The materials to prepare students for practical lessons of inorganic chemistry

The materials to prepare students for practical lessons of inorganic chemistry

LESSON № 20.

THEME. р-elements of the VІ group. Oxygen and properties of it compounds. р-elements of the VІA group. Sulfur, Selenium, Tellurium and polonium

 

The elements in this group have six electrons in their outer quantum level, and can thus achieve a noble gas configuration by acquiring two electrons.

PROPERTIES OF THE ELEMENTS OF THE VІ GROUP

Some of the more important physical properties of the elements are given in Table 1. Melting and boiling points increase with increasing atomic number from oxygen to tellurium, with oxygen showing the deviation typical of a group head element. The expected decrease in ionisation energy with increase in atomic number and size of the atoms should be noted.

Although the electron affinities do not change regularly with increasing atomic number, the increasing ionic radii imply decreasing lattice and hydration enthalpies. Hence, although oxygen forms a large number of wholly or partly ionic oxides with metals, containing O^{2 ~}, sulphur forms ionic compounds only with the more electropositive elements such as sodium, and most of its compounds are partly or wholly covalent. All the elements are able to share two electrons forming two covalent bonds. The two covalent bonds formed by oxygen can be separate bonds, for example

O 1s²2s²2p⁴

S 2s²2p⁶3s²3p⁴

Se 3s²3p⁶3d¹⁰4s²4p⁴

Te 4s²4p⁶4d¹⁰5s²5p⁴

Po 4s²4p⁶4d¹⁰4f¹⁴5s²5p⁶5d¹⁰6s²6p⁴

Table 1.

As we approach the right-hand side of the periodic table, similarities among the elements within a group become greater again. This is true of group VIA. Except polonium, which is radioactive and usually omitted from discussion, all members of the group form X^2– ions when combined with highly electropositive metals. The tendency to be reduced to the –2 oxidation state decreases significantly from top to bottom of the group, however, and tellurium shows some metallic properties. The group VIA elements are called chalcogens because most ores of copper (Greek chalkos) are oxides or sulfides, and such ores contain traces of selenium and tellurium.

 (The chalcogens (/ˈkælkədʒɨnz/) are the chemical elements in group 16 of the periodic table. This group is also known as the oxygen family. It consists of the elements oxygen (O), sulfur (S), selenium (Se), tellurium Te), and the radioactive element polonium (Po). The synthetic element livermorium (Lv) is predicted to be a chalcogen as well. Often, oxygen is treated separately from the other chalcogens, sometimes even excluded from the scope of the term “chalcogen” altogether, due to its very different chemical behavior from sulfur, selenium, tellurium and polonium. The word “chalcogen” is derived from a combination of the Greek word khalkόs (χαλκός) principally meaning copper (the term was also used for bronze/brass, any metal in the poetic sense, ore or coin), and the Latinised Greek word genēs, meaning born or produced. Sulfur has been known since antiquity, and oxygen was recognized as an element in the 18th century. Selenium, tellurium and polonium were discovered in the 19th century, and livermorium in 2000.

All of the chalcogens have six valence electrons, leaving them two electrons short of a full outer shell. Their most common oxidation states are −2, +2, +4, and +6. They have relatively low atomic radii, especially the lighter ones.

Lighter chalcogens are typically nontoxic in their elemental form, and are often critical to life, while the heavier chalcogens are typically toxic. All of the chalcogens have some role in biological functions, either as a nutrient or a toxin. The lighter chalcogens, such as oxygen and sulfur, are rarely toxic and usually helpful in their pure form. Selenium is an important nutrient but is also commonly toxic. Tellurium often has unpleasant effects (although some organisms can use it), and polonium is always extremely harmful, both in its chemical toxicity and its radioactivity.

Sulfur has more than 20 allotropes, oxygen has nine, selenium has at least five, polonium has two, and only one crystal structure of tellurium has so far been discovered. There are numerous organic chalcogen compounds. Not counting oxygen, organic sulfur compounds are generally the most common, followed by organic selenium compounds and organic tellurium compounds. This trend also occurs with chalcogen pnictides and compounds containing chalcogens and carbon group elements.

Oxygen is generally extracted from air and sulfur is extracted from oil and natural gas. Selenium and tellurium are produced as byproducts of copper refining. Polonium and livermorium are most available in particle accelerators. The primary use of elemental oxygen is in steelmaking. Sulfur is mostly converted into sulfuric acid, which is heavily used in the chemical industry. Selenium’s most common application is glassmaking. Tellurium compounds are mostly used in optical disks, electronic devices, and solar cells. Some of polonium’s applications are due to its radioactivity.

History

Early discoveries

Sulfur was known in the ancient history and is mentioned in Bible 15 times. Sulfur was known to the ancient Greeks and commonly mined by the ancient Romans. Sulfur was also historically used as a component of Greek fire. In the Middle Ages, sulfur was a key part of alchemical experiments. In the 1700s and 1800s, scientists Joseph Louis Gay-Lussac and Louis-Jacques Thénard proved sulfur to be a chemical element.

Early attempts to discover oxygen from air were hampered by the fact that air was thought of as a single element up to the 17th and 18th centuries. Robert Hooke, Mikhail Lomonosov, Ole Borch, and Pierre Bayden all successfully created oxygen, but did not realize it at the time. Oxygen was discovered by Joseph Priestley in 1774 when he focused sunlight on a sample of mercuric oxide and collected the resulting gas. Carl Wilhelm Scheele had also created oxygen in 1771 by the same method, but Scheele did not publish his results until 1777.

Tellurium was first discovered in 1783 by Franz Joseph Müller von Reichenstein. He discovered tellurium in a sample of what is now known as calaverite. Müller assumed at first that the sample was pure antimony, but tests he ran on the sample did not agree with this. Muller then guessed that the sample was bismuth sulfide, but tests confirmed that the sample was not that. For some years, Muller pondered the problem. Eventually he realized that the sample was gold bonded with an unknown element. In 1796, Müller sent part of the sample to the German chemist Martin Klaproth, who purified the undiscovered element. Klaproth decided to call the element tellurium after the Latin word for earth.

Selenium was discovered in 1817 by Jöns Jacob Berzelius. Berzelius discovered a reddish-brown sediment at a sulfuric acid manufacturing plant. The sample was thought to contain arsenic. Berzelius initially thought that the sediment contained tellurium, but came to realize that the sample also contained a new element, which he named selenium after the Greek word for moon.

Periodic table placing

Mendeleev’s periodic system proposed in 1871 showing oxygen, sulfur, selenium and tellurium part of his group VI.

Three of the chalcogens (sulfur, selenium, and tellurium) were part of the discovery of periodicity, as they are among a series of triads of elements in the same group that were noted by Johann Wolfgang Döbereiner as having similar properties. Around 1865 John Newlands produced a series of papers where he listed the elements in order of increasing atomic weight and similar physical and chemical properties that recurred at intervals of eight; he likened such periodicity to the octaves of music. His version included a “group b” consisting of oxygen, sulfur, selenium, tellurium, and osmium.

After 1869, Dmitri Mendeleev proposed his periodic table placing oxygen at the top of “group VI” above sulfur, selenium, and tellurium. Chromium, molybdenum, tungsten, and uranium were sometimes included in this group, but they would be later rearranged as part of group VIB; uranium would later be moved to the actinide series. Oxygen, along with sulfur, selenium, tellurium, and later polonium would be grouped in group VIA, until the group’s name was changed to group 16 in 1988.

Modern discoveries

In the late 19th century, Marie Curie and Pierre Curie discovered that a sample of pitchblende was emitting four times as much radioactivity as could be explained by the presence of uranium alone. The Curies gathered several tons of pitchblende and refined it for several months until they had a pure sample of polonium. The discovery officially took place in 1898. Prior to the invention of particle accelerators, the only way to create polonium was to extract it over several months from uranium ore.

The first attempt at creating livermorium was from 1976 to 1977 at the LBNL, who bombarded curium-248 with calcium-48, but were not successful. After several failed attempts in 1977, 1998, and 1999 by research groups in Russia, Germany, and the USA, livermorium was created successfully in 2000 at the Joint Institute for Nuclear Research by bombarding curium-248 atoms with calcium-48 atoms. The element was known as ununhexium until it was officially named livermorium in 2012.

Atomic properties of the chalcogens are summarized in the table.

Table 2.

Properties of the Group VIA Elements

 

Table 3.

At ordinary temperatures and pressures, oxygen is a gas. It exists in either of two allotropic forms: O₂, which makes up 21 percent of the earth’s atmosphere, or O₃ (ozone), which slowly decomposes to O₂. O₃ can be prepared by passing an electrical discharge through O₂ or air:

3O₂(g) 2O₃(g)

This reaction occurs naturally as a result of lightning bolts. O₃ is also produced by any device which produces electrical sparks. You may have noticed its distinctive odor in the vicinity of an electric motor, for example.

Ozone is formed in the earth’s stratosphere (between altitudes of 10 and. 50 km) by ultraviolet rays whose wavelengths are shorter than 250 nm:

O₂ 2O

O + O₂ → O₃      (1)

The ozone itself absorbs longer-wavelength ultraviolet radiation (up to 340 nm), preventing these harmful rays fom reaching the earth’s surface. Otherwise these rays would increase the incidence of human skin cancer and cause other environmental problems. In recent years convincing evidence has been obtained to show that nitrogen oxide emissions from supersonic transport (SST) airplanes (which fly in the stratosphere) can reduce the concentration of ozone. Similar conclusions have been drawn regarding chlorofluorocarbons(sometimes referred to as CFCs) used as propellants in aerosol hair sprays and deodorants. Once in the atmosphere, a photochemical reaction causes atomic chlorine to be broken off from CFCs. This atomic chlorine can then participate in a catalytic ozone depleting reaction:

Cl + O₃ → ClO + O₂

ClO + O₃ → Cl + 2 O₂

Atomic chlorine is regenerated, meaning that each CFC molecule has the potential to deplete large amounts of ozone. In the 1980s, it was determined that use of chemicals such as CFCs were thinning stratospheric ozone. This is also when the “ozone hole” over Antarctica was discovered. In response to the depletion of ozone, and the danger presented by it, the Montreal Protocol on Substances that Deplete the Ozone Layer was signed by leaders of multiple countries, with the goal to phase out production and use of CFCs and other chemicals harmful to the ozone layer. Today, 191 countries have signed the protocol, and while it is projected to take until 2075 for ozone levels to return to normal, the Montreal Protocol has so far proven a success.

O₃ is also an important component of photochemical smog. It is produced when O atoms (formed by breaking N—O bonds in NO₂) react with molecules according to Eq. (1). O₃ is a stronger oxidizing agent than O₂. It reacts with unsaturated hydrocarbons ( alkenes) in evaporated gasoline to produce aldehydes and ketones which are eye irritants. Rubber is a polymeric material which contains bonds, and so it too reacts with O₃. Further, ground level ozone and the accompanying smog has proven a significant health concern, irritating and damaging the respiratory system and also having links to asthma. So ozone is beneficial when in the upper atmosphere, but has adverse effects when at ground level.

Sulfur occurs in a variety of allotropic forms. At room temperature the most stable form is rhombic sulfur. This yellow solid consists of S₈ molecules (seen in the Jmol below) packed in a crystal lattice which belongs to the orthorhombic system (listed on the page discussing crystal systems).

When heated to 96°C, solid rhombic sulfur changes very slowly into monoclinic sulfur, in which one-third of the S₈ molecules are randomly oriented in the crystal lattice. When either form of sulfur melts, the liquid is at first pale yellow and flows readily, but above 160°C it becomes increasingly viscous. Only near the boiling point of 444.6°C does it thin out again. This unusual change in viscosity with temperature is attributed to opening of the eight-membered ring of S₈ and formation of long chains of sulfur atoms. These intertwine and prevent the liquid from flowing. This explanation is supported by the fact that if the viscous liquid is cooled rapidly by pouring it into water, the amorphous sulfur produced can be shown experimentally to consist of long chains of sulfur atoms.

Both selenium and tellurium have solid structures in which the atoms are bonded in long spiral chains. Both are semiconductors, and the electrical conductivity of selenium depends on the intensity of light falling on the element. This property is utilized in selenium photocells, which are often used in photographic exposure meters.

Selenium is also used in rectifiers to convert alternating electrical current to direct current. Compounds of selenium and tellurium are of little commercial importance, and they often are toxic. Moreover, many of them have foul odors, are taken up by the body, and are given off in perspiration and on the breath. These properties have inhibited study of tellurium and selenium compounds.

 

Chemical Reactions and Compounds

Oxygen

Since oxygen has the second largest electronegativity among all the elements, it is found in the –2 oxidation state in most compounds. Important oxides have already been discussed in sections dealing with the elements from which they form, and so we will deal only with unusual oxidation states of oxygen here. One of these is the +2 state found in OF₂, the most common compound in which oxygen is combined with the more electronegative fluorine. We have already mentioned the –½ and –1 states observed in alkali-metal superoxides and peroxides, but one important peroxide, hydrogen peroxide (H₂O₂), has not yet been discussed.

H₂O₂ can be prepared by electrolysis of solutions containing sulfate ions. H₂O₂ is a weak acid, and it can serve as an oxidizing agent (oxygen being reduced to the –2 state) or as a reducing agent (oxygen being oxidized to the 0 state). Like the peroxide ion, the H₂O₂ molecule contains an O—O single bond. This bond is rather weak compared with many other single bonds, and this contributes to the reactivity of H₂O₂. The compound decomposes easily, especially if exposed to light or contaminated with traces of transition metals. The decomposition

2H₂O₂(l) → 2H₂O(l) + O₂(g)

can occur explosively in the case of the pure liquid.

 

Sulfur

Although this element is only sixteenth in abundance at the surface of the earth, it is one of the few that has been known and used throughout history. Deposits of elemental sulfur are not uncommon, and, because they were stones that would burn, were originally called brimstone. Burning sulfur produces sulfur dioxide,

S₈(s) + 😯₂(g) → 8SO₂(g)

This colorless gas has a choking odor and is more poisonous than carbon monoxide. It is the anhydride of sulfurous acid, a weak diprotic acid:

SO₂(g) + H₂O(l) → H₂SO₃(aq)

SO₂ is also produced when almost any sulfur-containing substance is burned in air. Coal, for example, usually contains from 1 to 4% sulfur, and so burning coal releases SO₂ to the atmosphere. Many metal ores are sulfides, and when they are heated in air, SO₂ is produced. Copper, for example, may be obtained as the element by heating copper(I) sulfide:

Cu₂S(s) + O₂(g) 2Cu(s) + SO₂(g)

Since SO₂ is so poisonous, its release to the atmosphere is a major pollution problem. Once in the air, SO₂ is slowly oxidized to sulfur trioxide, SO₃:

2SO₂(g) + O₂(g) → 2SO₃(g)

This compound is the anhydride of sulfuric acid, H₂SO₄:

SO₃(g) + H₂O(l) → H₂SO₄(aq)

Thus if air is polluted with SO₂ and SO₃, a fine mist of dilute droplets of can form. All three substances are very irritating to the throat and lungs and are responsible for considerable damage to human health.

The natural mechanism for removal of sulfur oxides from the air is solution in raindrops, followed by precipitation. This makes the rainwater more acidic than it would otherwise be, and acid rain is now common in industrialized areas of the United States and Europe. Acid rain can slowly dissolve limestone and marble, both of which consist of CaCO₃:

CaCO₃(s) + H₃O⁺(aq) → Ca²⁺(aq) + HCO₃^–(aq) + H₂O(l)

Thus statues and buildings made of these materials may be damaged.

Despite the fact that a tremendous amount of sulfur is released to the environment by coal combustion and ore smelting, this element is not usually recovered from such processes. Instead it is obtained commercially from large deposits along the U.S. Gulf Coast and from refining of sour petroleum. Sour petroleum contains numerous sulfur compounds, including H₂S, which smells like rotten eggs. The deposits of elemental sulfur in Texas and Louisiana are mined by the Frasch process. Water at 170°C is pumped down a pipe to melt the sulfur, and the latter is forced to the surface by compressed air. Most of the H₂S or S₈ obtained from these sources is oxidized to SO₂, passed over a vanadium catalyst to make SO₃, and dissolved in water to make H₂SO₄. In 2005 an estimated 190 billion kg of H₂SO₄ was produced in the world, making H₂SO₄ one of the most important industrial chemicals. About half of it is used in phosphate fertilizer production.

Pure H₂SO₄ is a liquid at room temperature and has a great affinity for H₂O. This is apparently due to the reaction

H₂SO₄ + H₂O → H₃O⁺ + HSO₄^–

Formation of H₃O⁺ releases energy, and the reaction is exothermic. Concentrated H₂SO₄ is 93% H₂SO₄ and 7% H₂O by mass, corresponding to more than twice as many H₂SO₄ as H₂O molecules. Since many H₂SO₄ molecules still have protons to donate, concentrated H₂SO₄ also has a great affinity for H₂O. It is often used as a drying agent and can be employed in condensation reactions which give off H₂O.

 

THE ELEMENTS: OCCURRENCE AND EXTRACTION

OXYGEN

Oxygen occurs free in the atmosphere (21% by volume. 23% by weight). The proportion is constant-over the earth’s surface; it is also constant for many miles upwards, because the turbulence of the atmosphere prevents the tendency for the lighter gases, for example helium, to increase in amount at higher altitudes.

Hoffman electrolysis apparatus used in electrolysis of water.

Water contains 89% by weight of oxygen, and the outer crust of the earth contains about 47%; hence air, earth and sea together contain about 50 % by weight of oxygen. On the industrial scale oxygen is obtained by the fractional distillation of air. A common laboratory method for the preparation of oxygen is by the decomposition of hydrogen peroxide, H₂O₂, a reaction catalysed by manganese(IV) oxide:

2H₂O₂ → 2H₂O + O₂↑

A similar decomposition of the chlorate(I) (hypochlorite) ion, OC1^~, catalysed by both light and cobalt(II) ions, is less commonly used:

2C1O^– →2Cl^– + O₂↑

Oxygen can also be prepared by the thermal decomposition of certain solid compounds containing it. These include oxides of the more noble metals, for example of mercury or silver:

2HgO → 2Hg + O₂↑

certain higher oxides, for example of lead(IV) and manganese(IV):

2PbO₂ →2PbO + O₂↑

peroxides, for example of barium:

2BaO₂ ® 2BaO + O₂­

and certain oxosalts, notably the nitrates, chlorates(V), iodates(V) and manganates(VII) of alkali metals.

Pure oxygen is conveniently prepared by the thermal decomposition of potassium manganate(VII):

2KMnO₄ = K₂MnO₄ + MnO₂ + O₂­

2KClO₃ = 2KCl + 3O₂­

Oxygen can be produced by certain reactions in solution, for example the oxidation of hydrogen peroxide by potassium manganate(VII) acidified with sulphuric acid:

2MnO₄^~ + 5H₂O₂ + 6H₃O⁺ → 2Mn²⁺ + 14H₂O + 5O₂↑

 

Biological role

Photosynthesis and respiration. In nature, free oxygen is produced by the light-driven splitting of water during oxygenic photosynthesis. Green algae and cyanobacteria in marine environments provide about 70% of the free oxygen produced on earth and the rest is produced by terrestrial plants.

A simplified overall formula for photosynthesis is:

6 CO₂ + 6 H₂O + photons → C₆H₁₂O₆ + 6 O₂↑

(or simply carbon dioxide + water + sunlight → glucose + dioxygen)

Toxicity

MAIN SYMPTOMS OF OXYGEN TOXICITY

Oxygen gas (O₂) can be toxic at elevated partial pressures, leading to convulsions and other health problems. Oxygen toxicity usually begins to occur at partial pressures more than 50 kilopascals (kPa), equal to about 50% oxygen composition at standard pressure or 2.5 times the normal sea-level O₂ partial pressure of about 21 kPa. This is not a problem except for patients on mechanical ventilators, since gas supplied through oxygen masks in medical applications is typically composed of only 30%–50% O₂ by volume (about 30 kPa at standard pressure). (although this figure also is subject to wide variation, depending on type of mask).

At one time, premature babies were placed in incubators containing O₂-rich air, but this practice was discontinued after some babies were blinded by the oxygen content being too high.

Breathing pure O₂ in space applications, such as in some modern space suits, or in early spacecraft such as Apollo, causes no damage due to the low total pressures used. In the case of spacesuits, the O₂ partial pressure in the breathing gas is, in general, about 30 kPa (1.4 times normal), and the resulting O₂ partial pressure in the astronaut’s arterial blood is only marginally more than normal sea-level O₂ partial pressure (for more information on this, see space suit and arterial blood gas).

Oxygen toxicity to the lungs and central nervous system can also occur in deep scuba diving and surface supplied diving. Prolonged breathing of an air mixture with an O₂ partial pressure more than 60 kPa can eventually lead to permanent pulmonary fibrosis. Exposure to a O₂ partial pressures greater than 160 kPa (about 1.6 atm) may lead to convulsions (normally fatal for divers). Acute oxygen toxicity (causing seizures, its most feared effect for divers) can occur by breathing an air mixture with 21% O₂ at 66 m or more of depth; the same thing can occur by breathing 100% O₂ at only 6 m.

Combustion and other hazards

Highly concentrated sources of oxygen promote rapid combustion. Fire and explosion hazards exist when concentrated oxidants and fuels are brought into close proximity; however, an ignition event, such as heat or a spark, is needed to trigger combustion. Oxygen itself is not the fuel, but the oxidant. Combustion hazards also apply to compounds of oxygen with a high oxidative potential, such as peroxides, chlorates, nitrates, perchlorates, and dichromates because they can donate oxygen to a fire.

Concentrated O₂ will allow combustion to proceed rapidly and energetically. Steel pipes and storage vessels used to store and transmit both gaseous and liquid oxygen will act as a fuel; and therefore the design and manufacture of O₂ systems requires special training to ensure that ignition sources are minimized. The fire that killed the Apollo 1 crew in a launch pad test spread so rapidly because the capsule was pressurized with pure O₂ but at slightly more than atmospheric pressure, instead of the 1⁄3 normal pressure that would be used in a mission.

Liquid oxygen spills, if allowed to soak into organic matter, such as wood, petrochemicals, and asphalt can cause these materials to detonate unpredictably on subsequent mechanical impact. As with other cryogenic liquids, on contact with the human body it can cause frostbites to the skin and the eyes.

SULPHUR

Large deposits of free sulphur occur in America, Sicily and Japan. Combined sulphur occurs as sulphides, for example galena, PbS, zinc blende, ZnS, and iron pyrites, FeS₂, and as sulphates, notably as gypsum or anhydrite, CaSO₄. In America, the sulphur deposits (mostly in Louisiana and Texas) are dome-shaped layers about 30cm thick, between limestone above and anhydrite below. From these, the sulphur is extracted by the Frasch process. A metal tube, about 15 cm diameter and containing two concentric inner tubes is sunk into the top of the deposit. Water, superheated to 450 K, is forced under pressure down the outer tube, and enters the sulphur layer through perforations. The sulphur melts (m.p. 388 K) and enters the inner pipe at the bottom, up which it flows for some distance. Compressed air is forced down the innermost pipe; this emulsifies the water and molten sulphur mixture, so lowering its density, and the emulsion rises to the top of the pipe, where it is run off into vats to solidify. The purity is usually 99.8 %. Large quantities of sulphur are recovered from petroleum and natural gas. Naturally occurring hydrogen sulphide, H₂S, and that produced in the cracking and catalytic hydrogenation of petroleum is first removed by absorption and the regenerated gas is converted to sulphur by partial combustion with air, the overall reaction being,

6H₂S + 3O₂ → 6H₂O + 6S↓

3S + 6NaOH Û 2Na₂S + Na₂SO₃ + 3H₂O

SELENIUM AND TELLURIUM

Selenium and tellurium occur naturally in sulphide ores, usually as an impurity in the sulphide of a heavy metal. They are recovered from the flue dust produced when the heavy metal sulphide is roasted.

Te + 2H₂O = TeO₂ + 2H₂­

POLONIUM

This is a radioactive element. It occurs in minute traces in barium and thorium minerals, but it can be produced by irradiation of bismuth in a nuclear reactor. (The study of its chemistry presents great difficulty because of its intense a radiation).

ALLOTROPES

Oxygen, sulphur and selenium are known to exist in more than one allotropic form.

Element	Allotropes
Carbon	·         Diamond ·         – an extremely hard, transparent crystal, with the carbon atoms arranged in a tetrahedral lattice. A poor electrical conductor. An excellent thermal conductor. ·         Lonsdaleite ·         – also called hexagonal diamond. ·         Graphite ·         – a soft, black, flaky solid, a moderate electrical conductor. The C atoms are bonded in flat hexagonal lattices (graphene), which are then layered in sheets. ·         Linear acetylenic carbon ·         (Carbyne) ·         Amorphous carbon ·         Fullerenes, including ·         Buckminsterfullerene, aka “buckyballs”, such as C60. ·         Carboanotubes ·         – allotropes of carbon with a cylindrical nanostructure.
Phosphorus	·         White phosphorus ·         – crystalline solid ·         P4 ·         Red phosphorus ·         – polymeric solid ·         Scarlet phosphorus ·         Violet phosphorus ·         Black phosphorus ·         – semiconductor, analogous to graphite ·         Diphosphorus
Oxygen	·         dioxygen, O2 ·         – colorless (faint blue) ·         Ozone, O3 ·         – blue ·         Tetraoxygen, O4 ·         – ·         metastable ·         Octaoxygen, ·         O8 ·         – red
Sulfur	·         Sulfur has a large number of allotropes, second only to carbon
Selenium	·         “Red selenium,” cyclo-Se8 ·         Gray selenium, polymeric Se ·         Black selenium

 

OXYGEN

This exists in two allotropic forms, oxygen, O₂ and ozone, O₃. Oxygen is a colourless gas which condenses to a pale blue liquid, b.p. 90 K, which is markedly paramagnetic indicating the presence of unpaired electrons. Simple valence bond theory would indicate the structure :’p: q: i.e. 0=0 which accounts for the high oxygen-oxygen bond strength (bond dissociation energy, 49 kJ mol l^-1 ). but does not explain the paramagnetism. The molecular orbital theory of bonding, however, suggests not only a doubly bonded structure but also two molecular orbitals (i.e. orbitals of the complete O₂ molecule) of equal energy each containing one electron, and this satisfactorily explains both the high bond strength and paramagnetism. Oxygen, like nitrogen oxide, NO, shows little tendency to dimerise although the presence of the unstable, weakly bonded species, tetratomic oxygen O₄, has been reported as a constituent of liquid oxygen.

Ozone, O₃, is found in trace quantities in the upper atmosphere where it is believed to be formed by the photochemical dissociation of oxygen molecules by the intense ultra-violet light from the sun; absorption of this light in the process prevents it from reaching the earth where it would destroy all living matter very rapidly. Small quantities of ozone are produced when oxygen and air are subjected to an electrical discharge and it is, therefore, found in the neighbourhood of working electrical machines. Probably a small quantity of atomic oxygen is initially produced; most of this reeombines quickly to give oxygen, O₂, but a few atoms react to form ozone:

O₂ + O → O₃

O₂ + hn ® O₂*

O₂* + O₂ ® O₃ + O

The ozone molecules also decompose by reaction with atomic oxygen, so that the actual concentration of ozone is small. Ozone is formed in certain chemical reactions, including the action of fluorine on water and the thermal decomposition of iodic(VII) (periodic) acid. It is also formed when dilute (about 1 M) sulphuric acid is electrolysed at high current density; at low temperatures the oxygen evolved at the anode can contain as much as 30% ozone.

Ozone is normally produced by the use of a silent electrical discharge and a number of ozonisers have been produced. Using a potential of approximately 20000 V the ozonised oxygen produced can contain up to 10% ozone and pure ozone can be obtained by liquifaction of the mixture followed by fractional distillation (O₂, b.p. 90 K; O₃, b.p. 161 K). At room temperature ozone is a slightly blue diamagnetic gas which condenses to a deep blue liquid. It has a characteristic smell, and is toxic. Ozone is a very endothermic compound:

3O₂ = 2O₃, DG= 326 kJ/mol

8Ag + O₃ = 4Ag₂O + O₂­

K + O₃ = KO₃.

It decomposes exothermically to oxygen, a reaction which can be explosive. Even dilute ozone decomposes slowly at room tempera ture; the decomposition is catalysed by various substances (for example manganese(IV) oxide and soda-lime) and occurs more rapidly on heating.

Ozone is very much more reactive than oxygen and is a powerful oxidising agent especially in acid solution. Some examples are:

1. the conversion of black lead(II) sulphide to white lead(II) sulphate (an example of oxidation by addition of oxygen):

PbS + 4O₃ → PbSO₄↓ + 4O₂↑

2. the oxidation of iron(II) to iron(III) in acid solution:

2Fe ²⁺ + O₃ + 2H₃O⁺ → 2Fe ^{3 +} + O₂↑ + 3H₂O

The adherence of mercury to glass, i.e. tailing’ in presence of ozone, is probably due to the formation of an oxide. The oxidation of the iodide ion to iodine in solution is used to determine ozone quantiatively.

2KI + O₃ + H₂O = I₂¯ + 2KOH + O₂­

The liberated iodine is titrated with standard sodium thiosulphate(VI) solution after acidification to remove the hydroxide ions. Addition compounds called ozonides are produced when alkenes react with ozone and reductive cleavage of these compounds is used extensively in preparative and diagnostic organic chemistry. The molecular formula of ozone was determined by comparing its rate of diffusion with that of a known gas. The geometric structure of the molecule is angular with two equal O—O distances, which are slightly greater than in the oxygen molecule, and an O—O—O angle of 116°. Ozone has long been used on a small scale for water purification since it destroys viruses, and recent developments suggest that this use will increase in importance.

SULPHUR

The structures of sulphur in solid, liquid and gaseous phases are complicated. Rhombic sulphur is the solid allotrope stable at room temperature. It is yellow, readily soluble in carbon disulphide, from which it can be crystallised, and has a density of 2.06 g cm³ .

Video (http://www.youtube.com/watch?v=ZCR1HAad4ww) Formation of Sulphur

Above 369 K, the transition temperature, rhombic sulphur is no longer stable, slowly changing to monoelinic sulphur, and if rhombic sulphur is melted, allowed to partly solidify, and the remaining molten sulphur is poured off, there remain long needle-like crystals (almost colourless) of monoelinic sulphur, density 1.96 g cm^~3 . A good specimen of monoelinic sulphur can be prepared by crystallising a concentrated solution of sulphur in xylene, taking care to keep the temperature above 368 K. On standing at room temperature, monoelinic sulphur slowly changes to the rhombic form. Both these allotropes contain S₈ molecules with rings of eight sulphur atoms.

When sulphur is melted viscosity changes occur as the temperature is raised. These changes are due to the formation of long-chain polymers (in very pure sulphur, chains containing about 100 000 atoms may be formed). The polymeric nature of molten sulphur can be recognised if molten sulphur is poured in a thin stream into cold water, when a plastic rubbery mass known as plastic sulphur is obtained. This is only slightly soluble in carbon disulphide, but on standing it loses its plasticity and reverts to the soluble rhombic form. If certain substances, for example iodine or oxides of arsenic, are incorporated into the plastic sulphur, the rubbery character can be preserved. Colloidal sulphur is produced by careful addition of acid to sodium thiosulphate solution.

SELENIUM

Like sulphur, selenium exists in a number of allotropic forms. These include both crystalline, rhombic and monoelinic modifications which almost certainly contain Se₈ ring structures. Selenium, however, also has a grey allotrope which is metallic in appearance. It is stable at room temperature and is made up of extended spiral chains of selenium atoms.

TELLURIUM

Only one form of tellurium is known with certainty. It has a silvery white metallic appearance.

CHEMICAL REACTIVITY

Video (http://www.youtube.com/watch?v=FYXlBAqnhDI&feature=BF&list=ULwx2gSk2-sz4&index=1)  Chemical properties of oxygen

1. REACTIONS WITH AIR

At high temperatures oxygen reacts with the nitrogen in the air form ing small amounts of nitrogen oxide. Sulphur burns with a blue flame when heated in air to form sulphur dioxide SO₂, and a little sulphur trioxide SO₃. Selenium and tellurium also burn with a blue flame when heated in air, but form only their dioxides, SeO₂ and TeO₂.

2. REACTIONS WITH ACIDS

Oxygen

Oxygen is unaffected by aqueous acids unless they have powerful reducing properties when the acid is oxidised. For example

2HNO₂ + O₂ → 2HNO₃

4HI + O₂ → 2I2 + 2H₂O

However, hydrogen chloride gas, obtained as a by-product in chlorination reactions, is commercially converted to chlorine by passing the hydrogen chloride mixed with air over a copper catalyst at a temperature of 600-670K when the following reaction occurs:

4HCl + O₂ → 2H₂O + 2C1₂↑

However, when oxygen gas is passed into a solution where oxidation might be expected, the reaction is often too slow to be observed —there is an adverse kinetic factor.

Sulphur, selenium and tellurium.

These elements are generally unaffected by non-oxidising acids (behaviour expected for non-metallic elements) but they do react when heated with concentrated sulphuric and nitric acids, both powerful oxidising agents. Sulphur is oxidised to sulphur dioxide by hot concentrated sulphuric acid,

S + 2H₂SO₄ → 2H₂O + 3SO₂↑

and to sulphuric acid by hot concentrated nitric acid,

S + 6HNO₃ → H₂SO₄ + 6NO₂↑ 4- 2H₂O

With concentrated nitric acid, selenium and tellurium form only their + 4 oxoacids, H₂SeO₃ and H₂TeO₃ respectively, indicating a tendency for the higher oxidation states to become less stable as the atomic number of the element is increased.

Polonium

The more metallic nature of polonium is shown by the fact that it dissolves not only in concentrated nitric and sulphuric acids but also in hydrofluoric and hydrochloric acids.

3. REACTIONS WITH ALKALIS

Oxygen does not read with alkalis. Sulphur dissolves slowly in strong alkalis to give a mixture of sulphite [sulphate(IV)] and sulphide initially:

3S + 6OH^– → 2S^2- + SO₃^2-↑ + 3H₂O

However, the sulphide ion can attach to itself further atoms of sulphur to give polysulphide ions, and so these are found in solution also. Further, the sulphite ion can add on a sulphur atom to give the thiosulphate ion, S₂O₃^2~ which is also found in the reaction mixture. Selenium and tellurium react similarly, forming selenides and selenates(IV), and tellurides and tellurates(IV) respectively. Like the sulphide ion, S^{2 ~}, the ions Se^{2 ~} and Te^{2 ~} form polyanions but to a much lesser extent.

4. REACTIONS WITH ELEMENTS

Oxygen is a very reactive element and many metals and non-metals burn in it to give oxides; these reactions are dealt with under the individual group headings. Sulphur is less reactive than oxygen but still quite a reactive element and when heated it combines directly with the non-metallic elements, oxygen, hydrogen, the halogens (except iodine), carbon and phosphorus, and also with many metals to give sulphides. Selenium and tellurium are less reactive than sulphur but when heated combine directly with many metals and non-metals.

USES OF THE ELEMENTS

OXYGEN

Video http://www.youtube.com/watch?v=8cG3Y36gPxs Discover the element oxygen, its discovery, extraction, abundance, uses and place in the Periodic Table

Very large quantities of oxygen are used in steel manufacture. Other important uses include organic oxidation reactions; the oxidation of ethene CH₂=CH₂ to epoxyethane, CH₂—CH₂, is particular importance. The high temperature flames obtained when hydrocarbons burn in oxygen have many uses. The oxygen-ethyne (acetylene) flame, for example, is used in the cutting and welding of metals. All these products of complete hydrocarbon-oxygen combustion are gases and considerable expansion therefore occurs on reaction. The thrust produced is the basis of the internal combustion and many rocket engines.

SULPHUR

Sulphur is used in the manufacture of matches and fireworks, as a dust insecticide and for vulcanising rubber. Most of the world supply of sulphur, however, is used for the manufacture of sulphuric acid.

SELENIUM

Like sulphur, selenium has been used in the vulcanisation of rubber. It is also used in photoelectric cells.

COMPOUNDS OF GROUP VI ELEMENTS

HYDRIDES

All Group VI elements form a hydride H₂X. With the notable exception of water, they are all poisonous gases with very unpleasant smells.

Water, H₂O

PHYSICAL PROPERTIES

The fact that water is a liquid at room temperature with high enthalpies of fusion and vaporisation can be attributed to hydrogenbond formation. Because of the presence of the lone pairs of electrons, the molecule has a dipole moment (and the liquid a high permittivity or dielectric constant). In ice, there is an infinite three-dimensional structure in which the oxygen atom of each water molecule is surrounded by four hydrogen atoms arranged approximately tetrahedrally, two (in the molecule) attached by covalent bonds, and two from adjacent molecules by longer hydrogen bonds. As the temperature is increased hydrogen bonds begin to break and at 273 K there are insufficient to maintain the crystalline lattice and the solid melts. The liquid formed at 273 K has a quasi-crystalline structure. Between 273 K and 277 K the hydrogen bonds rearrange and the ‘crystal’ structure changes; the molecules pack more closely together so that the density increases. But above 277 K (where the density reaches a maximum value) the effect of thermal agitation of the ‘molecules’ becomes increasingly important and there is an overall expansion.

SOLVENT PROPERTIES

The high permittivity (dielectric constant) makes water a highly effective solvent for ionic crystals, since the electrostatic attractive forces between oppositely charged ions are reduced when the crystal is placed in water. Moreover, since water is not composed of randomly arranged molecules but has some degree of ‘structure’, the introduction of charged ions which attract the polar water molecules, produces a new ‘structure’, and a fraction of the water molecules become associated with the ions ‑ the process known as hydration. Energy is evolved in this process ‑ hydration energy ‑ and this assists the solution of both ionic and partly covalent substances: in the latter case hydrolysis may also occur. There are, however, many non-ionic substances for which water is a good solvent; this is because the molecules of such substances almost always contain hydrogen and oxygen atoms which can form hydrogen bonds with water molecules. Hence, for example, substances containing the ‑ OH group, for example alcohols, carboxylic acids and some carbohydrates, are soluble in water, provided that the rest of the molecule is not too large. As expected from the enthalpy of formation, water is thermally very stable but when steam is heated to above 1300 K slight dissociation to the elements does occur. Pure water is almost a nonconductor of electricity but slight ionic dissociation occurs :

2H₂O→H₃O ⁺ + OH^–. K₂₉₈ = l0^-14 mol ² l ^-2

Thus water can behave as an acid towards bases stronger than itself, for example

H₂O + NH₃ → NH₄⁺  + OH^–

and as a base to acids stronger than itself, for example

H₂O + HCI → H₃O⁺ + Cl^–

H₂O + HNO₃ → H₃O⁺ + NO₃^–

Water can also behave as both an oxidising and a reducing agent :

2H₂O + 2e → H₂↑(g) + 2OH^– (aq); E = – 0.83 V

Many metals are oxidised by water. At ordinary temperatures the more electropositive metals, for example, sodium, calcium (or their amalgams with mercury), react to give hydrogen, for example :

2Na + 2H₂O → 2NaOH + H₂↑

HATURAL WATER

Because of its excellent solvent properties naturally-occurring water is never pure. During its passage through the air, rain water absorbs carbon dioxide, small amounts of oxygen and nitrogen, and in urban areas, small quantities of other gaseous oxides such as those of sulphur. On reaching the ground it can absorb more carbon dioxide from decaying animals and vegetable material and dissolve any soluble salts. The dissolved carbon dioxide can attack limestone or other rock containing the carbonates of calcium and magnesium:

CaCO₃(s) + CO₂(aq) + H₂O → Ca²⁺ (aq) + 2HCO₃^– (aq)

Such water, and also that containing salts of multipositive metals, (usually sulphates), is said to be hard since it does not readily produce a lather with soap. Experiments with alkali metal salts can be performed to verify that the hardness is due to the presence of the multipositive metal ions and not to any of the anions present. The hardness due to calcium and magnesium hydrogencarbonates is said to be temporary since it can be removed by boiling:

Ca ²⁺ + 2HCO₃^– → CaCO₃↓ + CO₂↑ + H₂O

Whilst that due to other salts is called permanent hardness and is unaffected by boiling. Soap, essentially sodium stearate C₁₇H₃₅COONa, gives stearate and sodium ions in solution. The metal ions causing hardness form insoluble stearates which appear as scum, using up soap needed to wsolubilise” the fats and oils mainly responsible for ‘dirt’. The metal stearate precipitates—scum—may be slightly coloured, and water for washing and laundering must be softened, or a detergent used as an alternative to soap.

METHODS FOR REMOVING THE METAL IONS RESPONSIBLE FOR

HARDNESS IN WATER

Temporary hardness only may be removed:

1. By boiling, as explained above; a method too expensive for use on a large scale.

2. By addition of slaked lime, in calculated quantity for the par ticular degree of hardness (Clark’s method):

Ca(HCO₃)₂ + Ca(OH)₂ → 2CaCO₃↓ + 2H₂O

For temporary hardness due to magnesium carbonate, more lime is required, since the magnesium precipitates as the hydroxide (less soluble than the carbonate):

Mg(HCO₃)₂ + 2Ca(OH)₂ → Mg(OH)₂↓ + 2CaCO₃↓ + 2H₂O

It is thus important to determine the relative amounts of calcium and magnesium, for addition of too much lime means that calcium ions are reintroduced into the water, i.e. it becomes hard again, the hardness being permanent.

Temporary or permanent hardness may be removed:

1. By addition of sodium carbonate, for example.

Ca(HCO₃)₂ + Na₂CO₃ → CaCO₃↓+ 2NaHCO₃

CaSO₄ + Na₂CO₃ → CaCO₃↓ + Na₂SO₄

2. By the use of an ion-exchanger. An ion-exchanger can be a naturally-occurring aluminatesilicate, called a zeolite, or its synthetic equivalent known by a trade name, for example ‘Permutit’. Such exchangers have large, open three-dimensional structured anions with the negative charges at intervals, and balancing cations capable of free movement throughout the open structure.

Pure water for use in the laboratory can be obtained from tap water (hard or soft) by distillation; if water of great purity is required, distillation must be carried out in special apparatus, usually made of quartz, not glass or metal; precautions must be taken to avoid any spray getting into the distillate. Water which is sufficiently pure for most laboratory purposes can, however, be obtained by passing tap water through cation-exchangers and anion-exchangers as described above, when the water is ‘deionised’.

Hydrogen peroxide, H₂O₂

Hydrogen peroxide is probably unique in the very large number of reactions by which it is formed. Some of these may be mentioned :

1. From hydrogen and oxygen, by

(a) Burning hydrogen in oxygen and cooling the flame rapidly, by directing against ice.

(b) By exposing hydrogen and oxygen to intense ultra-violet light

(c) By exposure to certain radioactive rays, for example neutrons or electrons.

2. By passage of a glow discharge through water vapour. This can produce good yields of highly concentrated hydrogen peroxide.

3. By oxidation processes, for example oxidation of hydrocarbons, fatty acids and even some metals.

4. By electrolytic oxidation.

2H₂O₂ = 2H₂O + O₂­

In the laboratory, hydrogen peroxide can be prepared in dilute aqueous solution by adding barium peroxide to ice-cold dilute sulphuric acid:

BaO₂ + H₂SO₄ = H₂O₂ + BaSO₄¯

H₂O₂ + 2NaOH = Na₂O₂ + 2H₂O

 

Structure and dimensions of the H₂O₂ molecule in the gas phase…

and in the solid (crystalline) phase.

ACIDITY

Hydrogen peroxide in aqueous solution is a weak dibasic acid; the dissociation constant Ka for H₂O₂ → H⁺ + HO₂^– is 2.4 x 10^-12 mol l^-1 , indicating the strength of the acid (pKa = 11.6). The salts, known as peroxides (e.g. Na₂O₂) yield hydrogen peroxide on acidify cation and this reaction provides a useful method of differentiating between peroxides which contain the O—O linkage, and dioxides.

OXIDISING AND REDUCING PROPERTIES

Hydrogen peroxide has both oxidising properties (when it is converted to water) and reducing properties (when it is converted to oxygen); the half-reactions are (acid solution):

oxidation: H₂O₂ + 2H⁺ + 2e = 2H₂O, E= 1,77B

reduction: O₂ + 2H⁺ + 2e = H₂O₂, E=0,68B.

The following reactions are examples of hydrogen peroxide used as an oxidising agent:

1. Lead(II) sulphide is oxidised to lead(II) sulphate; this reaction has been used in the restoration of old pictures where the white lead pigment has become blackened by conversion to lead sulphide due to hydrogen sulphide in urban air:

PbS + 4H₂O₂ → PbSO₄↓ + 4H₂O

black                   white

2. Iron(II) is oxidised to iron(III) in acid solutions:

2Fe ²⁺  + H₂O₂ + 2H⁺ → 2Fe ³⁺ + 2H₂O

3. Iodide ions are oxidised to iodine in acid solution :

2KI + Na₂O₂ + 2H₂SO₄ = I₂ + Na₂SO₄ + K₂SO₄ + 2H₂O

As the above redox potentials indicate, only in the presence of very powerful oxidising agents does hydrogen peroxide behave as a reducing agent. For example:

1. Chlorine water is reduced to hydrochloric acid:

HC1O + H₂O₂ → H₂O + HC1 + O2↑

2. Manganate(VII) is reduced to manganese(II) ion in acid solution (usually sulphuric acid):

2KMnO₄ + 5H₂O₂ + 3H₂SO₄ = 2MnSO₄ + 5O₂ + K₂SO₄ + 8H₂O

It has been shown in reaction (3) that all the evolved oxygen comes from the hydrogen peroxide and none from the manganate(VII) or water, by using H₂ ¹⁸ O₂ and determining the  ¹⁸ O isotope in the evolved gas.

¹⁸O in the He-shell.

The reaction with acidified potassium manganate(VII) is used in the quantitative estimation of hydrogen peroxide.

TWO TESTS FOR HYDROGEN PEROXIDE

1. The oxidation of black lead(II) sulphide to the white sulphate is a very sensitive test if the black sulphide is used as a stain on filter paper.

PbS + 4H₂O₂ → PbSO₄↓ + 4H₂O

2. Addition of dilute potassium dichromate(VI) solution, K₂Cr₂O₇, to a solution of hydrogen peroxide produces chromium peroxide, CrO₅, as an unstable blue coloration; on adding a little ether and shaking this compcund transfers to the organic layer in which it is rather more stable.

4H₂O₂ + K₂Cr₂O₇ + H₂SO₄ → 2CrO₅ + K₂SO₄ + 5H₂O

USES

Pure hydrogen peroxide (or highly concentrated solution) is used together with oil as an under-water fuel. The fuel is ignited by inducing the strongly exothermic decomposition reaction by spraying it with a finely-divided solid catalyst. Mixtures of hydrazine and hydrogen peroxide are used for rocket propulsion. Hydrogen peroxide in aqueous solution has many uses, because the products from its reaction are either water or oxygen, which are generally innocuous. The chief use is bleaching of textiles, both natural and synthetic, and of wood pulp for paper. Other uses are the oxidation of dyestuffs, in photography and in the production of porous concrete and foam rubber where the evolved oxygen leavens’ the product. Hydrogen peroxide is a useful antiseptic (for example toothpaste). It is increasingly used to prepare organic peroxocompounds, which are used as catalysts in, for example, polymerisation reactions, and to prepare epoxy-compounds (where an oxygen atom adds on across a carbon-carbon double bond); these are used as plasticisers.

Hydrogen Peroxide – Internal Use

Over the last one hundred years many thousands of people have reported a number of health benefits from the internal use of Hydrogen Peroxide. Bio-oxidative therapy (the treatment of the body with extra oxygen) has been used by many medical practitioners to treat a variety of conditions from cancer to gangrene. This extra oxygen is delivered in a number of different ways including adding ozone (O₃) to the blood, the use of hyper-barric oxygen chambers (breathing oxygen (O₂) under pressure) and through either the intravenous or oral use of hydrogen peroxide.

We have learnt that many people purchase our food grade hydrogen peroxide to take orally, so thought it best to provide as much information as possible about the oral use of hydrogen peroxide. Although we do not recommend any particular method (because we’re not allowed to!!) we have set out below a number of ‘protocols’ that have been used safely by people wishing to try and rid themselves of various ailments.

Normally we would include the caveat here that you should consult your doctor before attempting any course of treatment. However, asking a doctor what he thinks about taking hydrogen peroxide will normally ellicit the usual ignorant response as no doubt, he/she knows very little if anything on the subject. So we repeat here our disclaimer, that we are all responsible for our own health, and as responsible retailers of hydrogen peroxide we shall do our utmost to provide you with as much information as we can so that you can make educated decisions about how you use our products.

Hydrogen peroxide occurs naturally in many of the bodies processes (in particular our immune systems) and in unpolluted regions of our planet hydrogen peroxide falls with the rain. If it was dangerous to have hydrogen peroxide in our bodies we would have died out as a species long ago.

As with so many things the toxicity is in the dose.

No one would consider it dangerous to take one aspirin, in fact it can save the life of someone who is having a heart attack.

Would it be considered dangerous to take 700 aspirins? Of course it would, it would kill you!!

35% hydrogen peroxide is 700 times more concentrated than the amount that is considered safe.

Always dilute your 35% food grade hydrogen peroxide in distilled water. Tap or mineral water contains dissolved minerals that will react with the extra oxygen, wasting it.

Never take hydrogen peroxide with food in your stomach. For most people this means an hour and a half either side of food. However if you feel a little nauseous after taking the H₂O₂ it could be because you haven’t left it long enough. Vitamin C, iron and fats in the stomach change hydrogen peroxide into super-oxide free radicals . This can severely damage the lining of your stomach . The same is true of iron, copper, silver or manganese so if you are supplementing these minerals

DO NOT TAKE HYDROGEN PEROXIDE INTERNALLY.

Hydrogen Peroxide is also thought to boost the immune system, so if you have recently had an organ transplant do not take hydrogen peroxide as it may cause your body to reject the transplant.

 

Hydrogen sulphide H₂S

Sulphur can be reduced directly to hydrogen sulphide by passing hydrogen through molten sulphur; the reversible reaction H₂ + S → H₂S occurs.

In the laboratory the gas is most conveniently prepared by the action of an acid on a metal sulphide, iron(II) and dilute hydrochloric acid commonly being used:

FeS + 2HC1 → FeCl₂ + H₂S↑

Na₂S + HOH Û NaHS + NaOH

SiS₂ + 3H₂O = H₂SiO₃ + 2H₂S

Al₂S₃ + 6H₂O = 2Al(OH)₃ + 3H₂S­

The gas is washed with water to remove any hydrogen chloride. Since iron(II) sulphide is a non-stoichiometric compound and always contains some free iron, the hydrogen sulphide always contains some hydrogen, liberated by the action of the iron on the acid. A sample of hydrogen sulphide of better purity can be obtained if antimony(HI) sulphide, (stibnite) Sb₂S₃, is warmed with concentrated hydrochloric acid:

Sb₂S₃ + 6HC1 → 2SbCl₃ + 3H₂S↑

2H₂S + 3O₂ → 2SO₂↑ + 2H₂O

2H₂S + O₂ → 2Si + 2H₂O

Hydrogen sulphide is a reducing agent in both acid and alkaline solution as shown by the following examples :

1. Its aqueous solution oxidises slowly on standing in air depositing sulphur.

2. It reduces the halogen elements in aqueous solution depositing sulphur :

C1₂ + H₂S → 2HC1 + S↓

3. It reduces sulphur dioxide, in aqueous solution :

2H₂S + SO₃^2- + 2H⁺ → 3H₂O + 3S↓

4. In acid solution, dichromates(VI) (and also chromates(VI) which are converted to dichromates) are reduced to chromium(HI) salts:

Cr₂O₇^2- + 8H⁺ + 3H₂S → 2Cr ³⁺ + 7H₂O

(Hence the orange colour of a dichromate is converted to the green colour of the hydra ted ehromium(III) ion, Cr ³⁺ , and sulphur is precipitated when hydrogen sulphide is passed through an acid solution.)

5. In acid solution, the manganate(VII) ion is reduced to the manganese(II) ion with decolorisation :

2KMnO₄ + 5H₂S + 3H₂SO₄ = 2MnSO₄ + 5S¯ + K₂SO₄ + 8H₂O,

H₂S + 4Br₂ + 4H₂O = H₂SO₄ + 8HBr

6. Iron(III) is reduced to iron(II) :

2Fe ³⁺  + H₂S → 2Fe ^{2 +} + 2H⁺ + S↓

 

TESTS FOR HYDROGEN SULPHIDE

1. Its smell.

2. The blackening of filter paper, moistened with a soluble lead(II) salt (e.g. the ethanoate or nitrate), by the formation of lead(II) sulphide. Hydrogen selenide (selenium hydride), H₂Se, and hydrogen telluride (tellurium hydride), H₂Te These two gases can readily be prepared by the action of acids on selenides and tellurides respectively, the reactions being analogous to that for the preparation of hydrogen sulphide. These gases have lower thermal stabilities than hydrogen sulphide as expected from their enthalpies of formation  and they are consequently more powerful reducing agents than hydrogen sulphide. Since the hydrogen-element bond energy decreases from sulphur to tellurium they are stronger acids than hydrogen sulphide in aqueous solution but are still classified as weak acids—similar change in acid strength is observed for Group VII hydrides. Many of the reactions of these acids, however, closely resemble those of hydrogen sulphide, the main difference being one of degree.

Polonium hydride, H₂Po

This has been made in trace quantities by the action of dilute hydrochloric acid on magnesium plated with polonium. As expected, it is extremely unstable and decomposes even at 100K.

OXIDES

Oxygen will unite with, i.e. oxidise (in the simplest sense), most elements other than the noble gases, forming oxides. With strongly electropositive metals, for example sodium or calcium, the oxides formed are ionic, for example sodium gives the oxide Na₂O, containing the ion O^2-.

2Mg + O₂ = 2MgO

S + O₂ = SO₂.

Mg(OH)₂ = MgO + H₂O

Ba(OH)₂ = BaO = H₂O

Hg(OH)₂ = HgO + H₂O

2AgOH = Ag₂O + H₂O

4HNO₃ = 4NO₂­ + 2H₂O + O₂­

2H₃BO₃ = B₂O₃ + 3H₂O

H₂CO₃ = H₂O + CO₂­

CaCO₃ ® CaO + CO₂­

Fe₂(SO₄)₃ ® Fe₂O₃ + 3SO₃­

With less electropositive metals or elements, for example aluminium, zinc, lead, the bond between element and oxygen may assume a partly covalent character, and the oxide becomes amphoteric, dissolving in both acids and bases, for example

A1₂O₃ + 6H⁺  + 9H₂O → 2[A1(H₂O)₆] ^{3 +}

A1₂O₃ + 6OH^– + 3H²O → 2[AI(OH)₆] ^3-

ZnO + 2HCl = ZnCl₂ + H₂O

ZnO + 2NaOH = Na₂ZnO₂ + H₂O

Notice that the acidic character is associated with the ability of aluminium to increase its covalency from three in the oxide to six in the hydroxoaluminate ion, [A1(OH)₆]^3-; the same ability to increase covalency is found in other metals whose oxides are amphoterie, for example

ZnO → [Zn(OH)₄]^{2 ~} or [Zn(OH)₆]^{4 ~}

PbO → [Pb(OH)₄] ^{2 –} or [Pb(OH)₆] ^{4 –}

HIGHER OXIDES

Variable oxidation state is also exhibited in the oxides themselves among metals in this region of electronegativity. Thus lead, for example, forms the monoxide PbO ( + 2) and the dioxide PbO₂ ( + 4) (the compound Pb₃O₄ is not a simple oxide but is sometimes called a ‘compound’ oxide). Similarly, manganese gives the oxides MnO and MnO₂.

Although the dioxides are oxidising agents, for example

PbO₂ + 4HC1 → PbCl₂ + 2H₂O + C1₂↑

PbO₂ + 2H₂SO₄ = Pb(SO₄)₂ + 2H₂O

TeO₂ + 2KOH = K₂TeO₃ + H₂O

SeO₂ + H₂O = H₂SeO₃

the oxidising power lies in the higher valency or oxidation state of the metal, not in the presence of more oxygen. The more noble metals (for example copper, mercury and silver) can form oxides, and exhibit variable oxidation state in such compounds (for example Cu₂O, CuO), but it is not easy to prepare such oxides by direct action of oxygen on the metal, and elevated temperatures are necessary. Moreover, in the case of silver and mercury, loss of oxygen from the oxide by heating is easy. The oxides are, however, basic (for example Ag₂O → Ag⁺, CuO → Cu^{2 +} in acids).

ACIDIC OXIDES

The other more electronegative elements are non-metals and form oxides which are entirely covalent and usually acidic. For example, sulphur yields the oxides SO₂ and SO₃, dissolving in bases to form the ions SO₃^2- and SO₄^2- respectively. A few non-metallic oxides are often described as neutral (for example carbon monoxide and dinitrogen oxide) because no directly related acid anion is known to exist.

SO₂ + 2NaOH = Na₂SO₄ + H₂O

P₂O₅ + 3Ca₃(PO₄)₂.

SULPHIDES

1. The alkali metal sulphides

These are ionic solids and can exist as the anhydrous salts (prepared by heating together sulphur with excess of the alkali metal) or as hydrates, for example Na₂S.9H₂O. Since hydrogen sulphide is a weak acid these salts are hydrolysed in water,

S^2- + H₂O → HS↑

Na₂S + HOH Û NaHS + NaOH

H^S- + H₂O → H₂S + OH^–

and smell of hydrogen sulphide. Aqueous solutions of these salts are conveniently prepared by the action of hydrogen sulphide on the alkali metal hydroxide ; if excess hydrogen sulphide is used the hydrogensulphide is formed, for example NaHS. Solutions of these sulphides can dissolve sulphur to give coloured polysulphides, for example Na₂S₄ containing anionic sulphur chains.

2. The sulphides of alkaline earth metals

These are similar to those of the alkali metals but are rather less soluble in water. However, calcium sulphide, for example, is not precipitated by addition of sulphide ions to a solution of a calcium salt, since in acid solution the equilibrium position

H₂S + Ca ^{2 +} → CaS↓ + 2H⁺

is very much to the left and ieutral, or alkaline solution the soluble hydrogensulphide is formed, for example

CaS + H₂O → Ca ^{2 +} + HS^– + OH^–

3. The sulphides of aluminium and chromium

These can be prepared by the direct combination of the elements. They are rapidly hydrolysed by water and the hydrolysis of solid aluminium sulphide can be used to prepare hydrogen sulphide:

Al₂S₃ + 6H₂O = 2Al(OH)₃ + 3H₂S­

SELENIDES AND TELLURIDES

These closely resemble the corresponding sulphides. The alkali metal selenides and tellurides are colourless solids, and are powerful reducing agents in aqueous solution, being oxidised by air to the elements selenium and tellurium respectively (cf. the reducing power of the hydrides).

OXIDES AND OXO-ACIDS AND THEIR SALTS

The elements, sulphur, selenium and tellurium form both di- and tri-oxides. The dioxides reflect the increasing metallic character of the elements. At room temperature, sulphur dioxide is a gas, boiling point 263 K, selenium dioxide is a volatile solid which sublimes at 588 K under 1 atmosphere pressure, and tellurium dioxide is a colourless, apparently ionic, crystalline dimorphic solid.

Sulphur

SULPHUR DIOXIDE, SO₂

 

Sulphur dioxide is formed together with a little of the trioxide when sulphur burns in air:

S + O₂ → SO₂↑

2S + 3O₂ → 2SO₃↑

It can be prepared by the reduction of hot concentrated sulphuric acid by a metal. Copper is used since it does not also liberate hydrogen from the acid:

Cu + 2H₂SO₄ = CuSO₄ + SO₂ + 2H₂O

The equation is not strictly representative of the reaction for the acid is reduced further and a black deposit consisting of copper(I) and copper(II) sulphides is also produced. Sulphur dioxide is also produced by the action of an acid (usually concentrated sulphuric since it is involatile) on a sulphite or hydrogensulphite, for example

2HSO₃^– + H₂SO₄ → SO₄²– + 2H₂O + 2SO₂↑

 Properties of sulphur dioxide

Sulphur dioxide is oxidised by chlorine in the presence of charcoal or camphor to give sulphur dichloride dioxide (sulphuryl chloride), SO₂C1₂;

SO₂ + C1₂ → SO₂C1₂

CuCl₂ + 2SO₃ = CuSO₄ + SO₂Cl₂

2SO₂ + O₂ Û 2SO₃,

Na₂S₂O₇ = Na₂SO₄ + SO₃.

Dioxides and peroxides oxidise it to yield sulphates:

PbO₂ + SO₂ → PbSO₄

Na₂O₂ + SO₂ → Na₂SO₄

Sulphur dioxide is an acidic oxide and dissolves readily in water, and in alkalis with which it forms salts:

NaOH + SO₂ = NaHSO₃;

sodium hydrogensuiphite

2NaOH + SO₂ = Na₂SO₃ + H₂O

sodium sulphite

Although sulphur dioxide, as a gas, is a reducing agent in the sense that it unites with oxygen, free or combined (for example in dioxides or peroxides) most of its reducing reactions in aqueous solution are better regarded as reactions of ‘sulphurous acid’ (in acid solution), or the sulphite ion (in alkaline solution).

Biochemical and biomedical roles

Sulfur dioxide is toxic in large amounts. It or its conjugate base bisulfite is produced biologically as an intermediate in both sulfate-reducing organisms and in sulfur oxidizing bacteria as well. The role of sulfur dioxide in mammalian biology is not yet well understood. Sulfur dioxide blocks nerve signals from the pulmonary stretch receptors (PSRs) and abolishes the Hering–Breuer inflation reflex.

Safety. Inhalation

Inhaling sulfur dioxide is associated with increased respiratory symptoms and disease, difficulty in breathing, and premature death. In 2008, the American Conference of Governmental Industrial Hygienists reduced the short-term exposure limit from 5 ppm to 0.25 ppm. The OSHA PEL is currently set at 5 ppm (13 mg/m3) time weighted average. NIOSH has set the IDLH at 100 ppm. A 2011 systematic review concluded that exposure to sulfur dioxide is associated with preterm birth.

Ingestion

In the United States, the Center for Science in the Public Interest lists the two food preservatives, sulfur dioxide and sodium bisulfite, as being safe for human consumption except for certain individuals who may be sensitive to them, especially in large amounts. Symptoms of sensitivity to sulfiting agents, including sulfur dioxide, manifest as potentially life-threatening trouble breathing within minutes of ingestion.

SULPHUROUS ACID

The solution obtained when sulphur dioxide dissolves in water has long been thought to contain unionised sulphurous acid, H₂SO₃, but more probably contains hydrated sulphur dioxide (cf. NH₃ solution). The solution behaves as a dibasic acid, i.e.

SO₂ + H₂O Û H₂SO₃ Û H⁺ + HSO₃^– Û 2H⁺ + SO₃^2-.

(К_а1 = 2 × 10^-2, К_а2 = 6 × 10^-8)

The sulphite ion, SO₃^2-, has a pyramidal structure and the short S — O bond length suggests the presence of double bonding, i.e.

Uses

The reducing action of sulphurous acid and sulphites in solution leads to their use as mild bleaching agents (for example magenta and some natural dyes, such as indigo, and the yellow dye in wool and straw are bleached). They are also used as a preservative for fruit and other foodstuffs for this reason. Other uses are to remove chlorine from fabrics after bleaching and in photography.

SULPHITES AND HYDROGENSULPHITES

When a saturated solution of sulphur dioxide is titrated against approximately 2 M sodium hydroxide solution the following pH curve is obtained Evaporation and crystallisation of the sodium sulphite solution gives crystals of the heptahydrate Na₂SO₃.7H₂O. However, on evaporation of the hydrogensulphite solution, the solid obtained is chiefly sodium pentaoxodisulphate(IV) ( s metabisulphite’) Na₂S₂O₅, and contains little if any of the hydrogensulphite. However, the hydrogen sulphite ion is obtained when the solid redissolves in water:

Properties

The redox properties have already been considered. A number of reactions of soluble (alkali metal) sulphites are noteworthy:

1. On boiling a solution of a sulphite with sulphur a thiosulphate(VI) is formed, and sulphur ‘dissolves’:

SO₃^2- + S → S₂O₃^2-  (e.g. Na₂S₂O₃)

Sodium thiosulphate is an important reducing agent used in volumetric analysis for the estimation of iodine:

I₂ + 2Na₂S₂O₃ → Na₂S₄O₆ + 2NaI

tetrathionate

It is used as the Tixer’ in photography under the name ‘hypo’.

2. Addition of barium chloride precipitates white barium sulphite:

Ba²⁺ + SO₃^2- →BaSO₃↓

SULPHUR TRIOXIDE

Sulphur trioxide was first prepared by heating iron(III) sulphate :

Fe₂(SO₄)₃ → Fe₂O₃ + 3SO₃

It is also obtained by the dehydration of concentrated sulphuric acid with phosphorus(V) oxide:

2H₂SO₄ + P₄O₁₀ → 4HPO₃ + 2SO₃↑

and the thermal decomposition of iron(II) sulphate :

2FeSO₄ → Fe₂O₃ + SO₂↑ + SO₃↑

iron(II) sulphate     iron(III) oxide

In the laboratory it is commonly prepared by the reaction between sulphur dioxide and oxygen at high temperature in the presence of a platinum catalyst :

2SO₂ + O₂ → 2SO₃↑

Sulphur trioxide can be collected as a white solid in a receiver surrounded by a freezing mixture of ice and salt.

Properties

In the vapour state, sulphur trioxide has the formula SO₃. The molecule is planar with all the S — O bonds short and of equal length. Solid sulphur trioxide reacts explosively with liquid water :

SO₃ + H₂O → H₂SO₄ ; ∆H = – 88 kJ mol^-1

and it fumes strongly in moist air. The gas sulphur trioxide does not readily dissolve in water, but it reacts with concentrated sulphuric acid, thus :

H₂SO₄ + SO₃ → H₂S₂O₇

H₂S₂O₇ + SO₃ → H₂S₃O₁₀ and so on.

Sulphur trioxide unites exothermically with basic oxides to give sulphates, for example

CaO + SO₃ → CaSO₄

Sulphur trioxide is used on an industrial scale for sulphonating organic compounds.

SULPHURIC ACID, H₂SO₄

Sulfuric acid

Space-filling model

Ball-and-stick model

Sulphuric acid is probably the most important chemical substance not found naturally. Its manufacture is therefore important ; the total world production is about 25 000 000 tons a year.

Grades of sulfuric acid

Although nearly 99% sulfuric acid can be made, the subsequent loss of SO₃ at the boiling point brings the concentration to 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:

“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.

Sulfuric acid reacts with its anhydride, SO₃, to form H₂S₂O₇, called pyrosulfuric acid, fuming sulfuric acid, Disulfuric acid or oleum or, less commonly,4 Nordhausen acid. 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.

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 Pharmacopeia (USP) grade are used for making pharmaceuticals and dyestuffs. Analytical grades are also available.

S + 2H₂SO₄ = 3SO₂­ + 2H₂O

Na₂S₂O₃ + H₂SO₄ = Na₂SO₄ + H₂S₂O₃

H₂S₂O₃ = H₂SO₃ + S⁰

H₂S₂O₈ + 2H₂O = 2H₂SO₄ + H₂O₂

 

 

Drops of concentrated sulfuric acid rapidly dehydrate a piece of cotton towel

 Homemade Sulfuric Acid Materials

Actually, this method starts with diluted sulfuric acid, which you boil to make concentrated sulfuric acid. This is the safest and easiest method of making sulfuric acid at home.

·  car battery acid

·  glass container

·  outdoor source of heat, like a grill

Battery acid, which may be purchased at an automotive supply store, is approximately 35% sulfuric acid. In many cases, this will be strong enough for your activities, but if you need concentrated sulfuric acid, all you need to do is remove the water. The resulting acid will not be as pure as reagent-grade sulfuric acid, so keep this in mind.

Safest Method

If you aren’t in a hurry, you can concentrate sulfuric acid by allowing the water to evaporate naturally. This takes several days.

1.                     Place an open container of sulfuric acid someplace with good circulation, safe from the possibility of a spill.

2.                     Loosely cover the container to minimize contamination with dust and other particulates.

3.                     Wait. The water will evaporate out of the solution, eventually leaving you with concentrated sulfuric acid. Note that sulfuric acid is highly hygroscopic, so it will retain a certain amount of water. You would need to heat the liquid to drive off the remaining water.

Quickest Method

The fastest method to concentrate sulfuric acid is to boil the water out of the acid. This is fast, but requires extreme care. You’ll want to do this outdoors so that you won’t be exposed to acid fumes, using borosilicate glass (e.g., Pyrex or Kimax). There is always a risk of shattering a glass container, no matter what you are heating, so you need to be prepared for that possibility. Do not leave this project unattended!

1.                     Heat the battery acid in a borosilicate glass pan.

2.                     When the liquid level stops dropping, you will have concentrated the acid as much as you can. At this point, the steam will be replaced by white vapor, too. Be careful to avoid inhaling the fumes.

3.                     Allow the liquid to cool before transferring it to another container.

4.                     Seal the container to prevent water from the air from getting into the acid. If the container is left open for too long, the sulfuric acid will become diluted.

Safety Notes

·  It’s advisable to keep baking soda (sodium bicarbonate) or another base on hand. If you spill some acid, you can quickly neutralize it by reacting it with the baking soda. Simply sprinkle baking soda on the spill.

·  Be careful to avoid contact with the sulfuric acid! Sulfuric acid is one of the strong acids. It is extremely corrosive and will react vigorously and unpleasantly with skin, mucous membranes, clothing and just about anything else it touches. Do not breathe the vapors, do not touch the acid and do not spill it. Tie long hair back, wear goggles and gloves and cover exposed skin.

·  Don’t use metal pans or utensils. Sulfuric acid reacts with metal. Also, it will attack some types of plastic. Glass is a good choice.

·  Sulfuric acid reacts with water in an exothermic reaction, but dilution with water is the best way to deal with an acid spill. Be sure to have copious amounts of water available, just in case something goes wrong. You can flood a small amount of acid with water. Caution: Sulfuric acid will splash when mixed with water! If you are going to work with this acid, know and respect its properties.

Notes About Battery Acid

Battery acid is about 35% sulfuric acid. You can purchase it at an automotive supply store. It may not be on the shelf, so ask for it. Battery acid may be sold in 5 gallon boxes, with the acid in a heavy duty plastic bag and a plastic tube to dispense the liquid. The box is heavy; it would be disastrous to drop it. Therefore, it’s a good idea to know what to expect.

It’s practical to dispense a working volume of acid, rather than try to deal with the entire container. Although the acid may come in a plastic container, it’s best to store this acid in a glass bottle. Sulfuric acid reacts with some types of plastic and may corrode a plastic container. I used a glass wine bottle that had a plastic screw-top cap. Whatever container you use, be certain to label it as sulfuric acid and poison and store it somewhere that children and pets can’t get to it. Also, don’t store acid with ammonia because the two chemicals mix to release toxic fumes.

 

Manufacture

The different methods of manufacturing sulphuric acid are essentially the same in principle and consist of three distinct processes :

I . Production of sulphur dioxide.

2. Conversion of sulphur dioxide to sulphur trioxide.

3. Conversion of sulphur trioxide to sulphuric acid.

1. Sulphur dioxide is obtained in the following three ways :

(a) By burning elemental sulphur (imported) :

S + O₂ → SO₂↑

(b) As a by-product of the roasting process in the extraction of certain metals from their sulphide ores, for example

2ZnS + 3O₂ → 2ZnO + 2SO₂↑

2. The combination of sulphur dioxide and oxygen to form the trioxide is slow and does not proceed to completion :

2SO₂ + O₂ → 2SO₃ : ∆H = – 94 kJ mol^-1

SO₂ + H₂O → H₂SO₃

3. The conversion of sulphur trioxide to sulphuric acid arises as a separate reaction only in the Contact process. Sulphur trioxide is not very soluble in water but dissolves readily in concentrated sulphuric acid. The sulphur trioxide from the Contact chamber is passed into concentrated sulphuric acid, to which water is added at the required rate:

SO₃ + H₂SO₄ → H₂S₂O₇

H₂S₂O₇ + H₂O → 2H₂SO₄

The 94 % acid from the sulphur dioxide drying towers (above) is used here and its strength brought up to 98 %. This is “concentrated1 sulphuric acid”. Stronger acid up to 106% may also be made. This concentration is suitable for sulphonating in, for example, the detergent industry.

Properties

Pure sulphuric acid is a colourless, viscous and rather heavy liquid (density 1.84g cm^-3). On heating, it decomposes near its boiling point, forming sulphur trioxide and a constant boiling (603 K) mixture of water and sulphuric acid containing 98 % of the latter. This is ‘concentrated’ sulphuric acid, which is usually used. Further heating gives complete dissociation into water and sulphur trioxide (K₁ = 1×10³, K₂ = 1,2×10^-2).

Oxidising properties

Concentrated sulphuric acid is an oxidising agent, particularly when hot, but the oxidising power of sulphuric acid decreases rapidly with dilution. The hot concentrated acid will oxidise nonmetals, for example carbon, sulphur and phosphorous to give, respectively, carbon dioxide, sulphur dioxide and phosphoric(V) acid. It also oxidises many metals to give their sulphates; cast iron, however, is not affected. The mechanisms of these reactions are complex and the acid gives a number of reduction products. Hot concentrated sulphuric acid is a useful reagent for differentiating between chloride, bromide and iodide salts, since it is able to oxidise (a) iodide, giving iodine (purple) and the reduction products, hydrogen sulphide, sulphur and sulphur dioxide together with a little hydrogen iodide; (b) bromide, giving bromine (red-brown) and the reduction product sulphur dioxide together with hydrogen bromide. It is unable to oxidise the chloride ion and steamy fumes of hydrogen chloride are evolved.

Acidic properties

Concentrated sulphuric acid displaces more volatile acids from their salts, for example hydrogen chloride from chlorides (see above) and nitric acid from nitrates. The dilute acid is a good conductor of electricity. It behaves as a strong dibasic acid :

H₂SO₄ + H₂O → H₃O⁺  + HSO₄ : Ka = 40 mol l^-1  at 298 K

HSO₄ + H₂O → H₃O⁺  +SO₄^2- : Ka=1.0 10^-2  mol l^-1  at 289K

the value of Ka for the first dissociation indicating that this reaction goes virtually to completion in dilute solution. The acid exhibits all the properties of the hydrogen ion, i.e. neutralising bases, giving hydrogen with many metals and so on. Dilute sulphuric acid attacks iron, but lead very soon becomes resistant due to the formation of a superficial layer of insoluble lead sulphate.

Cu + 2H₂SO_{4 conc} = CuSO₄ + SO₂­ + 2H₂O

Zn + 2H₂SO_{4 conc} = ZnSO₄ + SO₂­ + 2H₂O

3Zn + 4H₂SO_{4 conc} = 3ZnSO₄ + S¯ + 4H₂O

4Zn + 5H₂SO_{4 conc} = 4ZnSO₄ + H₂S­ + 4H₂O

Reaction with water and dehydrating property

Drops of concentrated sulfuric acid dehydrate a piece of cotton towel rapidly.

Because the hydration reaction of sulfuric acid is highly exothermic, dilution should always be performed by adding the acid to the water rather than the water to the acid. Because the reaction is in an equilibrium that favors the rapid protonation of water, addition of acid to the water ensures that the acid is the limiting reagent. This reaction is best thought of as the formation of hydronium ions:

H₂SO₄ + H₂O → H₃O⁺ + HSO₄⁻    K₁ = 2.4×10⁶   (strong acid)

HSO−4 + H₂O → H₃O⁺ + SO2−4     K₂ = 1.0×10⁻²

HSO−4 is the bisulfate anion and SO2−4 is the sulfate anion. K₁ and K₂ are the acid dissociation constants.

Because the hydration of sulfuric acid is thermodynamically favorable and the affinity of it for water is sufficiently strong, sulfuric acid is an excellent dehydrating agent. Concentrated sulfuric acid has a very powerful dehydrating property, removing water (H₂ O) from other compounds including sugar and other carbohydrates and producing carbon, heat, steam, and a more dilute acid containing increased amounts of hydronium and bisulfate ions.

In laboratory, this is often demonstrated by mixing table sugar (sucrose) into sulfuric acid. The sugar changes from white to dark brown and then to black as carbon is formed. A rigid column of black, porous carbon will emerge as well. The carbon will smell strongly of caramel due to the heat generated.

C₁₂H₂₂O₁₁ (white sucrose) + sulfuric acid → 12 C_{(black graphitic foam)} + 11 H₂O (steam) + sulfuric acid/water mixture

Similarly, mixing starch into concentrated sulfuric acid will give elemental carbon and water as absorbed by the sulfuric acid (which becomes slightly diluted). The effect of this can be seen when concentrated sulfuric acid is spilled on paper which is composed of cellulose; the cellulose reacts to give a burnt appearance, the carbon appears much as soot would in a fire. Although less dramatic, the action of the acid on cotton, even in diluted form, will destroy the fabric.

(C₆H₁₀O₅)n + sulfuric acid → 6n C + 5n H₂O

The reaction with copper(II) sulfate can also demonstrate the dehydration property of sulfuric acid. The blue crystal is changed into white powder as water is removed.

CuSO₄·5H₂O (blue crystal) + sulfuric acid → CuSO₄ (white powder) + 5 H₂O

Acid-base properties

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 iitration 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.

Solid state structure of the [D₃SO₄]⁺ ion present in [D₃SO₄]⁺[SbF₆]^–, synthesized by using DF in place of HF

When allowed to react with superacids, sulfuric acid can act as a base and be protonated, forming the [H₃SO₄]⁺ ion. Salt of [H₃SO₄]⁺ have been prepared using the following reaction in liquid HF:

((CH₃)₃SiO)₂SO₂ + 3 HF + SbF₅ → [H₃SO₄]⁺[SbF₆]^– + 2 (CH₃)₃SiF

The above reaction is thermodynamically favored due to the high bond enthalpy of the Si–F bond in the side product. Protonation using simply HF/SbF₅, however, have met with failure, as pure sulfuric acid undergoes self-ionization to give [H₃O]⁺ ions, which prevents the conversion of H₂SO₄ to [H₃SO₄]⁺ by the HF/SbF₅ system:

2 H₂SO₄ [H₃O]⁺ + [HS₂O₇]^–

Reactions with metals and strong oxidizing property

Dilute sulfuric acid reacts with metals via a single displacement reaction as with other typical acids, producing hydrogen gas and salts (the metal sulfate). It attacks reactive metals (metals at positions above copper in the reactivity series) such as iron, aluminium, zinc, manganese, magnesium and nickel.

Fe (s) + H₂SO₄ (aq) → H₂ (g) + FeSO₄ (aq)

However, concentrated sulfuric acid is a strong oxidizing agent and does not react with metals in the same way as other typical acids. Sulfur dioxide, water and SO₄^2- ions are evolved instead of the hydrogen and salts.

2 H₂SO₄ + 2 e^– → SO₂ + 2 H₂O + SO₄^2-

It can oxidize non-active metals such as tin and copper, depending upon the temperature of it like the nitric acid.

Cu + 2 H₂SO₄ → SO₂ + 2 H₂O + SO₄^2- + Cu²⁺

Lead and tungsten, however, are resistant to sulfuric acid.

Reactions with non-metals

Hot concentrated sulfuric acid oxidizes non-metals such as carbon and sulfur.

C + 2 H₂SO₄ → CO₂ + 2 SO₂ + 2 H₂O

S + 2 H₂SO₄ → 3 SO₂ + 2 H₂O

Reaction with sodium chloride

It reacts with sodium chloride, and gives hydrogen chloride gas and sodium bisulfate:

NaCl + H₂SO₄ → NaHSO₄ + HCl

Electrophilic aromatic substitution

Benzene undergoes electrophilic aromatic substitution with sulfuric acid to give the corresponding sulfonic acids:

Domestic uses

Concentrated sulfuric acid is frequently the major ingredient in acidic drain cleaners which are used to remove grease, hair, tissue paper, etc. Similar to their alkaline versions, such drain openers can dissolve fats and proteins via hydrolysis. Moreover, as concentrated sulfuric acid has a strong dehydrating property, it can remove tissue paper via dehydrating process as well. Since the acid may react with water vigorously, such acidic drain openers should be added slowly into the pipe to be cleaned.

Health

Sulfuric acid and sulfonated phenolics are the primary ingredients in Debacterol, a liquid topical agent that is used in the treatment of recurrent aphthous stomatitis (canker sores/mouth ulcers) or for any procedures in the oral cavity which require controlled, focal debridement of necrotic tissues.

Safety

Acidic drain cleaners usually contain sulfuric acid at a high concentration which turns a piece of pH paper red and chars it instantly, demonstrating the strong acidic nature and dehydrating property.

 

Laboratory hazards

Drops of 98% sulfuric acid char a piece of tissue paper instantly. Carbons are left after the dehydration reaction staining the paper black.

Sulfuric acid is capable of causing very severe burns, especially when it is at high concentrations. In common with other corrosive strong acids and strong alkalis, it readily decomposes proteins and lipids through amide hydrolysis and ester hydrolysis upon contact with living tissues. In addition, it exhibits a strong dehydrating property on carbohydrates, liberating extra heat and causing secondary thermal burns. The strong oxidizing property may also extend its corrosiveness on the tissue. Because of such reasons, damage posed by sulfuric acid is potentially more severe than that caused by other comparable strong acids, such as hydrochloric acid and nitric acid. Accordingly, it rapidly attacks the cornea and can induce permanent blindness if splashed onto eyes. If ingested, it damages internal organs irreversibly and may even be fatal. Protective equipment should always be used when handling this acid. Moreover, its strong oxidizing property makes it highly corrosive to many metals.

Sulfuric acid must be stored carefully in containers made of nonreactive material (such as glass). 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”. However, even the normal laboratory “dilute” grade (approximately 1 M, 10%) will char paper by dehydration if left in contact for a sufficient time.

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 around, 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.

The main occupational of risks posed by this acid are skin contact leading to burns (see above) 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.

FUMING SULPHURIC ACID (OLEUM)

When sulphur trioxide is dissolved in concentrated sulphuric acid the pure 100% acid is first formed; then a further molecule of the trioxide adds on:

H₂SO₄ + SO₃ → H₂S₂O₇

The formation of other polysulphuric acids H₂S₃O₁₀ up to H₂O(SO₃)n, by the addition of more sulphur trioxide, have been reported.

THE SULPHATES AND HYDROGENSULPHATES

The hydrogensulphates (or bisulphates) containing the ion are only known in the solid state for the alkali metals and ammonium. Sodium hydrogensulphate is formed when sodium chloride is treated with cold concentrated sulphuric acid:

NaCl + H₂SO₄ → NaHSO₄ + HCl↑

It may also be obtained by crystallising sodium sulphate from a dilute sulphuric acid solution:

Na₂SO₄ + H₂SO₄ → 2NaHSO₄

Fe₂(SO₄)₃ → Fe₂O₃ + 3SO₃↑

2Ag₂SO₄ → 4Ag↓ + 2SO₃↑ + O₂↑

OTHER ACIDS

In addition to the simple acids discussed above, sulphur forms two peroxosulphuric acids containing the —O—O— linkage and a number of thionic acids containing more than one sulphur atom. Oxides and oxo-acids of selenium Selenium dioxide is a volatile solid obtained when selenium is burnt in air or oxygen. It is very soluble in water, forming a solution of selenic(IV) (selenious) acid H₂SeO₃, a dibasic acid forming two series of salts. Both the acid and its salts are fairly good oxidizing agents, oxidising (for example) sulphur dioxide and hydrogen iodide. Selenium trioxide, SeO₃, is a white deliquescent solid which has never been obtained completely pure. When selenic acid(Vl), H₂SeO₄, is dehydrated a mixture of selenium dioxide and trioxide is obtained and oxygen is evolved. Selenic(VI) acid H₂SeO₄ is formed when selenium trioxide is dissolved in water and is a strong dibasic acid. It is a more powerful oxidising agent than sulphuric acid and will, for example, oxidise hydrochloric acid evolving chlorine.

Oxides and oxo-acids of tellurium

Tellurium dioxide, TeO₂, is a white non-volatile solid obtained when tellurium is burnt in air. It is only slightly soluble in water but dissolves in alkalis to form salts. Tellurium trioxide, TeO₃, is an orange yellow powder made by thermal decomposition of telluric(VI) acid Te(OH)_6. It is a strong oxidising agent which will, like H₂SeO₄, oxidise hydrogen chloride to chlorine. It dissolves in hot water to give telluric(VI) acid. This is a weak acid and quite different from sulphuric and selenic acids. Two series of salts are known.

HALIDES

Sulphur, selenium and tellurium form many halides, and only a brief introduction to the subject is given here.

Fluorides

All three elements form gaseous hexafluorides by the direct combination of the elements. They all have octahedral structures Sulphur and tellurium form a chloride of formula XC1_2. Sulphur dichloride SC1₂ is a red liquid at room temperature whilst the corresponding tellurium compound is a black solid. A number of bromides and iodides are known but there are no sulphur iodides.

TESTS FOR SULPHUR

Oxidation of a sulphur compound with concentrated nitric acid yields sulphuric acid or a sulphate, which can be tested for with barium chloride. This can be used to estimate the sulphur.

Practical skills

Qualitative tests on VI grpoup elements

1. Sulfide ion S^2-

Pb²⁺ + S^2- → PbS↓ – black precipitation

2. Sulfate ion SO₄^2-

Ba²⁺ + SO₄^2-→ BaSO₄ – white precipitation. It does not dissolved in strong acids.

3. Sulfite ion SO₃^2-

Na₂SO₃ + 2HCl → SO₂↑ + 2NaCl + H₂O

Identification of SO₂, according to the smell of the singed bone or colourless the indicator paper wetted in potassium permanganat solution

SO₂ + 2 KMnO₄ + H₂O → 2MnSO₄ + K₂SO₄ +2H₂SO₄

4. Thiosulfate ion S₂O₃^2-:

a) Na₂S₂O₃ + 2HCl → H₂S₂O₃ + 2NaCl

H₂S₂O₃ → H₂O + SO₂↑ + S↓

b) 2Na₂S₂O₃ + I₂ → Na₂S₄O₆ + 2 NaI

colourless of iodine solution.

5. Hydrogen peroxide

4H₂O₂ + K₂Cr₂O₇ + H₂SO₄ →^ether 2CrO₅ + K₂SO₄ + 5H₂O

Dark blue colour of organic layer

 

References:

1. The abstract of the lecture.

2. intranet.tdmu.edu.ua/auth.php

3. Atkins P. W. Physical chemistry / P.W. Atkins. – New York, 1994. – P.299‑307.

4. Cotton F. A. Chemical Applications of Group Theory / F. A. Cotton. ‑ John Wiley & Sons : New York, 1990.

5. Girolami G. S. Synthesis and Technique in Inorganic Chemistry / G. S. Girolami, T. B. Rauchfuss, R. J. Angelici. ‑ University Science Books : Mill Valley, CA, 1999.

6. Russell J. B. General chemistry / J B. Russell. New York.1992. – P. 550‑599.

7. Lawrence D. D. Analytical chemistry / D. D. Lawrence. –New York, 1992. – P. 218–224.

8. http://www.lsbu.ac.uk/water/ionish.html

9. http://en.wikipedia.org/wiki

The following website shows the reaction of VIA group elements. It’s cool stuff! Check it out!

www.youtube.com/watch?v=q6eIqNpdvlw

www.youtube.com/watch?v=NnFzHt6l4z8

www.youtube.com/watch?v=mjkuSm__G7s

www.youtube.com/watch?v=hxrQrOlWGH0

www.youtube.com/watch?v=J904M8YUdSQ

 

Prepared by PhD Falfushynska H.