

–114.2 °C (158.8 K) –85.1 °C (187.9 K) 720 g/L (20 °C) –7.0 (in water), 10.3 (in acetonitrile) Structure Linear 1.05 D Thermochemistry

Std enthalpy of formation ΔfHo298 Std enthalpy of combustion ΔcHo298 Specific heat capacity, C

–2.351 kJ/g –2.614 kJ/g

0.7981 J/g K Hazards MSDS JT Baker MSDS EU Index 017-002-00-2 Toxic (T) EU classification Corrosive (C) R-phrases R23, R35 S-phrases (S1/2), S9, S26, S36/37/39, S45 Flash point Non-flammable Related compounds Hydrogen fluoride Other anions Hydrogen bromide Hydrogen iodide Other cations Sodium chloride Related compounds Hydrochloric acid

The compound hydrogen chloride has the formula HCl. At room temperature, it is a colorless gas, which forms white fumes of hydrochloric acid upon contact with atmospheric humidity. Hydrogen chloride gas and hydrochloric acid are important in technology and industry. The formula HCl is often used to refer, somewhat misleadingly, to hydrochloric acid, an aqueous solution that can be derived from hydrogen chloride.

Chemistry

Hydrochloric acid fumes turning pH paper red showing that the fumes are acidic Hydrogen chloride is composed of diatomic molecules, each consisting of a hydrogen atom H and a chlorine atom Cl connected by a covalent single bond. Since the chlorine atom is much more electronegative than the hydrogen atom, the covalent bond between the two atoms is quite polar. Consequently, the molecule has a large dipole moment with a negative partial charge δ– at the chlorine atom and a positive partial charge δ+ at the hydrogen atom. In part due to its high polarity, HCl is very soluble in water (and in other polar solvents). Upon contact, H2O and HCl combine to form hydronium cations H3O+ and chloride anions Cl– through a reversible chemical reaction: HCl + H2O → H3O+ + Cl– The resulting solution is called hydrochloric acid and is a strong acid. The acid dissociation or ionization constant, Ka, is large, which means HCl dissociates or ionizes practically completely in water. Even in the absence of water, hydrogen chloride can still act as an acid. For example, hydrogen chloride can dissolve in certain other solvents such as methanol, protonate molecules or ions, and serve as an acid-catalyst for chemical reactions where anhydrous (water-free) conditions are desired. HCl + CH3OH → CH3O+H2 + Cl– Because of its acidic nature, hydrogen chloride is corrosive, particularly in the presence of moisture.

Structure and properties

Infrared (IR) absorption spectrum The infrared spectrum of gaseous hydrogen chloride consists of a number of sharp absorption lines grouped around 2886 cm−1 (wavelength ~3.47 µm). At room temperature, almost all molecules in the ground vibrational state v = 0. To promote an HCl molecule to the v = 1 state, we would expect to see an infrared absorption about 2880 cm−1. This absorption corresponding to the Q-branch is not observed due to it being forbidden due to symmetry. Instead, two sets of signals (P- and R-branches) are seen due to rotation of the molecules. Due to quantum mechanical rules, only certain rotational modes are permitted. They are characterized by the rotational quantum number J = 0, 1, 2, 3, ... ΔJ can only take values of ± 1. E(J) = h·B·J(J+1) The value of B is much smaller than ν e, such that a much smaller amount of energy is required to rotate the molecule; for a typical molecule, this lies within the microwave region. However, due to the vibrational energy of this molecule, the set of absorptions lie within the infrared region, allowing a spectrum showing the rovibrational modes of this molecule to be easily collected using an ordinary infrared spectrometer with a conventional gas cell.

One doublet in the IR spectrum due to isotopic composition of Chlorine. Naturally abundant chlorine consists of two isotopes, 35Cl and 37Cl, in a ratio of approximately 3:1. While the spring constants are very similar, the reduced masses are different causing significant differences in the rotational energy, thus doublets are observed on close inspection of each absorption line, weighted in the same ratio of 3:1.

Production Most hydrogen chloride produced on an industrial scale is used for hydrochloric acid production.

Direct synthesis

Flame inside HCl Oven.

In the chlor-alkali industry, brine (mixture of sodium chloride and water) solution is electrolyzed producing chlorine (Cl2), sodium hydroxide, and hydrogen (H2). The pure chlorine gas can be combined with hydrogen to produce hydrogen chloride. Cl2 + H2 → 2HCl As the reaction is exothermic, the installation is called an HCl oven or HCl burner. The resulting hydrogen chloride gas is absorbed in deionized water, resulting in chemically pure hydrochloric acid. This reaction can give a very pure product, e.g. for use in the food industry.

Organic synthesis The largest production of hydrochloric acid is integrated with the formation of chlorinated and fluorinated organic compounds, e.g., Teflon, Freon, and other CFCs, as well as chloroacetic acid, and PVC. Often this production of hydrochloric acid is integrated with captive use of it on-site. In the chemical reactions, hydrogen atoms on the hydrocarbon are replaced by chlorine atoms, whereupon the released hydrogen atom recombines with the spare atom from the chlorine molecule, forming hydrogen chloride. Fluorination is a subsequent chlorine-replacement reaction, producing again hydrogen chloride. R-H + Cl2 → R-Cl + HCl R-Cl + HF → R-F + HCl The resulting hydrogen chloride gas is either reused directly, or absorbed in water, resulting in hydrochloric acid of technical or industrial grade.

Laboratory methods Small amounts of HCl gas for laboratory use can be generated in a HCl generator by dehydrating hydrochloric acid with either sulfuric acid or anhydrous calcium chloride. Alternatively, HCl can be generated by the reaction of sulfuric acid with sodium chloride: NaCl + H2SO4 → NaHSO4 + HCl This reaction occurs at room temperature. Provided there is salt remaining in the generator and it is heated above 200 degrees Celsius, the reaction proceeds to; NaCl + NaHSO4 → HCl + Na2SO4 For such generators to function, the reagents should be dry. HCl can also be prepared by the hydrolysis of certain reactive chloride compounds such as phosphorus chlorides, thionyl chloride (SOCl2), and acyl chlorides. For example, cold

water can be gradually dripped onto phosphorus pentachloride (PCl5) to give HCl in this reaction: PCl5 + H2O → POCl3 + 2

HCl

High purity streams of the gas require lecture bottles or cylinders, both of which can be expensive. In comparison, the use of a generator requires only apparatus and materials commonly available in a laboratory.

Applications Most hydrogen chloride is used in the production of hydrochloric acid. It is also an important reagent in other industrial chemical transformations, e.g.:  

Hydrochlorination of rubber Production of vinyl and alkyl chlorides

In the semiconductor industry, it is used to both etch semiconductor crystals and to purify silicon via trichlorosilane (SiHCl3). It may also be used to treat cotton to delint it, and to separate it from wool. In the laboratory, anhydrous forms of the gas are particularly useful for generating chloride-based Lewis acids, which must be absolutely dry for their Lewis sites to function. It can also be used to dry the corresponding hydrated forms of these materials by passing it over as they are heated; the materials would otherwise fume HCl(g) themselves and decompose. Neither can these hydrates be dried using standard desiccator methods.

History Alchemists of the Middle Ages recognized that hydrochloric acid (then known as spirit of salt or acidum salis) released vaporous hydrogen chloride, which was called marine acid air. In the 17th century, Johann Rudolf Glauber used salt (sodium chloride) and sulfuric acid for the preparation of sodium sulfate, releasing hydrogen chloride gas. In 1772, Carl Wilhelm Scheele also reported this reaction and is sometimes credited with its discovery. Joseph Priestley prepared hydrogen chloride in 1772, and in 1810 Humphry Davy established that it is composed of hydrogen and chlorine. During the Industrial Revolution, demand for alkaline substances such as soda ash increased, and Nicolas Leblanc developed a new industrial-scale process for producing the soda ash. In the Leblanc process, salt was converted to soda ash, using sulfuric acid, limestone, and coal, giving hydrogen chloride as by-product. Initially, this gas was vented to air, but the Alkali Act of 1863 prohibited such release, so then soda ash producers absorbed the HCl waste gas in water, producing hydrochloric acid on an industrial scale. Later, the Hargreaves process was developed, which is similar to the Leblanc process

except sulfur dioxide, water, and air are used instead of sulfuric acid in a reaction which is exothermic overall. In the early 20th century the Leblanc process was effectively replaced by the Solvay process, which did not produce HCl. However, hydrogen chloride production continued as a step in hydrochloric acid production. Historical uses of hydrogen chloride in the 20th century include hydrochlorinations of alkynes in producing the chlorinated monomers chloroprene and vinyl chloride, which are subsequently polymerized to make polychloroprene (Neoprene) and polyvinyl chloride (PVC), respectively. In the production of vinyl chloride, acetylene (C2H2) is hydrochlorinated by adding the HCl across the triple bond of the C2H2 molecule, turning the triple into a double bond, yielding vinyl chloride. The "acetylene process", used until the 1960s for making chloroprene, starts out by joining two acetylene molecules, and then adds HCl to the joined intermediate across the triple bond to convert it to chloroprene as shown here:

This "acetylene process" has been replaced by a process which adds Cl2 to one of the double bonds in 1,3-butadiene instead, and subsequent elimination produces HCl instead, as well as chloroprene.

Safety Hydrogen chloride forms corrosive hydrochloric acid on contact with water found in body tissue. Inhalation of the fumes can cause coughing, choking, inflammation of the nose, throat, and upper respiratory tract, and in severe cases, pulmonary edema, circulatory system failure, and death. Skin contact can cause redness, pain, and severe skin burns. Hydrogen chloride may cause severe burns to the eye and permanent eye damage. The gas, being strongly hydrophilic, can be easily scrubbed from the exhaust gases of a reaction by bubbling it through water, producing useful hydrochloric acid as a byproduct. Any equipment handling hydrogen chloride gas must be checked on a routine basis; particularly valve stems and regulators. The gas requires the use of specialized materials on all whetted parts of the flow path, as it will interact with or corrode numerous materials hydrochloric acid alone will not; such as stainless and regular polymers.

Chapter-3

Hypochlorous Acid

Hypochlorous acid

IUPAC name hypochlorous acid, chloranol, hydroxidochlorine Other names Hydrogen hypochlorite, Chlorine hydroxide Identifiers CAS number 7790-92-3 PubChem 24341 ChemSpider 22757 UNII 712K4CDC10 EC number 232-232-5 Properties Molecular formula HClO Molar mass 52.46 g/mol Appearance Colorless aqueous solns Density Variable Solubility in water Soluble Acidity (pKa) 7.53 Hazards Main hazards Oxidizer

Related compounds Chlorine Related compounds Calcium hypochlorite Sodium hypochlorite

Hypochlorous acid is a weak acid with the chemical formula HClO. It forms when chlorine dissolves in water. It cannot be isolated in pure form due to rapid equilibration with its precursor. HClO is an oxidizer, and as its sodium salt sodium hypochlorite, NaClO, or its calcium salt calcium hypochlorite, is used as a bleach, a deodorant, and a disinfectant.

Uses In organic synthesis, HCIO converts alkenes to chlorohydrins. In biology, hypochlorous acid is generated in activated neutrophils by myeloperoxidasemediated peroxidation of chloride ions, and contributes to the destruction of bacteria. In water treatment, hypochlorous acid is the active sanitizer in hypochlorite-based products (e.g. used in swimming pools). In food service and water distribution, specialized equipment to generate weak solutions of HOCl from water and salt is sometimes used to generate adequate quantities of safe (unstable) disinfectant to treat food preparation surfaces and water supplies.

Formation, stablity and reactions Addition of chlorine to water gives both hydrochloric acid (HCl) and hypochlorous acid: Cl2 + H2O

HClO + HCl

When acids are added to aqueous salts of hypochlorous acid (such as sodium hypochlorite in commercial bleach solution), the resultant reaction is driven to the left, and chlorine gas is evolved. Thus, the formation of stable hypochlorite bleaches is facilitated by dissolving chlorine gas into basic water solutions, such as sodium hydroxide. The acid can also be prepared by dissolving dichlorine monoxide in water; under standard aqueous conditions, anhydrous hypochlorous acid is impossible to prepare due to the readily reversible equilibrium between it and its anhydride: 2 HOCl

Cl2O + H2O K(0°C) = 3.55×10−3 dm3mol−1

The presence of light or transition metal oxides of copper, nickel, or cobalt accelerates the exothermic decomposition into hydrochloric acid and oxygen:

2 Cl2 + 2 H2O → 4 HCl + O2

Chemical reactions In aqueous solution, hypochlorous acid partially dissociates into the anion hypochlorite OCl−: HClO

OCl− + H+

Salts of hypochlorous acid are called hypochlorites. One of the best-known hypochlorites is NaClO, the active ingredient in bleach. HClO is a stronger oxidant than chlorine under standard conditions. 2 HClO(aq) + 2 H+ + 2 e−

Cl2(g) + 2 H2O E = +1.63 V

HClO reacts with HCl to form chlorine gas: HClO + HCl → H2O + Cl2

Reactivity of HClO with biomolecules Hypochlorous acid reacts with a wide variety of biomolecules including DNA, RNA, fatty acid groups, cholesterol and proteins.

Reaction with protein sulfhydryl groups Knox et al. first noted that HClO is a sulfhydryl inhibitor that, in sufficient quantity, could completely inactivate proteins containing sulfhydryl groups. This is because HClO oxidises sulfhydryl groups, leading to the formation of disulfide bonds that can result in crosslinking of proteins. The HClO mechanism of sulfhydryl oxidation is similar to that of chloramine, and may only be bacteriostatic, because, once the residual chlorine is dissipated, some sulfhydryl function can be restored. One sulfhydryl-containing amino acid can scavenge up to four molecules of HOCl. Consistent with this, it has been proposed that sulfhydryl groups of sulfur-containing amino acids can be oxidized a total of three times by three HClO molecules, with the fourth reacting with the α-amino group. The first reaction yields sulfenic acid (R-SOH) then sulfinic acid (R-SO2H) and finally RSO3H. Each of those intermediates can also condense with another sulfhydryl group, causing cross-linking and aggregation of proteins. Sulfinic acid and R-SO3H derivatives are produced only at high molar excesses of HClO, and disulfides are formed primarily at bacteriocidal levels. Disulfide bonds can also be oxidized by HClO to sulfinic acid. Because the oxidation of sulfhydryls and disulfides evolves hydrochloric acid, this process results in the depletion HClO.

Reaction with protein amino groups Hypochlorous acid reacts readily with amino acids that have amino group side-chains, with the chlorine from HClO displacing a hydrogen, resulting in an organic chloramine. Chlorinated amino acids rapidly decompose, but protein chloramines are longer-lived and retain some oxidative capacity. Thomas et al. concluded from their results that most organic chloramines decayed by internal rearrangement and that fewer available NH2 groups promoted attack on the peptide bond, resulting in cleavage of the protein. McKenna and Davies found that 10 mM or greater HClO is necessary to fragment proteins in vivo. Consistent with these results, it was later proposed that the chloramine undergoes a molecular rearrangement, releasing HCl and ammonia to form an amide. The amide group can further react with another amino group to form a Schiff base, causing cross-linking and aggregation of proteins.

Reaction with DNA and nucleotides Hypochlourous acid reacts slowly with DNA and RNA as well as all nucleotides in vitro. GMP is the most reactive because HClO reacts with both the heterocyclic NH group and the amino group. In similar manner, TMP with only a heterocyclic NH group that is reactive with HClO is the second-most reactive. AMP and CMP, which have only a slowly reactive amino group, are less reactive with HClO. UMP has been reported to be reactive only at a very slow rate. The heterocyclic NH groups are more reactive than amino groups, and their secondary chloramines are able to donate the chlorine. These reactions likely interfere with DNA base pairing, and, consistent with this, Prütz has reported a decrease in viscosity of DNA exposed to HClO similar to that seen with heat denaturation. The sugar moieties are unreactive and the DNA backbone is not broken. NADH can react with chlorinated TMP and UMP as well as HClO. This reaction can regenerate UMP and TMP and results in the 5-hydroxy derivative of NADH. The reaction with TMP or UMP is slowly reversible to regenerate HClO. A second slower reaction that results in cleavage of the pyridine ring occurs when excess HClO is present. NAD+ is inert to HClO.

Reaction with lipids Hypochlorous acid reacts with unsaturated bonds in lipids, but not saturated bonds, and the OCl− ion does not participate in this reaction. This reaction occurs by hydrolysis with addition of chlorine to one of the carbons and a hydroxyl to the other. The resulting compound is a chlorhydrin. The polar chlorine disrupts lipid bilayers and could increase permeability. When chlorhydrin formation occurs in lipid bilayers of red blood cells, increased permeability occurs. Disruption could occur if enough chlorhydrin is formed. The addition of preformed chlorhydrins to red blood cells can affect permeability as well. Cholesterol chlorhydrins have also been observed, but do not greatly affect permeability, and it is believed that Cl2 is responsible for this reaction.

Mode of disinfectant action Escherichia coli exposed to hypochlorous acid lose viability in less than 100 ms due to inactivation of many vital systems. Hypochlorous acid has a reported LD50 of 0.0104– 0.156 ppm and 2.6 ppm caused 100% growth inhibition in 5 minutes. However it should be noted that the concentration required for bactericidal activity is also highly dependent on bacterial concentration.

Inhibition of glucose oxidation In 1948, Knox et al. proposed the idea that inhibition of glucose oxidation is a major factor in the bacteriocidal nature of chlorine solutions. He proposed that the active agent or agents diffuse across the cytoplasmic membrane to inactivate key sulfhydrylcontaining enzymes in the glycolytic pathway. This group was also the first to note that chlorine solutions (HOCl) inhibit sulfhydryl enzymes. Later studies have shown that, at bacteriocidal levels, the cytosol components do not react with HOCl. In agreement with this, McFeters and Camper found that aldolase, an enzyme that Knox et al. proposes would be inactivated, was unaffected by HOCl in vivo. It has been further shown that loss of sulfhydryls does not correlate with inactivation. That leaves the question concerning what causes inhibition of glucose oxidation. The discovery that HOCl blocks induction of β-galactosidase by added lactose led to a possible answer to this question. The uptake of radiolabeled substrates by both ATP hydrolysis and proton co-transport may be blocked by exposure to HOCl preceding loss of viability. From this observation, it proposed that HOCl blocks uptake of nutrients by inactivating transport proteins. The question of loss of glucose oxidation has been further explored in terms of loss of respiration. Venkobachar et al. found that succinic dehydrogenase was inhibited in vitro by HOCl, which led to the investigation of the possibility that disruption of electron transport could be the cause of bacterial inactivation. Albrich et al. subsequently found that HOCl destroys cytochromes and iron-sulfur clusters and observed that oxygen uptake is abolished by HOCl and adenine nucleotides are lost. It was also observed that irreversible oxidation of cytochromes paralleled the loss of respiratory activity. One way of addressing the loss of oxygen uptake was by studying the effects of HOCl on succinate-dependent electron transport. Rosen et al. found that levels of reductable cytochromes in HOCl-treated cells were normal, and these cells were unable to reduce them. Succinate dehydrogenase was also inhibited by HOCl, stopping the flow of electrons to oxygen. Later studies revealed that Ubiquinol oxidase activity ceases first, and the still-active cytochromes reduce the remaining quinone. The cytochromes then pass the electrons to oxygen, which explains why the cytochromes cannot be reoxidized, as observed by Rosen et al. However, this line of inquiry was ended when Albrich et al. found that cellular inactivation precedes loss of respiration by using a flow mixing system that allowed evaluation of viability on much smaller time scales. This group found that cells capable of respiring could not divide after exposure to HOCl.

Depletion of adenine nucleotides Having eliminated loss of respiration Albrich et al. proposes that the cause of death may be due to metabolic dysfunction caused by depletion of adenine nucleotides. Barrette et al. studied the loss of adenine nucleotides by studying the energy charge of HOClexposed cells and found that cells exposed to HOCl were unable to step up their energy charge after addition of nutrients. The conclusion was that exposed cells have lost the ability to regulate their adenylate pool, based on the fact that metabolite uptake was only 45% deficient after exposure to HOCl and the observation that HOCl causes intracellular ATP hydrolysis. It was also confirmed that, at bacteriocidal levels of HOCl, cytosolic components are unaffected. So it was proposed that modification of some membranebound protein results in extensive ATP hydrolysis, and this, coupled with the cells inability to remove AMP from the cytosol, depresses metabolic function. One protein involved in loss of ability to regenerate ATP has been found to be ATP synthetase. Much of this research on respiration reconfirms the observation that relevant bacteriocidal reactions take place at the cell membrane.

Inhibition of DNA replication Recently it has been proposed that bacterial inactivation by HOCl is the result of inhibition of DNA replication. When bacteria are exposed to HOCl, there is a precipitous decline in DNA synthesis that precedes inhibition of protein synthesis, and closely parallels loss of viability. During bacterial genome replication, the origin of replication (oriC in E. coli) binds to proteins that are associated with the cell membrane, and it was observed that HOCl treatment decreases the affinity of extracted membranes for oriC, and this decreased affinity also parallels loss of viability. A study by Rosen et al. compared the rate of HOCl inhibition of DNA replication of plasmids with different replication origins and found that certain plasmids exhibited a delay in the inhibition of replication when compared to plasmids containing oriC. Rosen’s group proposed that inactivation of membrane proteins involved in DNA replication are the mechanism of action of HOCl.

Protein unfolding and aggregation HOCl is known to cause post-translational modifications to proteins, the notable ones being cysteine and methionine oxidation. A recent examination of HOCl's bactericidal role revealed it to be a potent inducer of protein aggregation. Hsp33, a chaperone known to be activated by oxidative heat stress, protects bacteria from the effects of HOCl by acting as a holdase, effectively preventing protein aggregation. Strains of E. coli and Vibrio cholerae lacking Hsp33 were rendered especially sensitive to HOCl. Hsp33 protected many essential proteins from aggregation and inactivation due to HOCl, which is a probable mediator of HOCl's bactericidal effects.

Hypochlorites Hypochlorites are the salts of hypochlorous acid; commercially important hypochlorites are calcium hypochlorite and sodium hypochlorite.

Production of hypochlorites using electrolysis Solutions of hypochlorites can be produced by electrolysis of an aqueous chloride solution. Chlorine gas is produced at the anode, while hydrogen forms at the cathode. Some of the chlorine gas produced will dissolve forming hypochlorite ions. Hypochlorites are also produced by the disproportionation of chlorine gas in alkaline solutions.

Safety HOCl is a strong oxidizer and can form explosive mixtures.

Chapter-4

Hydrogen Cyanide

Hydrogen cyanide

IUPAC name formonitrile Other names hydrogen cyanide; methanenitrile; hydrocyanic acid; prussic acid; Zyklon B Identifiers CAS number 74-90-8 PubChem 768 ChemSpider 19951400 UNII 2WTB3V159F EC number 200-821-6 1051 (anhydrous) UN number 1613 (aqueous soln., <20%) 1614 (adsorbed) RTECS number MW6825000 Properties Molecular formula HCN Molar mass 27.0253 g/mol Colorless gas or pale blue Appearance highly volatile liquid Density 0.687 g/cm3, liquid. Melting point −13.4 °C, 260 K, 8 °F Boiling point 25.6 °C, 299 K, 78 °F Solubility in water completely miscible Acidity (pKa) 9.21 Refractive index (nD) 1.2675 Viscosity 0.201 cP

Molecular shape Dipole moment

Structure Linear 2.98 D Thermochemistry

Std enthalpy of formation ΔfHo298 Std enthalpy of combustion ΔcHo298 Specific heat capacity, C EU Index EU classification R-phrases S-phrases Flash point Autoignition temperature

-4.999 kJ/g -24.6 kJ/g 1.328 J/(g·K) (gas) 2.612 J/(g·K) (liquid) Hazards 006-006-00-X Extremely flammable (F+) Very toxic (T+) Dangerous for the environment (N) R12, R26, R50/53 (S1/2), S7/9, S16, S36/37, S38, S45, S60, S61 −17.8 °C (−64 °F) 538 °C

Related compounds Cyanogen Cyanogen chloride Related compounds Trimethylsilyl cyanide Methylidynephosphane

Hydrogen cyanide (with the historical common name of Prussic acid) is a chemical compound with chemical formula HCN. Hydrogen cyanide is a colorless, extremely poisonous liquid that boils slightly above room temperature at 26 °C (79 °F). Hydrogen cyanide is a linear molecule, with a triple bond between carbon and nitrogen. A minor tautomer of HCN is HNC, hydrogen isocyanide. Hydrogen cyanide is weakly acidic with a pKa of 9.2. It partly ionizes in water solution to give the cyanide anion, CN–. A solution of hydrogen cyanide in water is called hydrocyanic acid. The salts of hydrogen cyanide are known as cyanides. HCN has a faint, bitter, burnt almond-like odor that some people are unable to detect owing to a genetic trait. The volatile compound has been used as inhalation rodenticide and human poison. Cyanide ions interfere with iron-containing respiratory enzymes. HCN is produced on an industrial scale and is a highly valuable precursor to many chemical compounds ranging from polymers to pharmaceuticals.

History of discovery Hydrogen cyanide was first isolated from a blue dye (Prussian blue) which had been known from 1704 but whose structure was unknown. It is now known to be a

coordination polymer with a complex structure and an empirical formula of hydrated ferric ferrocyanide. In 1752, the French chemist Pierre Macquer made the important step of showing that Prussian blue could be converted to iron oxide plus a volatile component and that these could be used to reconstitute the dye. The new component was what we now know as hydrogen cyanide. Following Macquer's lead, it was first isolated from Prussian blue in pure form and characterized about 1783 by the Swedish chemist Carl Wilhelm Scheele, and was eventually given the German name Blausäure (literally "Blue acid") because of its acidic nature in water and its derivation from Prussian blue. In English it became known popularly as Prussic acid.

Anion of Prussian blue In 1787 the French chemist Claude Louis Berthollet showed that Prussic acid did not contain oxygen, an important contribution to acid theory, which had hitherto postulated that acids must contain oxygen (hence the name of oxygen itself, which is derived from Greek elements that mean "acid-former" and are likewise calqued into German as Sauerstoff). In 1815 Joseph Louis Gay-Lussac deduced Prussic acid's chemical formula. The radical cyanide in hydrogen cyanide was given its name from the Greek word for blue, again owing to its derivation from Prussian blue.

Production and synthesis Hydrogen cyanide forms in at least limited amounts from many combinations of hydrogen, carbon, and ammonia. Hydrogen cyanide is currently produced in great quantities by several processes, as well as being a recovered waste product from the manufacture of acrylonitrile. In the year 2000, 732,552 tons were produced in the US.

The most important process is the Andrussov oxidation invented by Leonid Andrussow at IG Farben in which methane and ammonia react in the presence of oxygen at about 1200 °C over a platinum catalyst: 2 CH4 + 2 NH3 + 3 O2 → 2 HCN + 6 H2O The energy needed for the reaction is provided by the partial oxidation of methane and ammonia. Of lesser importance is the Degussa process (BMA process) in which no oxygen is added and the energy must be transferred indirectly through the reactor wall: CH4 + NH3 → HCN + 3H2 This reaction is akin to steam reforming, the reaction of methane and water to give carbon monoxide and hydrogen. In the Shawinigan Process, ammonia and natural gas are passed over coke. As practiced at BASF, formamide is heated and split into hydrogen cyanide and water: CH(O)NH2 → HCN + H2O In the laboratory, small amounts of HCN are produced by the addition of acids to cyanide salts of alkali metals: H+ + NaCN → HCN + Na+ This reaction is sometimes the basis of accidental poisonings because the acid converts a nonvolatile cyanide salt into the gaseous HCN.

Historical methods of production The demand for cyanides for mining operations in the 1890s was met by George Thomas Beilby, who patented a method to produce hydrogen cyanide by passing ammonia over glowing coal in 1892. This method was used until Hamilton Castner in 1894 developed a synthesis starting from coal, ammonia, and sodium yielding sodium cyanide, which reacts with acid to form gaseous HCN.

Applications HCN is the precursor to sodium cyanide and potassium cyanide, which are used mainly in mining. Via the intermediacy of cyanohydrins, a variety of useful organic compounds are prepared from HCN including the monomer methyl methacrylate, from acetone, the amino acid methionine, via the Strecker synthesis, and the chelating agents EDTA and NTA. Via the hydrocyanation process, HCN is added to butadiene to give adiponitrile, a precursor to Nylon 66.

Occurrence HCN is obtainable from fruits that have a pit, such as cherries, apricots, apples, and bitter almonds, from which almond oil and flavoring are made. Many of these pits contain small amounts of cyanohydrins such as mandelonitrile and amygdalin, which slowly release hydrogen cyanide. One hundred grams of crushed apple seeds can yield about 10 mg of HCN. Some millipedes release hydrogen cyanide as a defense mechanism, as do certain insects, such as some burnet moths. Hydrogen cyanide is contained in the exhaust of vehicles, in tobacco and wood smoke, and in smoke from burning nitrogencontaining plastics. So-called "bitter" roots of the cassava plant may contain up to 1 gram of HCN per kilogram.

HCN and the Scientific Theory of the origin of life Hydrogen cyanide has been discussed as a precursor to amino acids and nucleic acids. It is believed by some, for example, that HCN played a part in the origin of life. Although the relationship of these chemical reactions to the origin of life theory remains speculative, studies in this area have led to discoveries of new pathways to organic compounds derived from condensation of HCN.

HCN in space HCN has been detected in the interstellar medium. Since then, extensive studies have probed formation and destruction pathways of HCN in various environments and examined its use as a tracer for a variety of astronomical species and processes. HCN can be observed from ground-based telescopes through a number of atmospheric windows. The J=1→0, J=3→2, J= 4→3, and J=10→9 pure rotational transitions have all been observed. HCN is formed in interstellar clouds through one of two major pathways: via a neutralneutral reaction (CH2 + N → HCN + H) and via dissociative recombination (HCNH+ + e→ HCN + H). The dissociative recombination pathway is dominant by 30%; however, the HCNH+ must be in its linear form. Dissociative recombination with its structural isomer, H2NC+ produces hydrogen isocyanide (HNC), exclusively. HCN is destroyed in interstellar clouds through a number of mechanisms depending on the location in the cloud. In photon-dominated regions (PDRs), photodissociation dominates, producing CN (HCN + ν → CN + H). At further depths, photodissociation by cosmic rays dominate, producing CN (HCN + cr → CN + H). In the dark core, two competing mechanisms destroy it, forming HCN+ and HCNH+ (HCN + H+ → HCN+ + H; HCN + HCO+ → HCNH+ + CO). The reaction with HCO+ dominates by a factor of ~3.5. HCN has been used to analyze a variety of species and processes in the interstellar medium. It has been suggested as a tracer for dense molecular gas and as a tracer of stellar inflow in high-mass star-forming regions. Further, the HNC/HCN ratio has been shown to be an excellent method for distinguishing between PDRs and X-ray-dominated regions (XDRs).

Hydrogen cyanide as a poison and chemical weapon A hydrogen cyanide concentration of 300 mg/m3 in air will kill a human within about 10 minutes. It is estimated that hydrogen cyanide at a concentration of 3500 ppm (about 3200 mg/m3) will kill a human in about 1 minute. The toxicity is caused by the cyanide ion, which halts cellular respiration by inhibiting an enzyme in mitochondria called cytochrome c oxidase. Hydrogen cyanide absorbed into a carrier for use as a pesticide (under IG Farben's brand name Cyclone B, or in German Zyklon B, with the B standing for Blausäure) was employed by Nazi Germany in the mid-20th century in extermination camps. The same product is currently made in the Czech Republic under the trademark "Uragan D2." Hydrogen cyanide is also the agent used in gas chambers employed in judicial execution in some U.S. states, where it is produced during the execution by the action of sulfuric acid on an egg-sized mass of potassium cyanide. Hydrogen cyanide is commonly listed amongst chemical warfare agents known as blood agents. As a substance listed under Schedule 3 of the Chemical Weapons Convention as a potential weapon which has large-scale industrial uses, manufacturing plants in signatory countries which produce more than 30 tonnes per year must be declared to, and can be inspected by, the Organisation for the Prohibition of Chemical Weapons. Under the name prussic acid, HCN has been used as a killing agent in whaling harpoons. Hydrogen cyanide gas in air is explosive at concentrations over 5.6%, equivalent to 56000 ppm.

Chapter-5

Chromic Acid and Perchloric Acid

Chromic acid Chromic acid

IUPAC name Chromic acid Systematic name Dihydroxidodioxidochromium Other names Chromic(VI) acid Tetraoxochromic acid Identifiers CAS number 7738-94-5 PubChem 24425 ChemSpider 22834 EC number 231-801-5 ChEBI CHEBI:33143 Gmelin Reference 25982 Properties Molecular formula H2CrO4 Molar mass 118.01 g mol−1 Exact mass 117.935820456 g mol-1 Appearance Red crystals

Density Melting point Boiling point Solubility in water

1.201 g cm-3 197 °C, 470 K, 387 °F 250 °C, 523 K, 482 °F (decomposes) 1666.6 g dm-3

The term chromic acid is usually used for a mixture made by adding concentrated sulfuric acid to a dichromate, which may contain a variety of compounds, including solid chromium trioxide. This kind of chromic acid may be used as a cleaning mixture for glass. Chromic acid may also refer to the molecular species, H2CrO4 of which the trioxide is the anhydride. Chromic acid features chromium in an oxidation state of +6 (or VI). It is a strong and corrosive oxidising agent.

Molecular chromic acid

Partial predominance diagram for chromate Molecular chromic acid, H2CrO4, has much in common with sulfuric acid, H2SO4. Both are classified as strong acids, though only the first proton is lost easily. H2CrO4

, [HCrO4]- + H+

The pKa for the equilibrium is not well characterized. Reported values vary between about -0.8 to 1.6. The value at zero ionic strength is difficult to determine because half

dissociation only occurs in very acidic solution, at about pH zero, that is, with an acid concentration of about 1 mol dm−3. A further complication is that the ion [HCrO4]- has a marked tendency to dimerize, with the loss of a water molecule, to form the dichromate ion, [Cr2O7]22 [HCrO4]-

[Cr2O7]2- + H2O, log KD = 2.05.

Furthermore, The dichromate can be protonated. [HCr2O7]-

[Cr2O7]2- + H+, pK = 1.8

The pK value for this reaction shows that is can be ignored at pH > 4. Loss of the second proton occurs in the pH range 4-8, making the ion [HCrO4]- a weak acid Molecular chromic acid could in principle be made by adding chromium trioxide to water (c.f. manufacture of sulfuric acid). CrO3 + H2O

H2CrO4

but in practice the reverse reaction occurs when molecular chromic acid is dehydrated. This is what happens when concentrated sulfuric acid is added to a dichromate solution. At first the colour changes from orange (dichromate) to red (chromic acid) and then deep red crystals of chromium trioxide precipitate from the mixture, without further colour change. The colours are due to LMCT charge transfer transitions. Chromium trioxide is the anydride of molecular chromic acid. It is a Lewis acid and can react with a Lewis base, such as pyridine in a non-aqueous medium such as dichloromethane (Collins reagent).

Dichromic acid Dichromic acid, H2Cr2O7, (structure illustrated top right) is the fully protonated form of the dichromate ion and also can be seen as the product of adding chromium trioxide to molecular chromic acid. [Cr2O7]2- + 2H+

H2Cr2O7

H2CrO4 + CrO3

It is probably present in chromic acid cleaning mixtures along with the mixed chromicsulfuric acid H2CrSO7

Uses Chromic acid is an intermediate in chromium plating, and is also used in ceramic glazes, and colored glass. Because a solution of chromic acid in sulfuric acid (also known as a

sulfochromic mixture or chromosulfuric acid) is a powerful oxidizing agent, it can be used to clean laboratory glassware, particularly of otherwise insoluble organic residues. This application has declined due to environmental concerns. Furthermore the acid leaves trace amounts of paramagnetic chromic ions Cr(III) that can interfere with certain applications, such as NMR spectroscopy. This is especially the case for NMR tubes. Chromic acid has also been widely used in the band instrument repair industry, due to its ability to "brighten" raw brass. A chromic acid dip leaves behind a bright yellow patina on the brass. Due to growing health and environmental concerns, many have discontinued use of this chemical in their repair shops.

Reactions Chromic acid is capable of oxidizing many kinds of organic compounds and many variations on this reagent have been developed: 



 

Chromic acid in aqueous sulfuric acid and acetone is known as the Jones reagent, which will oxidize primary and secondary alcohols to carboxylic acids and ketones respectively, while rarely affecting unsaturated bonds. Pyridinium chlorochromate is generated from chromium trioxide and pyridinium chloride. This reagent converts primary alcohols to the corresponding aldehydes (R-CHO). Collins reagent is an adduct of chromium trioxide and pyridine used for diverse oxidations. Chromyl chloride, CrO2Cl2 is a well-defined molecular compound that is generated from chromic acid.

Illustrative transformations   

Oxidation of methylbenzenes to benzoic acids. Oxidative scission of indene to homophthalic acid. Oxidation of secondary alcohol to ketone (cyclooctanone) and nortricyclanone.

Use in qualitative organic analysis In organic chemistry, dilute solutions of hexavalent chromium can be used to oxidize primary or secondary alcohols to the corresponding aldehydes and ketones. Tertiary alcohol groups are unaffected. Because of the oxidation is signaled by a color change from orange to a blue-green, chromic acid is used as a qualitative analytical test for the presence of primary or secondary alcohols.

Alternative reagents In oxidations of alcohols or aldehydes into carboxylic acids, chromic acid is one of several reagents, including several that are catalytic. For example nickel(II) salts catalyze oxidations by bleach. Aldehydes are relatively easily oxidised to carboxylic acids, and

mild oxidising agents are sufficient. Silver(I) compounds have been used for this purpose. Each oxidant offers advantages and disadvantages.

Safety Hexavalent chromium compounds are toxic and carcinogenic. For this reason, chromic acid oxidation is not used on an industrial scale.

Perchloric acid Perchloric acid

CAS number PubChem ChemSpider EC number UN number ChEMBL RTECS number Molecular formula Molar mass Appearance Density Melting point

Identifiers 7601-90-3 24247 22669 231-512-4 1873 CHEMBL1161634 SC7500000 Properties HClO4 100.46 g/mol colorless liquid 1.67 g/cm3 -17 C (azeotrope)

-112 °C (anhydrous) Boiling point 203 C (azeotrope) Solubility in water miscible Acidity (pKa) ≈ −8 Hazards MSDS ICSC 1006 EU Index 017-006-00-4 Oxidant (O) EU classification Corrosive (C) R-phrases R5, R8, R35 S-phrases (S1/2), S23, S26, S36, S45 Related compounds Hydrochloric acid Hypochlorous acid Related compounds Chlorous acid Chloric acid

Perchloric acid is the inorganic compound with the formula HClO4. Usually encountered as an aqueous solution, this colourless compound is a strong acid comparable in strength to sulfuric and nitric acids, as well as a powerful oxidizing agent. It is useful for preparing perchlorate salts, especially ammonium perchlorate, an important rocket fuel. Perchloric acid is also dangerously corrosive and readily forms explosive mixtures.

Production Perchloric acid is produced industrially by two routes. The traditional method exploits the very 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 involves no salts, entails anodic oxidation of aqueous chlorine at a platinum electrode.

Laboratory preparations Treatment of barium perchlorate with sulfuric acid precipitates barium sulfate, leaving perchloric acid. It also can be made by mixing nitric acid with ammonium perchlorate. The reaction gives nitrous oxide and perchloric acid due to a concurrent reaction involving the ammonium ion.

Properties Anhydrous perchloric acid is an 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 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, which is even more dangerous: 2 HClO4 + P4O10 → Cl2O7 + "H2P4O11"

Uses Perchloric acid is mainly produced as a precursor to ammonium perchlorate, which is used as rocket fuel. The growth in rocketry has led to increased production of perchloric acid. Several million kilograms are produced annually.

As an acid Perchloric acid, a superacid, is one of the strongest Brønsted-Lowry acids. Its pKa is −10. It provides strong acidity without interference from potential nucleophiles such as sulfate or chloride that complicate the use of sulfuric and hydrochloric acids. 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. For similar reasons, it is a useful eluent in ion-exchange chromatography. It is also used for electropolishing/etching of aluminum, molybdenum, and other metals.

Safety Anhydrous and monohydrated perchloric acid are explosive, but the usual aqueous solutions are stable in the absence of organic compounds. It is very corrosive to skin and eyes. Upon contact with concentrated perchloric acid, organic materials such as cloth and wood ignite. Salts of perchloric acid are also powerful oxidizers that can be explosive. Perchlorate salts tend to be more stable than their chlorate counterparts, which has led to their increased use in pyrotechnic compositions due to safety concerns. Because of these hazards, perchloric acid is usually handled under fume hoods with wash-down and air scrubbing capabilities, which are not available on standard laboratory fume hoods. The crystalline form of the acid, which is explosive and shock sensitive, can precipitate on hood surfaces; washing down the hood interior solves this problem.

O'Connor Electro-Plating Company Disaster 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 plant, 25 other buildings and 40 automobiles were obliterated and 250 nearby homes were damaged. The bath was being used to electro-polish aluminum furniture, and despite knowing the explosive dangers of the bath, the plant chemist, Robert M. McGee, allowed production to continue after the refrigeration system, to keep the batch cool, failed. 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.

Chapter-6

Carbonic Acid and Hydrogen Fluoride

Carbonic acid Carbonic acid

IUPAC name Carbonic acid Other names Carbon dioxide solution; Dihydrogen carbonate; acid of air; Aerial acid; Hydroxymethanoic acid Identifiers CAS number 463-79-6 ChemSpider 747 KEGG C01353 ChEMBL CHEMBL1161632 Properties Molecular formula H2CO3 Molar mass 62.03 g/mol Density 1.0 g/cm3 (dilute soln.) Melting point n/a Solubility in water Exists only in solution 6.352 (pKa1) Acidity (pKa)

Carbonic acid is the inorganic compound with the formula H2CO3 (equivalently OC(OH)2). It is also a name sometimes given to solutions of carbon dioxide in water, because such solutions contain small amounts of H2CO3. Carbonic acid salts forms two kinds of salts, the carbonates and the bicarbonates. It is a weak acid.

Chemical equilibria When dissolved in water, carbon dioxide exists in equilibrium with carbonic acid: CO2 + H2O

H2CO3

The hydration equilibrium constant at 25 °C is called Kh, which in the case of carbonic acid is [H2CO3]/[CO2] = 1.70×10−3: hence, the majority of the carbon dioxide is not converted into carbonic acid, remaining as CO2 molecules. In the absence of a catalyst, the equilibrium is reached quite slowly. The rate constants are 0.039 s−1 for the forward reaction (CO2 + H2O → H2CO3) and 23 s−1 for the reverse reaction (H2CO3 → CO2 + H2O). Carbonic acid is used in the making of soft drinks, inexpensive and artificially carbonated sparkling wines, and other bubbly drinks. The addition of two equivalents of water to CO2 would give orthocarbonic acid, C(OH)4, which is unimportant in aqueous solution. Addition of base to an excess of carbonic acid gives bicarbonate. With excess base, carbonic acid reacts to give carbonate salts.

Role of carbonic acid in blood Carbonic acid is an intermediate step in the transport of CO2 out of the body via respiratory gas exchange. The hydration reaction of CO2 is generally very slow in the absence of a catalyst, but red blood cells contain carbonic anhydrase, which both increases the reaction rate and dissociates a hydrogen ion (H+) from the resulting carbonic acid, leaving bicarbonate (HCO3-) dissolved in the blood plasma. This catalysed reaction is reversed in the lungs, where it converts the bicarbonate back into CO2 and allows it to be expelled. This equilibration plays an important role as a buffer in mammalian blood.

Role of carbonic acid in ocean chemistry The oceans of the world have absorbed almost half of the CO2 emitted by humans from the burning of fossil fuels. The extra dissolved carbon dioxide has caused the ocean's average surface pH to shift by about 0.1 unit from pre-industrial levels. This process is known as ocean acidification.

Acidity of carbonic acid Carbonic acid is diprotic: it has two protons, which may dissociate from the parent molecule. Thus there are two dissociation constants, the first one for the dissociation into the bicarbonate (also called hydrogen carbonate) ion HCO3−: H2CO3 HCO3− + H+ Ka1 = 4.45×10−7 ; pKa1 = 6.352 at 25 °C.

With a pKa1 of 6.352, carbonic acid H2CO3 is almost 10x weaker acid than acetic acid. The second for the dissociation of the bicarbonate ion into the carbonate ion CO32−: HCO3− CO32− + H+ Ka2 = 4.69×10−11 ; pKa2 = 10.329 at 25 °C and Ionic Strength = 0.0. Care must be taken when quoting and using the first dissociation constant of carbonic acid. In aqueous solution carbonic acid only exists in equilibrium with carbon dioxide, and the concentration of H2CO3 is much lower than the dissolved CO2 concentration. Since it is not possible to distinguish between H2CO3 and dissolved CO2 (referred to as CO2(aq)) by conventional methods, H2CO3* is used to represent the two species when writing the aqueous chemical equilibrium equation. The equation may be rewritten as follows (cf. sulfurous acid): H2CO3* HCO3− + H+ Ka = 4.6×10−7(General Chemistry: An Integrated Approach Third Edition); pKa = 6.352 at 25 °C and Ionic Strength = 0.0.(NIST CRITICAL Database) Whereas this pKa is quoted as the dissociation constant of carbonic acid, it is ambiguous: it might better be referred to as the acidity constant of dissolved carbon dioxide, as it is particularly useful for calculating the pH of CO2-containing solutions.

pH and composition of carbonic acid solutions At a given temperature, the composition of a pure carbonic acid solution (or of a pure CO2 solution) is completely determined by the partial pressure of carbon dioxide above the solution. To calculate this composition, account must be taken of the above equilibria between the three different carbonate forms (H2CO3, HCO3− and CO32−) as well as of the hydration equilibrium between dissolved CO2 and H2CO3 with constant and of the following equilibrium between the dissolved CO2 and the gaseous CO2 above the solution: CO2(gas) CO2(dissolved) with (Henry constant)

where kH=29.76 atm/(mol/L) at 25°C

relation and The corresponding equilibrium equations together with the the charge neutrality condition result in six equations for − + − the six unknowns [CO2], [H2CO3], [H ], [OH ], [HCO3 ] and [CO32−], showing that the composition of the solution is fully determined by . The equation obtained for [H+] is a cubic whose numerical solution yields the following values for the pH and the different species concentrations:

pH

(atm)

[CO2]

[H2CO3]

(mol/L)

−8

7.00

3.36 × 10

1.0 × 10−7

6.94

1.0 × 10−6

(mol/L)

(mol/L)

(mol/L)

5.71 × 10

1.42 × 10

7.90 × 10−13

3.36 × 10−09

5.71 × 10−12

5.90 × 10−09

1.90 × 10−12

6.81

3.36 × 10−08

5.71 × 10−11

9.16 × 10−08

3.30 × 10−11

1.0 × 10−5

6.42

3.36 × 10−07

5.71 × 10−09

3.78 × 10−07

4.53 × 10−11

1.0 × 10−4

5.92

3.36 × 10−06

5.71 × 10−09

1.19 × 10−06

5.57 × 10−11

3.5 × 10−4

5.65

1.18 × 10−05

2.00 × 10−08

2.23 × 10−06

5.60 × 10−11

1.0 × 10−3

5.42

3.36 × 10−05

5.71 × 10−08

3.78 × 10−06

5.61 × 10−11

1.0 × 10−2

4.92

3.36 × 10−04

5.71 × 10−07

1.19 × 10−05

5.61 × 10−11

1.0 × 10−1

4.42

3.36 × 10−03

5.71 × 10−06

3.78 × 10−05

5.61 × 10−11

1.0 × 10+0

3.92

3.36 × 10−02

5.71 × 10−05

1.20 × 10−04

5.61 × 10−11

2.5 × 10+0

3.72

8.40 × 10−02

1.43 × 10−04

1.89 × 10−04

5.61 × 10−11

1.0 × 10+1

3.42

3.36 × 10−01

5.71 × 10−04

3.78 × 10−04

5.61 × 10−11



 





−13

[CO32−]

−09

1.0 × 10

−10

[HCO3−]

We see that in the total range of pressure, the pH is always largely lower than pKa2 so that the CO32− concentration is always negligible with respect to HCO3− concentration. In fact CO32− plays no quantitative role in the present calculation. For vanishing , the pH is close to the one of pure water (pH = 7) and the dissolved carbon is essentially in the HCO3− form. For normal atmospheric conditions ( atm), we get a slightly acid solution (pH = 5.7) and the dissolved carbon is now essentially in the CO2 form. From this pressure on, [OH−] becomes also negligible so that the ionized part of the solution is now an equimolar mixture of H+ and HCO3−. For a CO2 pressure typical of the one in soda drink bottles ( ~ 2.5 atm), we get a relatively acid medium (pH = 3.7) with a high concentration of dissolved CO2. These features contribute to the sour and sparkling taste of these drinks. Between 2.5 and 10 atm, the pH crosses the pKa1 value (3.60) giving a dominant H2CO3 concentration (with respect to HCO3−) at high pressures.

Remark As noted above, [CO32−] may be neglected for this specific problem, resulting in the following very precise analytical expression for [H+]:

Spectroscopic studies of carbonic acid Theoretical calculations show that the presence of even a single molecule of water causes carbonic acid to revert to carbon dioxide and water. In the absence of water, the

dissociation of gaseous carbonic acid is predicted to be very slow, with a half-life of 180,000 years. It has long been recognized that pure carbonic acid cannot be obtained at room temperatures (about 20 °C or about 70 °F). It can be generated by exposing a frozen mixture of water and carbon dioxide to high-energy radiation, and then warming to remove the excess water. The carbonic acid that remained was characterized by infrared spectroscopy. The fact that the carbonic acid was prepared by irradiating a solid H2O + CO2 mixture may suggest that H2CO3 might be found in outer space, where frozen ices of H2O and CO2 are common, as are cosmic rays and ultraviolet light, to help them react. The same carbonic acid polymorph (denoted beta-carbonic acid) was prepared by heating alternating layers of glassy aqueous solutions of bicarbonate and acid in vacuo, which causes protonation of bicarbonate, followed by removal of the solvent. Alpha-carbonic acid was prepared by the same technique using methanol rather than water as a solvent.

Hydrogen fluoride Hydrogen fluoride

CAS number PubChem ChemSpider UNII KEGG RTECS number Molecular formula Molar mass Appearance Density Melting point Boiling point Solubility in water

Identifiers 7664-39-3 16211014 14214 RGL5YE86CZ C16487 MW7875000 Properties HF 20.00634 g/mol colorless gas 1.15 g/l, gas (25 °C) 0.99 g/mL, liquid (19.5oC) −83.6 °C, 190 K, -118 °F 19.5 °C, 293 K, 67 °F miscible

Acidity (pKa) Refractive index (nD) Molecular shape Dipole moment Std enthalpy of formation ΔfHo298 Standard molar entropy So298

3.2 1.00001 Structure Linear 1.86 D Thermochemistry −13.66 kJ/g (gas) −14.99 kJ/g (liquid) 8.687 J/g K (gas)

Related compounds Hydrogen chloride Other anions Hydrogen bromide Hydrogen iodide Other cations Sodium fluoride Related compounds Hydrofluoric acid

Hydrogen fluoride is a chemical compound with the formula HF. It is the principal industrial source of fluorine, often in the aqueous form as hydrofluoric acid, and thus is the precursor to many important compounds including pharmaceuticals and polymers (e.g. Teflon). HF is widely used in the petrochemical industry and a component of many superacids. HF boils just below room temperature whereas the other hydrogen halides condense at much lower temperatures. Unlike the other hydrogen halides, HF is lighter than air and it is particularly penetrating, which can damage the lungs. Aqueous solutions of HF, called hydrofluoric acid, are strongly corrosive.

Structure

HF forms orthorhombic crystals, consisting of zig-zag chains of HF molecules. The HF molecules, with a short H–F bond of 0.95 Å, are linked to neighboring molecules by intermolecular H–F distances of 1.55 Å. Liquid HF also consists of chains of HF molecules, but the chains are shorter, consisting on average of only five or six molecules. The higher boiling point of HF relative to

analogous species, such as HCl, is attributed to hydrogen bonding between HF molecules, as indicated by the existence of chains even in the liquid state.

Acidity The acidity of hydrofluoric acid solutions vary with concentration owing to hydrogenbond interactions of the fluoride ion. Dilute solutions are weakly acidic with an acid ionization constant Ka = 6.6 × 10−4 (or pKa = 3.18), in contrast to corresponding solutions of the other hydrogen halides which are strong acids. Concentrated solutions of hydrogen fluoride are much more strongly acid than implied by this value, as shown by measurements of the Hammett acidity function H0 (or “effective pH”). For 100%, HF has an H0, estimated to be between −10.2 and −11, which is comparable to the value −12 for sulfuric acid. In thermodynamic terms, HF solutions are highly non-ideal, with the activity of HF increasing much more rapidly than its concentration. The weak acidity in dilute solution is sometimes attributed to the high H—F bond strength, which combines with the high dissolution enthalpy of HF to outweigh the more negative enthalpy of hydration of the fluoride ion However Giguère and Turrell have shown by infrared spectroscopy that the predominant solute species is the hydrogen-bonded ion-pair [H3O+•F−], which suggests that the ionization can be described as a double equilibrium: H2O + HF

[H3O+•F−]

H3O+ + F−

The first equilibrium to the right and the second to the left, meaning that HF is extensively dissociated, but that the tight ion pairs reduce the thermodynamic activity coefficient of H3O+, so that the solution is effectively less acidic. In concentrated solution, the additional HF causes the ion pair to dissociate with formation of the hydrogen-bonded hydrogen difluoride ion. [H3O+•F−] + HF

H3O+ + HF2−

The increase in free H3O+ due to this reaction accounts for the rapid increase in acidity, while fluoride ions are stabilized (and become less basic) by strong hydrogen bonding to HF to form HF2−. This interaction between the acid and its own conjugate base is an example of homoconjugation. At the limit of 100% liquid HF, there is autoionization 2 HF

H2F+ + F−

that forms an extremely acidic solution (H0 = −11). The acidity of anhydrous HF can be increased even further by the addition of Lewis acids such as SbF5, which can reduce H0 to −21.

Production and uses Hydrogen fluoride is produced as by the action of sulfuric acid on pure grades of the mineral fluorite and also as a side-product of the extraction of the fertilizer precursor phosphoric acid from various minerals. The anhydrous compound hydrogen fluoride is more commonly used than its aqueous solution, hydrofluoric acid. HF serves as a catalyst in alkylation processes in oil refineries. A component of high-octane gasoline called "alkylate" is generated in Alkylation units that combine C3 and C4 olefins and isobutane to generate gasoline. HF is a reactive solvent in the electrochemical fluorination of organic compounds. In this approach, HF is oxidized in the presence of a hydrocarbon and the fluorine replaces C–H bonds with C–F bonds. Perfluorinated carboxylic acids and sulfonic acids are produced in this way. Hydrogen fluoride is an important catalyst used in the majority of the installed linear alkyl benzene production in the world. The process involves dehydrogenation of nparaffins to olefins, and subsequent reaction with benzene using HF as catalyst. Elemental fluorine, F2, is prepared by electrolysis of a solution of HF and potassium bifluoride. The potassium bifluoride is needed because anhydrous hydrogen fluoride does not conduct electricity. Several million kilograms of F2 are produced annually.

Health effects Upon contact with moisture, including tissue, hydrogen fluoride immediately converts to hydrofluoric acid, which is highly corrosive and toxic, and requires immediate medical attention upon exposure.

Chapter-7

Hydrogen Isocyanide and Isocyanic acid

Hydrogen isocyanide Hydrogen isocyanide

IUPAC name hydrogen isocyanide or methanidylidyneazanium Other names isocyanic acid; hydroisocyanic acid; nitriliomethanide Identifiers PubChem 6432654 ChemSpider 4937885 Properties Molecular formula HNC Molar mass 27.03 g/mol

Hydrogen isocyanide is a chemical with the molecular formula HNC. It is a minor tautomer of hydrogen cyanide, HCN). Its importance in the field of astrochemistry is linked to its ubiquity in the interstellar medium.

Nomenclature Both 'hydrogen isocyanide' and 'methanidylidyneazanium' are correct IUPAC names for HNC. Currently there is no Preferred IUPAC name. The second one is according to the substitutive nomenclature rules, derived from the parent hydride azane (NH3) and the anion methanide (C-).

Molecular properties Hydrogen isocyanide (HNC) is a linear triatomic molecule with C∞v point group symmetry. It is a zwitterion and an isomer of hydrogen cyanide (HCN). Both HNC and HCN have large, similar dipole moments, with respectively μHNC=3.05 Debye and

μHCN=2.98 Debye. These large dipole moments facilitate the easy observation of these species in the interstellar medium.

HNC−HCN tautomerism As HNC is higher in energy than HCN by 3920 cm−1 (46.9 kJ/mol), the naïve expectation would be that the two would have an equilibrium ratio at T<100K of ([HNC]/[HCN])eq,T<100K<10−25. However, observations show a very different conclusion; ([HNC]/[HCN])observed is much higher than 10−25, and is in fact on the order of unity in cold environments. This is because of the potential energy path of the tautomerization reaction; there is an activation barrier on the order of roughly 12,000 cm−1 for the tautomerization to occur, which corresponds to a temperature at which HNC would already have been destroyed by neutral-neutral reactions.

Spectral properties In practice, HNC is almost exclusively observed astronomically using the J=1→0 transition. This transition occurs at ~90.66 GHz, which is a point of good visibility in the atmospheric window, thus making astronomical observations of HNC particularly simple. Many other related species (including HCN) are observed in roughly the same window.

Significance in the interstellar medium HNC is intricately linked to the formation and destruction of numerous other molecules of importance in the interstellar medium - aside from the obvious partners HCN, HCNH+, and CN, HNC is linked to the abundances of many other compounds, either directly of through few degrees of separation. As such, an understanding of the chemistry of HNC leads to an understanding of countless other species - HNC is an integral piece in the complex puzzle representing interstellar chemistry. Furthermore, HNC (alongside HCN) is a commonly used tracer of dense gas in molecular clouds, as referenced in this paper. Aside from the potential to use HNC to investigate gravitational collapse as the means of star formation, HNC abundance (relative to the abundance of other nitrogenous molecules) can be used to determine the evolutionary stage of protostellar cores. This is demonstrated in the aforementioned paper by Tennekes et al. In the same paper, the authors also elaborate on the HNC/HCN abundance ratio as a means of determining the temperature of the environment. This paper demonstrates a myriad of uses for knowledge of the abundance of HNC. In it, the HCO+/HNC line ratio is used to good effect as a measure of density of gas. This information provides great insight into the mechanisms of the formation of (Ultra)Luminous Infrared Galaxies ((U)LIRGs), as it provides data on the nuclear environment, star formation, and even black hole fueling. Furthermore, the HNC/HCN line ratio is used to distinguish between photon-dissociation regions (PDRs) and X-ray-dissociation regions (XDRs) on the basis that [HNC]/[HCN] is roughly unity in PDR sources, but greater than unity in XDR sources.

The study of HNC is a relatively simple pursuit, and this is one of the greatest motivations for its study. Aside from having its J=1→0 transition in a clear portion of the atmospheric window, as well as having numerous isotopomers also available for easy study, and in addition to having a large dipole moment that makes observations particularly simple, HNC is, in its molecular nature, a quite simple molecule. This makes the study of the reaction pathways that lead to its formation and destruction a good means of obtaining insight to the workings of these reactions in space. Furthermore, the study of the tautomerization of HNC to HCN (and vice versa), which has been studied extensively, has been suggested as a model by which more complicated isomerization reactions can be studied.

Chemistry in the interstellar medium HNC is found primarily in dense molecular clouds, though it is ubiquitous in the interstellar medium. HNC is formed primarily through the dissociative recombination of HNCH+ and H2NC+, and it is destroyed primarily through ion-neutral reactions with H3+ and C+. Rate constants are taken from udfa.net, and data on fractional abundances is taken from here. Rate calculations were done at 3.16x105 years, which is considered early time, and at 20K, which is a typical temperature for dense molecular clouds. Formation Reactions Reactant Reactant Product Product Rate 1 2 1 2 constant

Rate/[H2 Relative ]2 Rate

HCNH+

e-

HNC

H

9.50e-8

4.76e-25

3.4

H2NC+

e-

HNC

H

1.80e-7

1.39e-25

1.0

Destruction Reactions Reactant Reactant Product Product Rate 1 2 1 2 constant

Rate/[H2 Relative ]2 Rate

H3+

HNC

HCNH+

H2

8.10e-9

1.26e-24

1.7

C+

HNC

C2N+

H

3.10e-9

7.48e-25

1.0

These four reactions are merely the four most dominant, and thus the most significant in the formation of the HNC abundances in dense molecular clouds; there are dozens more reactions for the formation and destruction of HNC. Though these reactions primarily lead to various protonated species, HNC is linked closely to the abundances of many other nitrogen containing molecules, for example, NH3 and CN. The pathways leading between these species can be found in the paper by Turner et al. that is linked above. The abundance HNC is also inexorably linked to the abundance of HCN, and the two tend to exist in a specific ratio based on the environment, as noted in the paper by Hiraoka et al. that is linked above. This is because the reactions that form HNC can often also form HCN, and vice versa, depending on the conditions in which the reaction occurs, and also

that there exist isomerization reactions for the two species. A simplified pathway showing many of the methods of HNC formation and destruction is available as Fig. 10 from Turner et al.

Astronomical detections HNC was first detected in June 1970 by L. E. Snyder and D. Buhl using the 36-foot radio telescope of the National Radio Astronomy Observatory (NRAO). The main molecular isotope, H12C14N, was observed via its J=1→0 transition at 88.6 GHz in six different sources: W3 (OH), Orion A, Sgr A(NH3A), W49, W51, DR 21(OH). A secondary molecular isotope, H13C14N, was observed via its J=1→0 transition at 86.3 GHz in only two of these sources: Orion A and Sgr A(NH3A). HNC was then later detected extragalactically in 1988 by C. Henkel, R. Mauersberger, and P. Schilke using the IRAM 30-m telescope at the Pico de Veleta in Spain. It was observed via its J=1→0 transition at 90.7 GHz toward IC342. A number of detections have been made towards the end of confirming the temperature dependence of the abundance ratio of [HNC]/[HCN]. A strong fit between temperature and the abundance ratio would allow observers to spectroscopically detect the ratio and then extrapolate the temperature of the environment, thus gaining great insight into the environment of the species. In 1986, Goldsmith et al. measured the abundances of rare isotopes of HNC and HCN along the OMC-1 and determined that the abundance ratio varies by more than an order of magnitude in warm regions versus cold regions. In 1992, Schilke et al. measured abundances of HNC, HCN, and deuterated analogs along the OMC-1 ridge and core and confirmed the temperature dependence of the abundance ratio. Helmich and van Dishoeck performed a survey of the W 3 Giant Molecular Cloud in 1997 in which they detected over 24 different molecular isotopes, comprising over 14 distinct chemical species, including HNC, HN13C, and H15NC. This survey further confirmed the temperature dependence of the abundance ratio, [HNC]/[HCN], this time ever confirming the dependence of the isotopomers. These are not the only detections of importance of HNC in the interstellar medium. In 1997, Pratap et al. observed HNC along the TMC-1 ridge and found that its abundance relative to HCO+ to be constant along the ridge – this led credence to the reaction pathway that posits that HNC is derived initially from HCO+. One significant astronomical detection that demonstrated the practical use of observing HNC occurred in 2006 by Tennekes et al., in which the authors detected and then used the abundances of various nitrogenous compounds (including HN13C and H15NC) to determine the stage of evolution of the protostellar core Cha-MMS1 based on the relative magnitudes of the abundances.

Isocyanic acid Isocyanic acid

CAS number PubChem ChemSpider Molecular formula Molar mass Appearance Density Melting point Boiling point Solubility in water Solubility Main hazards

IUPAC name Isocyanic acid Identifiers 75-13-8 , 420-05-3 (cyanic acid) 6347 6107 Properties HNCO 43.03 g/mol Colorless liquid or gas (b.p. near room temperature) 1.14 g/cm3 (20 °C) -86 °C 23.5 °C Dissolves Soluble in benzene, toluene, ether Hazards Poisonous

Isocyanic acid is an inorganic compound with the formula HNCO, discovered in 1830 by Liebig and Wöhler. This colourless substance is volatile and poisonous, with a boiling point of 23.5 °C. Isocyanic acid is the simplest stable chemical compound that contains carbon, hydrogen, nitrogen, and oxygen, the four most commonly-found elements in organic chemistry and biology.

Preparation and reactions Isocyanic acid can be made by protonation of the cyanate anion, such as from salts like potassium cyanate, by either gaseous hydrogen chloride or acids such as oxalic acid.

H+ + NCO → HNCO

HNCO also can be made by the high-temperature thermal decomposition of cyanuric acid, a trimer. C3H3N3O3 → 3 HNCO Isocyanic acid hydrolyses to carbon dioxide and ammonia: HNCO + H2O → CO2 + NH3 At sufficiently high concentrations, isocyanic acid oligomerizes to give cyanuric acid and cyamelide, a polymer. These species usually are easily separated from liquid- or gasphase reaction products. Dilute solutions of isocyanic acid are stable in inert solvents, e.g. ether and chlorinated hydrocarbons. Isocyanic acid reacts with amines to give ureas (carbamides): HNCO + RNH2 → RNHC(O)NH2. This reaction is called carbamylation. HNCO adds across electron-rich double bonds, such as vinylethers, to give the corresponding isocyanates.

Isomers: Cyanic acid and fulminic acid Low-temperature photolysis of solids containing HNCO has been shown to make H-OC≡N, known as cyanic acid or hydrogen cyanate; it is a tautomer of isocyanic acid. Pure cyanic acid has not been isolated, and isocyanic acid is the predominant form in all solvents. Note that sometimes information presented for cyanic acid in reference books is actually for isocyanic acid. Cyanic and isocyanic acids are isomers of fulminic acid (H-C=N-O), an unstable compound.

Chapter-8

Nitric Acid

Nitric acid

IUPAC name Nitric acid Identifiers CAS number 7697-37-2 PubChem 944 ChemSpider 919 UNII 411VRN1TV4 EC number 231-714-2 UN number 2031 KEGG D02313 MeSH Nitric+acid ChEBI CHEBI:48107 ChEMBL CHEMBL1352 RTECS number QU5775000 Gmelin Reference 1576 3DMet B00068 Properties Molecular formula HNO3

63.01 g mol−1 62.995642903 g mol-1 Colorless liquid 1.5129 g cm-3 -42 °C, 231 K, -44 °F 83 °C, 356 K, 181 °F (68% solution boils Boiling point at 121 °C) Solubility in water Completely miscible -1.4 Acidity (pKa) Refractive index 1.397 (16.5 °C) (nD) Dipole moment 2.17 ± 0.02 D Hazards ICSC 0183 MSDS PCTL Safety Website EU Index 007-004-00-1 Toxic (T) EU classification Corrosive (C) Oxidant (O) R-phrases R8 R35 S-phrases (S1/2) S23 S26 S36 S45 Flash point Non-flammable Related compounds Other anions Nitrous acid Sodium nitrate Other cations Potassium nitrate Ammonium nitrate Related compounds Dinitrogen pentoxide Molar mass Exact mass Appearance Density Melting point

Nitric acid (HNO3), also known as aqua fortis and spirit of nitre, is a highly corrosive and toxic strong acid. Colorless when pure, older samples tend to acquire a yellow cast due to the accumulation of oxides of nitrogen. If the solution contains more than 86% nitric acid, it is referred to as fuming nitric acid. Depending on the amount of nitrogen dioxide present, fuming nitric acid is further characterized as white fuming nitric acid or red fuming nitric acid, at concentrations above 95%. At room temperature, nitric acid tends to rapidly develop a yellow color due to decomposition. Nitric acid is also commonly used as a strong oxidizing agent.

Properties Pure anhydrous nitric acid (100%) is a colorless mobile liquid with a density of 1.512 g/cm3 which solidifies at −42 °C to form white crystals and boils at 83 °C. When boiling in light, and slowly even at room temperature, there is a partial decomposition with the formation of nitrogen dioxide following the reaction:

4 HNO3 → 2 H2O + 4 NO2 + O2 Thus, anhydrous nitric acid should be stored below 0 °C to avoid decomposition. The nitrogen dioxide (NO2) remains dissolved in the nitric acid coloring it yellow, or red at higher temperatures. While the pure acid tends to give off white fumes when exposed to air, acid with dissolved nitrogen dioxide gives off reddish-brown vapors, leading to the common name "red fuming acid" or "fuming nitric acid". Fuming nitric acid is also referred to as 16 molar nitric acid. It is the most concentrated form of nitric acid at Standard Temperature and Pressure (STP). Nitric acid is miscible with water and distillation gives a maximum-boiling azeotrope with a concentration of 68% HNO3 and a boiling temperature of 120.5 °C at 1 atm, which is the ordinary concentrated nitric acid of commerce. Two solid hydrates are known; the monohydrate (HNO3·H2O) and the trihydrate (HNO3·3H2O). Nitrogen oxides (NOx) are soluble in nitric acid and this property influences more or less all the physical characteristics depending on the concentration of the oxides. These mainly include the vapor pressure above the liquid and the boiling temperature, as well as the color mentioned above. Nitric acid is subject to thermal or light decomposition with increasing concentration and this may give rise to some non-negligible variations in the vapor pressure above the liquid because the nitrogen oxides produced dissolve partly or completely in the acid.

Acid-base properties Nitric acid is normally considered to be a strong acid at ambient temperatures. There is some disagreement over the value of the acid dissociation constant, though the pKa value is usually reported as less than –1. This means that the nitric acid in solution is fully dissociated except in extremely acidic solutions. The pKa value rises to 1 at a temperature of 250 °C. Nitric acid can act as a base with respect to an acid such as sulfuric acid. HNO3 + 2H2SO4

NO2+ + H3O+ + 2HSO4–; K ~ 22

The nitronium ion, NO2+, is the active reagent in aromatic nitration reactions. Since nitric acid has both acidic and basic properties it can undergo an autoprotolysis reaction, similar to the self-ionization of water 2HNO3

NO2+ + NO3– + H2O

Oxidizing properties

Reactions with metals Nitric acid reacts with most metals. This characteristic has made it a common agent to be used in acid tests. Some precious metals, such as pure gold do not react with nitric acid, though pure gold does react with aqua regia, a mixture of concentrated nitric acid and hydrochloric acid. However, some less noble metals (Ag, Cu, ...) present in some gold alloys relatively poor in gold such as colored gold can be easily oxidized and dissolved by nitric acid, leading to colour changes of the gold-alloy surface. Nitric acid is used as a cheap means in jewelry shops to quickly spot low-gold alloys (< 14 carats) and to rapidly assess the gold purity. Strongly electropositive metals, such as magnesium react with nitric acid as with other acids, reducing the hydrogen ion. Mg + 2 H+ → Mg2+ + H2 With less electropositve metals the products depend on temperature and the acid concentration. For example, copper reacts with dilute nitric acid at ambient temperatures with a 3:8 stoichiometry. 3 Cu + 8 HNO3 → 3 Cu2+ + 2 NO + 4 H2O + 6 NO3The nitric oxide produced may react with atmospheric oxygen to give nitrogen dioxide. With more concentrated nitric acid, nitrogen dioxide is produced directly in a reaction with 1:4 stoichiometry. Cu + 4 H+ + 2 NO3− → Cu2+ + 2 NO2 + 2 H2O

Passivation Although chromium (Cr), iron (Fe) and aluminium (Al) readily dissolve in dilute nitric acid, the concentrated acid forms a metal oxide layer that protects the metal from further oxidation, which is called passivation. Typical passivation concentrations range from 18% to 22% by weight.

Reactions with non-metals Being a powerful oxidizing acid, nitric acid reacts violently with many organic materials and the reactions may be explosive. Reaction with non-metallic elements, with the exceptions of nitrogen, oxygen, noble gases, silicon and halogens, usually oxidizes them to their highest oxidation states as acids with the formation of nitrogen dioxide for concentrated acid and nitric oxide for dilute acid.

C + 4 HNO3 → CO2 + 4 NO2 + 2 H2O or 3 C + 4 HNO3 → 3 CO2 + 4 NO + 2 H2O

Xanthoproteic test Nitric acid reacts with proteins to form yellow nitrated products. This reaction is known as the xanthoproteic reaction. This test is carried out by adding concentrated nitric acid to the substance being tested, and then heating the mixture. If proteins are present that contains amino acids with aromatic rings, the mixture turns yellow. Upon adding a strong base such as liquid ammonia, the color turns orange. These color changes are caused by nitrated aromatic rings in the protein. Xanthoproteins are formed when the acid contacts epithelial cells and are indicative of inadequate safety precautions when handling nitric acid.

Grades The concentrated nitric acid of commerce consists of the maximum boiling azeotrope of nitric acid and water. Technical grades are normally 68% HNO3, (approx 15 molar), while reagent grades are specified at 70% HNO3. The density of concentrated nitric acid is 1.42 g/mL. An older density scale is occasionally seen, with concentrated nitric acid specified as 42° Baumé. White fuming nitric acid, also called 100% nitric acid or WFNA, is very close to anhydrous nitric acid. One specification for white fuming nitric acid is that it has a maximum of 2% water and a maximum of 0.5% dissolved NO2. Anhydrous nitric acid has a density of 1.513 g/mL and has the approximate concentration of 24 molar. A commercial grade of fuming nitric acid, referred to in the trade as "strong nitric acid" contains 90% HNO3 and has a density of 1.50 g/mL. This grade is much used in the explosives industry. It is not as volatile nor as corrosive as the anhydrous acid and has the approximate concentration of 21.4 molar. Red fuming nitric acid, or RFNA, contains substantial quantities of dissolved nitrogen dioxide (NO2) leaving the solution with a reddish-brown color. One formulation of RFNA specifies a minimum of 17% NO2, another specifies 13% NO2. Because of the dissolved nitrogen dioxide, the density of red fuming nitric acid is lower at 1.490 g/mL. An inhibited fuming nitric acid (either IWFNA, or IRFNA) can be made by the addition of 0.6 to 0.7% hydrogen fluoride (HF). This fluoride is added for corrosion resistance in metal tanks. The fluoride creates a metal fluoride layer that protects the metal.

Industrial production Nitric acid is made by reacting nitrogen dioxide (NO2) with water. 3 NO2 + H2O → 2 HNO3 + NO Normally, the nitric oxide produced by the reaction is reoxidized by the oxygen in air to produce additional nitrogen dioxide. Bubbling nitrogen dioxide through hydrogen peroxide can help to improve acid yield. 2 NO2 + H2O2 → 2 HNO3 Almost pure nitric acid can be made by adding sulfuric acid to a nitrate salt, and heating the mixture with an oil bath. A condenser is used to condense the nitric acid fumes that bubble out of the solution. 2 NaNO3 + H2SO4 → 2 HNO3 + Na2SO4 Dilute nitric acid may be concentrated by distillation up to 68% acid, which is a maximum boiling azeotrope containing 32% water. In the laboratory, further concentration involves distillation with either sulfuric acid or magnesium nitrate which act as dehydrating agents. Such distillations must be done with all-glass apparatus at reduced pressure, to prevent decomposition of the acid. Industrially, strong nitric acid is produced by dissolving additional nitrogen dioxide in 68% nitric acid in an absorption tower. Dissolved nitrogen oxides are either stripped in the case of white fuming nitric acid, or remain in solution to form red fuming nitric acid. More recently, electrochemical means have been developed to produce anhydrous acid from concentrated nitric acid feedstock. Commercial grade nitric acid solutions are usually between 52% and 68% nitric acid. Production of nitric acid is via the Ostwald process, named after German chemist Wilhelm Ostwald. In this process, anhydrous ammonia is oxidized to nitric oxide, which is then reacted with oxygen in air to form nitrogen dioxide. This is subsequently absorbed in water to form nitric acid and nitric oxide. The nitric oxide is cycled back for reoxidation. By using ammonia derived from the Haber process, the final product can be produced from nitrogen, hydrogen, and oxygen which are derived from air and natural gas as the sole feedstocks. Prior to the introduction of the Haber process for the production of ammonia in 1913, nitric acid was produced using the Birkeland–Eyde process, also known as the arc process. This process is based upon the oxidation of atmospheric nitrogen by atmospheric oxygen to nitric oxide at very high temperatures. An electric arc was used to provide the high temperatures, and yields of up to 4% nitric oxide were obtained. The nitric oxide was cooled and oxidized by the remaining atmospheric oxygen to nitrogen dioxide, and this was subsequently absorbed in dilute nitric acid. The process was very energy

intensive and was rapidly displaced by the Ostwald process once cheap ammonia became available.

Laboratory synthesis In laboratory, nitric acid can be made from copper(II) nitrate or by reacting approximately equal masses of a nitrate salt with 96% sulfuric acid (H2SO4), and distilling this mixture at nitric acid's boiling point of 83 °C until only a white crystalline mass, a metal sulfate, remains in the reaction vessel. The red fuming nitric acid obtained may be converted to the white nitric acid. H2SO4 + NO−3 → HSO−4(s) + HNO3(g) The dissolved NOx are readily removed using reduced pressure at room temperature (1030 min at 200 mmHg or 27 kPa) to give white fuming nitric acid. This procedure can also be performed under reduced pressure and temperature in one step in order to produce less nitrogen dioxide gas.

Uses

Nitric acid in a laboratory. The main use of nitric acid is for the production of fertilizers; other important uses include the production of explosives, etching and dissolution of metals, especially as a

component of aqua regia for the purification and extraction of gold, and in chemical synthesis.

Rocket fuel Nitric acid has been used in various forms as the oxidizer in liquid-fueled rockets. These forms include red fuming nitric acid, white fuming nitric acid, mixtures with sulfuric acid, and these forms with HF inhibitor. IRFNA (inhibited red fuming nitric acid) was one of 3 liquid fuel components for the BOMARC missile.

Chemical reagent In elemental analysis by ICP-MS, ICP-AES, GFAA, and Flame AA, dilute nitric acid (0.5 to 5.0 %) is used as a matrix compound for determining metal traces in solutions. Ultrapure trace metal grade acid is required for such determination, because small amounts of metal ions could affect the result of the analysis. It is also typically used in the digestion process of turbid water samples, sludge samples, solid samples as well as other types of unique samples which require elemental analysis via ICP-MS, ICP-OES, ICP-AES, GFAA and FAA. Typically these digestions use a 50% solution of the purchased HNO3 mixed with Type 1 DI Water. In organic synthesis, nitric acid may be used to introduce the nitro group. When used with sulfuric acid, it generates the nitronium ion, which electrophilically reacts with aromatic compounds such as benzene.

Woodworking In a low concentration (approximately 10%), nitric acid is often used to artificially age pine and maple. The color produced is a grey-gold very much like very old wax or oil finished wood (wood finishing).

Other uses A solution of nitric acid and alcohol, Nital, is used for etching of metals to reveal the microstructure. ISO 14104 is one of the standards detailing this well known procedure. Commercially available aqueous blends of 5-30% nitric acid and 15-40% phosphoric acid are commonly used for cleaning food and dairy equipment primarily to remove precipitated calcium and magnesium compounds (either deposited from the process stream or resulting from the use of hard water during production and cleaning).

Safety Nitric acid is a powerful oxidizing agent, and the reactions of nitric acid with compounds such as cyanides, carbides, and metallic powders can be explosive. Reactions of nitric

acid with many organic compounds, such as turpentine, are violent and hypergolic (i.e., self-igniting). Due to its properties it is stored away from bases and organics. Concentrated nitric acid dyes human skin yellow due to a reaction with the keratin. These yellow stains turn orange when neutralized.

History The first mention of nitric acid is in Pseudo-Geber´s De Inventione Veritatis, wherein it is obtained by calcining a mixture of niter, alum and blue vitriol. It was again described by Albert the Great in the 13th century and by Ramon Lull, who prepared it by heating niter and clay and called it "eau forte" (aqua fortis). Glauber devised the process still used today to obtain it, namely by heating niter with strong sulfuric acid. Its true nature was determined by Lavoisier in (1776), when he showed that it contained oxygen, whilst in 1785 Henry Cavendish determined its constitution and showed that it could be synthesized by passing a stream of electric sparks through moist air.

Chapter-9

Hydrogen Peroxide

Hydrogen peroxide

CAS number PubChem ChemSpider UNII EC number UN number KEGG ChEMBL IUPHAR ligand RTECS number Molecular formula Molar mass Appearance Density Melting point Boiling point Solubility in water Solubility Acidity (pKa)

IUPAC name dihydrogen dioxide Other names Dioxidane Identifiers 7722-84-1 784 763 BBX060AN9V 231-765-0 2015 (>60% soln.) 2014 (20–60% soln.) 2984 (8–20% soln.) D00008 CHEMBL71595 2448 MX0900000 (>90% soln.) MX0887000 (>30% soln.) Properties H2O2 34.0147 g/mol Very light blue color; colorless in solution 1.110 g/cm3 (20 °C, 30-percent) 1.450 g/cm3 (20 °C, pure) -0.43 °C, 273 K, 31 °F 150.2 °C, 423 K, 302 °F Miscible soluble in ether 11.62

Refractive index (nD) Viscosity Dipole moment

1.34 1.245 cP (20 °C) 2.26 D Thermochemistry

Std enthalpy of -4.007 kJ/g formation ΔfHo298 Specific heat capacity, 1.267 J/g K (gas) C 2.619 J/g K (liquid) Hazards MSDS ICSC 0164 (>60% soln.) EU Index 008-003-00-9 Oxidant (O) EU classification Corrosive (C) Harmful (Xn) R-phrases R5, R8, R20/22, R35 (S1/2), S17, S26, S28, S36/37/39, S-phrases S45 Flash point Non-flammable 1518 mg/kg LD50 Related compounds Water Ozone Related compounds Hydrazine Hydrogen disulfide

Hydrogen peroxide (H2O2) is an oxidizer commonly used as a bleach. It is the simplest peroxide (a compound with an oxygen-oxygen single bond). Hydrogen peroxide is a clear liquid, slightly more viscous than water, that appears colorless in dilute solution. It is used as a disinfectant, antiseptic, oxidizer, and in rocketry as a propellant. The oxidizing capacity of hydrogen peroxide is so strong that it is considered a highly reactive oxygen species. Hydrogen peroxide is naturally produced in organisms as a by-product of oxidative metabolism. Nearly all living things (specifically, all obligate and facultative aerobes) possess enzymes known as peroxidases, which harmlessly and catalytically decompose low concentrations of hydrogen peroxide to water and oxygen.

Structure and properties

H2O2 adopts a nonplanar structure of C2 symmetry. Although chiral, the molecule undergoes rapid racemization. The flat shape of the anti conformer would minimize steric repulsions, the 90° torsion angle of the syn conformer would optimize mixing between the filled p-type orbital of the oxygen (one of the lone pairs) and the LUMO of the vicinal O-H bond. The observed anticlinal "skewed" shape is a compromise between the two conformers. Despite the fact that the O-O bond is a single bond, the molecule has a high barrier to complete rotation of 29.45 kJ/mol (compared with 12.5 kJ/mol for the rotational barrier of ethane). The increased barrier is attributed to repulsion between one lone pair and other lone pairs. The bond angles are affected by hydrogen bonding, which is relevant to the structural difference between gaseous and crystalline forms; indeed a wide range of values is seen in crystals containing molecular H2O2.

Comparison with analogues Analogues of hydrogen peroxide include the chemically identical deuterium peroxide and hydrogen disulfide. Hydrogen disulfide has a boiling point of only 70.7°C despite having a higher molecular weight, indicating that hydrogen bonding increases the boiling point of hydrogen peroxide.

Physical properties of hydrogen peroxide solutions The properties of aqueous solutions of hydrogen peroxide differ from those of the neat material, reflecting the effects of hydrogen bonding between water and hydrogen peroxide molecules. Hydrogen peroxide and water form a eutectic mixture, exhibiting freezing-point depression. Whereas pure water melts and freezes at approximately 273K, and pure hydrogen peroxide just 0.4K below that, a 50% (by volume) solution melts and freezes at 221 K.

pH of H2O2 Pure hydrogen peroxide has a pH of 6.2, making it a weak acid. The pH can be as low as 4.5 when diluted at approximately 60%.

History Hydrogen peroxide was first isolated in 1818 by Louis Jacques Thénard by reacting barium peroxide with nitric acid. An improved version of this process used hydrochloric acid, followed by sulfuric acid to precipitate the barium sulfate byproduct. Thénard's process was used from the end of the 19th century until the middle of the 20th century. Modern production methods are discussed below. For a long time, pure hydrogen peroxide was believed to be unstable, because attempts to separate the hydrogen peroxide from the water, which is present during synthesis, failed. This instability was however due to traces of impurities (transition metals salts) that catalyze the decomposition of the hydrogen peroxide. One hundred percent pure hydrogen peroxide was first obtained through vacuum distillation by Richard Wolffenstein in 1894. At the end of 19th century, Petre Melikishvili and his pupil L. Pizarjevski showed that of the many proposed formulas of hydrogen peroxide, the correct one was H-O-O-H. The use of H2O2 sterilization in biological safety cabinets and barrier isolators is a popular alternative to ethylene oxide (EtO) as a safer, more efficient decontamination method. H2O2 has long been widely used in the pharmaceutical industry. In aerospace research, H2O2 is used to sterilize artificial satellites and space probes. The U.S. FDA has recently granted 510(k) clearance to use H2O2 in individual medical device manufacturing applications. EtO criteria outlined in ANSI/AAMI/ISO 14937 may be used as a validation guideline. Sanyo was the first manufacturer to use the H2O2 process in situ in a cell culture incubator, which is a faster and more efficient cell culture sterilization process.

Manufacture Formerly, hydrogen peroxide was prepared by the electrolysis of an aqueous solution of sulfuric acid or acidic ammonium bisulfate (NH4HSO4), followed by hydrolysis of the peroxodisulfate ((SO4)2)2− that is formed. Today, hydrogen peroxide is manufactured almost exclusively by the autoxidation of a 2-alkyl anthrahydroquinone (or 2-alkyl-9,10dihydroxyanthracene) to the corresponding 2-alkyl anthraquinone in the so called anthraquinone process. Major producers commonly use either the 2-ethyl or the 2-amyl derivative. The cyclic reaction depicted below shows the 2-ethyl derivative, where 2ethyl-9,10-dihydroxyanthracene (C16H14O2) is oxidized to the corresponding 2ethylanthraquinone (C16H12O2) and hydrogen peroxide. Most commercial processes achieve this by bubbling compressed air through a solution of the anthracene, whereby the oxygen present in the air reacts with the labile hydrogen atoms (of the hydroxy

group), giving hydrogen peroxide and regenerating the anthraquinone. Hydrogen peroxide is then extracted and the anthraquinone derivative is reduced back to the dihydroxy (anthracene) compound using hydrogen gas in the presence of a metal catalyst. The cycle then repeats itself.

This process is known as the Riedl-Pfleiderer process, having been first discovered by them in 1936. The overall equation for the process is deceptively simple: H2 + O2 → H2O2 The economics of the process depend heavily on effective recycling of the quinone (which is expensive) and extraction solvents, and of the hydrogenation catalyst. In 1994, world production of H2O2 was around 1.9 million tonnes and grew to 2.2 million in 2006, most of which was at a concentration of 70% or less. In that year bulk 30% H2O2 sold for around US $0.54 per kg, equivalent to US $1.50 per kg (US $0.68 per lb) on a "100% basis".

New developments A new, so-called "high-productivity/high-yield" process, based on an optimized distribution of isomers of 2-amyl anthraquinone, has been developed by Solvay. In July 2008, this process allowed the construction of a "mega-scale" single-train plant in Zandvliet (Belgium). The plant has an annual production capacity more than twice that of the world's next-largest single-train plant. An even-larger plant is scheduled to come onstream at Map Ta Phut (Thailand) in 2011. It is likely that this will lead to a reduction in the cost of production due to economies of scale. A process to produce hydrogen peroxide directly from the elements has been of interest for many years. The problem with the direct synthesis process is that, in terms of thermodynamics, the reaction of hydrogen with oxygen favors production of water. It had been recognized for some time that a finely dispersed catalyst is beneficial in promoting

selectivity to hydrogen peroxide, but, while selectivity was improved, it was still not sufficiently high to permit commercial development of the process. However, an apparent breakthrough was made in the early 2000s by researchers at Headwaters Technology. The breakthrough revolves around development of a minute (nanometer-size) phasecontrolled noble metal crystal particles on carbon support. This advance led, in a joint venture with Evonik Industries, to the construction of a pilot plant in Germany in late 2005. It is claimed that there are reductions in investment cost because the process is simpler and involves less equipment; however, the process is also more corrosive and unproven. This process results in low concentrations of hydrogen peroxide (about 5–10 wt% versus about 40 wt% through the anthraquione process). In 2009, another catalyst development was announced by workers at Cardiff University. This development also relates to the direct synthesis, but, in this case, using gold– palladium nanoparticles. Under normal circumstances, the direct synthesis must be carried out in an acid medium to prevent immediate decomposition of the hydrogen peroxide once it is formed. Whereas hydrogen peroxide tends to decompose on its own (which is why, even after production, it is often necessary to add stabilisers to the commercial product when it is to be transported or stored for long periods), the nature of the catalyst can cause this decomposition to accelerate rapidly. It is claimed that the use of this gold-palladium catalyst reduces this decomposition and, as a consequence, little to no acid is required. The process is in a very early stage of development and currently results in very low concentrations of hydrogen peroxide being formed (less than about 1– 2 wt%). Nonetheless, it is envisaged by the inventors that the process will lead to an inexpensive, efficient, and environmentally friendly process. A novel electrochemical process for the production of alkaline hydrogen peroxide has been developed by Dow. The process employs a monopolar cell to achieve an electrolytic reduction of oxygen in a dilute sodium hydroxide solution.

Availability Hydrogen peroxide is most commonly available as a solution in water. For consumers, it is usually available from pharmacies at 3 and 6 wt% concentrations. The concentrations are sometimes described in terms of the volume of oxygen gas generated, one milliliter of a 20-volume solution generates twenty milliliters of oxygen gas when completely decomposed. For laboratory use, 30 wt% solutions are most common. Commercial grades from 70% to 98% are also available, but due to the potential of solutions of >68% hydrogen peroxide to be converted entirely to steam and oxygen (with the temperature of the steam increasing as the concentration increases above 68%) these grades are potentially far more hazardous, and require special care in dedicated storage areas. Buyers must typically submit to inspection by the small number of commercial manufacturers.

Reactions Decomposition

Manganese dioxide decomposing a very dilute solution of hydrogen peroxide Hydrogen peroxide decomposes (disproportionates) exothermically into water and oxygen gas spontaneously: 2 H2O2 → 2 H2O + O2 This process is thermodynamically favorable. It has a ΔHo of −98.2 kJ·mol−1 and a ΔGo of −119.2 kJ·mol−1 and a ΔS of 70.5 J·mol−1·K−1. The rate of decomposition is dependent on the temperature and concentration of the peroxide, as well as the pH and the presence of impurities and stabilizers. Hydrogen peroxide is incompatible with many substances that catalyse its decomposition, including most of the transition metals and their compounds. Common catalysts include manganese dioxide, silver, and platinum. The same reaction is catalysed by the enzyme catalase, found in the liver, whose main function in the body is the removal of toxic byproducts of metabolism and the reduction of oxidative stress. The decomposition occurs more rapidly in alkali, so acid is often added as a stabilizer. The liberation of oxygen and energy in the decomposition has dangerous side-effects. Spilling high concentrations of hydrogen peroxide on a flammable substance can cause an immediate fire, which is further fueled by the oxygen released by the decomposing hydrogen peroxide. High test peroxide, or HTP (also called high-strength peroxide) must be stored in a suitable, vented container to prevent the buildup of oxygen gas, which would otherwise lead to the eventual rupture of the container.

In the presence of certain catalysts, such as Fe2+ or Ti3+, the decomposition may take a different path, with free radicals such as HO· (hydroxyl) and HOO· being formed. A combination of H2O2 and Fe2+ is known as Fenton's reagent. A common concentration for hydrogen peroxide is 20-volume, which means that, when 1 volume of hydrogen peroxide is decomposed, it produces 20 volumes of oxygen. A 20volume concentration of hydrogen peroxide is equivalent to 1.667 mol/dm3 (Molar solution) or about 6%.

Redox reactions In acidic solutions, H2O2 is one of the most powerful oxidizers known—stronger than chlorine, chlorine dioxide, and potassium permanganate. Also, through catalysis, H2O2 can be converted into hydroxyl radicals (.OH), which are highly reactive. Oxidant/Reduced product Oxidation potential, V [[fuck shit cunt|2.1 1.8 Potassium permanganate/Manganese dioxide 1.7 Chlorine dioxide/HClO 1.5 Chlorine/Chloride 1.4 In aqueous solutions, hydrogen peroxide can oxidize or reduce a variety of inorganic ions. When it acts as a reducing agent, oxygen gas is also produced. In acidic solutions Fe2+ is oxidized to Fe3+ (hydrogen peroxide acting as an oxidizing agent), 2 Fe2+(aq) + H2O2 + 2 H+(aq) → 2 Fe3+(aq) + 2H2O(l) and sulfite (SO32−) is oxidized to sulfate (SO42−). However, potassium permanganate is reduced to Mn2+ by acidic H2O2. Under alkaline conditions, however, some of these reactions reverse; for example, Mn2+ is oxidized to Mn4+ (as MnO2). Other examples of hydrogen peroxide's action as a reducing agent are reaction with sodium hypochlorite or potassium permanganate, which is a convenient method for preparing oxygen in the laboratory. NaOCl + H2O2 → O2 + NaCl + H2O 2 KMnO4 + 3 H2O2 → 2 MnO2 + 2 KOH + 2 H2O + 3 O2 Hydrogen peroxide is frequently used as an oxidizing agent in organic chemistry. One application is for the oxidation of thioethers to sulfoxides. For example, methyl phenyl sulfide was oxidized to methyl phenyl sulfoxide in 99% yield in methanol in 18 hours (or 20 minutes using a TiCl3 catalyst):

Ph-S-CH3 + H2O2 → Ph-S(O)-CH3 + H2O Alkaline hydrogen peroxide is used for epoxidation of electron-deficient alkenes such as acrylic acids, and also for oxidation of alkylboranes to alcohols, the second step of hydroboration-oxidation.

Formation of peroxide compounds Hydrogen peroxide is a weak acid, and it can form hydroperoxide or peroxide salts or derivatives of many metals. For example, on addition to an aqueous solution of chromic acid (CrO3) or acidic solutions of dichromate salts, it will form an unstable blue peroxide CrO(O2)2. In aqueous solution it rapidly decomposes to form oxygen gas and chromium salts. It can also produce peroxoanions by reaction with anions; for example, reaction with borax leads to sodium perborate, a bleach used in laundry detergents: Na2B4O7 + 4 H2O2 + 2 NaOH → 2 Na2B2O4(OH)4 + H2O H2O2 converts carboxylic acids (RCOOH) into peroxy acids (RCOOOH), which are themselves used as oxidizing agents. Hydrogen peroxide reacts with acetone to form acetone peroxide, and it interacts with ozone to form hydrogen trioxide, also known as trioxidane. Reaction with urea produces carbamide peroxide, used for whitening teeth. An acid-base adduct with triphenylphosphine oxide is a useful "carrier" for H2O2 in some reactions.

Alkalinity Hydrogen peroxide can still form adducts with very strong acids. The superacid HF/SbF5 forms unstable compounds containing the [H3O2]+ ion.

Uses Industrial applications

ISO tank container for hydrogen peroxide transportation. About 50% of the world's production of hydrogen peroxide in 1994 was used for pulpand paper-bleaching. Other bleaching applications are becoming more important as hydrogen peroxide is seen as an environmentally benign alternative to chlorine-based bleaches. Other major industrial applications for hydrogen peroxide include the manufacture of sodium percarbonate and sodium perborate, used as mild bleaches in laundry detergents. It is used in the production of certain organic peroxides such as dibenzoyl peroxide, used in polymerisations and other chemical processes. Hydrogen peroxide is also used in the production of epoxides such as propylene oxide. Reaction with carboxylic acids produces a corresponding peroxy acid. Peracetic acid and meta-chloroperoxybenzoic acid (commonly abbreviated mCPBA) are prepared from acetic acid and meta-chlorobenzoic acid, respectively. The latter is commonly reacted with alkenes to give the corresponding epoxide. In the PCB manufacturing process, hydrogen peroxide mixed with sulfuric acid was used as the microetch chemical for copper surface roughening preparation.

A combination of a powdered precious metal-based catalyst, hydrogen peroxide, methanol and water can produce superheated steam in one to two seconds, releasing only CO2 and high-temperature steam for a variety of purposes. Recently, there has been increased use of vaporized hydrogen peroxide in the validation and bio-decontamination of half-suit and glove-port isolators in pharmaceutical production. Nuclear pressurized water reactors (PWRs) use hydrogen peroxide during the plant shutdown to force the oxidation and dissolution of activated corrosion products deposited on the fuel. The corrosion products are then removed with the cleanup systems before the reactor is disassembled. Hydrogen peroxide is also used in the oil and gas exploration industry to oxidize rock matrix in preparation for micro-fossil analysis.

Chemical applications A method of producing propylene oxide from hydrogen peroxide has been developed. The process is claimed to be environmentally friendly, since the only significant byproduct is water. It is also claimed the process has significantly lower investment and operating costs. Two of these "HPPO" (hydrogen peroxide to propylene oxide) plants came onstream in 2008: One of them located in Belgium is a Solvay, Dow-BASF joint venture, and the other in Korea is a EvonikHeadwaters, SK Chemicals joint venture. A caprolactam application for hydrogen peroxide has been commercialized. Potential routes to phenol and epichlorohydrin utilizing hydrogen peroxide have been postulated.

Biological function Hydrogen peroxide is also one of the two chief chemicals in the defense system of the bombardier beetle, reacting with hydroquinone to discourage predators. A study published in Nature found that hydrogen peroxide plays a role in the immune system. Scientists found that hydrogen peroxide is released after tissues are damaged in zebra fish, which is thought to act as a signal to white blood cells to converge on the site and initiate the healing process. When the genes required to produce hydrogen peroxide were disabled, white blood cells did not accumulate at the site of damage. The experiments were conducted on fish; however, because fish are genetically similar to humans, the same process is speculated to occur in humans. The study in Nature suggested asthma sufferers have higher levels of hydrogen peroxide in their lungs than healthy people, which could explain why asthma sufferers have inappropriate levels of white blood cells in their lungs.

Domestic uses

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Diluted H2O2 (between 3% and 8%) is used to bleach human hair when mixed with ammonium hydroxide, hence the phrase "peroxide blonde". It is absorbed by skin upon contact and creates a local skin capillary embolism that appears as a temporary whitening of the skin. It is used to whiten bones that are to be put on display. 3% H2O2 is used medically for cleaning wounds, removing dead tissue, and as an oral debriding agent. Peroxide stops slow (small vessel) wound bleeding/oozing, as well. However, recent studies have suggested that hydrogen peroxide impedes scarless healing as it destroys newly formed skin cells. Most over-the-counter peroxide solutions are not suitable for ingestion. If a dog has swallowed a harmful substance (e.g., rat poison), small amounts of hydrogen peroxide can be given to induce vomiting. 35% hydrogen peroxide is used to prevent infection transmission in the hospital environment, hydrogen peroxide vapor is registered with the US EPA as a sporicidal sterilant. 3% H2O2 is effective at treating fresh (red) blood-stains in clothing and on other items. It must be applied to clothing before blood stains can be accidentally "set" with heated water. Cold water and soap are then used to remove the peroxide treated blood. The United States Food and Drug Administration (FDA) has classified hydrogen peroxide as a Low Regulatory Priority (LRP) drug for use in controlling fungus on fish and fish eggs.

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Some horticulturalists and users of hydroponics advocate the use of weak hydrogen peroxide solution in watering solutions. Its spontaneous decomposition releases oxygen that enhances a plant's root development and helps to treat root rot (cellular root death due to lack of oxygen) and a variety of other pests. There is some peer-reviewed academic research to back up some of the claims. Laboratory tests conducted by fish culturists in recent years have demonstrated that common household hydrogen peroxide can be used safely to provide oxygen for small fish. Hydrogen peroxide releases oxygen by decomposition when it is exposed to catalysts such as manganese dioxide. Hydrogen peroxide is a strong oxidizer effective in controlling sulfide and organic-related odors in wastewater collection and treatment systems. It is typically applied to a wastewater system where there is a retention time of 30 minutes to 5 hours before hydrogen sulfide is released. Hydrogen peroxide oxidizes the hydrogen sulfide and promotes bio-oxidation of organic odors. Hydrogen peroxide decomposes to oxygen and water, adding dissolved oxygen to the system, thereby negating some Biochemical Oxygen Demand (BOD). Mixed with baking soda and a small amount of hand soap, hydrogen peroxide is effective at removing skunk odor. Hydrogen peroxide is used with phenyl oxalate ester and an appropriate dye in glow sticks as an oxidizing agent. It reacts with the ester to form an unstable CO2 dimer, which excites the dye to an excited state; the dye emits a photon (light) when it spontaneously relaxes back to the ground state. Hydrogen peroxide can be combined with vinegar and table salt to form a substitute for industrial chemicals such as ferric chloride, ammonium persulfate, or hydrochloric acid as a hobbyist's printed circuit board etchant.

Use as propellant

Rocket Belt hydrogen peroxide propulsion system High concentration H2O2 is referred to as HTP or High test peroxide. It can be used either as a monopropellant (not mixed with fuel) or as the oxidizer component of a bipropellant rocket. Use as a monopropellant takes advantage of the decomposition of 70–98+% concentration hydrogen peroxide into steam and oxygen. The propellant is pumped into a reaction chamber where a catalyst, usually a silver or platinum screen, triggers decomposition, producing steam at over 600 °C, which is expelled through a nozzle, generating thrust. H2O2 monopropellant produces a maximum specific impulse (Isp) of 161 s (1.6 kN·s/kg), which makes it a low-performance monopropellant. Peroxide generates much less thrust than hydrazine, but is not toxic. The Bell Rocket Belt used hydrogen peroxide monopropellant.

As a bipropellant H2O2 is decomposed to burn a fuel as an oxidizer. Specific impulses as high as 350 s (3.5 kN·s/kg) can be achieved, depending on the fuel. Peroxide used as an oxidizer gives a somewhat lower Isp than liquid oxygen, but is dense, storable, noncryogenic and can be more easily used to drive gas turbines to give high pressures using an efficient closed cycle. It can also be used for regenerative cooling of rocket engines. Peroxide was used very successfully as an oxidizer in World-War-II German rockets (e.g. T-Stoff, containing oxyquinoline stabilizer, for the Me-163), and for the low-cost British Black Knight and Black Arrow launchers. In the 1940s and 1950s, the Walter turbine used hydrogen peroxide for use in submarines while submerged; it was found to be too noisy and require too much maintenance compared to diesel-electric power systems. Some torpedoes used hydrogen peroxide as oxidizer or propellant, but this was dangerous and has been discontinued by most navies. Hydrogen peroxide leaks were blamed for the sinkings of HMS Sidon and the Russian submarine Kursk. It was discovered, for example, by the Japanese Navy in torpedo trials, that the concentration of H2O2 in right-angle bends in HTP pipework can often lead to explosions in submarines and torpedoes. SAAB Underwater Systems is manufacturing the Torpedo 2000. This torpedo, used by the Swedish navy, is powered by a piston engine propelled by HTP as an oxidizer and kerosene as a fuel in a bipropellant system. While rarely used now as a monopropellant for large engines, small hydrogen peroxide attitude control thrusters are still in use on some satellites. They are easy to throttle, and safer to fuel and handle before launch than hydrazine thrusters. However, hydrazine is more often used in spacecraft because of its higher specific impulse and lower rate of decomposition.

Therapeutic use Hydrogen peroxide is generally recognized as safe (GRAS) as an antimicrobial agent, an oxidizing agent and for other purposes by the FDA. Hydrogen peroxide has been used as an antiseptic and anti-bacterial agent for many years due to its oxidizing effect. While its use has decreased in recent years with the popularity of readily available over the counter products, it is still used by many hospitals, doctors and dentists. 

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Like many oxidative antiseptics, hydrogen peroxide causes mild damage to tissue in open wounds, but it also is effective at rapidly stopping capillary bleeding (slow blood oozing from small vessels in abrasions), and is sometimes used sparingly for this purpose, as well as cleaning. Hydrogen peroxide can be used as a toothpaste when mixed with correct quantities of baking soda and salt. Hydrogen peroxide and benzoyl peroxide are sometimes used to treat acne. Hydrogen peroxide is used as an emetic in veterinary practice.

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The American Cancer Society states that "there is no scientific evidence that hydrogen peroxide is a safe, effective or useful cancer treatment", and advises cancer patients to "remain in the care of qualified doctors who use proven methods of treatment and approved clinical trials of promising new treatments." Another controversial alternative medical procedure is inhalation of hydrogen peroxide at a concentration of about 1%. Intravenous usage of hydrogen peroxide has been linked to several deaths.

Improvised explosive device / home-made bomb precursor Hydrogen peroxide was the main ingredient in the 7 July 2005 London bombings that killed 52 London Underground and bus passengers. The bomb-making ingredients are reported to be easier to buy than large numbers of aspirin pills.

Safety Regulations vary, but low concentrations, such as 3%, are widely available and legal to buy for medical use. Higher concentrations may be considered hazardous and are typically accompanied by a Material Safety Data Sheet (MSDS). In high concentrations, hydrogen peroxide is an aggressive oxidizer and will corrode many materials, including human skin. In the presence of a reducing agent, high concentrations of H2O2 will react violently. High-concentration hydrogen peroxide streams, typically above 40%, should be considered a D001 hazardous waste, due to concentrated hydrogen peroxide's meeting the definition of a DOT oxidizer according to U.S. regulations, if released into the environment. The EPA Reportable Quantity (RQ) for D001 hazardous wastes is 100 pounds, or approximately ten gallons, of concentrated hydrogen peroxide. Hydrogen peroxide should be stored in a cool, dry, well-ventilated area and away from any flammable or combustible substances. It should be stored in a container composed of non-reactive materials such as stainless steel or glass (other materials including some plastics and aluminium alloys may also be suitable). Because it breaks down quickly when exposed to light, it should be stored in an opaque container, and pharmaceutical formulations typically come in brown bottles that filter out light. Hydrogen peroxide, either in pure or diluted form, can pose several risks: 

Explosive vapors. Above roughly 70% concentrations, hydrogen peroxide can give off vapor that can detonate above 70 °C (158 °F) at normal atmospheric pressure. This can then cause a boiling liquid expanding vapor explosion (BLEVE) of the remaining liquid. Distillation of hydrogen peroxide at normal pressures is thus highly dangerous.

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Hazardous reactions. Hydrogen peroxide vapors can form sensitive contact explosives with hydrocarbons such as greases. Hazardous reactions ranging from ignition to explosion have been reported with alcohols, ketones, carboxylic acids (particularly acetic acid), amines and phosphorus. Spontaneous ignition. Concentrated hydrogen peroxide, if spilled on clothing (or other flammable materials), will preferentially evaporate water until the concentration reaches sufficient strength, at which point the material may spontaneously ignite. Corrosive. Concentrated hydrogen peroxide (>50%) is corrosive, and even domestic-strength solutions can cause irritation to the eyes, mucous membranes and skin. Swallowing hydrogen peroxide solutions is particularly dangerous, as decomposition in the stomach releases large quantities of gas (10 times the volume of a 3% solution) leading to internal bleeding. Inhaling over 10% can cause severe pulmonary irritation. Bleach agent. Low concentrations of hydrogen peroxide, on the order of 3% or less, will chemically bleach many types of clothing to a pinkish hue. Caution should be exercised when using common products that may contain hydrogen peroxide, such as facial cleaner or contact lens solution, which easily splatter upon other surfaces. Internal ailments. Large oral doses of hydrogen peroxide at a 3% concentration may cause "irritation and blistering to the mouth, (which is known as Black hairy tongue) throat, and abdomen", as well as "abdominal pain, vomiting, and diarrhea". Vapor pressure. Hydrogen peroxide has a significant vapor pressure (1.2 kPa at 50 oC[CRC Handbook of Chemistry and Physics, 76th Ed, 1995-1996]) and exposure to the vapor is potentially hazardous. Hydrogen peroxide vapor is a primary irritant, primarily affecting the eyes and respiratory system and the NIOSH Immediately dangerous to life and health limit (IDLH) is only 75 ppm. Documentation for Immediately Dangerous to Life or Health Concentrations (IDLH): NIOSH. Chemical Listing and Documentation of Revised IDLH Values (as of 3/1/95). Long term exposure to low ppm concentrations is also hazardous and can result in permanent lung damage and OSHA Occupational Safety and Health Administration has established a permissible exposure limit of 1.0 ppm calculated as an eight hour time weighted average (29 CFR 1910.1000, Table Z-1) and hydrogen peroxide has also been classified by the ACGIH American Conference of Industrial Hygienists (ACGIH) as a "known animal carcinogen, with unknown relevance on humans.[2008 Threshold Limit Values for Chemical Substances and Physical Agents & Biological Exposure Indices, ACGIH] In applications where high concentrations of hydrogen peroxide are used, suitable personal protective equipment should be worn and it is prudent in situations where the vapor is likely to be generated, such as hydrogen peroxide gas or vapor sterilization, to ensure that there is adequate ventilation and the vapor concentration monitored with a continuous gas monitor for hydrogen peroxide. Continuous gas monitors for hydrogen peroxide are available from several suppliers. Further information on the hazards of hydrogen peroxide is available

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from OSHA Occupational Safety and Health Guideline for Hydrogen Peroxide and from the ATSDR. Agency for Toxic Substances and Disease Registry Skin disorders. Vitiligo is an acquired skin disorder with the loss of native skin pigment, which affects about 0.5-1% of the world population. Recent studies have discovered increased H2O2 levels in the epidermis and in blood are one of many hallmarks of this disease.

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On July 16, 1934 in Kummersdorf, Germany a rocket engine using hydrogen peroxide exploded, killing three people. As a result of this incident, Werner von Braun decided not to use hydrogen peroxide as an oxidizer in the rockets he developed afterward. Several people received minor injuries after a hydrogen peroxide spill on board Northwest Airlines flight 957 from Orlando to Memphis on October 28, 1998 and subsequent fire on Northwest Airlines flight 7. During the Second World War, doctors in Nazi concentration camps experimented with the use of hydrogen peroxide injections in the killing of human subjects. Hydrogen peroxide was said to be one of the ingredients in the bombs that failed to explode in the July 21, 2005 London bombings. The Russian submarine K-141 Kursk sailed out to sea to perform an exercise of firing dummy torpedoes at the Pyotr Velikiy, a Kirov class battlecruiser. On August 12, 2000 at 11:28 local time (07:28 UTC), there was an explosion while preparing to fire the torpedoes. The only credible report to date is that this was due to the failure and explosion of one of the Kursk's hydrogen peroxide-fueled torpedoes. It is believed that HTP, a form of highly concentrated hydrogen peroxide used as propellant for the torpedo, seeped through rust in the torpedo casing. A similar incident was responsible for the loss of HMS Sidon in 1955 On August 16, 2010 a spill of about 10 gallons of cleaning fluid spilled on the 53rd floor of 1515 Broadway, in Times Square, New York City. The spill, which a spokesperson for the New York City fire department said was of Hydrogen Peroxide, shut down Broadway between West 42nd and West 48th streets as a number of fire engines responded to the hazmat situation. There were no reported injuries.

Chapter-10

Properties of Water

Water (H2O)

IUPAC name Water Oxidane Other names Hydrogen oxide Dihydrogen monoxide Hydrogen monoxide Identifiers CAS number 7732-18-5 PubChem 962 ChemSpider 937 UNII 059QF0KO0R ChEBI CHEBI:15377 ChEMBL CHEMBL1098659 RTECS number ZC0110000 Properties Molecular formula H2O Molar mass 18.01528(33) g/mol white solid or almost colorless, Appearance transparent, with a slight hint of blue, crystalline solid or liquid 1000 kg/m3, liquid (4 °C) (62.4 lb/cu. Density ft) 917 kg/m3, solid Melting point 0 °C, 32 °F (273.15 K) Boiling point 99.98 °C, 212 °F (373.13 K)

15.74 ~35–36 15.74 Basicity (pKb) Refractive index (nD) 1.3330 Viscosity 0.001 Pa s at 20 °C Structure Crystal structure Hexagonal Molecular shape Bent Dipole moment 1.85 D Hazards Main hazards Drowning Related compounds Hydrogen sulfide Hydrogen selenide Other cations Hydrogen telluride Hydrogen polonide Hydrogen peroxide acetone Related solvents methanol water vapor Related compounds ice heavy water Acidity (pKa)

Water (H2O) is the most abundant compound on Earth's surface, covering about 70%. In nature, it exists in liquid, solid, and gaseous states. It is in dynamic equilibrium between the liquid and gas states at standard temperature and pressure. At room temperature, it is a tasteless and odorless liquid, nearly colorless with a hint of blue. Many substances dissolve in water and it is commonly referred to as the universal solvent. Because of this, water in nature and in use is rarely pure and some of its properties may vary slightly from those of the pure substance. However, there are also many compounds that are essentially, if not completely, insoluble in water. Water is the only common substance found naturally in all three common states of matter and it is essential for all life on Earth. Water usually makes up 55% to 78% of the human body.

Forms of water Like many substances, water can take numerous forms that are broadly categorized by phase of matter. The liquid phase is the most common among water's phases (within the Earth's atmosphere and surface) and is the form that is generally denoted by the word "water." The solid phase of water is known as ice and commonly takes the structure of hard, amalgamated crystals, such as ice cubes, or loosely accumulated granular crystals, like snow. The gaseous phase of water is known as water vapor (or steam), and is characterized by water assuming the configuration of a transparent cloud. (Note that the visible steam and clouds are, in fact, water in the liquid form as minute droplets suspended in the air.) The fourth state of water, that of a supercritical fluid, is much less common than the other three and only rarely occurs in nature, in extremely uninhabitable conditions. When water achieves a specific critical temperature and a specific critical

pressure (647 K and 22.064 MPa), liquid and gas phase merge to one homogeneous fluid phase, with properties of both gas and liquid. One example of naturally occurring supercritical water is found in the hottest parts of deep water hydrothermal vents, in which water is heated to the critical temperature by scalding volcanic plumes and achieves the critical pressure because of the crushing weight of the ocean at the extreme depths at which the vents are located. Additionally, anywhere there is volcanic activity below a depth of 2.25 km (1.4 miles) can be expected to have water in the supercritical phase. Vienna Standard Mean Ocean Water is the current international standard for water isotopes. Naturally occurring water is almost completely composed of the neutron-less hydrogen isotope protium. Only 155 ppm include deuterium (2H or D), a hydrogen isotope with one neutron, and less than 20 parts per quintillion include tritium (3H or T), which has two. Heavy water is water with a higher-than-average deuterium content, up to 100%. Chemically, it is similar but not identical to normal water. This is because the nucleus of deuterium is twice as heavy as protium, and this causes noticeable differences in bonding energies. Because water molecules exchange hydrogen atoms with one another, hydrogen deuterium oxide (DOH) is much more common in low-purity heavy water than pure dideuterium monoxide (D2O). Humans are generally unaware of taste differences, but sometimes report a burning sensation or sweet flavor. Rats, however, are able to avoid heavy water by smell. Toxic to many animals, heavy water is used in the nuclear reactor industry to moderate (slow down) neutrons. Light water reactors are also common, where "light" simply designates normal water. Light water more specifically refers to deuterium-depleted water (DDW), water in which the deuterium content has been reduced below the standard 155ppm level. Light water has been found to be beneficial for improving cancer survival rates in mice and humans undergoing chemotherapy.

Physics and chemistry Water is the chemical substance with chemical formula H2O: one molecule of water has two hydrogen atoms covalently bonded to a single oxygen atom. Water is a tasteless, odorless liquid at ambient temperature and pressure, and appears colorless in small quantities, although it has its own intrinsic very light blue hue. Ice also appears colorless, and water vapor (steam) is essentially invisible as a gas. Water is primarily a liquid under standard conditions, which is not predicted from its relationship to other analogous hydrides of the oxygen family in the periodic table, which are gases such as hydrogen sulfide. The elements surrounding oxygen in the periodic table, nitrogen, fluorine, phosphorus, sulfur and chlorine, all combine with hydrogen to produce gases under standard conditions. The reason that water forms a liquid is that oxygen is more electronegative than all of these elements with the exception of fluorine. Oxygen attracts electrons much more strongly than hydrogen, resulting in a net positive

charge on the hydrogen atoms, and a net negative charge on the oxygen atom. The presence of a charge on each of these atoms gives each water molecule a net dipole moment. Electrical attraction between water molecules due to this dipole pulls individual molecules closer together, making it more difficult to separate the molecules and therefore raising the boiling point. This attraction is known as hydrogen bonding. The molecules of water are constantly moving in relation to each other, and the hydrogen bonds are continually breaking and reforming at timescales faster than 200 femtoseconds. However, this bond is sufficiently strong to create many of the peculiar properties of water, such as the those that make it integral to life. Water can be described as a polar liquid that slightly dissociates disproportionately into the hydronium ion (H3O+(aq)) and an associated hydroxide ion (OH−(aq)). 2 H2O (l)

H3O+ (aq) + OH− (aq)

Water, ice and vapor

Heat capacity and heats of vaporization and fusion Temperature (°C)

Heat of vaporization Hv (kJ/mol)

0

45.054

25

43.99

40

43.35

60

42.482

80

41.585

100

40.657

120

39.684

140

38.643

160

37.518

180

36.304

200

34.962

220

33.468

240

31.809

260

29.93

280

27.795

300

25.3

320

22.297

340

18.502

360

12.966

374

2.066

Water has the second highest specific heat capacity of all known substances, after ammonia, as well as a high heat of vaporization (40.65 kJ/mol or 2257 kJ/kg), both of which are a result of the extensive hydrogen bonding between its molecules. These two unusual properties allow water to moderate Earth's climate by buffering large fluctuations in temperature. According to Josh Willis, of NASA's Jet Propulsion Laboratory, the oceans absorb one thousand times more heat than the atmosphere (air) and are holding 80 to 90% of global warming heat.

The specific enthalpy of fusion of water is 333.55 kJ/kg at 0 °C. Of common substances, only that of ammonia is higher. This property confers resistance to melting on the ice of glaciers and drift ice. Before and since the advent of mechanical refrigeration, ice was and still is in common use for retarding food spoilage.

Temperature (°C)

Constant-pressure heat capacity Cp (J/(g·K) at 100 kPa)

0

4.2176

10

4.1921

20

4.1818

30

4.1784

40

4.1785

50

4.1806

60

4.1843

70

4.1895

80

4.1963

90

4.205

100

4.2159

Note that the specific heat capacity of ice at –10 °C is about 2.05 J/(g·K) and that the heat capacity of steam at 100 °C is about 2.080 J/(g·K).

Density of water and ice Density of liquid water Temp (°C) Density (kg/m3) +100 958.4 +80 971.8 +60 983.2 +40 992.2 +30 995.6502 +25 997.0479 +22 997.7735 +20 998.2071 +15 999.1026 +10 999.7026 +4 999.9720 0 999.8395 −10 998.117 −20 993.547 −30 983.854 The values below 0 °C refer to supercooled water.

The density of water is approximately one gram per cubic centimeter. More precisely, it is dependent on its temperature, but the relation is not linear and is not even monotonic. When cooled from room temperature liquid water becomes increasingly dense, just like other substances. But at approximately 4 °C, pure water reaches its maximum density. As it is cooled further, it expands to become less dense. This unusual negative thermal expansion is attributed to strong, orientation-dependent, intermolecular interactions and is also observed in molten silica. The solid form of most substances is denser than the liquid phase; thus, a block of most solids will sink in the liquid. However, a block of ice floats in liquid water because ice is less dense. Upon freezing, the density of water decreases by about 9%. The reason for this is the 'cooling' of intermolecular vibrations allowing the molecules to form steady hydrogen bonds with their neighbors and thereby gradually locking into positions reminiscent of the hexagonal packing achieved upon freezing to ice Ih. Whereas the hydrogen bonds are shorter in the crystal than in the liquid, this locking effect reduces the average coordination number of molecules as the liquid approaches nucleation. Other substances that expand on freezing are silicon, gallium, germanium, antimony, bismuth, plutonium and other compounds that form spacious crystal lattices with tetrahedral coordination. Only ordinary hexagonal ice is less dense than the liquid. Under increasing pressure, ice undergoes a number of transitions to other allotropic forms with higher density than

liquid water, such as high density amorphous ice (HDA) and very high density amorphous ice (VHDA). Water also expands significantly as the temperature increases. Its density decreases by 4% from its highest value when approaching its boiling point. The melting point of ice is 0 °C (32 °F, 273 K) at standard pressure, however, pure liquid water can be supercooled well below that temperature without freezing if the liquid is not mechanically disturbed. It can remain in a fluid state down to its homogeneous nucleation point of approximately 231 K (−42 °C). The melting point of ordinary hexagonal ice falls slightly under moderately high pressures, but as ice transforms into its allotropes above 209.9 MPa (2,072 atm), the melting point increases markedly with pressure, i.e., reaching 355 K (82 °C) at 2.216 GPa (21,870 atm) (triple point of Ice VII). A significant increase of pressure is required to lower the melting point of ordinary ice — the pressure exerted by an ice skater on the ice only reduces the melting point by approximately 0.09 °C (0.16 °F). These properties of water have important consequences in its role in the ecosystem of Earth. Water at a temperature of 4 °C will always accumulate at the bottom of fresh water lakes, irrespective of the temperature in the atmosphere. Since water and ice are poor conductors of heat (good insulators) it is unlikely that sufficiently deep lakes will freeze completely, unless stirred by strong currents that mix cooler and warmer water and accelerate the cooling. In warming weather, chunks of ice float, rather than sink to the bottom where they might melt extremely slowly. These phenomena thus may help to preserve aquatic life.

Density of saltwater and ice

WOA surface density. The density of water is dependent on the dissolved salt content as well as the temperature of the water. Ice still floats in the oceans, otherwise they would freeze from the bottom up. However, the salt content of oceans lowers the freezing point by about 2 °C and lowers the temperature of the density maximum of water to the freezing point. This is why, in ocean water, the downward convection of colder water is not blocked by an expansion of water as it becomes colder near the freezing point. The oceans' cold water near the freezing point continues to sink. For this reason, any creature attempting to survive at the bottom of such cold water as the Arctic Ocean generally lives in water that is 4 °C colder than the temperature at the bottom of frozen-over fresh water lakes and rivers in the winter. In cold countries, when the temperature of the ocean reaches 4°C, the layers of water near the top in contact with cold air continue to lose heat energy and their temperature falls below 4°C. On cooling below 4°C, these layers do not sink but may rise up as water has a maximum density at 4°C. (Refer: Polarity and hydrogen bonding ) Due to this, the layer of water at 4°C remains at the bottom and above this layers of water 3°C, 2°C, 1°C and 0°C are formed. Since ice a poor conductor of heat, it does not allow heat energy from the water beneath the layer of ice which prevents the water freezing. Thus, aquatic creatures survive in such places. As the surface of salt water begins to freeze (at −1.9 °C for normal salinity seawater, 3.5%) the ice that forms is essentially salt free with a density approximately equal to that

of freshwater ice. This ice floats on the surface and the salt that is "frozen out" adds to the salinity and density of the seawater just below it, in a process known as brine rejection. This denser saltwater sinks by convection and the replacing seawater is subject to the same process. This provides essentially freshwater ice at −1.9 °C on the surface. The increased density of the seawater beneath the forming ice causes it to sink towards the bottom. On a large scale, the process of brine rejection and sinking cold salty water results in ocean currents forming to transport such water away from the Poles. One potential consequence of global warming is that the loss of Arctic and Antarctic ice could result in the loss of these currents as well, which could have unforeseeable consequences on near and distant climates.

Miscibility and condensation

Red line shows saturation Water is miscible with many liquids, for example ethanol in all proportions, forming a single homogeneous liquid. On the other hand, water and most oils are immiscible usually forming layers according to increasing density from the top.

As a gas, water vapor is completely miscible with air. On the other hand the maximum water vapor pressure that is thermodynamically stable with the liquid (or solid) at a given temperature is relatively low compared with total atmospheric pressure. For example, if the vapor partial pressure is 2% of atmospheric pressure and the air is cooled from 25 °C, starting at about 22 °C water will start to condense, defining the dew point, and creating fog or dew. The reverse process accounts for the fog burning off in the morning. If the humidity is increased at room temperature, for example, by running a hot shower or a bath, and the temperature stays about the same, the vapor soon reaches the pressure for phase change, and then condenses out as minute water droplets, commonly referred to as steam. A gas in this context is referred to as saturated or 100% relative humidity, when the vapor pressure of water in the air is at the equilibrium with vapor pressure due to (liquid) water; water (or ice, if cool enough) will fail to lose mass through evaporation when exposed to saturated air. Because the amount of water vapor in air is small, relative humidity, the ratio of the partial pressure due to the water vapor to the saturated partial vapor pressure, is much more useful. Water vapor pressure above 100% relative humidity is called super-saturated and can occur if air is rapidly cooled, for example, by rising suddenly in an updraft.

Vapor pressure

Vapor pressure diagrams of water Temperature Pressure °C K °F Pa atm torr in Hg psi 0 273 32 611 0.00603 4.58 0.180 0.0886 5 278 41 872 0.00861 6.54 0.257 0.1265 10 283 50 1,228 0.01212 9.21 0.363 0.1781

12 14 16 17 18 19 20 21 22 23 24 25

285 287 289 290 291 292 293 294 295 296 297 298

54 1,403 0.01385 10.52 57 1,599 0.01578 11.99 61 1,817 0.01793 13.63 63 1,937 0.01912 14.53 64 2,064 0.02037 15.48 66 2,197 0.02168 16.48 68 2,338 0.02307 17.54 70 2,486 0.02453 18.65 72 2,644 0.02609 19.83 73 2,809 0.02772 21.07 75 2,984 0.02945 22.38 77 3,168 0.03127 23.76

0.414 0.2034 0.472 0.2318 0.537 0.2636 0.572 0.2810 0.609 0.2993 0.649 0.3187 0.691 0.3392 0.734 0.3606 0.781 0.3834 0.830 0.4074 0.881 0.4328 0.935 0.4594

Compressibility The compressibility of water is a function of pressure and temperature. At 0 °C, at the limit of zero pressure, the compressibility is 5.1×10−10 Pa−1. At the zero-pressure limit, the compressibility reaches a minimum of 4.4×10−10 Pa−1 around 45 °C before increasing again with increasing temperature. As the pressure is increased, the compressibility decreases, being 3.9×10−10 Pa−1 at 0 °C and 100 MPa. The bulk modulus of water is 2.2 GPa. The low compressibility of non-gases, and of water in particular, leads to their often being assumed as incompressible. The low compressibility of water means that even in the deep oceans at 4 km depth, where pressures are 40 MPa, there is only a 1.8% decrease in volume.

Triple point The various triple points of water Phases in stable equilibrium Pressure Temperature liquid water, ice Ih, and water vapor 611.73 Pa 273.16 K (0.01 °C) liquid water, ice Ih, and ice III 209.9 MPa 251 K (−22 °C) liquid water, ice III, and ice V 350.1 MPa −17.0 °C liquid water, ice V, and ice VI 632.4 MPa 0.16 °C ice Ih, Ice II, and ice III 213 MPa −35 °C ice II, ice III, and ice V 344 MPa −24 °C ice II, ice V, and ice VI 626 MPa −70 °C The temperature and pressure at which solid, liquid, and gaseous water coexist in equilibrium is called the triple point of water. This point is used to define the units of

temperature (the kelvin, the SI unit of thermodynamic temperature and, indirectly, the degree Celsius and even the degree Fahrenheit). As a consequence, water's triple point temperature is a prescribed value rather than a measured quantity.

water phase diagram: Y-axis = Pressure in pascals (10n), X-axis = temperature in kelvins, S = solid, L = liquid, V = vapor, CP = critical point, TP = triple point of water The triple point is at a temperature of 273.16 K (0.01 °C) by convention, and at a pressure of 611.73 Pa. This pressure is quite low, about 1⁄166 of the normal sea level barometric pressure of 101,325 Pa. The atmospheric surface pressure on planet Mars is remarkably close to the triple point pressure, and the zero-elevation or "sea level" of Mars is defined by the height at which the atmospheric pressure corresponds to the triple point of water. Although it is commonly named as "the triple point of water", the stable combination of liquid water, ice I, and water vapor is but one of several triple points on the phase diagram of water. Gustav Heinrich Johann Apollon Tammann in Göttingen produced data on several other triple points in the early 20th century. Kamb and others documented further triple points in the 1960s.

Electrical properties

Electrical conductivity Pure water containing no ions is an excellent insulator, but not even "deionized" water is completely free of ions. Water undergoes auto-ionization in the liquid state. Further, because water is such a good solvent, it almost always has some solute dissolved in it, most frequently a salt. If water has even a tiny amount of such an impurity, then it can conduct electricity readily, as impurities such as salt separate into free ions in aqueous solution by which an electric current can flow. It is known that the theoretical maximum electrical resistivity for water is approximately 182 kΩ·m at 25 °C. This figure agrees well with what is typically seen on reverse osmosis, ultra-filtered and deionized ultra-pure water systems used, for instance, in semiconductor manufacturing plants. A salt or acid contaminant level exceeding even 100 parts per trillion (ppt) in ultra-pure water begins to noticeably lower its resistivity level by up to several kOhm·m (or hundreds of nanosiemens per meter). The low electrical conductivity of water increases significantly upon solvation of a small amount of ionic material, such as hydrogen chloride or any salt. Any electrical conductivity observable in water is the result of ions of mineral salts and carbon dioxide dissolved in it. Carbon dioxide forms carbonate ions in water. Water selfionizes, where two water molecules form one hydroxide anion and one hydronium cation, but not enough to carry enough electric current to do any work or harm for most operations. In pure water, sensitive equipment can detect a very slight electrical conductivity of 0.055 µS/cm at 25 °C. Water can also be electrolyzed into oxygen and hydrogen gases but in the absence of dissolved ions this is a very slow process, as very little current is conducted. While electrons are the primary charge carriers in water (and metals), in ice the primary charge carriers are protons.

Electrolysis Water can be split into its constituent elements, hydrogen and oxygen, by passing an electric current through it. This process is called electrolysis. Water molecules naturally dissociate into H+ and OH− ions, which are attracted toward the cathode and anode, respectively. At the cathode, two H+ ions pick up electrons and form H2 gas. At the anode, four OH− ions combine and release O2 gas, molecular water, and four electrons. The gases produced bubble to the surface, where they can be collected. The standard potential of the water electrolysis cell is 1.23 V at 25 °C.

Dielectric constant dielectric constant of water 10 20 30 40 50 60 70 80 90 100 temperature /°C 0 87.9 83.95 80.18 76.58 73.18 69.88 66.76 63.78 60.93 58.2 55.58 ε

Polarity and hydrogen bonding

Model of hydrogen bonds between molecules of water An important feature of water is its polar nature. The water molecule forms an angle, with hydrogen atoms at the tips and oxygen at the vertex. Since oxygen has a higher electronegativity than hydrogen, the side of the molecule with the oxygen atom has a partial negative charge. An object with such a charge difference is called a dipole meaning two poles. The oxygen end is partially negative and the hydrogen end is partially positive, because of this the direction of the dipole moment points towards the oxygen. The charge differences cause water molecules to be attracted to each other (the relatively positive areas being attracted to the relatively negative areas) and to other polar molecules. This attraction contributes to hydrogen bonding, and explains many of the properties of water, such as solvent action. A water molecule can form a maximum of four hydrogen bonds because it can accept two and donate two hydrogen atoms. Other molecules like hydrogen fluoride, ammonia, methanol form hydrogen bonds but they do not show anomalous behavior of thermodynamic, kinetic or structural properties like those observed in water. The answer

to the apparent difference between water and other hydrogen bonding liquids lies in the fact that apart from water none of the hydrogen bonding molecules can form four hydrogen bonds either due to an inability to donate/accept hydrogens or due to steric effects in bulky residues. In water local tetrahedral order due to the four hydrogen bonds gives rise to an open structure and a 3-dimensional bonding network, resulting in the anomalous decrease of density when cooled below 4 °C. Although hydrogen bonding is a relatively weak attraction compared to the covalent bonds within the water molecule itself, it is responsible for a number of water's physical properties. One such property is its relatively high melting and boiling point temperatures; more energy is required to break the hydrogen bonds between molecules. The similar compound hydrogen sulfide (H2S), which has much weaker hydrogen bonding, is a gas at room temperature even though it has twice the molecular mass of water. The extra bonding between water molecules also gives liquid water a large specific heat capacity. This high heat capacity makes water a good heat storage medium (coolant) and heat shield.

Cohesion and adhesion

Dew drops adhering to a spider web Water molecules stay close to each other (cohesion), due to the collective action of hydrogen bonds between water molecules. These hydrogen bonds are constantly

breaking, with new bonds being formed with different water molecules; but at any given time in a sample of liquid water, a large portion of the molecules are held together by such bonds. Water also has high adhesion properties because of its polar nature. On extremely clean/smooth glass the water may form a thin film because the molecular forces between glass and water molecules (adhesive forces) are stronger than the cohesive forces. In biological cells and organelles, water is in contact with membrane and protein surfaces that are hydrophilic; that is, surfaces that have a strong attraction to water. Irving Langmuir observed a strong repulsive force between hydrophilic surfaces. To dehydrate hydrophilic surfaces—to remove the strongly held layers of water of hydration—requires doing substantial work against these forces, called hydration forces. These forces are very large but decrease rapidly over a nanometer or less. They are important in biology, particularly when cells are dehydrated by exposure to dry atmospheres or to extracellular freezing.

Surface tension

This paper clip is under the water level, which has risen gently and smoothly. Surface tension prevents the clip from submerging and the water from overflowing the glass edges.

Temperature dependence of the surface tension of pure water Water has a high surface tension of 72.8 mN/m at room temperature, caused by the strong cohesion between water molecules, the highest of the non-metallic liquids. This can be seen when small quantities of water are placed onto a sorption-free (non-adsorbent and non-absorbent) surface, such as polyethylene or Teflon, and the water stays together as drops. Just as significantly, air trapped in surface disturbances forms bubbles, which sometimes last long enough to transfer gas molecules to the water. Another surface tension effect is capillary waves, which are the surface ripples that form around the impacts of drops on water surfaces, and sometimes occur with strong subsurface currents flowing to the water surface. The apparent elasticity caused by surface tension drives the waves.

Capillary action Due to an interplay of the forces of adhesion and surface tension, water exhibits capillary action whereby water rises into a narrow tube against the force of gravity. Water adheres to the inside wall of the tube and surface tension tends to straighten the surface causing a surface rise and more water is pulled up through cohesion. The process continues as the water flows up the tube until there is enough water such that gravity balances the adhesive force.

Surface tension and capillary action are important in biology. For example, when water is carried through xylem up stems in plants, the strong intermolecular attractions (cohesion) hold the water column together and adhesive properties maintain the water attachment to the xylem and prevent tension rupture caused by transpiration pull.

Water as a solvent

Presence of colloidal calcium carbonate from high concentrations of dissolved lime turns the water of Havasu Falls turquoise. Water is also a good solvent due to its polarity. Substances that will mix well and dissolve in water (e.g. salts) are known as hydrophilic ("water-loving") substances, while those that do not mix well with water (e.g. fats and oils), are known as hydrophobic

("water-fearing") substances. The ability of a substance to dissolve in water is determined by whether or not the substance can match or better the strong attractive forces that water molecules generate between other water molecules. If a substance has properties that do not allow it to overcome these strong intermolecular forces, the molecules are "pushed out" from the water, and do not dissolve. Contrary to the common misconception, water and hydrophobic substances do not "repel", and the hydration of a hydrophobic surface is energetically, but not entropically, favorable. When an ionic or polar compound enters water, it is surrounded by water molecules (Hydration). The relatively small size of water molecules typically allows many water molecules to surround one molecule of solute. The partially negative dipole ends of the water are attracted to positively charged components of the solute, and vice versa for the positive dipole ends. In general, ionic and polar substances such as acids, alcohols, and salts are relatively soluble in water, and non-polar substances such as fats and oils are not. Non-polar molecules stay together in water because it is energetically more favorable for the water molecules to hydrogen bond to each other than to engage in van der Waals interactions with non-polar molecules. An example of an ionic solute is table salt; the sodium chloride, NaCl, separates into Na+ cations and Cl− anions, each being surrounded by water molecules. The ions are then easily transported away from their crystalline lattice into solution. An example of a nonionic solute is table sugar. The water dipoles make hydrogen bonds with the polar regions of the sugar molecule (OH groups) and allow it to be carried away into solution.

Water in acid-base reactions Chemically, water is amphoteric: it can act as either an acid or a base in chemical reactions. According to the Brønsted-Lowry definition, an acid is defined as a species which donates a proton (a H+ ion) in a reaction, and a base as one which receives a proton. When reacting with a stronger acid, water acts as a base; when reacting with a stronger base, it acts as an acid. For instance, water receives an H+ ion from HCl when hydrochloric acid is formed: HCl (acid) + H2O (base)

H3O+ + Cl−

In the reaction with ammonia, NH3, water donates a H+ ion, and is thus acting as an acid: NH3 (base) + H2O (acid)

NH+4 + OH−

Because the oxygen atom in water has two lone pairs, water often acts as a Lewis base, or electron pair donor, in reactions with Lewis acids, although it can also react with Lewis bases, forming hydrogen bonds between the electron pair donors and the hydrogen atoms of water. HSAB theory describes water as both a weak hard acid and a weak hard base, meaning that it reacts preferentially with other hard species:

H+ (Lewis acid) + H2O (Lewis base) → H3O+ Fe3+ (Lewis acid) + H2O (Lewis base) → Fe(H2O)3+6 Cl− (Lewis base) + H2O (Lewis acid) → Cl(H2O)−6 When a salt of a weak acid or of a weak base is dissolved in water, water can partially hydrolyze the salt, producing the corresponding base or acid, which gives aqueous solutions of soap and baking soda their basic pH: Na2CO3 + H2O

NaOH + NaHCO3

Ligand chemistry Water's Lewis base character makes it a common ligand in transition metal complexes, examples of which range from solvated ions, such as Fe(H2O)3+ 6, to perrhenic acid, which contains two water molecules coordinated to a rhenium atom, to various solid hydrates, such as CoCl2·6H2O. Water is typically a monodentate ligand, it forms only one bond with the central atom.

Organic chemistry As a hard base, water reacts readily with organic carbocations, for example in hydration reaction, in which a hydroxyl group (OH−) and an acidic proton are added to the two carbon atoms bonded together in the carbon-carbon double bond, resulting in an alcohol. When addition of water to an organic molecule cleaves the molecule in two, hydrolysis is said to occur. Notable examples of hydrolysis are saponification of fats and digestion of proteins and polysaccharides. Water can also be a leaving group in SN2 substitution and E2 elimination reactions, the latter is then known as dehydration reaction.

Acidity in nature Pure water has the concentration of hydroxide ions (OH−) equal to that of the hydronium (H3O+) or hydrogen (H+) ions, which gives pH of 7 at 298 K. In practice, pure water is very difficult to produce. Water left exposed to air for any length of time will dissolve carbon dioxide, forming a dilute solution of carbonic acid, with a limiting pH of about 5.7. As cloud droplets form in the atmosphere and as raindrops fall through the air minor amounts of CO2 are absorbed and thus most rain is slightly acidic. If high amounts of nitrogen and sulfur oxides are present in the air, they too will dissolve into the cloud and rain drops producing acid rain.

Water in redox reactions Water contains hydrogen in oxidation state +1 and oxygen in oxidation state −2. Because of that, water oxidizes chemicals with reduction potential below the potential of H+/H2, such as hydrides, alkali and alkaline earth metals (except for beryllium), etc. Some other reactive metals, such as aluminum, are oxidized by water as well, but their oxides are not

soluble, and the reaction stops because of passivation. Note, however, that rusting of iron is a reaction between iron and oxygen, dissolved in water, not between iron and water. 2 Na + 2 H2O → 2 NaOH + H2 Water can be oxidized itself, emitting oxygen gas, but very few oxidants react with water even if their reduction potential is greater than the potential of O2/O2−. Almost all such reactions require a catalyst 4 AgF2 + 2 H2O → 4 AgF + 4 HF + O2

Geochemistry Action of water on rock over long periods of time typically leads to weathering and water erosion, physical processes that convert solid rocks and minerals into soil and sediment, but under some conditions chemical reactions with water occur as well, resulting in metasomatism or mineral hydration, a type of chemical alteration of a rock which produces clay minerals in nature and also occurs when Portland cement hardens. Water ice can form clathrate compounds, known as clathrate hydrates, with a variety of small molecules that can be embedded in its spacious crystal lattice. The most notable of these is methane clathrate, 4CH4·23H2O, naturally found in large quantities on the ocean floor.

Transparency Water is relatively transparent to visible light, near ultraviolet light, and far-red light, but it absorbs most ultraviolet light, infrared light, and microwaves. Most photoreceptors and photosynthetic pigments utilize the portion of the light spectrum that is transmitted well through water. Microwave ovens take advantage of water's opacity to microwave radiation to heat the water inside of foods. The very weak onset of absorption in the red end of the visible spectrum lends water its intrinsic blue hue.

Heavy water and isotopologues Several isotopes of both hydrogen and oxygen exist, giving rise to several known isotopologues of water. Hydrogen occurs naturally in three isotopes. The most common (1H) accounting for more than 99.98% of hydrogen in water, consists of only a single proton in its nucleus. A second, stable isotope, deuterium (chemical symbol D or 2H), has an additional neutron. Deuterium oxide, D2O, is also known as heavy water because of its higher density. It is used in nuclear reactors as a neutron moderator. The third isotope, tritium, has 1 proton and 2 neutrons, and is radioactive, decaying with a half-life of 4500 days. T2O exists in nature only in minute quantities, being produced primarily via cosmic ray-induced nuclear reactions in the atmosphere. Water with one deuterium atom HDO occurs

naturally in ordinary water in low concentrations (~0.03%) and D2O in far lower amounts (0.000003%). The most notable physical differences between H2O and D2O, other than the simple difference in specific mass, involve properties that are affected by hydrogen bonding, such as freezing and boiling, and other kinetic effects. The difference in boiling points allows the isotopologues to be separated. Consumption of pure isolated D2O may affect biochemical processes - ingestion of large amounts impairs kidney and central nervous system function. Small quantities can be consumed without any ill-effects, and even very large amounts of heavy water must be consumed for any toxicity to become apparent. Oxygen also has three stable isotopes, with 16O present in 99.76%, 17O in 0.04%, and 18O in 0.2% of water molecules.

Liquid crystal state in the exclusion zone Near hydrophilic surfaces, water exists in a liquid crystal state. This liquid crystal state has the following properties:      

the water molecules are constrained in movement (as shown by nuclear magnetic resonance imagery) it is more stable (as shown by infrared radiation imagery) it has a negative charge (as shown by a test of its electric potential) it abosrbs at 270 nm (as shown by light absorption imagery) it is more viscous than liquid water (as shown by falling ball viscometry) the molecules are aligned (as shown by polarizing microscopy)

Gerald Pollack speculated that this liquid crystal zone remained relatively unexplored recently, despite extensive writing on this topic up through 1949, because of the polywater and water memory debacles.

History The first decomposition of water into hydrogen and oxygen, by electrolysis, was done in 1800 by an English chemist William Nicholson. In 1805, Joseph Louis Gay-Lussac and Alexander von Humboldt showed that water is composed of two parts hydrogen and one part oxygen. Gilbert Newton Lewis isolated the first sample of pure heavy water in 1933. The properties of water have historically been used to define various temperature scales. Notably, the Kelvin, Celsius, Rankine, and Fahrenheit scales were, or currently are, defined by the freezing and boiling points of water. The less common scales of Delisle,

Newton, Réaumur and Rømer were defined similarly. The triple point of water is a more commonly used standard point today.

Systematic naming The accepted IUPAC name of water is oxidane or simply water, or its equivalent in different languages, although there are other systematic names which can be used to describe the molecule. The simplest systematic name of water is hydrogen oxide. This is analogous to related compounds such as hydrogen peroxide, hydrogen sulfide, and deuterium oxide (heavy water). Another systematic name, oxidane, is accepted by IUPAC as a parent name for the systematic naming of oxygen-based substituent groups, although even these commonly have other recommended names. For example, the name hydroxyl is recommended over oxidanyl for the –OH group. The name oxane is explicitly mentioned by the IUPAC as being unsuitable for this purpose, since it is already the name of a cyclic ether also known as tetrahydropyran. The polarized form of the water molecule, H+OH−, is also called hydron hydroxide by IUPAC nomenclature. Dihydrogen monoxide (DHMO) is a rarely used name of water. This term has been used in various hoaxes that call for this "lethal chemical" to be banned, such as in the dihydrogen monoxide hoax. Other systematic names for water include hydroxic acid, hydroxylic acid, and hydrogen hydroxide. Both acid and alkali names exist for water because it is amphoteric (able to react both as an acid or an alkali). None of these exotic names are used widely.

Chapter-11

Phosphorous Acid and Sulfamic Acid

Phosphorous acid Phosphorous acid

IUPAC name phosphonic acid Other names Dihydroxyphosphine oxide Dihydroxy(oxo)-λ5-phosphane Dihydroxy-λ5-phosphanone Orthophosphorous acid Oxo-λ5-phosphanediol Oxo-λ5-phosphonous acid Identifiers CAS number 13598-36-2 ChemSpider 10449259 , 10459438 (17O3) KEGG C06701 RTECS number SZ6400000 Properties Molecular formula H3PO3 Molar mass 82.00 g/mol Appearance white solid

deliquescent Density 1.651 g/cm3 (21 °C) Melting point 73.6 °C, 347 K, 164 °F Boiling point 200 °C (decomp) Solubility in water 310 g/100 mL Solubility soluble in alcohol Structure Molecular shape tetrahedral Hazards R-phrases 22-35 S-phrases 26-36/37/39-45 Related compounds H3PO4 (i.e., PO(OH)3) Related compounds H3PO2 (i.e., H2PO(OH))

Phosphorous acid is the compound described by the formula H3PO3. This acid is diprotic (readily ionizes two protons), not triprotic as might be suggested by this formula. Phosphorous acid is as an intermediate in the preparation of other phosphorus compounds.

Nomenclature and tautomerism H3PO3 is more clearly described with the structural formula HPO(OH)2. This species exists in equilibrium with a minor tautomer P(OH)3. IUPAC recommendations, 2005, are that the latter is called phosphorous acid, whereas the dihydroxy form is called phosphonic acid. Only the reduced phosphorus compounds are spelled with an "ous" ending. Other important oxyacids of phosphorus are phosphoric acid (H3PO4) and hypophosphorous acid (H3PO2). The reduced phosphorus acids are subject to similar tautomerism involving shifts of H between O and P. The P(OH)3 tautomer has been observed as a ligand bonded to molybdenum.

Structure and oxidation state In the solid state, HP(O)(OH)2 is tetrahedral with one shorter P=O bond of 148 pm and two longer P-O(H) bonds of 154 pm. Because the electronegativity of H and P are similar, the covalent P-H bond does not alter oxidation state of phosphorus, which is assigned the formal oxidation state P(II).

Preparation HPO(OH)2 is the product of the hydrolysis of its acid anhydride: P4O6 + 6 H2O → 4 HPO(OH)2 (An analogous relationship connects H3PO4 and P4O10).

On an industrial scale, the acid is prepared by hydrolysis of phosphorus trichloride with water or steam: PCl3 + 3 H2O → HPO(OH)2 + 3 HCl Potassium phosphite is also a convenient precursor to phosphorous acid: K2HPO3 + 2 HCl → 2 KCl + H3PO3 In practice aqueous potassium phosphite is treated with excess hydrochloric acid. By concentrating the solution and precipitations with alcohols, the pure acid can be separated from the salt.

Reactions Phosphorous acid on heating at 200°C converts to phosphoric acid and phosphine: 4H3PO3 → 3H3PO4 + PH3 although in practice the reaction yields a number of brownish undefined phosphorus suboxides as well. Phosphorous acid is a moderately strong dibasic acid. It reacts with alkalis forming acid phosphites and normal phosphites. Thus, reaction with sodium hydroxide gives sodium dihydrogen phosphite and disodium hydrogen phosphite, but not trisodium phosphite, Na3PO3 as the third (P-bound) hydrogen is not acidic. H3PO3 + NaOH → NaH2PO3 + H2O H3PO3 + 2NaOH → Na2HPO3 + 2H2O Phosphorous acid is a powerful reducing agent. When treated with a cold solution of mercuric chloride, a white precipitate of mercurous chloride forms: H3PO3 + 2HgCl2 + H2O → Hg2Cl2 + H3PO4 + 2HCl Mercurous chloride is reduced further by phosphorous acid to mercury on heating or on standing: H3PO3 + Hg2Cl2 + H2O → 2Hg + H3PO4 + 2HCl

Acid-base properties Phosphorous acid is a diprotic acid, since the hydrogen bonded directly to the central phosphorus atom is not readily ionizable. Chemistry examinations often test students' appreciation of the fact that all three hydrogen atoms are not acidic under aqueous conditions, in contrast with phosphoric acid. HP(O)2(OH)− is a moderately strong acid. HP(O)(OH)2 → HP(O)2(OH)− + H+ pKa = 1.3

HP(O)2(OH)− → HPO32− + H+ pKa = 6.7 The HP(O)2(OH)− species is called the hydrogenphosphite, and the HPO32− the phosphite ion.(Note that the IUPAC recommendations are dihydrogenphosphite and hydrogenphosphite respectively) The IUPAC (mostly organic) name is phosphonic acid. This nomenclature is commonly reserved for substituted derivatives, that is, organic group bonded to phosphorus, not simply an ester. For example, (CH3)PO(OH)2 is "methylphosphonic acid", which may of course form "methylphosphonate" esters. Both phosphorous acid and its deprotonated forms are good reducing agents, although not necessarily quick to react. They are oxidized to phosphoric acid or its salts. It reduces solutions of noble metal cations to the metals.

Uses In industry and agriculture The most important use of phosphorous acid is the production of phosphonates which are used in water treatment. Phosphorous acid is also used for preparing phosphite salts, such as potassium phosphite. These salts, as well as aqueous solutions of pure phosphorous acid, have shown effectiveness in controlling a variety of microbial plant diseases, in particular, treatment using either trunk injection or foliar containing phosphorous acid salts is indicated in response to infections by phytophthora and pythium-type plant pathogens (both within class oomycetes, known as water molds), such as dieback/root rot and downy mildew. Anti-microbial products containing salts of phosphorous acid are marketed in Australia as 'Yates Anti-Rot'; and in the United States of America, for example, aluminum salts of the diethyl ester of phosphorous acid (known generically as 'Fosetyl-Al') are sold under the trade name 'Aliette'. Phosphorous acid and its salts, unlike phosphoric acid, are somewhat toxic and should be handled carefully.

As a chemical reagent Phosphorous acid is used in chemical reactions as a reducing agent that is somewhat less vigorous than the related hypophosphorous acid.

Sulfamic acid Sulfamic acid

IUPAC name Sulfamic acid Identifiers CAS number 5329-14-6 PubChem 5987 ChemSpider 5767 EC number 226-218-8 UN number 2967 ChEMBL CHEMBL68253 RTECS number WO5950000 Properties Molecular formula H3NSO3 Molar mass 97.10 g/mol Density 2.15 g/cm3 Melting point 205 °C decomp. Solubility in water moderate, with slow hydrolysis Hazards MSDS ICSC 0328 EU Index 016-026-00-0 EU classification Irritant (Xi) R-phrases R36/38, R52/53 S-phrases (S2), S26, S28, S61 Related compounds Other cations Ammonium sulfamate

Sulfamic acid, also known as amidosulfonic acid, amidosulfuric acid, aminosulfonic acid, and sulfamidic acid, is a molecular compound with the formula H3NSO3. This colorless, water-soluble compound finds many applications. Sulfamic acid (H3NSO3) may be considered an intermediate compound between sulfuric acid (H2SO4), and sulfamide (H4N2SO2), effectively - though see below - replacing an OH group with an -NH2 group at each step. This pattern can extend no further in either direction without breaking down the -SO2 group.

Structure and reactivity First, it should be noticed that the compound is well described by the formula H3NSO3, not the tautomer H2NSO2(OH). The relevant bond distances are S=O, 1.44 and S-N 1.77 Å. The greater length of the S-N distance is consistent with a single bond. Furthermore, a neutron diffraction study located the hydrogen atoms, all three of which are 1.03 Å distant from nitrogen. In the solid state, the molecule of sulfamic acid is well described by a zwitterionic form :

Ball-and-stick model of a sulfamic acid zwitterion in the crystal Sulfamic acid is a moderately strong acid, Ka = 1.01 x 10−1. Because the solid is nonhygroscopic, it is used as a standard in acidimetry (quantitative assays of acid content). H3NSO3 + NaOH → NaH2NSO3 + H2O Double deprotonation can be effected in NH3 solution to give [HNSO3]2−. Sulfamic acid melts at 205 °C before decomposing at higher temperatures to H2O, SO3, SO2, and N2. Water solutions are unstable and slowly hydrolyze to ammonium bisulfate, but the crystalline solid is indefinitely stable under ordinary storage conditions.

With HNO2, sulfamic acid reacts to give N2, while with HNO3, it affords N2O. HNO2 + H3NSO3 → H2SO4 + N2 + H2O HNO3 + H3NSO3 → H2SO4 + N2O + H2O The behavior of H3NSO3 resembles that of urea, (H2N)2CO, in some ways. Both feature amino groups linked to electron-withdrawing centers that can participate in delocalized bonding. Both liberate ammonia upon heating in water.

Applications The most famous application of sulfamic acid is in the synthesis of sweet-tasting compounds. Reaction with cyclohexylamine followed by addition of NaOH gives C6H11NHSO3Na, sodium cyclamate. Sulfamates (O-substituted-, N-substituted-, or di-/tri-substituted derivatives of sulfamic acid) have been used in the design of many types of therapeutic agents such as antibiotics, nucleoside/nucleotide human immunodeficiency virus (HIV) reverse transcriptase inhibitors, HIV protease inhibitors (PIs), anti-cancer drugs (steroid sulfatase and carbonic anhydrase inhibitors), anti-epileptic drugs, and weight loss drugs. Sulfamic acid is used as an acidic cleaning agent, sometimes pure or as a component of proprietary mixtures, typically for metals and ceramics. It is frequently used for removing rust and limescale, replacing the more volatile and irritating hydrochloric acid. It is often a component of household descaling agents, for example, Lime-A-Way Thick Gel contains up to 8% sulfamic acid and pH 2 - 2.2 , or detergents used for removal of limescale. When compared to most of the common strong mineral acids, Sulfamic acid has desirable water descaling properties, low volatility, low toxicity and is a water soluble solid forming soluble calcium and iron-III salts. Its also finds applications in the industrial cleaning of dairy and brew-house equipment. Although it is considered less corrosive than hydrochloric acid due to its lower pKa, corrosion inhibitors are often added to commercial cleansers of which it is a component. It is possible that the amino group could act as a ligand under certain circumstances, as does the chloride ion for FeIII, when hydrochloric acid is used in rust removal. Sulfamic acid is used in the S.C. Johnson & Sons, Inc. "Scrubbing Bubbles Fizz-Its Toilet Tablets", and in the Saeco Dezcal™ descaling powder for home coffee and espresso equipment.     

Catalyst for esterification process Dye and pigment manufacturing Herbicide Ingredient in Denture Tablets Coagulator for urea-formaldehyde resins

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Ingredient in fire extinguishing media. Sulfamic acid is the main raw material for Ammonium sulfamate which is a widely used herbicide and fire retardant material for household product. Pulp and paper industry as a chloride stabilizer Synthesis of nitrous oxide by reaction with nitric acid In household cleaning chemical products such as Cameo. The deprotonated form (sulfamate) is a common counterion for Nickel(II) in electroplating. Easy-Off

Silver polishing According to the label on the consumer product, the liquid silver cleaning product TarnX contains thiourea, a detergent, and sulfamic acid.

Chapter-12

Phosphoric Acid

Phosphoric acid

IUPAC name trihydroxidooxidophosphorus phosphoric acid Other names Orthophosphoric acid Identifiers 7664-38-2 , CAS number 16271-20-8 (hemihydrate) PubChem 1004 ChemSpider 979 UNII E4GA8884NN EC number 231-633-2 UN number 1805 KEGG D05467 ChEMBL CHEMBL1187 RTECS number TB6300000 Properties Molecular formula H3PO4 Molar mass 98.00 g/mol Appearance white solid or colourless, viscous

liquid (>42 °C) 1.885 g/mL (liquid) Density 1.685 g/mL (85 % solution) 2.030 g/mL (crystal at 25 °C) 42.35 °C (anhydrous) Melting point 29.32 °C (hemihydrate) Boiling point 158 °C (decomp) Solubility in water 5.48 g/mL 2.148, 7.198, 12.375 Acidity (pKa) 2.4–9.4 cP (85% aq. soln.) Viscosity 147 cP (100 %) Hazards MSDS ICSC 1008 EU Index 015-011-00-6 EU classification Corrosive (C) R-phrases R34 S-phrases (S1/2) S26 S45 Flash point Non-flammable Related compounds Hypophosphorous acid Phosphorous acid Related phosphorus Pyrophosphoric acid oxoacids Triphosphoric acid Perphosphoric acid Permonophosphoric acid

Phosphoric acid, also known as orthophosphoric acid or phosphoric(V) acid, is a mineral (inorganic) acid having the chemical formula H3PO4. Orthophosphoric acid molecules can combine with themselves to form a variety of compounds which are also referred to as phosphoric acids, but in a more general way. The term phosphoric acid can also refer to a chemical or reagent consisting of phosphoric acids, usually orthophosphoric acid.

Orthophosphoric acid chemistry Pure anhydrous phosphoric acid is a white solid that melts at 42.35 °C to form a colorless, viscous liquid. Most people and even chemists refer to orthophosphoric acid as phosphoric acid, which is the IUPAC name for this compound. The prefix ortho is used to distinguish the acid from other phosphoric acids, called polyphosphoric acids two(ii). Orthophosphoric acid is a non-toxic, inorganic, rather weak triprotic acid, which, when pure, is a solid at room temperature and pressure. The chemical structure of orthophosphoric acid is shown above in the data table. Orthophosphoric acid is a very polar molecule; therefore it is highly soluble in water. The oxidation state of phosphorus (P) in ortho- and other phosphoric acids is +5; the oxidation state of all the oxygen atoms (O) is −2 and all the hydrogen atoms (H) is +1. Triprotic means that an orthophosphoric acid molecule can dissociate up

to three times, giving up an H+ each time, which typically combines with a water molecule, H2O, as shown in these reactions: H3PO4(s) + H2O(l) H2PO4−(aq)+ H2O(l) HPO42−(aq)+ H2O(l)

H3O+(aq) + H2PO4−(aq) H3O+(aq) + HPO42−(aq) H3O+(aq) + PO43−(aq)

Ka1= 7.25×10−3 Ka2= 6.31×10−8 Ka3= 3.98×10−13

The anion after the first dissociation, H2PO4−, is the dihydrogen phosphate anion. The anion after the second dissociation, HPO42−, is the hydrogen phosphate anion. The anion after the third dissociation, PO43−, is the phosphate or orthophosphate anion. For each of the dissociation reactions shown above, there is a separate acid dissociation constant, called Ka1, Ka2, and Ka3 given at 25 °C. Associated with these three dissociation constants are corresponding pKa1=2.12 , pKa2=7.21 , and pKa3=12.67 values at 25 °C. Even though all three hydrogen (H ) atoms are equivalent on an orthophosphoric acid molecule, the successive Ka values differ since it is energetically less favorable to lose another H+ if one (or more) has already been lost and the molecule/ion is more negatively-charged. Because the triprotic dissociation of orthophosphoric acid, the fact that its conjugate bases (the phosphates mentioned above) cover a wide pH range, and, because phosphoric acid/phosphate solutions are, in general, non-toxic, mixtures of these types of phosphates are often used as buffering agents or to make buffer solutions, where the desired pH depends on the proportions of the phosphates in the mixtures. Similarly, the non-toxic, anion salts of triprotic organic citric acid are also often used to make buffers. Phosphates are found pervasively in biology, especially in the compounds derived from phosphorylated sugars, such as DNA, RNA, and adenosine triphosphate (ATP). There is a separate article on phosphate as an anion or its salts. Upon heating orthophosphoric acid, condensation of the phosphoric units can be induced by driving off the water formed from condensation. When one molecule of water has been removed for each two molecules of phosphoric acid, the result is pyrophosphoric acid (H4P2O7). When an average of one molecule of water per phosphoric unit has been driven off, the resulting substance is a glassy solid having an empirical formula of HPO3 and is called metaphosphoric acid. Metaphosphoric acid is a singly anhydrous version of orthophosphoic acid and is sometimes used as a water- or moisture-absorbing reagent. Further dehydrating is very difficult, and can be accomplished only by means of an extremely strong desiccant (and not by heating alone). It produces phosphoric anhydride, which has an empirical formula P2O5, although an actual molecule has a chemical formula of P4O10. Phosphoric anhydride is a solid, which is very strongly moistureabsorbing and is used as a desiccant.

pH and composition of a phosphoric acid aqueous solution For a given total acid concentration [A] = [H3PO4] + [H2PO4−] + [HPO42−] + [PO43−] ([A] is the total number of moles of pure H3PO4 which have been used to prepare 1 liter of solution), the composition of an aqueous solution of phosphoric acid can be calculated using the equilibrium equations associated with the three reactions described above

together with the [H+][OH−] = 10−14 relation and the electrical neutrality equation. Possible concentrations of polyphosphoric molecules and ions is neglected. The system may be reduced to a fifth degree equation for [H+] which can be solved numerically, yielding: [A] (mol/L)

pH

[H3PO4]/[A] (%)

[H2PO4−]/[A] (%)

[HPO42−]/[A] (%)

[PO43−]/[A] (%)

1

1.08

91.7

8.29

6.20×10−6

1.60×10−17

10−1

1.62

76.1

23.9

6.20×10−5

5.55×10−16

10−2

2.25

43.1

56.9

6.20×10−4

2.33×10−14

10−3

3.05

10.6

89.3

6.20×10−3

1.48×10−12

10−4

4.01

1.30

98.6

6.19×10−2

1.34×10−10

10−5

5.00

0.133

99.3

0.612

1.30×10−8

10−6

5.97

1.34×10−2

94.5

5.50

1.11×10−6

10−7

6.74

1.80×10−3

74.5

25.5

3.02×10−5

10−10

7.00

8.24×10−4

61.7

38.3

8.18×10−5

For large acid concentrations, the solution is mainly composed of H3PO4. For [A] = 10−2, the pH is closed to pKa1, giving an equimolar mixture of H3PO4 and H2PO4−. For [A] below 10−3, the solution is mainly composed of H2PO4− with [HPO42−] becoming non negligible for very dilute solutions. [PO43−] is always negligible.

Chemical reagent Pure 75–85% aqueous solutions (the most common) are clear, colourless, odourless, nonvolatile, rather viscous, syrupy liquids, but still pourable. Phosphoric acid is very commonly used as an aqueous solution of 85% phosphoric acid or H3PO4. Because it is a concentrated acid, an 85% solution can be corrosive, although nontoxic when diluted. Because of the high percentage of phosphoric acid in this reagent, at least some of the orthophosphoric acid is condensed into polyphosphoric acids in a temperature-dependent equilibrium, but, for the sake of labeling and simplicity, the 85% represents H3PO4 as if it were all orthophosphoric acid. Other percentages are possible too, even above 100%, where the phosphoric acids and water would be in an unspecified equilibrium, but the overall elemental mole content would be considered specified. When aqueous solutions of phosphoric acid and/or phosphate are dilute, they are in or will reach an equilibrium after a while where practically all the phosphoric/phosphate units are in the ortho- form.

Preparation of hydrogen halides Phosphoric acid reacts with halides to form the corresponding hydrogen halide gas (steamy fumes are observed on warming the reaction mixture). This is a common practice for the laboratory preparation of hydrogen halides. NaCl(s) + H3PO4(l) → NaH2PO4(s) + HCl(g) NaBr(s) + H3PO4(l) → NaH2PO4(s) + HBr(g) NaI(s) + H3PO4(l) → NaH2PO4(s) + HI(g)

Rust removal Phosphoric acid may be used as a "rust converter", by direct application to rusted iron, steel tools, or surfaces. The phosphoric acid converts reddish-brown iron(III) oxide (rust) to black ferric phosphate, FePO4. "Rust converter" is sometimes a greenish liquid suitable for dipping (in the same sort of acid bath as is used for pickling metal), but it is more often formulated as a gel, commonly called naval jelly. It is sometimes sold under other names, such as "rust remover" or "rust killer". As a thick gel, it may be applied to sloping, vertical, or even overhead surfaces. After treatment, the black ferric-phosphate coating can be scrubbed off, leaving a fresh metal surface. Multiple applications of phosphoric acid may be required to remove all rust. The black phosphate coating can also be left in place, where it will provide moderate further corrosion resistance. (Such protection is also provided by the superficially similar Parkerizing and blued electrochemical conversion coating processes.)

Processed food use Food-grade phosphoric acid (additive E338) is used to acidify foods and beverages such as various colas, but not without controversy regarding its health effects. It provides a tangy or sour taste and, being a mass-produced chemical, is available cheaply and in large quantities. The low cost and bulk availability is unlike more expensive seasonings that give comparable flavors, such as citric acid which is obtainable from lemons and limes. However, most citric acid in the food industry is not extracted from citrus fruit, but fermented by Aspergillus niger mold from scrap molasses, waste starch hydrolysates and phosphoric acid.

Biological effects on bone calcium and kidney health Phosphoric acid, used in many soft drinks (primarily cola), has been linked to lower bone density in epidemiological studies. For example, a study using dual-energy X-ray absorptiometry rather than a questionnaire about breakage, provides reasonable evidence to support the theory that drinking cola results in lower bone density. This study was

published in the American Journal of Clinical Nutrition. A total of 1672 women and 1148 men were studied between 1996 and 2001. Dietary information was collected using a food frequency questionnaire that had specific questions about the number of servings of cola and other carbonated beverages and that also made a differentiation between regular, caffeine-free, and diet drinks. The paper cites significant statistical evidence to show that women who consume cola daily have lower bone density. Total phosphorus intake was not significantly higher in daily cola consumers than in nonconsumers; however, the calcium-to-phosphorus ratios were lower. The study also suggests that further research is needed to confirm the findings. On the other hand, a study funded by Pepsi suggests that insufficient intake of phosphorus leads to lower bone density. The study does not examine the effect of phosphoric acid, which binds with magnesium and calcium in the digestive tract to form salts that are not absorbed, but rather studies general phosphorus intake. However, a well-controlled clinical study by Heaney and Rafferty using calcium-balance methods found no impact of carbonated soft drinks containing phosphoric acid on calcium excretion. The study compared the impact of water, milk, and various soft drinks (two with caffeine and two without; two with phosphoric acid and two with citric acid) on the calcium balance of 20- to 40-year-old women who customarily consumed ~3 or more cups (680 mL) of a carbonated soft drink per day. They found that, relative to water, only milk and the two caffeine-containing soft drinks increased urinary calcium, and that the calcium loss associated with the caffeinated soft drink consumption was about equal to that previously found for caffeine alone. Phosphoric acid without caffeine had no impact on urine calcium, nor did it augment the urinary calcium loss related to caffeine. Because studies have shown that the effect of caffeine is compensated for by reduced calcium losses later in the day, Heaney and Rafferty concluded that the net effect of carbonated beverages—including those with caffeine and phosphoric acid—is negligible, and that the skeletal effects of carbonated soft drink consumption are likely due primarily to milk displacement. Other chemicals such as caffeine (also a significant component of popular common cola drinks) were also suspected as possible contributors to low bone density, due to the known effect of caffeine on calciuria. One other study, involving 30 women over the course of a week, suggests that phosphoric acid in colas has no such effect, and postulates that caffeine has only a temporary effect, which is later reversed. The authors of this study conclude that the skeletal effects of carbonated beverage consumption are likely due primarily to milk displacement. (Another possible confounding factor may be an association between high soft drink consumption and sedentary lifestyle.) Cola consumption has also been associated with chronic kidney disease and kidney stones through medical research. The preliminary results suggest that cola consumption may increase the risk of chronic kidney disease.

Medical use Phosphoric acid is used in dentistry and orthodontics as an etching solution, to clean and roughen the surfaces of teeth where dental appliances or fillings will be placed. Phosphoric acid is also an ingredient in over-the-counter anti-nausea medications that also contain high levels of sugar (glucose and fructose). This acid is also used in many teeth whiteners to eliminate plaque that may be on the teeth before application.

Preparation Phosphoric acid can be prepared by three routes – the thermal process, the wet process and the dry kiln process.

Thermal phosphoric acid This very pure phosphoric acid is obtained by burning elemental phosphorus to produce phosphorus pentoxide and dissolving the product in dilute phosphoric acid. This produces a very pure phosphoric acid, since most impurities present in the rock have been removed when extracting phosphorus from the rock in a furnace. The end result is food-grade, thermal phosphoric acid; however, for critical applications, additional processing to remove arsenic compounds may be needed.

Wet phosphoric acid Wet process phosphoric acid is prepared by adding sulfuric acid to tricalcium phosphate rock, typically found in nature as apatite. The reaction is: Ca5(PO4)3X + 5 H2SO4 + 10 H2O → 3 H3PO4 + 5 CaSO4·2H2O + HX where X may include OH, F, Cl, and Br The initial phosphoric acid solution may contain 23–33% P2O5, but can be concentrated by the evaporation of water to produce commercial- or merchant-grade phosphoric acid, which contains about 54% P2O5. Further evaporation of water yields superphosphoric acid with a P2O5 concentration above 70%. Digestion of the phosphate ore using sulfuric acid yields the insoluble calcium sulfate (gypsum), which is filtered and removed as phosphogypsum. Wet-process acid can be further purified by removing fluorine to produce animal-grade phosphoric acid, or by solvent extraction and arsenic removal to produce food-grade phosphoric acid.

Kiln Phosphoric Acid Kiln phosphoric acid (KPA) process technology is the most recent technology. Called the “Improved Hard Process”, this technology will both make low grade phosphate rock

reserves commercially viable and will increase the P2O5 recovery from existing phosphate reserves. This may significantly extend the commercial viability of phosphate reserves.

Other applications 

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Phosphoric acid is used as the electrolyte in phosphoric acid fuel cells. It is also used as an external standard for phosphorus-31 Nuclear magnetic resonance (NMR). Phosphoric acid is used as a cleaner by construction trades to remove mineral deposits, cementitious smears, and hard water stains. It is also used as a chelant in some household cleaners aimed at similar cleaning tasks. Hot phosphoric acid is used in microfabrication to etch silicon nitride (Si3N4). It is highly selective in etching Si3N4 instead of SiO2, silicon dioxide. Phosphoric acid is used as a flux by hobbyists (such as model railroaders) as an aid to soldering. Phosphoric acid is also used in hydroponics pH solutions to lower the pH of nutrient solutions. While other types of acids can be used, phosphorus is a nutrient used by plants, especially during flowering, making phosphoric acid particularly desirable. General Hydroponics pH Down liquid solution contains phosphoric acid in addition to citric acid and ammonium bisulfate with buffers to maintain a stable pH in the nutrient reservoir. Phosphoric acid is used as an electrolyte in copper electropolishing for burr removal and circuit board planarization. Phosphoric acid is used with distilled water (2–3 drops per gallon) as an electrolyte in oxyhydrogen (HHO) generators. Phosphoric acid is used as a pH adjuster in cosmetics and skin-care products. Phosphoric acid is used as a chemical oxidizing agent for activated carbon production, as used in the Wentworth Process. Phosphoric acid is also used for high-performance liquid chromatography. Phosphoric acid can be used as a dispersing agent in detergents and leather treatment.. Phosphoric acid can be used as an additive to stabilize acidic aqueous solutions within a wanted and specified pH range Phosphoric acid is the key ingredient that gives the bite taste in Coca-Cola and Pepsi sodas. In compound semiconductor processing, phosphoric acid is a common wet etching agent: for example, in combination with hydrogen peroxide and water it is used to etch InGaAs selective to InP.

Chapter-13

Hydrogen Sulfide

Hydrogen sulfide

Systematic name Hydrogen sulfide Other names Dihydrogen monosulfide Dihydrogen sulfide Sewer gas Stink damp Sulfane Sulfurated hydrogen Sulfureted hydrogen Sulfuretted hydrogen

CAS number PubChem ChemSpider UNII EC number UN number KEGG MeSH ChEBI ChEMBL RTECS number

Sulfur hydride Identifiers 7783-06-4 402 391 YY9FVM7NSN 231-977-3 1053 C00283 Hydrogen+sulfide CHEBI:16136 CHEMBL1200739 MX1225000

Beilstein Reference Gmelin Reference 3DMet

3535004 303 B01206 Properties Molecular formula H2S Molar mass 34.08 g mol−1 Exact mass 33.987720754 g mol-1 Appearance Colorless gas Density 0.001363 g cm-3 Melting point -82 °C, 191 K, -116 °F Boiling point -60 °C, 213 K, -76 °F Solubility in water 4 g dm-3 (at 20 °C) Vapor pressure 1740 kPa (at 21 °C) Acidity (pKa) 7.05 6.95 Basicity (pKb) Refractive index (nD) 1.000644 (0 °C) Structure Molecular shape Bent Dipole moment 0.97 D Thermochemistry Std enthalpy of -20.599 kJ mol-1 formation ΔfHo298 Specific heat capacity, C 1.003 J K-1 g-1 Hazards EU Index 016-001-00-4 EU classification F+

R-phrases S-phrases Autoignition temperature Explosive limits

T+

N

R12, R26, R50 (S1/2), S9, S16, S36, S38, S45, S61 300 °C

4.3–46% Related compounds Water Hydrogen selenide Related hydrogen Hydrogen telluride chalcogenides Hydrogen polonide Hydrogen disulfide Related compounds Phosphine

Hydrogen sulfide (British English: hydrogen sulphide) is the chemical compound with the formula H2S. It is a colorless, very poisonous, flammable gas with the characteristic foul odor of rotten eggs at concentrations up to 100 parts per million. It often results from the bacterial breakdown of organic matter in the absence of oxygen, such as in swamps and sewers (anaerobic digestion). It also occurs in volcanic gases, natural gas, and some

well waters. The human body produces small amounts of H2S and uses it as a signaling molecule.

Properties Hydrogen sulfide is slightly heavier than air; a mixture of H2S and air is explosive. Hydrogen sulfide and oxygen burn with a blue flame to form sulfur dioxide (SO2) and water. In general, hydrogen sulfide acts as a reducing agent. At high temperature or in the presence of catalysts, sulfur dioxide can be made to react with hydrogen sulfide to form elemental sulfur and water. This is exploited in the Claus process, the main way to convert hydrogen sulfide into elemental sulfur. Hydrogen sulfide is slightly soluble in water and acts as a weak acid, giving the hydrosulfide ion HS− (pKa=6.9 in 0.01-0.1 mol/litre solutions at 18°C) and the sulfide ion S2− (pKa=11.96). A solution of hydrogen sulfide in water is initially clear but over time turns cloudy. This is due to the slow reaction of hydrogen sulfide with the oxygen dissolved in water, yielding elemental sulfur, which precipitates out. Hydrogen sulfide reacts with metal ions to form metal sulfides, which may be considered the salts of hydrogen sulfide. Some ores are sulfides. Metal sulfides often have a dark color. Lead(II) acetate paper is used to detect hydrogen sulfide because it turns grey in the presence of the gas as lead(II) sulfide is produced. Reacting metal sulfides with strong acid liberates hydrogen sulfide. If gaseous hydrogen sulfide is put into contact with concentrated nitric acid, it explodes. Hydrogen sulfide reacts with alcohols to form thiols.

Production Hydrogen sulfide is most commonly obtained by its separation from sour gas, which is natural gas with high content of H2S. It can also be produced by reacting hydrogen gas with molten elemental sulfur at about 450 °C. Hydrocarbons can replace hydrogen in this process. Sulfate-reducing bacteria (resp. sulfur-reducing bacteria) generate usable energy under low-oxygen conditions by using sulfates (resp. elemental sulfur) to oxidize organic compounds or hydrogen; this produces hydrogen sulfide as a waste product. The standard lab preparation is to gently heat ferrous sulfide (FeS) with a strong acid in a Kipp generator. FeS + 2 HCl → FeCl2 + H2S

A less well-known and more convenient alternative is to react aluminium sulfide with water: 6 H2O + Al2S3 → 3 H2S + 2 Al(OH)3 This gas is also produced by heating sulfur with solid organic compounds and by reducing sulfurated organic compounds with hydrogen. Hydrogen sulfide is also a byproduct of some reactions and caution should be taken when production is likely as exposure can be fatal.

Occurrence

Deposit of sulfur on a rock, caused by volcanic gases Small amounts of hydrogen sulfide occur in crude petroleum, but natural gas can contain up to 90%. Volcanoes and some hot springs (as well as cold springs) emit some H2S, where it probably arises via the hydrolysis of sulfide minerals, i.e. MS + H2O → MO + H2S. About 10% of total global emissions of H2S is due to human activity. By far the largest industrial route to H2S occurs in petroleum refineries: The hydrodesulfurization process liberates sulfur from petroleum by the action of hydrogen. The resulting H2S is converted

to elemental sulfur by partial combustion via the Claus process, which is a major source of elemental sulfur. Other anthropogenic sources of hydrogen sulfide include coke ovens, paper mills (using the sulfate method), and tanneries. H2S arises from virtually anywhere where elemental sulfur comes in contact with organic material, especially at high temperatures. Hydrogen sulfide can be present naturally in well water. In such cases, ozone is often used for its removal; an alternative method uses a filter with manganese dioxide. Both methods oxidize sulfides to much less toxic sulfates.

Uses Production of thioorganic compounds Several organosulfur compounds are produced using hydrogen sulfide. These include methanethiol, ethanethiol, and thioglycolic acid.

Alkali metal sulfides Upon combining with alkali metal bases, hydrogen sulfide converts to alkali hydrosulfides such as sodium hydrosulfide and sodium sulfide, which are used in the degradation of biopolymers. The depilation of hides and the delignification of pulp by the Kraft process both are effected by alkali sulfides.

Analytical chemistry Hydrogen sulfide used to have importance in analytical chemistry for well over a century, in the qualitative inorganic analysis of metal ions. In these analyses, heavy metal (and nonmetal) ions (e.g., Pb(II), Cu(II), Hg(II), As(III)) are precipitated from solution upon exposure to H2S. The components of the resulting precipitate redissolve with some selectivity. For small-scale laboratory use in analytic chemistry, the use of thioacetamide has superseded H2S as a source of sulfide ions.

Precursor to metal sulfides As indicated above, many metal ions react with hydrogen sulfide to give the corresponding metal sulfides. This conversion is widely exploited. For example, gases or waters contaminated by hydrogen sulfide can be cleaned with metal sulfides. In the purification of metal ores by flotation, mineral powders are often treated with hydrogen sulfide to enhance the separation. Metal parts are sometimes passivated with hydrogen sulfide. Catalysts used in hydrodesulfurization are routinely activated with hydrogen sulfide, and the behavior of metallic catalysts used in other parts of a refinery is also modified using hydrogen sulfide.

Miscellaneous applications Hydrogen sulfide is used to separate deuterium oxide, or heavy water, from normal water via the Girdler Sulfide process.

Removal from fuel gases Hydrogen sulfide is commonly found in natural gas, biogas, and LPG. It can be removed in a number of ways.

Reaction with iron oxide Gas is pumped through a container of hydrated iron(III) oxide, which combines with hydrogen sulfide. Fe2O3(s) + H2O(l) + 3 H2S(g) → Fe2S3(s) + 4 H2O(l) In order to regenerate iron(III) oxide, the container must be taken out of service, flooded with water and aerated. 2 Fe2S3(s) + 3 O2(g) + 2 H2O(l) → 2 Fe2O3(s) + 2H2O(l) + 6 S(s) On completion of the regeneration reaction the container is drained of water and can be returned to service. The advantage of this system is that it is completely passive during the extraction phase.

Hydrodesulfurization Hydrodesulfurization is a more complex method of removing sulfur from fuels.

Safety Hydrogen sulfide is a highly toxic and flammable gas (explosion limits: 4.3 % - 46 %). Being heavier than air, it tends to accumulate at the bottom of poorly ventilated spaces. Although very pungent at first, it quickly deadens the sense of smell, so potential victims may be unaware of its presence until it is too late. For safe handling procedures, a hydrogen sulfide material safety data sheet (MSDS) should be consulted.

Toxicity Hydrogen sulfide is considered a broad-spectrum poison, meaning that it can poison several different systems in the body, although the nervous system is most affected. The toxicity of H2S is comparable with that of hydrogen cyanide. It forms a complex bond with iron in the mitochondrial cytochrome enzymes, thus preventing cellular respiration.

Since hydrogen sulfide occurs naturally in the body, the environment and the gut, enzymes exist in the body capable of detoxifying it by oxidation to (harmless) sulfate. Hence, low levels of hydrogen sulfide may be tolerated indefinitely. At some threshold level, believed to average around 300–350 ppm, the oxidative enzymes become overwhelmed. Many personal safety gas detectors, such as those used by utility, sewage and petrochemical workers, are set to alarm at as low as 5 to 10 ppm and to go into high alarm at 15 ppm. An interesting diagnostic clue of extreme poisoning by H2S is the discoloration of copper coins in the pockets of the victim. Treatment involves immediate inhalation of amyl nitrite, injections of sodium nitrite, inhalation of pure oxygen, administration of bronchodilators to overcome eventual bronchospasm, and in some cases hyperbaric oxygen therapy (HBO). HBO therapy has anecdotal support and remains controversial. Exposure to lower concentrations can result in eye irritation, a sore throat and cough, nausea, shortness of breath, and fluid in the lungs. These effects are believed to be due to the fact that hydrogen sulfide combines with alkali present in moist surface tissues to form sodium sulfide, a caustic. These symptoms usually go away in a few weeks. Long-term, low-level exposure may result in fatigue, loss of appetite, headaches, irritability, poor memory, and dizziness. Chronic exposure to low level H2S (around 2 ppm) has been implicated in increased miscarriage and reproductive health issues among Russian and Finnish wood pulp workers, but the reports have not (as of circa 1995) been replicated. 

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0.00047 ppm is the recognition threshold, the concentration at which 50% of humans can detect the characteristic odor of hydrogen sulfide, normally described as resembling "a rotten egg". Less than 10 ppm has an exposure limit of 8 hours per day. 10–20 ppm is the borderline concentration for eye irritation. 50–100 ppm leads to eye damage. At 100–150 ppm the olfactory nerve is paralyzed after a few inhalations, and the sense of smell disappears, often together with awareness of danger. 320–530 ppm leads to pulmonary edema with the possibility of death. 530–1000 ppm causes strong stimulation of the central nervous system and rapid breathing, leading to loss of breathing. 800 ppm is the lethal concentration for 50% of humans for 5 minutes exposure (LC50). Concentrations over 1000 ppm cause immediate collapse with loss of breathing, even after inhalation of a single breath.

Hydrogen sulfide was used by the British as a chemical agent during World War I. It was not considered to be an ideal war gas, but, while other gases were in short supply, it was used on two occasions in 1916. The gas, produced by mixing certain household ingredients, was used in a suicide wave in 2008, primarily but not exclusively in Japan.

As of 2010, this has occurred in a number of US cities (and in Putney West London, England), prompting warnings to first responders who can be exposed when responding to a suicide. A dump of toxic waste containing hydrogen sulfide is believed to have caused 17 deaths and thousands of illnesses in Abidjan, on the West Africa coast, in the 2006 Côte d'Ivoire toxic waste dump. In 1975, a hydrogen sulfide explosion in Denver City, located in Yoakum and Gaines counties, Texas, caused the state legislature to focus on the deadly hazards of the gas. State Representative E L Short of Tahoka in Lynn County, took the lead in endorsing an investigation by the Texas Railroad Commission and urged that residents be warned "by knocking on doors if necessary" of the imminent danger stemming from the gas. One may die from the second inhalation of the gas, and a warning itself may be too late.

Function in the body Hydrogen sulfide is produced in small amounts by some cells of the mammalian body and has a number of biological signaling functions. (Only two other such gases are currently known: nitric oxide (NO) and carbon monoxide (CO).) The gas is produced from cysteine by the enzymes cystathionine beta-synthase and cystathionine gamma-lyase. It acts as a relaxant of smooth muscle and as a vasodilator and is also active in the brain, where it increases the response of the NMDA receptor and facilitates long term potentiation, which is involved in the formation of memory. Eventually the gas is converted to sulfite in the mitochondria by thiosulfate reductase, and the sulfite is further oxidized to thiosulfate and sulfate by sulfite oxidase. The sulfates are excreted in the urine. Due to its effects similar to nitric oxide (without its potential to form peroxides by interacting with superoxide), hydrogen sulfide is now recognized as potentially protecting against cardiovascular disease. The cardioprotective role effect of garlic is caused by catabolism of the polysulfide group in allicin to H2S, a reaction that could depend on reduction mediated by glutathione. Though both nitric oxide and hydrogen sulfide have been shown to relax blood vessels, their mechanisms of action are different: while NO activates the enzyme guanylyl cyclase, H2S activates ATP-sensitive potassium channel in smooth muscle cells. Researchers are not clear how the vessel-relaxing responsibilities are shared between nitric oxide and hydrogen sulfide. However there exists some evidence to suggest that nitric oxide does most of the vessel-relaxing work in large vessels and hydrogen sulfide is responsible for similar action in smaller blood vessels.

Like nitric oxide, hydrogen sulfide is involved in the relaxation of smooth muscle that causes erection of the penis, presenting possible new therapy opportunities for erectile dysfunction. In Alzheimer's disease the brain's hydrogen sulfide concentration is severely decreased. In trisomy 21 (the most common form of Down syndrome) the body produces an excess of hydrogen sulfide. Hydrogen sulfide is also involved in the disease process of type 1 diabetes. The beta cells of the pancreas in type 1 diabetes produce an excess of the gas, leading to the death of beta cells and to a reduced production of insulin by those that remain.

Induced hypothermia/suspended animation In 2005, it was shown that mice can be put into a state of suspended animation-like hypothermia by applying a low dosage of hydrogen sulfide (81 ppm H2S) in the air. The breathing rate of the animals sank from 120 to 10 breaths per minute and their temperature fell from 37 °C to just 2 °C above ambient temperature (in effect, they had become cold-blooded). The mice survived this procedure for 6 hours and afterwards showed no negative health consequences. In 2006 it was shown that the blood pressure of mice treated in this fashion with hydrogen sulfide did not significantly decrease. A similar process known as hibernation occurs naturally in many mammals and also in toads, but not in mice. (Mice can fall into a state called clinical torpor when food shortage occurs). If the H2S-induced hibernation can be made to work in humans, it could be useful in the emergency management of severely injured patients, and in the conservation of donated organs. In 2008, hypothermia induced by hydrogen sulfide for 48 hours was shown to reduce the extent of brain damage caused by experimental stroke in rats. As mentioned above, hydrogen sulfide binds to cytochrome oxidase and thereby prevents oxygen from binding, which leads to the dramatic slowdown of metabolism. Animals and humans naturally produce some hydrogen sulfide in their body; researchers have proposed that the gas is used to regulate metabolic activity and body temperature, which would explain the above findings. Two recent studies cast doubt that the effect can be achieved in larger mammals. A 2008 study failed to reproduce the effect in pigs, concluding that the effects seen in mice were not present in larger mammals. Likewise a paper by Haouzi et al. noted that there is no induction of hypometabolism in sheep, either. However, at a February 2010 TED conference, Mark Roth announced that hydrogen sulfide induced hypothermia had completed Phase I clinical trials. He estimated that further trials would take 'a few years.'

Participant in the sulfur cycle Hydrogen sulfide is a central participant in the sulfur cycle, the biogeochemical cycle of sulfur on Earth. In the absence of oxygen, sulfur-reducing and sulfate-reducing bacteria derive energy from oxidizing hydrogen or organic molecules by reducing elemental sulfur or sulfate to hydrogen sulfide. Other bacteria liberate hydrogen sulfide from sulfur-containing amino acids; this gives rise to the odor of rotten eggs and contributes to the odor of flatulence.

Sludge from a pond; the black color is due to metal sulfides As organic matter decays under low-oxygen (or hypoxic) conditions (such as in swamps, eutrophic lakes or dead zones of oceans), sulfate-reducing bacteria will use the sulfates present in the water to oxidize the organic matter, producing hydrogen sulfide as waste. Some of the hydrogen sulfide will react with metal ions in the water to produce metal sulfides, which are not water soluble. These metal sulfides, such as ferrous sulfide FeS, are often black or brown, leading to the dark color of sludge. Several groups of bacteria can use hydrogen sulfide as fuel, oxidizing it to elemental sulfur or to sulfate by using dissolved oxygen, metal oxides (e.g., Fe oxyhydroxides and Mn oxides) or nitrate as oxidant.

The purple sulfur bacteria and the green sulfur bacteria use hydrogen sulfide as electron donor in photosynthesis, thereby producing elemental sulfur. (In fact, this mode of photosynthesis is older than the mode of cyanobacteria, algae, and plants, which uses water as electron donor and liberates oxygen.)

Mass extinctions Hydrogen sulfide could have been implicated in several mass extinctions that have occurred in the Earth's past. In particular, a buildup of hydrogen sulfide in the atmosphere may have caused the Permian-Triassic extinction event 252 million years ago. Organic residues from these extinction boundaries indicate that the oceans were anoxic (oxygen-depleted) and had species of shallow plankton that metabolized H2S. The formation of H2S may have been initiated by massive volcanic eruptions, which emitted CO2 and methane into the atmosphere, which warmed the oceans, lowering their capacity to absorb oxygen that would otherwise oxidize H2S. The increased levels of hydrogen sulfide could have killed oxygen-generating plants as well as depleted the ozone layer, causing further stress. Small H2S blooms have been detected in modern times in the Dead Sea and in the Atlantic ocean off the coast of Namibia.

Chapter-14

Sulfuric Acid

Sulfuric acid/ Sulphuric acid

CAS number ChemSpider UNII EC number UN number KEGG ChEMBL RTECS number Molecular formula Molar mass

IUPAC name Sulfuric acid Other names Oil of vitriol Identifiers 7664-93-9 1086 O40UQP6WCF 231-639-5 1830 D05963 CHEMBL572964 WS5600000 Properties H2SO4 98.086 g/mol

Appearance Density Melting point Boiling point Solubility in water Acidity (pKa) Viscosity

Clear, colorless, odorless liquid 1.84 g/cm3, liquid 10 °C, 283 K, 50 °F 337 °C, 610 K, 639 °F miscible −3 26.7 cP (20 °C) Hazards MSDS External MSDS EU Index 016-020-00-8 Toxic (T) EU classification Corrosive (C) Dangerous for the environment (N) R-phrases R35 S-phrases (S1/2) S26 S30 S45 Flash point Non-flammable Related compounds Selenic acid Related strong acids Hydrochloric acid Nitric acid Sulfurous acid Peroxymonosulfuric acid Related compounds Sulfur trioxide Oleum Supplementary data page Structure and n, εr, etc. properties Thermodynamic Phase behaviour data Solid, liquid, gas Spectral data UV, IR, NMR, MS

Sulfuric acid (alternative spelling sulphuric acid) is a strong mineral acid with the molecular formula H2SO4. Its historical name is vitriol. The salts of sulfuric acid are called sulfates. Sulfuric acid is soluble in water at all concentrations. Sulfuric acid has many applications, and is a central substance in the chemical industry. Principal uses include lead-acid batteries for cars and other vehicles, ore processing, fertilizer manufacturing, oil refining, wastewater processing, and chemical synthesis.

History

John Dalton's 1808 sulfuric acid molecule shows a central sulfur atom bonded to three oxygen atoms. The study of vitriol began in ancient times. Sumerians had a list of types of vitriol that they classified according to substance's color. Some of the earliest discussions on the origin and properties of vitriol are in the works of the Greek physician Dioscorides (first century AD) and the roman naturalist Pliny the Elder (23–79 AD). Galen also discussed its medical use. Metallurgical uses for vitriolic substances were recorded in the Hellenistic alchemical works of Zosimos of Panopolis, in the treatise Phisica et Mystica, and the "Leyden Papyrus x". Iranian alchemists like Jabir Ibn Hayyan (c. 721 – c. 815 AD), Al-Razi (865 – 925 AD), and Jamal Din al-Watwat (d. 1318, wrote the book Mabāhij al-fikar wa-manāhij al'ibar), included vitriol in their mineral classification lists. Ibn Sina focused on its medical uses and different varieties of vitriol. Sulfuric acid was called "oil of vitriol" by medieval European alchemists. There are mentions to it in the works of Vincent of Beauvais and in the Compositum de Compositis ascribed to Albertus Magnus. A passage from Pseudo-Geber´s Summa Perfectionis was long considered to be the first recipe for sulphuric acid, but this was a misinterpretation. In the 17th century, the German-Dutch chemist Johann Glauber prepared sulfuric acid by burning sulfur together with saltpeter (potassium nitrate, KNO3), in the presence of steam. As saltpeter decomposes, it oxidizes the sulfur to SO3, which combines with water to produce sulfuric acid. In 1736, Joshua Ward, a London pharmacist, used this method to begin the first large-scale production of sulfuric acid. In 1746 in Birmingham, John Roebuck adapted this method to produce sulfuric acid in lead-lined chambers, which were stronger, less expensive, and could be made larger than the previously used glass containers. This lead chamber process allowed the effective industrialization of sulfuric acid production. After several refinements, this method remained the standard for sulfuric acid production for almost two centuries. Sulfuric acid created by John Roebuck's process only approached a 35–40% concentration. Later refinements to the lead-chamber process by French chemist JosephLouis Gay-Lussac and British chemist John Glover improved the yield to 78%. However, the manufacture of some dyes and other chemical processes require a more concentrated product. Throughout the 18th century, this could only be made by dry distilling minerals in a technique similar to the original alchemical processes. Pyrite (iron disulfide, FeS2) was heated in air to yield iron (II) sulfate, FeSO4, which was oxidized by further heating in air to form iron(III) sulfate, Fe2(SO4)3, which, when heated to 480 °C, decomposed to iron(III) oxide and sulfur trioxide, which could be passed through water to yield sulfuric acid in any concentration. However, the expense of this process prevented the large-scale use of concentrated sulfuric acid.

In 1831, British vinegar merchant Peregrine Phillips patented the contact process, which was a far more economical process for producing sulfur trioxide and concentrated sulfuric acid. Today, nearly all of the world's sulfuric acid is produced using this method.

Physical properties Grades of sulfuric acid Although nearly 100% sulfuric acid can be made, this loses SO3 at the boiling point to produce 98.3% acid. The 98% grade is more stable in storage, and is the usual form of what is described as "concentrated sulfuric acid." Other concentrations are used for different purposes. Some common concentrations are: Mass fraction Density Concentration Common name H2SO4 (kg/L) (mol/L) 10% 1.07 ~1 dilute sulfuric acid battery acid 29–32% 1.25–1.28 4.2–5 (used in lead–acid batteries) chamber acid 62–70% 1.52–1.60 9.6–11.5 fertilizer acid tower acid 78–80% 1.70–1.73 13.5–14 Glover acid 95–98% 1.83 ~18 concentrated sulfuric acid "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 lead chamber itself (<70% to avoid contamination with nitrosylsulfuric acid) 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 (the modern equivalent of chamber acid, used in many titrations) is prepared by slowly adding 98% sulfuric acid to an equal volume of water, with good stirring: the temperature of the mixture can rise to 80 °C (176 °F) or higher. When high concentrations of SO3 gas are added to sulfuric acid, H2S2O7, called pyrosulfuric acid, fuming sulfuric acid or oleum or, less commonly, Nordhausen acid, is formed. Concentrations of oleum are either expressed in terms of % SO3 (called % oleum) or as % H2SO4 (the amount made if H2O were added); common concentrations are 40% oleum (109% H2SO4) and 65% oleum (114.6% H2SO4). Pure H2S2O7 is a solid with melting point 36°C. Pure sulfuric acid is a viscous clear liquid, like oil, and this explains the old name of the acid ('oil of vitriol').

Commercial sulfuric acid is sold in several different purity grades. Technical grade H2SO4 is impure and often colored, but is suitable for making fertilizer. Pure grades such as United States Pharmacopoeia (USP) grade are used for making pharmaceuticals and dyestuffs. Analytical grades are also available.

Polarity and conductivity Anhydrous H2SO4 is a very 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−4

The equilibrium constant for the autoprotolysis is Kap(25°C)= [H3SO+4][HSO−4] = 2.7×10−4. The comparable equilibrium constant for water, Kw is 10−14, a factor of 1010 (10 billion) 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 intra-molecular proton-switch mechanism (analogous to the Grotthuss mechanism in water), making sulfuric acid a good conductor. It is also an excellent solvent for many reactions. The equilibrium is actually more complex than shown above; 100% H2SO4 contains the following species at equilibrium (figures shown as millimoles per kilogram of solvent): HSO−4 (15.0), H3SO+4 (11.3), H3O+ (8.0), HS2O−7 (4.4), H2S2O7 (3.6), H2O (0.1).

Chemical properties Reaction with water The hydration reaction of sulfuric acid is highly exothermic. One should always add 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+ + HSO4− K1 = 2.4 x 106 (strong acid) HSO4− + H2O → H3O+ + SO42− K2 = 1.0 x 10-2 HSO4- is the bisulfate anion and SO42- is the sulfate anion. K1 and K2 are the acid dissociation constants. Because the hydration of sulfuric acid is thermodynamically favorable, sulfuric acid is an excellent dehydrating agent, and is used to prepare many

dried fruits. The affinity of sulfuric acid for water is sufficiently strong that it will remove hydrogen and oxygen atoms from other compounds; for example, mixing starch (C6H12O6)n and concentrated sulfuric acid will give elemental carbon and water which is absorbed by the sulfuric acid (which becomes slightly diluted): (C6H12O6)n → 6n C + 6n H2O The effect of this can be seen when concentrated sulfuric acid is spilled on paper; the cellulose reacts to give a burnt appearance, the carbon appears much as soot would in a fire. A more dramatic reaction occurs when sulfuric acid is added to a tablespoon of white sugar in a beaker; a rigid column of black, porous carbon will quickly emerge. The carbon will smell strongly of caramel due to the heat generated. Although less dramatic, the action of the acid on cotton, even in diluted form, will destroy the fabric.

Other reactions As an acid, sulfuric acid reacts with most bases to give the corresponding sulfate. For example, the blue copper salt copper(II) sulfate, commonly used for electroplating and as a fungicide, is prepared by the reaction of copper(II) oxide with sulfuric acid: CuO (s) + H2SO4 (aq) → CuSO4 (aq) + H2O (l) Sulfuric acid can also be used to displace weaker acids from their salts. Reaction with sodium acetate, for example, displaces acetic acid, CH3COOH, and forms sodium bisulfate: H2SO4 + CH3COONa → NaHSO4 + CH3COOH Similarly, reacting sulfuric acid with potassium nitrate can be used to produce nitric acid and a precipitate of potassium bisulfate. When combined with nitric acid, sulfuric acid acts both as an acid and a dehydrating agent, forming the nitronium ion NO+ 2, which is important in nitration reactions involving electrophilic aromatic substitution. This type of reaction, where protonation occurs on an oxygen atom, is important in many organic chemistry reactions, such as Fischer esterification and dehydration of alcohols. Concentrated sulfuric acid reacts with sodium chloride, and gives hydrogen chloride gas and sodium bisulfate: NaCl + H2SO4 → NaHSO4 + HCl Concentrated sulfuric acid also dehydrates sugar, leaving a porous black carbon mass behind. Sulfuric acid does not take part in this reaction, but it decomposes the sugar. During this reaction heat is produced and water vapor is given off. Reaction is as follows: C12H22O11 → 12C + 11H2O

Sulfuric acid reacts with most metals via a single displacement reaction to produce hydrogen gas and the metal sulfate. Dilute H2SO4 attacks iron, aluminium, zinc, manganese, magnesium and nickel, but reactions with tin and copper require the acid to be hot and concentrated. Lead and tungsten, however, are resistant to sulfuric acid. The reaction with iron shown below is typical for most of these metals, but the reaction with tin produces sulfur dioxide rather than hydrogen. Fe (s) + H2SO4 (aq) → H2 (g) + FeSO4 (aq) Sn (s) + 2 H2SO4 (aq) → SnSO4 (aq) + 2 H2O (l) + SO2 (g) These reactions may be taken as typical: the hot concentrated acid generally acts as an oxidizing agent whereas the dilute acid acts a typical acid. Hence hot concentrated acid reacts with tin, zinc and copper to produce the salt, water and sulfur dioxide, whereas the dilute acid reacts with metals high in the reactivity series (such as Zn) to produce a salt and hydrogen. This is explained more fully in 'A New Certificate Chemistry' by Holderness and Lambert. Benzene undergoes electrophilic aromatic substitution with sulfuric acid to give the corresponding sulfonic acids:

Occurrence Pure sulfuric acid is not encountered naturally on Earth, due to its great affinity for water. Apart from that, sulfuric acid is a constituent of acid rain, which is formed by atmospheric oxidation of sulfur dioxide in the presence of water – i.e., oxidation of sulfurous acid. Sulfur dioxide is the main byproduct produced when sulfur-containing fuels such as coal or oil are burned. Sulfuric acid is formed naturally by the oxidation of sulfide minerals, such as iron sulfide. The resulting water can be highly acidic and is called acid mine drainage (AMD) or acid rock drainage (ARD). This acidic water is capable of dissolving metals present in sulfide ores, which results in brightly colored, toxic streams. The oxidation of pyrite (iron sulfide) by molecular oxygen produces iron(II), or Fe2+: 2 FeS2 (s) + 7 O2 + 2 H2O → 2 Fe2+ (aq) + 4 SO2− + 4 (aq) + 4 H The Fe2+ can be further oxidized to Fe3+:

4 Fe2+ + O2 + 4 H+ → 4 Fe3+ + 2 H2O The Fe3+ produced can be precipitated as the hydroxide or hydrous oxide: Fe3+ (aq) + 3 H2O → Fe(OH)3 (s) + 3 H+ The iron(III) ion ("ferric iron") can also oxidize pyrite: FeS2 (s) + 14 Fe3+ + 8 H2O → 15 Fe2+ (aq) + 2 SO2−4 (aq) + 16 H+ When iron(III) oxidation of pyrite occurs, the process can become rapid. pH values below zero have been measured in ARD produced by this process. ARD can also produce sulfuric acid at a slower rate, so that the acid neutralizing capacity (ANC) of the aquifer can neutralize the produced acid. In such cases, the total dissolved solids (TDS) concentration of the water can be increased from the dissolution of minerals from the acid-neutralization reaction with the minerals.

Extraterrestrial sulfuric acid

Venus Sulfuric acid is produced in the upper atmosphere of Venus by the Sun's photochemical action on carbon dioxide, sulfur dioxide, and water vapor. Ultraviolet photons of wavelengths less than 169 nm can photodissociate carbon dioxide into carbon monoxide and atomic oxygen. Atomic oxygen is highly reactive. When it reacts with sulfur dioxide, a trace component of the Venusian atmosphere, the result is sulfur trioxide, which can combine with water vapor, another trace component of Venus's atmosphere, to yield sulfuric acid. In the upper, cooler portions of Venus's atmosphere, sulfuric acid exists as a liquid, and thick sulfuric acid clouds completely obscure the planet's surface when viewed from above. The main cloud layer extends from 45–70 km above the planet's surface, with thinner hazes extending as low as 30 km and as high as 90 km above the surface. The permanent Venusian clouds produce a concentrated acid rain, as the clouds in the atmosphere of Earth produce water rain. The atmosphere exhibits a sulfuric acid cycle. As sulfuric acid rain droplets fall down through the hotter layers of the atmosphere's temperature gradient, they are heated up and release water vapor, becoming more and more concentrated. When they reach temperatures above 300°C, sulfuric acid begins to decompose into sulfur trioxide and water, both in the gas phase. Sulfur trioxide is highly reactive and dissociates into sulfur dioxide and atomic oxygen, which oxidizes traces of carbon monoxide to form carbon dioxide. Sulfur dioxide and water vapor rise on convection currents from the mid-level atmospheric layers to higher altitudes, where they will be transformed again into sulfuric acid, and the cycle repeats.

Europa Infrared spectra from NASA's Galileo mission show distinct absorptions on Jupiter's moon Europa that have been attributed to one or more sulfuric acid hydrates. Sulfuric acid in solution with water causes significant freezing-point depression of water's melting point, down to 210 K, and this would make more likely the existence of liquid solutions beneath Europa's icy crust.The interpretation of the spectra is somewhat controversial. Some planetary scientists prefer to assign the spectral features to the sulfate ion, perhaps as part of one or more minerals on Europa's surface.

Manufacture Sulfuric acid is produced from sulfur, oxygen and water via the conventional contact process (DCDA) or the wet sulfuric acid process (WSA).

Contact process (DCDA) In the first step, sulfur is burned to produce sulfur dioxide. S (s) + O2 (g) → SO2 (g) This is then oxidized to sulfur trioxide using oxygen in the presence of a vanadium(V) oxide catalyst. 2 SO2 (g) + O2 (g) → 2 SO3 (g) (in presence of V2O5) The sulfur trioxide is absorbed into 97–98% H2SO4 to form oleum (H2S2O7), also known as fuming sulfuric acid. The oleum is then diluted with water to form concentrated sulfuric acid. H2SO4 (l) + SO3 → H2S2O7 (l) H2S2O7 (l) + H2O (l) → 2 H2SO4 (l) Note that directly dissolving SO3 in water is not practical due to the highly exothermic nature of the reaction between sulfur trioxide and water. The reaction forms a corrosive aerosol that is very difficult to separate, instead of a liquid. SO3 (g) + H2O (l) → H2SO4 (l)

Wet sulfuric acid process (WSA) In the first step, sulfur is burned to produce sulfur dioxide: S(s) + O2(g) → SO2(g) or, alternatively, hydrogen sulfide (H2S) gas is incinerated to SO2 gas:

2 H2S + 3 O2 → 2 H2O + 2 SO2 (−518 kJ/mol) This is then oxidized to sulfur trioxide using oxygen with vanadium(V) oxide as catalyst. 2 SO2 + O2 → 2 SO3 (−99 kJ/mol) The sulfur trioxide is hydrated into sulfuric acid H2SO4: SO3 + H2O → H2SO4(g) (−101 kJ/mol) The last step is the condensation of the sulfuric acid to liquid 97–98% H2SO4: H2SO4(g) → H2SO4(l) (−69 kJ/mol)

Other methods Another method is the less well-known metabisulfite method, in which metabisulfite is placed at the bottom of a beaker, and 12.6 molar concentration hydrochloric acid is added. The resulting gas is bubbled through nitric acid, which will release brown/red vapors. The completion of the reaction is indicated by the ceasing of the fumes. This method does not produce an inseparable mist, which is quite convenient. Sulfuric acid can be produced in the laboratory by burning sulfur in air and dissolving the gas produced in a hydrogen peroxide solution. SO2 + H2O2 → H2SO4 Prior to 1900, most sulfuric acid was manufactured by the chamber process. As late as 1940, up to 50% of sulfuric acid manufactured in the United States was produced by chamber process plants.

Uses

Sulfuric acid production in 2000

Sulfuric acid is a very important commodity chemical, and indeed, a nation's sulfuric acid production is a good indicator of its industrial strength. World production in 2001 was 165 million tons, with an approximate value of US$8 billion. The major use (60% of total production worldwide) for sulfuric acid is in the "wet method" for the production of phosphoric acid, used for manufacture of phosphate fertilizers as well as trisodium phosphate for detergents. In this method, phosphate rock is used, and more than 100 million tonnes are processed annually. This raw material is shown below as fluorapatite, though the exact composition may vary. This is treated with 93% sulfuric acid to produce calcium sulfate, hydrogen fluoride (HF) and phosphoric acid. The HF is removed as hydrofluoric acid. The overall process can be represented as: Ca5F(PO4)3 + 5 H2SO4 + 10 H2O → 5 CaSO4·2 H2O + HF + 3 H3PO4 Sulfuric acid is used in large quantities by the iron and steelmaking industry to remove oxidation, rust and scale from rolled sheet and billets prior to sale to the automobile and white goods (appliances) industry. Used acid is often recycled using a Spent Acid Regeneration (SAR) plant. These plants combust spent acid with natural gas, refinery gas, fuel oil or other fuel sources. This combustion process produces gaseous sulfur dioxide (SO2) and sulfur trioxide (SO3) which are then used to manufacture "new" sulfuric acid. SAR plants are common additions to metal smelting plants, oil refineries, and other industries where sulfuric acid is consumed in bulk, as operating a SAR plant is much cheaper than the recurring costs of spent acid disposal and new acid purchases. Ammonium sulfate, an important nitrogen fertilizer, is most commonly produced as a byproduct from coking plants supplying the iron and steel making plants. Reacting the ammonia produced in the thermal decomposition of coal with waste sulfuric acid allows the ammonia to be crystallized out as a salt (often brown because of iron contamination) and sold into the agro-chemicals industry. Another important use for sulfuric acid is for the manufacture of aluminum sulfate, also known as paper maker's alum. This can react with small amounts of soap on paper pulp fibers to give gelatinous aluminum carboxylates, which help to coagulate the pulp fibers into a hard paper surface. It is also used for making aluminum hydroxide, which is used at water treatment plants to filter out impurities, as well as to improve the taste of the water. Aluminum sulfate is made by reacting bauxite with sulfuric acid: Al2O3 + 3 H2SO4 → Al2(SO4)3 + 3 H2O Sulfuric acid is used for a variety of other purposes in the chemical industry. For example, it is the usual acid catalyst for the conversion of cyclohexanone oxime to caprolactam, used for making nylon. It is used for making hydrochloric acid from salt via the Mannheim process. Much H2SO4 is used in petroleum refining, for example as a catalyst for the reaction of isobutane with isobutylene to give isooctane, a compound that raises the octane rating of gasoline (petrol). Sulfuric acid is also important in the manufacture of dyestuffs solutions and is the "acid" in lead-acid (car) batteries.

Sulfuric acid is also used as a general dehydrating agent in its concentrated form.

Sulfur-iodine cycle The sulfur-iodine cycle is a series of thermo-chemical processes used to obtain hydrogen. It consists of three chemical reactions whose net reactant is water and whose net products are hydrogen and oxygen. 2 H2SO4 → 2 SO2 + 2 H2O + O2 I2 + SO2 + 2 H2O → 2 HI + H2SO4 2 HI → I2 + H2

(830 °C) (120 °C) (320 °C)

The sulfur and iodine compounds are recovered and reused, hence the consideration of the process as a cycle. This process is endothermic and must occur at high temperatures, so energy in the form of heat has to be supplied. The sulfur-iodine cycle has been proposed as a way to supply hydrogen for a hydrogenbased economy. It does not require hydrocarbons like current methods of steam reforming. The sulfur-iodine cycle is currently being researched as a feasible method of obtaining hydrogen, but the concentrated, corrosive acid at high temperatures poses currently insurmountable safety hazards if the process were built on a large scale.

Safety Laboratory hazards

Drops of 98% sulfuric acid char a piece of tissue paper instantly The corrosive properties of sulfuric acid are accentuated by its highly exothermic reaction with water. Burns from sulfuric acid are potentially more serious than those of comparable strong acids (e.g. hydrochloric acid), as there is additional tissue damage due to dehydration and particularly secondary thermal damage due to the heat liberated by the reaction with water. The danger is greater with more concentrated preparations of sulfuric acid, but even the normal laboratory "dilute" grade (approximately 1 M, 10%) will char paper by dehydration if left in contact for a sufficient time. Therefore, solutions equal to or stronger than 1.5 M are labeled "CORROSIVE", while solutions greater than 0.5 M but less than 1.5 M are labeled "IRRITANT". Fuming sulfuric acid (oleum) is not recommended for use in schools as it is quite hazardous. The standard first aid treatment for acid spills on the skin is, as for other corrosive agents, irrigation with large quantities of water. Washing is continued for at least ten to fifteen minutes to cool the tissue surrounding the acid burn and to prevent secondary damage.

Contaminated clothing is removed immediately and the underlying skin washed thoroughly. Preparation of the diluted acid can also be dangerous due to the heat released in the dilution process. The concentrated acid is always added to water and not the other way round, to take advantage of the relatively high heat capacity of water. Addition of water to concentrated sulfuric acid leads to the dispersal of a sulfuric acid aerosol or worse, an explosion. Preparation of solutions greater than 6 M (35%) in concentration is most dangerous, as the heat produced may be sufficient to boil the diluted acid: efficient mechanical stirring and external cooling (such as an ice bath) are essential. On a laboratory scale, sulfuric acid can be diluted by pouring concentrated acid onto crushed ice made from de-ionized water. The ice melts in an endothermic process while dissolving the acid. The amount of heat needed to melt the ice in this process is greater than the amount of heat evolved by dissolving the acid so the solution remains cold. After all the ice has melted, further dilution can take place using water.

Industrial hazards Although sulfuric acid is non-flammable, contact with metals in the event of a spillage can lead to the liberation of hydrogen gas. The dispersal of acid aerosols and gaseous sulfur dioxide is an additional hazard of fires involving sulfuric acid. Sulfuric acid is not considered toxic besides its obvious corrosive hazard, and the main occupational risks are skin contact leading to burns and the inhalation of aerosols. Exposure to aerosols at high concentrations leads to immediate and severe irritation of the eyes, respiratory tract and mucous membranes: this ceases rapidly after exposure, although there is a risk of subsequent pulmonary edema if tissue damage has been more severe. At lower concentrations, the most commonly reported symptom of chronic exposure to sulfuric acid aerosols is erosion of the teeth, found in virtually all studies: indications of possible chronic damage to the respiratory tract are inconclusive as of 1997. In the United States, the permissible exposure limit (PEL) for sulfuric acid is fixed at 1 mg/m³: limits in other countries are similar. There have been reports of sulfuric acid ingestion leading to vitamin B12 deficiency with subacute combined degeneration. The spinal cord is most often affected in such cases, but the optic nerves may show demyelination, loss of axons and gliosis.

Legal restrictions International commerce of sulfuric acid is controlled under the United Nations Convention Against Illicit Traffic in Narcotic Drugs and Psychotropic Substances, 1988, which lists sulfuric acid under Table II of the convention as a chemical frequently used in the illicit manufacture of narcotic drugs or psychotropic substances. In the US sulfuric acid is included in List II of the list of essential or precursor chemicals established pursuant to the Chemical Diversion and Trafficking Act. Accordingly,

transactions of sulfuric acid—such as sales, transfers, exports from and imports to the United States—are subject to regulation and monitoring by the Drug Enforcement Administration.