The materials to prepare students for practical lessons of inorganic chemistry

The materials to prepare students for practical lessons of inorganic chemistry

LESSON 17.

Themes: GROUP IV (CARBON, SILICON, GERMANIUM, TIN, LEAD)

 

SUMMARY OF GENERAL CHARACTERISTICS

In this group the outer quantum level has a full s level and two electrons in the corresponding p level. As the size of the atom increases the ionisation energy changes (see Table 1) and these changes are reflected in the gradual change from a typical non- metallic element, carbon, to the weakly metallic element, lead.

Hence the oxides of carbon and silicon are acidic whilst those of tin and lead are amphoteric.

Near the middle of the periodic table there is greatest variability of properties among elements of the same group. This is certainly true of group IVA, which contains carbon, a nonmetal, silicon and germanium, both semi- metals, and tin and lead, which are definitely metallic. Elemental carbon exists in two allotropic forms, diamond and graphite, whose structures are shown below.

The crystal structure of (a) diamond and (b) graphite.

In diamond there is a three-dimensional network of covalent bonds, while graphite consists of two-dimensional layers covalently bonded. Silicon, germanium, and one allotrope of tin (gray tin) also have the diamond structure—each atom is surrounded by four others arranged tetrahedrally. White tin has an unusual structure in which there are four nearest-neighbor atoms at a distance of 302 pm and two others at 318 pm. Only lead has a typical closest-packed metallic structure in which each atom is surrounded by 12 others.

Properties of the Group IVA Elements

 

Some properties of the group IVA elements are summarized in the table. As in the case of group IIIA, there is a large decrease in ionization energy and electronegativity from carbon to silicon, but little change farther down the group. This occurs for the same reason in both groups, namely, that elements farther down the group have filled d subshells. Note also that ionization energies, especially the third and fourth, are rather large. Formation of true +4 ions is very difficult, and in their +4 oxidation states all group IVA elements form predominantly covalent bonds. The +2 oxidation state, corresponding to use of the np², but not the ns², electrons for bonding, occurs for all elements. It is most important in the case of tin and especially lead, the latter having an inert pair like that of thallium. In the +4 oxidation state lead is a rather strong oxidizing agent, gaining two electrons (6s²) and being reduced to the +2 state.

Chemical Reactions and Compounds

Carbon’s ability to form strong bonds with other carbon atoms and the tremendous variety of organic compounds have already been discussed extensively in the section on organic compounds. You may want to review the subsections dealing with hydrocarbons and the other organic compounds. The most important inorganic carbon compounds are carbon monoxide and carbon dioxide. Both are produced by combustion of any fuel containing carbon:

C + ½O₂ → CO      (1)

CO + ½O₂ → CO₂      (2)

The triple bond in is the strongest chemical bond known, and contains two double bonds, and so both molecules are quite stable. Equations (1) and (2) occur stepwise when a fuel is burned, and the strong bond makes Eq. (2) slow unless the temperature is rather high. If there is insufficient O₂ or if the products of combustion are cooled rapidly, significant quantities of CO can be produced. This is precisely what happens in an automobile engine, and the exhaust contains between 3 and 4% CO unless pollution controls have been installed.

CO is about 200 times better than O₂ at bonding to hemoglobin, the protein which transports O₂ through the bloodstream from the lungs to the tissues. Consequently a small concentration of CO in the air you breathe can inhibit transport of O₂ to the brain, causing drowsiness, loss of consciousness, and death. (After a few minutes of breathing undiluted auto exhaust, more than half your hemoglobin will be incapable of transporting O₂, and you will faint.) CO in automobile exhaust can be used to put animals to sleep. Because CO is colorless and odorless, your senses cannot detect it, and people must constantly be cautioned not to run cars in garages or other enclosed spaces. With the large number of cars and the great number of miles driven, it is important to limit CO emissions from automobiles. In the early 1970s new EPA standards led to the adoption of catalytic converters, which convert the poisonous CO into CO₂. Implementation and increasing effectiveness of these converters has caused CO levels to drop since the 1970s, despite the increase in automobiles on the road.

Like the organic compounds of carbon, the oxygen compounds of silicon which make up most of the earth’s crust have already been described. These substances illustrate a major contrast between the chemistry of carbon and silicon. The latter element does form a few compounds, called silanes, which are analogous to the alkanes, but the Si—Si bonds in silanes are much weaker than Si—O bonds. Consequently the silanes combine readily with oxygen from air, forming Si—O—Si linkages. Unlike the alkanes, which must be ignited with a spark or a match before they will burn, silanes catch fire of their own accord in air:

2Si₄H₁₀ + 13O₂ → 4SiO₂ + 5H₂O

Another important group of silicon compounds is the silicones. These polymeric substances contain Si—O—Si linkages and may be thought of as derived from silicon dioxide, SiO₂. To make silicones, one must first reduce silicon dioxide to silicon. This can be done using carbon as the reducing agent in a high-temperature furnace:

SiO₂(s) + 2C(s) Si(l) + 2CO(g)

The silicon is then reacted with chloromethane:

Si(s) + 2CH₃Cl(g) (CH₃)₂SiCl₂(g)

The dichlorodimethylsilane obtained in this reaction polymerizes when treated with water:

n(CH₃)₂SiCl₂ + nH₂O → + 2nHCl

The silicone polymer consists of a strongly bonded —Si—O—Si—O—Si—O chain, called a siloxane chain, with two methyl groups (or other organic groups) on each silicon atom. The strong backbone of a silicone polymer makes it stable to heat and difficult to decompose. Silicone oils make good lubricants and heat-transfer fluids, and rubber made from silicone remains flexible at low temperatures.

Besides the metals themselves, some tin and lead compounds are of commercial importance. Tin(II) fluoride (stannous fluoride), SnF₂, is added to some toothpastes to inhibit dental caries. Tooth decay involves dissolving of dental enamel [mainly Ca₁₀(PO₄)₆(OH)₂] in acids synthesized by bacteria in the mouth. Fluoride ions from SnF₂ inhibit decay by transforming tooth surfaces into Ca₁₀(PO₄)₆F₂, which is less soluble in acid:

Ca₁₀(PO₄)₆(OH)₂ + SnF₂ → Ca₁₀(PO₄)₆F₂ + Sn(OH)₂

Since F^– is a weaker base than OH^–, the F^– compound has less tendency to react with acids. Note that when tin or lead are in the +2 oxidation state and are combined with a highly electronegative element like fluorine, the compounds formed are rather ionic.

Lead is found in two main commercial applications. One, the lead-acid storage battery is used to start cars and power golf carts. The other is the lead found in automobile fuel. In the +4 oxidation state lead forms primarily covalent compounds and bonds strongly to carbon. The compound tetraethyllead may be synthesized by reacting with a sodium-lead alloy:

4NaPb + 4CH₃CH₂Cl → (CH₃CH₂)Pb + 4Pb + 4NaCl

Sodium dissolved in the lead makes the latter more reactive. Tetraethyl-lead prevents gasoline from igniting too soon or burning unevenly in an automobile engine, circumstances which cause the engine to “knock” or “ping.” This is where the term leaded gasoline comes from. A major problem connected to using tetraethyl-lead is the introduction of lead into the atmosphere. Lead is toxic, and thus use of TEL as an antiknock agent has been phased out in favor of other agents less dangerous to public health.

Most Glasses are Silicates

Quartz is a crystalline material with the composition SiO₂(s) and the crystalline structure shown in Figure. When crystalline quartz is melted and then cooled quickly to prevent the formation of crystals, a disordered three-dimensional array of polymeric chains, sheets, and other three-dimensional clusters forms. The resulting material is called quartz glass. All glass consists of such a random array.

Glass manufacturing is a 25-billion-dollar-per-year industry in the United States. The major component in glass is almost pure quartz sand. Among the other components of glass, soda, Na₂O(s), comes from soda ash, Na₂CO₃(s); lime, CaO(s), comes from limestone, CaCO₃(s); and aluminum oxide, Al₂O₃ (s), comes from feldspars, which have the general formula M₂O∙Al₂O₃∙6  SiO₂(s), where M is K or Na. All the components of glass are fairly inexpensive chemicals.

A wide variety of glass properties can be produced by varying the glass composition. For example, partial replacement of CaO(s) and Na₂O(s) by B₂O₃(s) gives a glass called borosilicate glass that expands less on heating (or contracts less on cooling) and is thus used in making glass utensils meant to be heated.

Colored glass is made by adding a few percent of a colored transition metal oxide, such as CoO(s) to make blue “cobalt” glass, Cr₂O₃(s) to make orange glass, and Au₂O₃ (s) to make ruby glass. Because it is made with gold, ruby glass is more expensive than other colored glass. Lead glass, which contains PbO(s), has attractive optical properties and is used to make decorative, cut-glass articles.

Photochromic eyeglasses have a small amount of added silver chloride dispersed throughout and trapped in the glass. When sunlight strikes this type of glass, the tiny AgCl(s) grains decompose into opaque clusters of silver atoms and chlorine atoms

according to . The chlorine atoms are trapped in the crystal lattice, so the silver and chlorine atoms can recombine in the dark to form silver chloride, a reaction that causes the glass to become clear.

Porcelain has a much higher percentage of Al₂O₃(s) than glass and as a result is a heterogeneous substance. Porcelain is stronger than glass because of this heterogeneity and is also more chemically resistant than glass. Earthenware is similar in composition to porcelain but is more porous because it is fired at a lower temperature

OXIDATION STATES

Gain of electrons

Only the carbon atom can gain four electrons; this only happens when it is combined with extremely electropositive elements and this state may be regarded as exceptional. Bonding in carbides is almost invariably predominantly covalent.

Loss of electrons

The oxidation state -1-4 involves both the s and p electrons. The oxidation state +2, involving only the p electrons, becomes increasingly important with increasing atomic size, and the two s.

Table 1.

SELECTED ROPERTIES OF THE ELEMENTS

 

Electrons are retained as an inert pair. There are no stable compounds of carbon and silicon in this + 2 oxidation state; it is uncommon (and strongly reducing) in germanium, less strongly reducing and commonly found in tin and it is the most stable oxidation state for lead. Only tin and lead are capable of forming + 2 ions which occur both in the solid state and in solution, where the ions are stabilised by solvation.

The oxidation state +4 is predominantly covalent and the stability of compounds with this oxidation state generally decreases with increasing atomic size. It is the most stable oxidation state for silicon, germanium and tin, but for lead the oxidation state +4 is found to be less stable than oxidation state + 2 and hence lead(IV) compounds have oxidising properties. The concept of oxidation states is best applied only to germanium, tin and lead, for the chemistry of carbon and silicon is almost wholly defined in terms of covalency with the carbon and silicon atoms sharing all their four outer quantum level electrons. These are often tetrahedrally arranged around the central atom. There are compounds of carbon in which the valency appears to be less than four but, with the exception of carbon monoxide, double or triple bonds are formed in such a way as to make the covalency of carbon always four. The exceptional structure of carbon monoxide makes the molecule an electron donor. Silicon does not form equivalent double- or triple-bonded molecules.

Silicon, germanium, tin and lead can make use of unfilled d-orbitals to expand their covalency beyond four and each of these elements is able (but only with a few ligands) to increase its covalency to six. Hence silicon in oxidation state +4 forms the octahedral hexafluorosilicate complex ion [SiF₆]^2- (but not [SiCl]^2– ).

Tin and lead in oxidation state 4-4 form the hexahydroxo complex ions, hexahydroxostannate(IV), [Sn(OH)₆]^2- and hexahydroxoplumbate(IV) respectively when excess alkali is added to an aqueous solution containing hydrated tin(IV) and lead(IV) ions.

Carbon, however, is unable to form similar complexes since the energy required to promote electrons to the next higher energy level, the 3s, is too great (or since carbon has no available d orbitals in its outer quantum level).

OCCURRENCE AND EXTRACTION OF THE ELEMENTS

Pure carbon occurs naturally in two modifications, diamond and graphite. In both these forms the carbon atoms are linked by covalent bonds to give giant molecules.

Diamonds

Diamonds are found in South Africa, India, South America and Russia. The largest ever found was the Cullinan diamond which weighed about 600 g. There are four possible crystalline arrangements all of which are found to occur naturally. The interatomic bonds are very strong. This high bond strength is reflected in the great hardness and high melting point of diamond. Diamond also has a high refractive index and is the densest form of carbon. The many uses of diamond are largely dependent on its great hardness, for example for cutting and grinding.

Very small synthetic diamonds have been made industrially by subjecting graphite to pressures in the range 5.5-6.9 GNm^–2 , at temperatures between 1500 and 2700 K. The diamonds produced are very small but competitive with natural diamonds for use in industrial cutting and grinding wheels.

Graphite

Graphite occurs naturally in Ceylon, Germany and the USA. It was formerly mined in Cumberland. Its name (Greek, grapho – I write) indicates its use in “lead” pencils. The structure of graphite: each carbon atom is joined to three others by six bonds, the arrangement being trigonal planar. The remaining electron on each carbon atom is in a p orbital. A sideways overlap of these orbitals occurs to give a delocalisedbond. It is this second bond which reduces the C—C bond distance in graphite compared with that found in diamond. The delocalisedbond readily explains the conductivity and colour of graphite, properties absent in diamond which has no such delocalised bonding. The planes of carbon atoms are held together by van der Waal’s forces which are much weaker than either a or bonding and allow the planes to slide over each other. Graphite is consequently anisotropic and much research has been carried out in attempts to produce large single crystals. Graphite manufactured on a large scale by the Acheson process, in which coke containing a little silica is heated in an electric furnace in the absence of air for many hours, does not produce large crystals. Single crystals of graphite, almost free from defects, have been produced by striking an electric arc between carbon rods. These “whiskers” have very high tensile strength along the planes of carbon atoms but are very brittle. So-called ‘carbon fibres’ have been produced by the controlled thermal degradation of certain acrylic textile fibres. The basic molecular orientation of the carbon atoms in the original fibre is retained. Plastics reinforced with carbon fibres are light in weight but have great strength, properties making them valuable to many industries and to the aero industry in particular. A process in which hydrocarbons are heated above 2300 K gives a material called pyrographite. This has properties indicating considerable ordering of the graphite crystals present. The thermal conductivity along the planes of carbon atoms is almost 100 times that at right angles to the planes, a property which makes the material valuable in rocket nose cones where rapid conduction from the hot zone is required and low conduction through to the interior. Electric conductance along the planes is 1000 times that found at right angles to the planes.

Amorphous carbon

Carbon Allotropes

Some allotropes of carbon: a) diamond; b) graphite; c) lonsdaleite; d–f) fullerenes (C₆₀, C₅₄₀, C₇₀); g) amorphous carbon; h) carbon nanotube.

In addition to diamond and graphite, carbon appears to exist in a number of other forms, collectively called amorphous carbon. Four common examples are coke, animal charcoal, lampblack and sugar carbon which can be prepared by heating coal, bones, oil and sugar respectively in the virtual absence of air. X-ray diffraction studies indicate that these and nearly all other forms of amorphous carbon are in fact microcrystallme graphite. Truly amorphous carbon, which gives random X-ray scattering, can be prepared by the low temperature decomposition of hexaiodobenzene, C₆I₆. Charcoal and lampblack have enormous surface areas for a small volume of sample, and are able to adsorb large amounts of gas or liquid. The effectiveness of the carbon can be greatly increased by heating the sample in a stream of steam to 1100-1300 K when impurities adsorbed during the initial preparation are driven off. This ‘activated’ charcoal has particularly good adsorption properties and is used as a catalyst. Lampblack is used in making printing ink, pigments and as a filler for rubber to be used in tyres.

In 1985 Harold Kroto, Robert Curl, and Richard Smalley, together with graduate students Jim Heath and Sean O’Brien, were attempting to create long carbon chains similar to those observed by radioastronomers near red dwarf stars. In their experiment they used a laser to vaporize a graphite disk, analyzing the products formed using a mass spectrometer. To their surprise, in addition to the sought-after compounds, they observed a particularly stable peak at a mass of 720 u, corresponding to a formula of C60. Neither the bonding in graphite nor that in diamond could account for a stable structure containing 60 carbon atoms. Although there wasn’t enough product available to perform a chemical analysis, the team suggested that the molecule was a  truncated icosahedron, a three-dimensional structure composed of 12 pentagons and 20 hexagons shaped like a soccer ball (Figure M.4)—a never-before-observed allotrope of carbon. The team named the new molecules buckminsterfullerenes, after architect Richard Buckminster Fuller, who designed the geodesic dome. They are also often called by their more colloquial name, buckyballs. Since their initial discovery, scientists have learned how to produce buckminsterfullerene in macroscopic quantities. Kroto, Curl, and Smalley shared the 1996 Nobel Prize in Chemistry for their discovery of fullerenes, a new allotrope of carbon (sidebar). Since this discovery, a number of other interesting stable compounds of similar structure, collectively called fullerenes, have been discovered, including C70, C76, C84, and so forth up to giant fullerene molecules containing more than half a million atoms. Fullerenes have been found to be a component of soot, observed spectroscopically in interstellar space, discovered in rocks hit by lightning, and found in meteors. One of the more intriguing findings was that of helium atoms with isotopic ratios not found on earth trapped inside fullerene cages isolated from meteorites, suggesting a possible origin outside our solar system.

Fullerenes are not the only new allotrope of carbon to form upon the vaporization of graphite; another cylindrical structure called carboanotubes has also been discovered and is the focus of much attention. The structure of carboanotubes is a chicken-wire-like mesh of carbon atoms composed of hexagons wrapped into a long cylinder 1–2 nm in diameter and up to a millimeter or longer in length (Frontispiece). Nanotubes may eventually find uses as new building materials for cars and aircraft, ultrastrong fibers for clothing, ultrathin flexible wires, and a variety of other innovative applications. The discovery of nanotubes has fueled the growth of a new field of research called nanotechnology, the goal of which is the production of molecular-sized machines and electronic devices that are several to a hundred nanometers in size.

Binary compounds in which carbon is combined with less electronegative elements are called  carbides. One of the most important carbides is calcium carbide. It is produced industrially by the reaction of lime, CaO(s), and coke (a form of solid carbon that can be derived from coal) according to

CaC₂(s) + CO(g) 2000°C→ CaO(s) + 3 C(s)

Calcium carbide is a gray-black, hard solid with a melting point over 2000°C. Calcium carbide reacts with water to produce acetylene. The equation for the reaction is:

CaC₂ (s) + 2 H₂O(l) → C₂H₂ (g) + Ca(OH)₂(s)

At one time, this reaction represented one of the major sources of acetylene for the chemical industry and for oxyacetylene welding, but its use has declined because of the high cost of the energy used in the production of the calcium carbide. Acetylene is now produced from petroleum and natural gas.

One other industrially important carbide is silicon carbide, SiC(s), also known as  carborundum. Carborundum is one of the hardest known materials and is used as an abrasive for cutting metals and polishing glass. Its structure is similar to the cubic crystal structure of diamond.

Carbon forms several sulfides, but only one of them, carbon disulfide, CS₂(l ), is stable at room temperature. Carbon disulfide is a colorless, poisonous, flammable liquid. The purified liquid has a sweet, pleasing odor, but the commonly occurring commercial and reagent grades have an extremely disagreeable odor that is due to organic impurities. Large quantities of CS₂(l ) are used in the manufacture of rayon, carbon tetrachloride, and cellophane. It is also used as a solvent for a number of substances.

 

TYPICAL REACTIONS OF THE ELEMENTS

1. THE REACTIONS WITH ACIDS

Dilute acids have no effect on any form of carbon, and diamond is resistant to attack by concentrated acids at room temperature, but is oxidised by both concentrated sulphuric and concentrated nitric acid at about 500 K, when an additional oxidising agent is present. Carbon dioxide is produced and the acids are reduced to gaseous oxides:

C + 4HNO₃ → CO₂ + 2H₂O + 4NO₂

C + 2H₂SO₄ → CO₂ + 2H₂O + 2SO²

Graphite reacts rather differently with mixtures of oxidising agents and concentrated oxoacids. A “graphite oxide’ is formed; the graphite swells because oxygen atoms become attached to some of the carbon atoms in the rings and distend the layer structure. ‘Graphite oxide’ is rather indefinite in composition. With concentrated sulphuric acid and an oxidising agent a blue solution called ‘graphite hydrogen sulphate‘ is formed; this has an approximate formula (Cx) ⁺ HSO₄ ^–. 2H₂SO₄.

Amorphous carbon, having a far greater effective surface area than either diamond or graphite, is the most reactive form of carbon. It reacts with both hot concentrated sulphuric and hot concentrated nitric acids in the absence of additional oxidising agents but is not attacked by hydrochloric acid.

2. THE REACTIONS WITH ALKALIS

Carbon does not react, even with molten alkali.

3. THE REACTIONS WITH OXYGEN

All forms of carbon, if heated to a sufficiently high temperature, give carbon dioxide in a plentiful supply of air, and carbon monoxide if the supply is limited:

C + O₂ → CO₂; ∆H= – 394 kJ mol^-1

C + ½O₂ → CO; ∆H = – 111 kJ mol ^-1

COMPOUNDS OF GROUP IV ELEMENTS

HYDRIDES

Carbon hydrides are commonly called hydrocarbons, its organic compounds (alkenes and alkynes). 

Structural formula of methane, the simplest possible organic compound.

 

Oxides of carbon

Carbon monoxide, CO

Carbon monoxide is a colourless, odourless gas. It is extremely poisonous, since the haemoglobin of the blood reacts with carbon monoxide in preference to oxygen so preventing the haemoglobin from acting in its normal capacity as an oxygen carrier.

Carbon monoxide is formed by the incomplete combustion of carbon. It is prepared in the laboratory by dropping methanoic (formic) acid into warm concentrated sulphuric acid; the latter dehydrates the methanoic acid:

H₂SO₄

HCOOH  →  CO↑ + H₂O

The gas is passed through caustic soda solution to remove any sulphur dioxide or carbon dioxide produced in side reactions.

The carbon dioxide is removed by passage of the gas through a mixture of sodium and calcium hydroxides. Very pure carbon monoxide is produced by heating nickel tetracarbonyl:

Ni(CO)₄ → Ni + 4CO↑

The commercial production of carbon monoxide in the form of water gas is now largely obsolete.The production by the reaction between steam and hydrocarbons is considered later.

The structure of carbon monoxide can be represented as a resonance hybrid between two structures (a) C ≡ O and (b) C = O

In structure (a) each atom has a complete octet; in the actual molecule, the carbon-oxygen bond length is greater than would be expected for a triple bond, and the molecule has a much smaller dipole moment than would be expected if the oxygen was donating electrons to the carbon as in (a); hence structure (b) must contribute to the actual structure. A simplified orbital picture of structure is shown at the top of the next page, where π₁ is formed by sharing electrons from both carbon and oxygen and π₂ is formed by electrons donated from oxygen only. This structure indicates that carbon monoxide should have donor properties, the carbon atom having a lone pair of electrons. Carbonmonoxide is in fact found to have donor properties and forms donor compounds, for example with diborane it splits the molecule by donating to the borane, BH₃.

2CO + B₂H₆ ↔ 2OC → BH₃

It also forms compounds known as carbonyls with many metals. The best known is nickel tetracarbonyl, Ni(CO)₄, a volatile liquid, clearly covalent.

Carbon monoxide burns with a characteristic blue flame in air or oxygen. The reaction

2CO + O₂ → 2CO₂ ; ∆H= – 283 kJ mol^-1

is very exothermic and as expected, therefore, carbon monoxide reacts with heated oxides of a number of metals, for example copper, lead, iron, reducing them to the metal. For example:

PbO + CO → Pb + CO₂↑

Carbon monoxide forms addition compounds. With chlorine in sunlight or in the presence of charcoal in the dark, carbonyl chloride (phosgene). COCl₂. is formed :

CO + C1₂ → COCl₂

Although carbon monoxide appears to be the anhydride of methanoic acid it does not react with water to give the acid; however, it will react with sodium hydroxide solution above 450 K, under pressure, to give sodium methanoate:

CO + NaOH → HCOONa

Carbon dioxide, CO₂.

Carbon dioxide is present in air and escapes from fissures in the earth in volcanic regions and where “mineral springs” occur. It may be prepared by:

(1) the action of dilute acid on any metal carbonate or hydrogencarbonate, for example

CaCO₃ + 2HC1 → CaCl₂ + CO₂↑ + H₂O

(2) the action of heat on a hydrogencarbonate,

2HCO₃^– → H₂O + CO₂↑ + CO₃^2-

(3) the action of heat on a metal carbonate, other than those of the alkali metals or barium. Industrially, carbon dioxide is obtained in large quantities by heating limestone:

CaCO₃ → CaO + CO₂↑

It is obtained as a by-product in the fermentation of sugars to give alcohols:

C₆H₁₂O₆ → 2C₂H₅OH + 2CO₂

CO + H₂O → CO₂ + H₂

In one process the carbon dioxide is removed using potassium carbonate solution, potassium hydrogencarbonate being produced:

K₂CO₃ + H₂O + CO₂ → 2KHCO₃

This reaction can be reversed by heat and the potassium carbonate and carbon dioxide recovered.

STRUCTURE

Carbon dioxide has a linear structure. The simple double-bonded formula, however, does not fully explain the structure since the measured carbon-oxygen bond lengths are equal but intermediate between those expected for a double and a triple bond. A more accurate representation is, therefore, obtained by considering carbon dioxide as a resonance hybrid of the three structures given below:

O⁺ ≡ C—O^– ↔ O = C = O ↔ O^–—C ≡ O⁺

  (a)                   (b)                  (c)

PROPERTIES

Carbon dioxide is a colourless gas which is virtually odourless and tasteless. Its density, relative to air, is 1.53; hence it accumulates at

CO + 3H₂ → CH₄ + H₂O

the bottom of towers or wells in which it is being prepared, and may reach dangerous concentrations there. (Carbon dioxide does not support respiration, but it is not toxic.) Its critical point is 304 K. i.e. it may be compressed to a liquid below this temperature. However, if carbon dioxide is cooled rapidly solid carbon dioxide is formed. This sublimes at 195 K and atmospheric pressure; it is a white solid, now much used as a refrigerant (known as “dry ice” or “Drikold”), since it leaves no residue after sublimation.

Chemically, carbon dioxide is not very reactive, and it is often used as an inactive gas to replace air when the latter might interact with a substance, for example in the preparation of chromium(II) salts. Very reactive metals, for example the alkali metals and magnesium can, however, continue to burn in carbon dioxide if heated sufficiently, for example

4K + 3CO₂ → 2K₂CO₃ + C

Carbon dioxide reacts with a solution of a metal hydroxide giving the carbonate, which may be precipitated, for example

Ca ²⁺  + 2OH^– + CO₂ → CaCO₃ + H₂O

This reaction is used as a test for carbon dioxide. Passage of an excess of carbon dioxide produces the soluble hydrogencarbonate :

CaCO₃ + CO₂ + H₂O → Ca ²⁺  + 2HCO₃^–

The hydrogencarbonate ion, produced iature by this reaction, is one of the main causes of temporary hardness in water. Carbon dioxide is fairly soluble in water, 1 cm³ dissolving 1.7 cm³ of the gas at stp. The variation of solubility with pressure does not obey Henry’s law, since the reaction

CO₂ + H₂O = H₂CO₃

takes place to a small extent, forming carbonic acid.

USES

Carbon dioxide is used in the manufacture of sodium carbonate by the ammonia-soda process, urea, salicyclic acid (for aspirin), fire extinguishers and aerated water. Lesser amounts are used to transfer heat generated by an atomic reactor to water and so produce steam and electric power, whilst solid carbon dioxide is used as a refrigerant, a mixture of solid carbon dioxide and alcohol providing a good low-temperature bath (195 K) in which reactions can be carried out in the laboratory. The following equilibria apply to a solution of carbon dioxide in water :

CO₂ + H₂O ↔ H₂CO₃  ↔H⁺  + HCO^– ↔ 2H⁺ + CO₃^2-

Carbonic acid

The amount of carbonic acid present, undissociated or dissociated, is only about 1 % of the total concentration of dissolved carbon dioxide. Carbonic acid, in respect of its dissociation into hydrogen and hydrogencarbonate ions, is actually a stronger acid than acetic acid. Since carbonic acid is a weak acid, its salts are hydrolysed in aqueous solution.

 

STABILITY OF CARBONATES AND HYDROGENCARBONATES TO HEAT

The stability to heat of metal carbonates is related to the size and charge of the cation present. Carbonates formed by metal ions with large radius: charge ratios, for example, Na ⁺ , K⁺ , Ba ²⁺ , are stable to heat at high temperatures whilst those ions with low radius : charge ratios, for example, Li⁺, Zn^{2 +} , Cu²⁺ form carbonates which are relatively easily decomposed by heat, the effect being so marked with Fe ³⁺ and A1³⁺ that neither of these ions is able to form a carbonate stable at room temperature. These changes in stability have been attributed to the amount of distortion of the carbonate ion that the metal ion causes; the greater this distortion the lower the stability of the carbonate. The hydrogencarbonate ion is unstable and decomposes on heating in either solid or solution thus:

2HCO₃^– → H₂O + CO₂↑ + CO₃^2-

(If the hydrogencarbonate is in solution and the cation is Ca ^{2 +} or Mg^{2 +} , the insoluble carbonate is precipitated; this reaction may be used, therefore, to remove hardness in water by precipitation of Ca^{2 +} or Mg^{2 +} ions.) The ease of decomposition of hydrogencarbonates affords a test to distinguish between a hydrogencarbonate and a carbonate; carbon dioxide is evolved by a hydrogencarbonate, but not by a carbonate, if it is heated, either as the solid or in solution, on a boiling water bath.

 

CHLORIDES AND OTHER IMPORTANT HALIDES OF

GROUP IV ELEMENTS

All Group IV elements form tetrachlorides, MX₄, which are predominantly tetrahedral and covalent. Germanium, tin and lead also form dichlorides, these becoming increasingly ionic in character as the atomic weight of the Group IV element increases and the element becomes more metallic. Carbon and silicon form catenated halides which have properties similar to their tetrahalides.

CARBON

When carbon forms four covalent bonds with halogen atoms the second quantum level on the carbon is completely filled with electrons. Carbon tetrachloride (tetrachloromethane) is a liquid, b.p. 350 1, and is prepared by the action of chlorine on carbon disulphide in the presence of a catalyst, usually manganese(II) chloride or iron(III) chloride :

CS₂ + 3C1₂ → CC1₄ + S₂C1₂

Further reaction then occurs between the disulphur dichloride and the carbon disulphide :

2S₂C1₂ + CS₂ → CC1₄ + 6S

Carbon tetrachloride is an excellent solvent for organic substances.

OTHER IMPORTANT COMPOUNDS

CARBIDES

These can be divided into three groups:

The salt-like carbides: Among these are aluminium tricarbide (methanide) A1₄C₃ (containing essentially C^{4 –} ions) in the crystal lattice and the rather more common dicarbides containing the C₂^2- ion, for example calcium dicarbide CaC₂; these carbides are hydrolysed by water yielding methane and ethyne respectively:

A1₄C₃ + 12H₂O → 4A1(OH)₃ + 3CH₄↑

2Al + 3H₂C₂ = Al₂(C₂)₃ + 3H₂­

CaC₂ + 2H₂O → Ca(OH)₂ + C₂H₂↑

The covalent carbides: These include boron carbide B4C and silicon carbide SiC; the latter is made by heating a mixture of silica and coke in an electric furnace to about 2000 K :

SiO₂ + 3C → SiC + 2CO↑

The interstitial carbides: These are formed by the transition metals (e.g. titanium, iron) and have the general formula M_XC. They are ofteon-stoichiometric—the carbon atoms can occupy some or all of the small spaces between the larger metal atoms, the arrangement of which remains essentially the same as in the pure metal (cf.the interstitial hydrides).

CARBON DISULPHIDE, CS₂

This was formerly manufactured by passing sulphur vapour over white hot coal or charcoal. An equilibrium was established and the carbon disulphide vapour was condensed, allowing the reaction to proceed:

C + 2S → CS₂

2CS₂ + 5O₂ → 2CO + 4SO₂

It is also decomposed by water above 420 K:

CS₂ + 2H₂O → CO₂ + 2H₂S

CHEMICAL TEST FOR GROUP IV ELEMENTS

All carbon compounds, if oxidised by either oxygen or an oxide (such as copper(II) oxide) yield carbon dioxide, which gives a precipitate of calcium carbonate when passed into aqueous calcium hydroxide.

Video Gas tests (Carbon dioxide) http://www.youtube.com/watch?v=LiAvDpl5aJA&feature=related

 SILICON

After oxygen, silicon is the most abundant element in the earth’s crust. It occurs extensively as the oxide, silica, in various forms, for example, flint, quartz, sand, and as silicates in rocks and clays, but not as the free element, silicon.

Silicon powder

Silicon is prepared by reduction of silica, SiO₂– Powdered “amorphous’ silicon can be obtained by heating dry powdered silica with either powdered magnesium or a mixture of powdered aluminium and sulphur (this supplies additional heat). After the reaction has ceased, magnesium (or aluminium) oxide and any unchanged silica is removed by washing with hydrofluoric acid in a polythene vessel:

SiO₂ + 2Mg → 2MgO + Si

(If an excess of magnesium is used, magnesium silicide, Mg₂Si, is also produced.) The silicon obtained is a light brown hygroscopic powder. Crystalline or ‘metallic’ silicon is obtained industrially by the reduction of silica with carbon in an electric arc furnace:

SiO₂ + 2C → 2CO + Si

 

Silicon crystallizes in the diamond cubic crystal structure

 

The formation of silicon carbide, SiC (carborundum), is prevented by the addition of a little iron; as much of the silicon is added to steel to increase its resistance to attack by acids, the presence of a trace of iron does not matter. (Addition of silicon to bronze is found to increase both the strength and the hardness of the bronze.) Silicon is also manufactured by the reaction between silicon tetrachloride and zinc at 1300K and by the reduction of trichlorosilane with hydrogen. Crystalline silicon has the tetrahedral diamond arrangement, but since the mean thermochemical bond strength between the silicon atoms is less than that found between carbon atoms (Si—Si, 226kJmol^–1 , C—C, 356kJmol^-1 ), silicon does not possess the great hardness found in diamond. Amorphous silicon (silicon powder) is microcrystalline silicon.

The chemistry of silicon differs markedly from that of carbon. Although carbon and silicon both have a normal valence of four and are similar in some ways, the differences between the chemical properties of their compounds are more striking than the similarities. There are three groups of compounds where the differences are easily seen.

1. The oxides

The oxides of the two elements have very different properties. Carbon dioxide, CO₂ is a gas at room temperature and atmospheric pressure. It consists of CO₂ molecules, in which there are C=O double bonds so as to satisfy the valences of both carbon and oxygen. The linear shape of the molecule is shown in this drawing of the molecule:

                                                 O=C=O

 

Silicon dioxide, usually known as silica, has the empirical formula SiO₂, but it is a solid with a very high melting point, not a gas. The contrast between the properties of CO₂ and SiO₂ could not be more striking. The reason for the difference is that silica consists of an extended network of silicon and oxygen atoms. It is difficult to give a clear impression of the three dimensional structure of silica, but the following two-dimensional structural formula gives the essential idea of the network:

 

 Each Si atom is bonded to four O atoms, and each O atom is bonded to two Si atoms. The bonds are all shown as single covalent bonds. In this way, both elements achieve their normal valence, four for Si and two for O. In the full three-dimensional structure of silica, the four O atoms around each Si atom form a tetrahedron. The network of atoms and bonds extends for a long distance in all directions, and this is indicated by the incomplete bonds at the edge of the above structure.

The network of bonds holds the atoms together in a strong solid, which can only be melted at very high temperature. In part, this is reflected in the DH_f^o for SiO₂(s), which is much more negative than DH_f^o for CO₂(g). (CHEMBOOK, Appendix I; Oxtoby, Appendix D)

To summarize, CO₂ molecules contain three atoms joined by C=O double bonds, whereas the SiO₂ lattice contains an enormous number of atoms joined by Si–O single bonds. This reminds us that, although second row elements form covalent double and triple bonds in many of their compounds, elements of the third and subsequent rows rarely form covalent double bonds. (CHEMBOOK page 82) As examples of this rule, carbon commonly forms double bonds, while silicon forms only single bonds.

2. The hydrides, or silanes

The hydrides of silicon have formulae such as SiH₄, Si₂H₆ etc, and are called silanes. These compounds are analogous to the alkanes CH₄, C₂H₆ etc.

The silanes are very reactive compounds. Silanes react vigorously with oxygen, and ignite or explode spontaneously in contact with air. Silanes react with water containing hydroxide ions OH^–, producing a form of form silica and hydrogen gas. By comparison, the alkanes are rather unreactive compounds. Alkanes can be mixed with oxygen without any reaction unless ignited by a spark or flame, and do not react with water at ordinary temperatures.

The reason that silanes are much more reactive than alkanes is not related to any “weakness” in the bonds. The Si–Si and Si–H bonds are almost as “strong” as the C–C and C–H bonds. The reactivity of silanes has been ascribed to several factors: the silicon atom is larger than the carbon atom, and has a lower electronegativity than carbon.

3. Silicates

Silicates are salts in which the anions contain silicon and oxygen. Silicates form a large fraction of the minerals and rocks of the earth’s crust (CHEMBOOK Figure 3-1). Many different silicate minerals are known, and they have bewildering variety of crystal structures. By comparison, there are relatively few carbonate salts.

Students studying geology will become experts in recognizing silicate minerals in the field, together with their crystal structures. There is an extensive display of minerals in the mineral museum in Miller Hall, which is open every weekday.

Most silicates are formed from tetrahedral SiO₄ structural units, which can be joined together by Si–O–Si groups of atoms. The SiO₄ tetrahedra can be joined in many different ways, each with a characteristic network of silicon and oxygen atoms.

The silicate minerals are the most characteristic compounds of silicon. In contrast, the compounds of carbon are the basis for the chemistry of life. Together, the two elements form the basis for understanding almost everything animal, vegetable and mineral (to use an old phrase).

 

GERMANIUM

Germanium is a greyish-white, brittle solid, obtained by reducing the dioxide, GeO₂, with hydrogen or carbon at red heat. Germanium is a rare element found in trace quantities in coke obtained from bituminous coal. When this coke is burnt, germanium dioxide, together with many other metal oxides, is deposited in the flue. The extraction of germanium dioxide from this mixture is a complex process. Impure germanium and silicon are both purified by zone refining and both can be obtained in a very high purity, for example silicon pure to one part in 1010 can be obtained*. Germanium, likesilicon, crystallises with a diamond structure, the mean thermochemical bond strength being Ge—Ge, 188 kJ mol ^-l.

TIN

The common ore of tin is tinstone or cassiterite, SnO₂, found in Cornwall and in Germany and other countries. The price of tin has risen so sharply in recent years that previously disregarded deposits in Cornwall are now being re-examined. Tin is obtained from the tin dioxide, SnO₂, by reducing it with coal in a reverbatory furnace:

SnO₂ + 2C → 2CO↑+ Sn

Before this treatment, the cassiterite content of the ore is increased by removing impurities such as clay, by washing and by roasting which drives off oxides of arsenic and sulphur. The crude tin obtained is often contaminated with iron and other metals. It is, therefore, remelted on an inclined hearth; the easily fusible tin melts away, leaving behind the less fusible impurities. The molten tin is finally stirred to bring it into intimate contact with air. Any remaining metal impurities are thereby oxidised to form a scum (“tin dross’) on the surface and this can be skimmed off. Very pure tin can be obtained by zone refining. Tin exists in three different forms (allotropes). “Grey tin” has a diamond structure, a density of 5.75gcm^~3  and is stable below 286 K. “White tin” exists as tetragonal crystals, has a density of 7.31 gcm^~3 and is stable between 286 and 434 K. Between 434 K and the melting point of tin, 505 K, tin has a rhombic structure, hence the name ‘rhombic tin’, and a density of 6.56 g cm^~3

LEAD

The principal ore of lead is galena, PbS. Although there are some galena deposits in Great Britain, much of this country’s requirements must be imported. In the extraction of lead, the sulphide ore is first roasted together with quartz in a current of air:

2PbS + 3O₂ → 2PbO + 2SO₂↑

Galena, lead ore

Any lead(II) sulphate formed in this process is converted to lead(II) silicate by reaction with the quartz. The oxide produced is then mixed with limestone and coke and heated in a blast furnace. The following reactions occur:

PbO + C → Pb + CO↑

PbO + CO → Pb + CO₂↑

PbSiO₃ + CaO + CO → Pb + CaSiO₃ + CO₂↑

The last equation explains the function of the limestone. An older process, in which the ore was partially roasted, the air shut off and the temperature raised so that excess sulphide reacted with the oxide produced to give lead, is now obsolete. Crude lead contains traces of a number of metals. The desilvering of lead is considered later under silver. Other metallic impurities are removed by remelting under controlled conditions when arsenic and antimony form a scum of lead(II) arsenate and antimonate on the surface while copper forms an infusible alloy which also takes up any sulphur, and also appears on the surface. The removal of bismuth, a valuable by-product, from lead is accomplished by making the crude lead the anode in an electrolytic bath consisting of a solution of lead in fluorosilicic acid. Gelatin is added so that a smooth coherent deposit of lead is obtained on the pure lead cathode when the current is passed. The impurities here (i.e. all other metals) form a sludge in the electrolytic bath and are not deposited on the cathode. Lead has only one form, a cubic metallic lattice. Thus we can see the change from non-metal to metal in the physical structure of these elements, occurring with increasing atomic weight of the elements carbon, silicon, germanium, tin and lead.

TYPICAL REACTIONS OF THE ELEMENTS

1. THE REACTIONS WITH ACIDS

Silicon

Silicon, like carbon, is unaffected by dilute acids. Powdered silicon dissolves incompletely in concentrated nitric acid to give insoluble silicon dioxide, SiO₂ :

3Si + 4HNO₃ → 3SiO₂ + 4NO + 2H₂O

Germanium

The gradual increase in electropositive character down the group is clearly shown in that, unlike both carbon and silicon, germanium very readily dissolves in both concentrated nitric and sulphuric acids; the hydrated germanium(IV) oxide is produced:

3Ge + 4HNO₃ → 3GeO₂ + 4NO + 2H₂O

Germanium, however, does not react with either dilute sulphuric or dilute hydrochloric acid, unlike tin, the next element in the group.

Tin

Tin slowly dissolves in dilute hydrochloric, nitric and sulphuric acids, and is in fact the only Group IV element to do so. The reactions with more concentrated acid are rapid. With hydrochloric acid, tin gives a solution of tin(II) chloride, there being no further oxidation to the + 4 oxidation state :

Sn + 2HCl → SnCl₂ + H₂↑

Droplet of melted tin

Concentrated nitric acid, however, is an oxidising agent and tin reacts to give hydrated tin(IV) oxide in a partly precipitated, partly colloidal form, together with a small amount of tin(II) nitrate, Sn(NO₃)₂ :

Sn + 4HNO₃ → SnO₂ + 4NO₂ + 2H₂O

A similar oxidation reaction occurs with concentrated sulphuric acid but in this case hydrated tin(IV) ions remain in solution :

Sn + 4H₂SO₄ → Sn(SO₄)₂ + 4H₂O + 2SO₂

Lead

Lead reacts only briefly with dilute hydrochloric and sulphuric acids for both lead(II) chloride and lead(II) sulphate are insoluble and form a film on the lead which effectively prevents further attack. Lead, however, does slowly dissolve in both concentrated sulphuric and hydrochloric acids. The sulphuric acid is reduced to sulphur dioxide :

Pb + 2H₂SO₄ → PbSO₄ + 2H₂O + SO₂↑

Lead reacts slowly with hot concentrated hydrochloric acid since the lead(II) chloride dissolves in an excess of the hot hydrochloric acid to form the acid H₂[PbCl₄]:

Pb + 4HC1 → H₂[PbCl₄] + H₂↑

Again, nitric acid readily dissolves lead but is unable to oxidise lead beyond the oxidation state +2. The reduction products of the nitric acid vary with the concentration of acid used, and a number of nitrogen oxides are usually obtained. Warm dilute nitric acid gives mainly nitrogen oxide, NO,

3Pb + 8HNO₃ → 3Pb(NO₃)₂ + 4H₂O + 2NO↑

whilst cold concentrated acid gives mainly nitrogen dioxide, NO₂ :

Pb + 4HNO₃ → Pb(NO₃)₂ + 2H₂O + 2NO₂↑

2. THE REACTIONS WITH ALKALIS

Silicon and germanium

Silicon and germanium readily react with even very dilute solutions of caustic alkali. Silicon is so sensitive to attack that it will dissolve when boiled with water which has been in contact with glass:

Si + 2OH^– + H₂O → SiO₃^2- + 2H₂↑

Ge + 2OH^– + H₂O → GeO₃^2- + 2H₂↑

Tin

Tin dissolves slowly in hot concentrated alkali forming a hexahydroxostannate(IV):

Sn + 4H₂O + 2OH^– → [Sn(OH)₆]^2- + 2H₂↑

Lead

Lead dissolves only very slowly in hot concentrated sodium hydroxide and forms hexahydroxoplumbate(II):

Pb + 4OH^– + 2H₂O → [Pb(OH)₆]^4- + H₂↑

Notice, again, that the lower oxidation state of lead is formed.

3. THE REACTIONS WITH OXYGEN

Silicon

Silicon burns when heated in air to red heat giving silicon dioxide, SiO₂. Several crystalline forms of SiO₂ are known.

Si + O₂ → SiO₂; ∆H = – 910 kJ mol^–l  (approximate)

Note the much larger enthalpy of formation of silicon dioxide as compared with carbon dioxide; this arises in part because of greater strength in the Si—O bonds and also because the Si—Si bond in silicon is much weaker than the C—C bond.

Tin

Ordinary white tin is not attacked by air at ordinary temperatures but on heating in air it forms tin(IY) oxide, SnO₂.

Sn + O₂ → SnO₂

Lead

Finely divided lead, when heated in air, forms first the lead(II) oxide, litharge’, PbO, and then on further heating in an ample supply of air, dilead(II) lead(IV) oxide, ‘red lead’, Pb₃O₄. Lead, in a very finely divided state, when allowed to fall through air, ignites and a shower of sparks is produced. Such finely divided powder is said to be “pyrophoric‘. It can be prepared by carefully heating lead tartrate.

COMPOUNDS OF GROUP IV ELEMENTS

HYDRIDES

Silicon

Silicon, unlike carbon, does not form a very large number of hydrides. A series of covalently bonded volatile hydrides called silanes analogous to the alkane hydrocarbons is known, with the general formula Si_nH_2n+2′ but less than ten members of the series have so far been prepared. Mono- and disilanes are more readily prepared by the reaction of the corresponding silicon chloride with lithium aluminium hydride in ether:

SiCl₄ + LiAlH₄ → SiH₄↑ + LiCl↓ + A1C1₃ and

2Si₂Cl₆ + 3LiAlH₄ → 2Si₂H₆t↑ + 3LiCl↓ + 3A1C1₃

Si₂H₆ + 8H₂O = 2Si(OH)₄ + 7H₂­,

SiH₄ + 2KOH + H₂O = K₂SiO₃ + 4H₂­,

SiH₄ + HBr = SiH₃Br + H₂­.

 

The Si — Si bond is weaker than the 3A1C1 3LiCli C bond (mean thermochemical bond energies are C — C in diamond, 356 kJ mol^–1 , Si — Si in silicon, 226kJmol^-1 ) and catenation (the phenomenon of self-linkage between atoms of the same element) is consequently less marked with silicon than with carbon; the higher silanes decompose slowly even at room temperature. Silanes are far more sensitive to oxygen than alkanes and all the silanes are spontaneously inflammable in air, for example

SiH₄ + 2O₂ → SiO₂ + 2H₂O

This greater reactivity of the silanes may be due to several factors, for example, the easier approach of an oxygen molecule (which may attach initially to the silane by use of the vacant silicon d orbitals) and the formation of strong Si—O bonds (stronger than C—O). Halogen derivatives of silanes can be obtained but direct halogenation often occurs with explosive violence; the halogen derivatives are usually prepared by reacting the silane at low temperature with a carbon compound such as tetrachloromethane, in the presence of the corresponding aluminium halide which acts as a catalyst. Silanes are very sensitive to attack by alkalis and will even react with water made alkaline by contact with glass; this reaction is in marked contrast to the reactions shown by alkanes. Unlike alkanes, silanes are found to have marked reducing properties and will reduce, for example, potassium manganate(VII) to manganese(IV) oxide, and iron(III) to iron(II). In addition to the volatile silanes, silicon also forms non-volatile hydrides with formulae (SiH₂)x but little is known about their structure. Silicon, however, does not form unsaturated hydrides corresponding to the simple alkenes.

Germanium

Germanium forms a series of hydrides of general formula Ge_nH_2n+2 which are quite similar to the corresponding silanes. Only a small number of germanes have so far been prepared. Germanes are not as inflammable as the corresponding silanes (the Ge—O bond is not as strong as the Si—O bond) and they are also less reactive towards alkalis, monogermane being resistant to quite concentrated alkali.

Germane is similar to methane.

Tin

The greater metallic nature of tin is clearly indicated here for tin forms only one hydride, stannane, SnH₄. It is best prepared by the reaction of lithium aluminium hydride and tin(IV) chloride in ether:

LiAlH₄ + SnCl₄ → SnH₄↑+ LiCl↓ + A1C1₃

It is a colourless gas which decomposes on heating above 420 K to give metallic tin, often deposited as a mirror, and hydrogen. It is a reducing agent and will reduce silver ions to silver and mercury(II) ions to mercury. SnSn bonding is unknown in hydrides but does exist in alkyl and aryl compounds, for example (CH₃)₃Sn-Sn(CH₃)₃.

Lead

Lead, like tin, forms only one hydride, plumbane. This hydride is very unstable, dissociating into lead and hydrogen with great rapidity. It has not been possible to analyse it rigorously or determine any of its physical properties, but it is probably PbH4. Although this hydride is unstable, some of its derivatives are stable; thus, for example, tetraethyllead, Pb(C₂H₅)₄, is one of the most stable compounds with lead in a formal oxidation state of + 4. It is used as an “antiknock’ in petrol.

OXIDES OF GROUP IV ELEMENTS

All Group IV elements form both a monoxide, MO, and a dioxide, MO₂. The stability of the monoxide increases with atomic weight of the Group IV elements from silicon to lead, and lead(II) oxide, PbO, is the most stable oxide of lead. The monoxide becomes more basic as the atomic mass of the Group IV elements increases, but no oxide in this Group is truly basic and even lead(II) oxide is amphoteric. Carbon monoxide has unusual properties and emphasises the different properties of the group head element and its compounds. The dioxides are all predominantly acidic but again acidity decreases with increasing atomic mass of the Group IV element and lead(IV) oxide, PbO₂, is amphoteric. The stability of the dioxides decreases with increasing atomic mass of the Group IV elements and although tin(IV) oxide, SnO₂, is the most stable oxide of tin, lead(IV) oxide is less stable than lead(II) oxide.

Oxides of silicon

SILICON MONOXIDE, SiO

When silica (silicon dioxide) and silicon are heated in vacuo to 1700 K, there is evidence for SiO in the gaseous state. On cooling, a brown powder is obtained which rapidly disproportionates:

2SiO → Si + SiO₂

SILICON DIOXIDE, SiO₂

Silica is found naturally in several crystalline forms (e.g. quartz, tridyniite. cristobalite) and as kieselguhr. a hydrated amorphous solid possessing great absorptive powers. It is not appropriate to refer to this oxide of silicon as a dioxide, since, in its crystalline forms, it forms “giant molecules” in which each silicon atom is linked tetrahedrally to four oxygen atoms: the structure can be represented diagrammatically thus, the linkages extending three-dimensionally. Pure silica may be obtained by hydrolysing silicon tetrafluoride or the tetrachlorid.

When so prepared, silica is hydrated ; it appears in fact as a gel i.e. a colloidal system in which a liquid is dispersed in a solid. This gel when filtered off and dried, loses much of its water, and on heating can be made anhydrous ; but formation of a solid gel takes place again when the anhydrous solid is exposed to a moist atmosphere, i.e. the solid absorbs water. Hence silica gel is a most useful drying agent, for it has a high capacity for absorbing water and it is also chemically inactive. Silica is attacked only by hydrofluoric acid, and by alkali to give silicates :

SiO₂ + 2OH^– → SiO₃^2- + H₂O

SiO₂ + 2Mg = Si + 2MgO

SiO₂ +C = Si + CO

SiO₂ + Ca(OH)₂ = CaSiO₃ + H₂O;

SiO₂ + 4HF = SiF₄­ + 2H₂O.

When silica is fused, silica glass is formed. This has advantages over ordinary glass in that it is much less easily fused, and has a very low coefficient of expansion. It is, therefore, used for crucibles and other articles required to be infusible and to resist chemical attack. It is also used for certain optical plates and lenses, since it transmits ultra-violet light better than ordinary glass.

“SILICIC ACID” AND THE SILICATES

When acid is added to any soluble silicate, the following reaction occurs:

SiO₃^2- (aq) + 2H⁺ → H₂SiO₃ (aq) → SiO₂* nH₂O

and the ‘silicic acid’ is converted to insoluble, hydra ted silica similar to that already described.

A soluble silicate—a trioxosilicate—is obtained when silica is fused with sodium carbonate:

SiO₂ + Na₂CO₃ → Na₂SiO₃ + CO₂↑

Clay and kaolin describe groups of substances with compositions which are similar chemically (they contain aluminium, silicon, oxygen and water) but with many different kinds of structure, the nature of which has been established by X-ray diffraction studies. The clays all possess a layer-like structure. When water is added to clay it enters between the layers and the clay swells and acquires plasticity thus enabling it to be moulded into bricks, pottery, and so on. On ignition or “firing”, these lose plasticity permanently acquiring thereby a fixed shape, hardness and strength. Kaolin is rather less ‘plastic’ than clay but can be moulded and then fired to give porcelain or “china”.

Glass is the name given to any amorphous solid produced when a liquid solidifies. Glasses are non-crystalline and isotropic, i.e. their physical properties are independent of the direction in which they are measured. When a glass is heated, it does not melt at a fixed temperature but gradually softens until a liquid is obtained. The word ‘glass’ commonly means the transparent substance obtained when white sand is fused with metal oxides or carbonates to give a mixture of silicates. Ordinary or ‘soda-glass’ has the approximate composition Na₂O . CaO . 6SiO₂. If sodium is replaced by potassium the melting point is raised (Jena glass) and the use of lithium gives added strength; replacement of calcium by lead gives a higher refracting power (flint glass), and the SiO₂ may be partly replaced by P₂O₅ (crown glass). Addition of aluminium and boron oxides gives a glass with a low coefficient of expansion suitable for vessels which are to be heated, e.g. “Pyrex”. Coloured glass is made by adding an oxide of a metal which gives a coloured silicate, e.g. cobalt (blue), iron(II) (green), copper(I) (red).

Oxides of germanium

GERMANIUM(II) OXIDE, GeO

The existence of germanium(II) oxide is well established. It is a solid which can be made, for example, by the action of water on germanium dichloride, GeCl₂:

GeCl₂ + H₂O → GeO + 2HC1

The product is a solid yellow hydrated oxide. If prepared by a method in the absence of water, a black anhydrous product is obtained. Germanium(II) oxide is stable in air at room temperature but is readily oxidised when heated in air or when treated at room temperature with, for example, nitric acid, hydrogen peroxide, or potassium manganate(VII). When heated in the absence of air it disproportionates at 800 K:

2GeO → Ge + GeO₂

The yellow hydrated oxide is slightly acidic and forms germanates(II) (germanites). The increased stability of germanium(II) oxide compared to silicon(II) oxide clearly indicates the more metallic nature of germanium.

GERMANIUM(IV) OXIDE

Germanium(IV) oxide occurs in two forms; one has a rutile lattice and melts at 1359K whilst the other has a quartz lattice and a melting point of 1389 K. It can be prepared by oxidation of germanium using, for example, concentrated nitric acid, or by the hydrolysis of germanium tetrachloride:

Ge + 4HNO₃ → GeO₂ + 4NO₂ + 2H₂O

GeCl₄ + 2H₂O → GeO₂ + 4HC1

Oxides of tin

TIN(II) OXIDE

If a solution of a tin(II) salt is treated with a small amount of an alkali, tin(II) hydroxide is precipitated, the reaction being represented by the equation:

Sn²⁺ + 2OH^– → Sn(OH)₂↓

Tin(II) oxide is a dark-coloured powder which oxidises spontaneously in air with the evolution of heat to give tin(IV) oxide, SnO₂:

2SnO + O₂ → SnO₂

It is amphoteric; it gives tin(II) salts with dilute acids and hydroxostannates(II) with alkalis, for example:

SnO + 2HC1 → SnCl₂ + H₂O

SnO + H₂O + OH^– → [Sn(OH)₃]^–

Stannate(II) ions are powerful reducing agents. Since, for tin, the stability of oxidation state +4 is greater than that of oxidation state +2, tin(II) always has reducing properties, but these are greater in alkaline conditions than in acid.

TIN(IV) OXIDE, SnO₂

Tin(IV) oxide occurs naturally, clearly indicating its high stability. It can be prepared either by heating tin in oxygen or by heating the hydrated oxide obtained when metallic tin reacts with concentrated nitric acid:

Sn + 4HNO₃ → SnO₂↓ + 4NO₂↑ + 2H₂O

Tin(IV) oxide is insoluble in water, but if fused with sodium hydroxide and the mass extracted with water, sodium hexahydroxostannate(IV) is formed in solution:

SnO₂ + 2NaOH + 2H₂O → Na₂[Sn(OH)₆]

Oxides of lead

LEAD(II) OXIDE, PbO

Lead(II) oxide is the most stable oxide of lead; it exists in two crystalline forms. One form is reddish yellow in colour, with a tetragonal lattice, and is called litharge. The other form, yellow in colour, has a rather greater density and a rhombic lattice ; it is called massicot. Litharge is obtained when molten lead is oxidised by a blast of air. By more careful heating, or by heating lead carbonate or lead nitrate, massicot is obtained. Litharge is the stable form at room temperature, but massicot changes only very slowly to litharge under ordinary conditions. Lead(II) oxide is the most basic oxide formed by a Group IV element. It dissolves easily in acids to give lead(II) salts but it also dissolves slowly in alkalis to give hydroxoplumbates(II) and must, therefore, be classed as an amphoteric oxide, for example :

PbO + 2H⁺ →Pb²⁺ + H₂O

PbO + 4OH^– + H₂O → [Pb(OH)₆]^4-

Lead(II) oxide is easily reduced to the metal when heated with a reducing agent such as hydrogen, carbon or carbon monoxide, for example :

PbO + H₂ → Pb + H₂O

LEAD(IV) OXIDE, PbO₂

Lead(IV) oxide can be prepared by the action of an alkaline chlorate(I) solution on a solution of a lead(II) salt. The reaction can be considered in two stages:

(1) Pb²⁺ + 2OH^– → Pb(OH)₂↓ – white

The white precipitate of lead hydroxide (or hydrated lead(II) oxide) is then oxidised by the chlorate(I) to the brown dioxide:

(2) Pb(OH)₂ + ClO^– → PbO₂↓ + Cl^– + H₂O

brown

Lead(IV) oxide is also obtained when ‘red lead’, Pb3O4 (see below), is treated with dilute nitric acid:

Pb₃O₄ + 4HNO₃ → 2Pb(NO₃)₂ + 2H₂O + PbO₂

When heated above 600 K lead(IV) oxide decomposes into the more stable lead(II) oxide and oxygen :

2PbO₂ → 2PbO + O₂

PbO₂ + 4HC1 → PbCl + C1₂ + 2H₂O

2PbO₂ + 2H₂SO₄ → 2PbSO₄ + 2H₂O + O₂

PbO₂ + S → Pb + SO₂

PbO₂ + SO₂ → PbSO₄

Lead dioxide is slightly soluble in concentrated nitric acid and concentrated sulphuric acid, and it dissolves in fused alkalis. It therefore has amphoteric properties, although these are not well characterised since it is relatively inert.

CHLORIDES AND OTHER IMPORTANT HALIDES OF

GROUP IV ELEMENTS

All Group IV elements form tetrachlorides, MX₄, which are predominantly tetrahedral and covalent. Germanium, tin and lead also form dichlorides, these becoming increasingly ionic in character as the atomic weight of the Group IV element increases and the element becomes more metallic. Carbon and silicon form catenated halides which have properties similar to their tetrahalides.

Silicon

Silicon tetrajluoride is formed when hydrogen fluoride reacts with silica or a silicate :

4HF + SiO₂ → SiF₄  + 2H₂O

The hydrogen fluoride is conveniently produced in situ by the action of concentrated sulphuric acid on calcium fluoride:

CaF₂ + H₂SO₄ → CaSO₄ + 2HF

Silicon tetrafluoride is a colourless gas, b.p. 203 K, the molecule having, like the tetrahalides of carbon, a tetrahedral covalent structure. It reacts with water to form hydrated silica (silica gel) and hexafluorosilicic acid, the latter product being obtained by a reaction between the hydrogen fluoride produced and excess silicon tetrafluoride :

SiF₄ + 2H₂O → SiO₂ + 4HF

SiF₄ + 2HF → H₂SiF₆

Silicon tetrachloride is a colourless liquid, b.p. 216.2 K, and again the molecule has a covalent structure. Silicon tetrachloride is hydrolysed by water :

SiCl₄ + 2H₂O → 4HC1 + SiO₂

Germanium

Germanium forms divalent compounds with all the halogens. Germanium(ll) chloride can be prepared by passing the vapour of germanium(IV) chloride (see below) over heated germanium. The reaction is reversible and disproportionation of germanium(II) chloride is complete at about 720 K at atmospheric pressure:

GeCl₄ + Ge → 2GeCl₂

Germanium(IV) chloride can be prepared by passing chlorine over germanium at a temperature of 350-450 K :

Ge + 2C1₂ → GeCl₄

Tin

TIN(ll) CHLORIDE

This chloride is prepared by dissolving tin in concentrated hydrochloric acid; on cooling, the solution deposits crystals of hydrated tin(II) chloride, SnCl2. 2H2O (‘tin salt’). The anhydrous chloride is prepared by heating tin in a current of hydrogen chloride:

Sn + 2HC1 → SnCl₂ + H₂

TIN(IV) CHLORIDE, SnCl₄

Stannic chloride is prepared by treating metallic tin with chlorine:

Sn + 2C1₂ → SnCl₄

SnCl₄ + 2H₂O → SnO₂ + 4HC1

Lead

LEAD(II) CHLORIDE

The solid is essentially ionic, made up of Pb²⁺  and Cl^– ions. The vapour contains bent molecules of PbCl₂. Lead chloride is precipitated when hydrochloric acid is added to a cold solution of a lead(II) salt. It dissolves in hot water.

LEAD(IV) CHLORIDE

PbCl₄ → PbCl₂ + C1₂↑

LEAD(II) IODIDE

The solid has a layer structure. Lead(II) iodide is soluble in hot water but on cooling, appears in the form of glistening golden ‘spangles’. This reaction is used as a test for lead(II) ions in solution.

OTHER IMPORTANT COMPOUNDS

Lead

LEAD(II) CARBONATE

Lead(II) carbonate occurs naturally as cerussite. It is prepared in the laboratory by passing carbon dioxide through, or adding sodium hydrogencarbonate to, a cold dilute solution of lead(II) nitrate or lead(II) ethanoate:

Pb²⁺ + 2HCO₃ ^–→ PbCO₃↓ + CO₂↑ + H₂O

If the normal carbonate is used, the basic carbonate or white lead, Pb(OH)₂. 2PbCO₃. is precipitated. The basic carbonate was used extensively as a base in paints but is now less common, having been largely replaced by either titanium dioxide or zinc oxide. Paints made with white lead are not only poisonous but blacken in urban atmospheres due to the formation of lead sulphide and it is hardly surprising that their use is declining.

LEAD(II) CHROMATE(Vl), PbCrO₄

Lead(II) chromate(VI) is precipitated when a soluble chromate(VI) or dichromate( VI) is added to a solution of a lead salt ieutral or slightly acid solution:

Pb²⁺ +CrO₄^2- →PbCrO↓

2Pb²⁺ + Cr₂O₇^2- + H₂O → 2PbCrO4↓ + 2H⁺

The precipitation of lead(II) chromate is used to estimate lead gravimetrically: the yellow precipitate of lead(II) chromate is filtered off, dried and weighed. Lead(II) chromate is used as a pigment under the name “chrome yellow”

THE LEAD ACCUMULATOR

The most widely-used storage battery is the lead accumulator. Each cell consists essentially of two lead plates immersed in an electrolyte of sulphuric acid. The lead plates are usually perforated and one is packed with lead(IV) oxide, the other with spongy lead. An inert porous insulator acts as a separator between the plates.

CHEMICAL TEST FOR GROUP IV ELEMENTS

Carbon

All carbon compounds, if oxidised by either oxygen or an oxide (such as copper(II) oxide) yield carbon dioxide, which gives a precipitate of calcium carbonate when passed into aqueous calcium hydroxide.

Silicon

All silicon compounds on oxidation yield silica or silicates; these are difficult to detect but silica (given by silicates after acid treatment) is insoluble in all acids except hydrofluoric acid.

Tin

In presence of hydrochloric acid, tin(II) in aqueous solution (1) is precipitated by hydrogen sulphide as brown SnS, and (2) will reduce mercury(II) chloride first to rnercury(I) chloride (white precipitate) and then to metallic mercury. Tin(IV) in aqueous acid gives a yellow precipitate with hydrogen sulphide, and no reaction with mercury(II) chloride.

Lead

Lead(II) in aqueous solution gives on addition of the appropriate anion (1) a white precipitate of lead(II) chloride, (2) a yellow precipitate of lead(II) chromate, and (3) a yellow precipitate of lead(II) iodide which dissolves on heating and reappears on cooling in the form of glistening ‘spangles’.

Characteristics of lead

A sample of freshly solidified lead (from molten state)

Lead is a bright and silvery metal with a very slight shade of blue in a dry atmosphere. Upon contact with air, it begins to tarnish by forming a complex mixture of compounds depending on the conditions. The color of the compounds can vary. The tarnish layer can contain significant amounts of carbonates and hydroxycarbonates. It has a few characteristic properties: high density, softness, ductility and malleability, poor electrical conductivity compared to other metals, high resistance to corrosion, and ability to react with organic chemicals.

Various traces of other metals change its properties significantly: the addition of small amounts of antimony or copper to lead increases the alloy’s hardness and improves corrosion resistance from sulfuric acid. A few other metals also improve only hardness and fight metal fatigue, such as cadmium, tin, or tellurium; metals like sodium or calcium also have this ability, but they weaken the chemical stability. Finally, zinc and bismuth simply impair the corrosion resistance (0.1% bismuth content is the industrial usage threshold).^[1] In return, lead impurities mostly worsen the quality of industrial materials, although there are exceptions: for example, small amounts of lead improve the ductility of steel.

Lead has only one common allotrope, which is face-centered cubic, with the lead–lead distance being 349 pm. At 327.5 °C (621.5 °F), lead melts; the melting point is above that of tin (232 °C, 449.5 °F), but significantly below that of germanium (938 °C, 1721 °F). The boiling point of lead is 1749 °C (3180 °F),^[7] which is below those of both tin (2602 °C, 4716 °F) and germanium (2833 °C, 5131 °F). Densities increase down the group: the Ge and Sn values (5.23 and 7.29 g·cm⁻³, respectively) are significantly below that of lead: 11.32 g·cm⁻³.

A lead atom has 82 electrons, having an electronic configuration of [Xe]4f¹⁴5d¹⁰6s²6p². In its compounds, lead (unlike the other group 14 elements) most commonly loses its two and not four outermost electrons, becoming lead(II) ions, Pb²⁺. Such unusual behavior is rationalized by considering the inert pair effect, which occurs because of the stabilization of the 6s-orbital due to relativistic effects, which are stronger closer to the bottom of the periodic table. Tin shows a weaker such effect: tin(II) is still a reducer.

The figures for electrode potential show that lead is only slightly easier to oxidize than hydrogen. Lead thus can dissolve in acids, but this is often impossible due to specific problems (such as the formation of insoluble salts). Powdered lead burns with a bluish-white flame. As with many metals, finely divided powdered lead exhibits pyrophoricity. Toxic fumes are released when lead is burned.

Isotopes

Lead occurs naturally on Earth exclusively in the form of four isotopes: lead-204, −206, −207, and −208. All four can be radioactive as the hypothetical alpha decay of any would be exothermic, but the lower half-life limit has been put only for lead-204: over 1.4×10¹⁷ years. This effect is, however, so weak that natural lead poses no radiation hazard. Three isotopes are also found in three of the four major decay chains: lead-206, −207 and −208 are final decay products of uranium-238, uranium-235, and thorium-232, respectively. Since the amounts of them iature depend also on other elements’ presence, the isotopic composition of natural lead varies by sample: in particular, the relative amount of lead-206 varies between 20.84% and 27.78%.

Aside from the stable ones, thirty-four radioisotopes have been synthesized: they have mass numbers of 178–215. Lead-205 is the most stable radioisotope of lead, with a half-life of over 10⁷ years. 47 nuclear isomers (long-lived excited nuclear states), corresponding to 24 lead isotopes, have been characterized. The most long-lived isomer is lead-204m2 (half-life of about 1.1 hours).

Chemical reactivity

Lead is classified as a post-transition metal and is also a member of the carbon group. Lead only forms a protective oxide layer although finely powdered highly purified lead can ignite in air. Melted lead is oxidized in air to lead monoxide. All chalcogens oxidize lead upon heating.

Fluorine does not oxidize cold lead. Hot lead can be oxidized, but the formation of a protective halide layer lowers the intensity of the reaction above 100 °C (210 °F). The reaction with chlorine is similar: thanks to the chloride layer, lead persistence against chlorine surpasses those of copper or steel up to 300 °C (570 °F).

Water in the presence of oxygen attacks lead to start an accelerating reaction. The presence of carbonates or sulfates results in the formation of insoluble lead salts, which protect the metal from corrosion. So does carbon dioxide, as the insoluble lead carbonate is formed; however, an excess of the gas leads to the formation of the soluble bicarbonate; this makes the use of lead pipes dangerous. Lead dissolves in organic acids (in the presence of oxygen) and concentrated (≥80%) sulfuric acid thanks to complexation; however, it is only weakly affected by hydrochloric acid and is stable against hydrofluoric acid, as the corresponding halides are weakly soluble. Lead also dissolves in quite concentrated alkalis (≥10%) because of the amphoteric character and solubility of plumbites.

Compounds

Lead compounds exist mainly in two main oxidation states, +2 and +4. The former is more common. Inorganic lead(IV) compounds are typically strong oxidants or exist only in highly acidic solutions.

Oxides and sulfides

Three oxides are known: lead(II) oxide or lead monoxide (PbO), lead tetroxide (Pb₃O₄) (sometimes called “minium”), and lead dioxide (PbO₂). The monoxide exists as two allotropes: α-PbO and β-PbO, both with layer structure and tetracoordinated lead. The alpha polymorph is red-colored and has the Pb–O distance of 230 pm; the beta polymorph is yellow-colored and has the Pb–O distance of 221 and 249 pm (due to asymmetry). Both polymorphs can exist under standard conditions (beta with small (10⁻⁵ relative) impurities, such as Si, Ge, Mo, etc.). PbO reacts with acids to form salts, and with alkalis to give plumbites, [Pb(OH)₃]⁻ or [Pb(OH)₄]²⁻. The monoxide oxidizes in air to trilead tetroxide, which at 550 °C (1020 °F) degrades back into PbO.

The dioxide may be prepared by, for example, halogenization of lead(II) salts. Regardless the polymorph, it has a black-brown color. The alpha allotrope is rhombohedral, and the beta allotrope is tetragonal. Both allotropes are black-brown in color and always contain some water, which cannot be removed, as heating also causes decomposition (to PbO and Pb₃O₄). The dioxide is a powerful oxidizer: it can oxidize hydrochloric and sulfuric acids. It does not react with alkaline solution, but reacts with solid alkalis to give hydroxyplumbates, or with basic oxides to give plumbates.

Reaction of lead salts with hydrogen sulfide yields lead monosulfide. The solid has the rocksalt-like simple cubic structure, which it keeps up to the melting point, 1114 °C (2037 °F). When heated in air, it oxidizes to the sulfate and then the monoxide. Lead monosulfide is almost insoluble in water, weak acids, and (NH₄)₂S/(NH₄)₂S₂ solution is the key for separation of lead from analytical groups I to III ions, tin, arsenic, and antimony. However, it dissolves iitric and hydrochloric acids, to give elemental sulfur and hydrogen sulfide, respectively. Upon heating under high pressures with sulfur, it gives the disulfide. In the compound, the lead atoms are linked octahedrally with the sulfur atoms. It is also a semiconductor. A mixture of the monoxide and the monosulfide when heated forms the metal.

2 PbO + PbS → 3 Pb + SO₂

Halides and other salts

A 3 kg lead weight used on a scuba diving weight belt.

Heating lead carbonate with hydrogen fluoride yields the hydrofluoride, which decomposes to the difluoride when it melts. This white crystalline powder is more soluble than the diiodide, but less than the dibromide and the dichloride. The tetrafluoride, a yellow crystalline powder, is unstable.

Other dihalides are obtained upon heating lead(II) salts with the halides of other metals; lead dihalides precipitate to give white orthorhombic crystals (diiodide forms yellow hexagonal crystals). They can also be obtained by direct reaction of their constituent elements at temperature exceeding melting points of dihalides. Their solubility increases with temperature; adding more halides first decreases the solubility, but then increases due to complexation, with the maximum coordinatioumber being 6. The complexation depends on halide ioumbers, atomic number of the alkali metal, the halide of which is added, temperature and solution ionic strength. The tetrachloride is obtained upon dissolving the dioxide in hydrochloric acid; to prevent the exothermic decomposition, it is kept under concentrated sulfuric acid. The tetrabromide may not, and the tetraiodide definitely does not exist. The diastatide has also been prepared.

The metal is not attacked by sulfuric or hydrochloric acids. It dissolves iitric acid with the evolution of nitric oxide gas to form dissolved Pb(NO₃)₂. It is a well-soluble solid in water; it is thus a key to receive the precipitates of halides, sulfate, chromate, carbonate, and basic carbonate Pb₃(OH)₂(CO₃)₂ salts of lead.

Organolead

The best-known compounds are the two simplest plumbane derivatives: tetramethyllead (TML) and tetraethyllead (TEL). The homologs of these, as well as hexaethyldilead (HEDL), are of lesser stability. The tetralkyl derivatives contain lead(IV), where the Pb–C bonds are covalent. They thus resemble typical organic compounds.

Lead readily forms an equimolar alloy with sodium metal that reacts with alkyl halides to form organometallic compounds of lead such as tetraethyllead. The Pb–C bond energies in TML and TEL are only 167 and 145 kJ/mol; the compounds thus decompose upon heating, with first signs of TEL composition seen at 100 °C (210 °F). The pyrolysis yields of elemental lead and alkyl radicals; their interreaction causes the synthesis of HEDL. TML and TEL also decompose upon sunlight or UV light. In presence of chlorine, the alkyls begin to be replaced with chlorides; the R₂PbCl₂ in the presence of HCl (a by-product of the previous reaction) leads to the complete mineralization to give PbCl₂. Reaction with bromine follows the same principle.^[29]

Roman lead pipes often bore the insignia of Roman emperors (see Roman lead pipe inscriptions). Lead plumbing in the Latin West may have been continued beyond the age of Theoderic the Great into the medieval period. Many Roman “pigs” (ingots) of lead figure in Derbyshire lead mining history and in the history of the industry in other English centers. The Romans also used lead in molten form to secure iron pins that held together large limestone blocks in certain monumental buildings.^[36] In alchemy, lead was thought to be the oldest metal and was associated with the planet Saturn. Alchemists accordingly used Saturn’s symbol (the scythe, ♄) to refer to lead.^[37]

Up to the 17th century, tin was ofteot distinguished from lead: lead was called plumbum nigrum (literally, “black lead”), while tin was called plumbum candidum (literally, “bright lead”). Their inherence through history can also be seen in other languages: the word “olovo” means lead in Czech, but in Russian it (“олово”) means tin. Lead’s symbol Pb is an abbreviation of its Latiame plumbum for soft metals; the English words “plumbing”, “plumber”, “plumb”, and “plumb-bob” also derive from this Latin root.

Lead production in the US commenced as early as the late 1600s by Indians in the The Southeast Missouri Lead District, commonly called the Lead Belt, is a lead mining district in the southeastern part of Missouri. Significant among Missouri’s lead mining concerns in the district was the Desloge Family and Desloge Consolidated Lead Company in Desloge, Missouri and Bonne Terre – having been active in lead trading, mining and lead smelting from 1823 in Potosi to 1929.

Occurrence

Lead and zinc bearing carbonate and clastic deposits. Source: USGS

Metallic lead does occur iature, but it is rare. Lead is usually found in ore with zinc, silver and (most abundantly) copper, and is extracted together with these metals. The main lead mineral is galena (PbS), which contains 86.6% lead by weight. Other common varieties are cerussite (PbCO₃) and anglesite (PbSO₄).

Ore processing

Galena, lead ore

Most ores contain less than 10% lead, and ores containing as little as 3% lead can be economically exploited. Ores are crushed and concentrated by froth flotation typically to 70% or more. Sulfide ores are roasted, producing primarily lead oxide and a mixture of sulfates and silicates of lead and other metals contained in the ore. Lead oxide from the roasting process is reduced in a coke-fired blast furnace to the metal. Additional layers separate in the process and float to the top of the metallic lead. These are slag (silicates containing 1.5% lead), matte (sulfides containing 15% lead), and speiss (arsenides of iron and copper). These wastes contain concentrations of copper, zinc, cadmium, and bismuth that can be recovered economically, as can their content of unreduced lead.^[42]

Metallic lead that results from the roasting and blast furnace processes still contains significant contaminants of arsenic, antimony, bismuth, zinc, copper, silver, and gold. The melt is treated in a reverberatory furnace with air, steam, and sulfur, which oxidizes the contaminants except silver, gold, and bismuth. The oxidized contaminants are removed by drossing, where they float to the top and are skimmed off. Since lead ores contain significant concentrations of silver, the smelted metal also is commonly contaminated with silver. Metallic silver as well as gold is removed and recovered economically by means of the Parkes process Desilvered lead is freed of bismuth according to the Betterton-Kroll process by treating it with metallic calcium and magnesium, which forms a bismuth dross that can be skimmed off. Very pure lead can be obtained by processing smelted lead electrolytically by means of the Betts process. The process uses anodes of impure lead and cathodes of pure lead in an electrolyte of silica fluoride.

At current use rates, the supply of lead is estimated to run out in 42 years. Environmental analyst Lester Brown has suggested lead could run out within 18 years based on an extrapolation of 2% growth per year. This may need to be reviewed to take account of renewed interest in recycling, and rapid progress in fuel cell technology. According to the International Resource Panel’s Metal Stocks in Society report, the global per capita stock of lead in use in society is 8 kg. Much of this is in more-developed countries (20–150 kg per capita) rather than less-developed countries (1–4 kg per capita).

Applications

Elemental form

Lead bricks are commonly used as radiation shielding.

Contrary to popular belief, pencil leads in wooden pencils have never been made from lead. The term comes from the Roman stylus, called the penicillus, a small brush used for painting. When the pencil originated as a wrapped graphite writing tool, the particular type of graphite being used was named plumbago (lit. act for lead, or lead mockup).

Lead is used in applications where its low melting point, ductility and high density are advantageous. The low melting point makes casting of lead easy, and therefore small arms ammunition and shotgun pellets can be cast with minimal technical equipment. It is also inexpensive and denser than other common metals.

Because of its high density and resistance from corrosion, lead is used for the ballast keel of sailboats. Its high density allows it to counterbalance the heeling effect of wind on the sails while at the same time occupying a small volume and thus offering the least underwater resistance. For the same reason it is used in scuba diving weight belts to counteract the diver’s natural buoyancy and that of his equipment It does not have the weight-to-volume ratio of many heavy metals, but its low cost increases its use in these and other applications.

Roman lead water pipes with taps

Lead pipe in Roman baths

Multicolor lead-glazing in a Tang dynasty Chinese sancai ceramic cup dating from the 8th century CE

More than half of the US lead production (at least 1.15 million tonnes in 2000) is used for automobiles, mostly as electrodes in the lead–acid battery, used extensively as a car battery.

Cathode (reduction)

PbO₂ + 4 H⁺ + SO2−

4 + 2e^– → PbSO₄ + 2 H₂O

Anode (oxidation)

Pb + SO²⁻₄ → PbSO₄ + 2e^–

Lead is used as electrodes in the process of electrolysis. It is used in solder for electronics, although this usage is being phased out by some countries to reduce the amount of environmentally hazardous waste, and in high voltage power cables as sheathing material to prevent water diffusion into insulation. Lead is one of three metals used in the Oddy test for museum materials, helping detect organic acids, aldehydes, and acidic gases. It is also used as shielding from radiation (e.g., in X-ray rooms). Molten lead is used as a coolant (e.g., for lead cooled fast reactors).

Lead is added to brass to reduce machine tool wear. In the form of strips, or tape, lead is used for the customization of tennis rackets. Tennis rackets of the past sometimes had lead added to them by the manufacturer to increase weight. It is also used to form glazing bars for stained glass or other multi-lit windows. The practice has become less common, not for danger but for stylistic reasons. Lead, or sheet-lead, is used as a sound deadening layer in some areas in wall, floor and ceiling design in sound studios where levels of airborne and mechanically produced sound are targeted for reduction or virtual elimination. It is the traditional base metal of organ pipes, mixed with varying amounts of tin to control the tone of the pipe.

Lead has many uses in the construction industry (e.g., lead sheets are used as architectural metals in roofing material, cladding, flashing, gutters and gutter joints, and on roof parapets). Detailed lead moldings are used as decorative motifs used to fix lead sheet. Lead is still widely used in statues and sculptures. Lead is often used to balance the wheels of a car; this use is being phased out in favor of other materials for environmental reasons. Owing to its half-life of 22.20 years, the radioactive isotope ²¹⁰Pb is used for dating material from marine sediment cores by radiometric methods.

Compounds

Lead compounds are used as a coloring element in ceramic glazes, notably in the colors red and yellow. Lead is frequently used in polyvinyl chloride (PVC) plastic, which coats electrical cords.

Lead is used in some candles to treat the wick to ensure a longer, more even burn. Because of the dangers, European and North American manufacturers use more expensive alternatives such as zinc. Lead glass is composed of 12–28% lead oxide. It changes the optical characteristics of the glass and reduces the transmission of radiation.

Some artists using oil-based paints continue to use lead carbonate white, citing its properties in comparison with the alternatives. Tetra-ethyl lead is used as an anti-knock additive for aviation fuel in piston-driven aircraft. Lead-based semiconductors, such as lead telluride, lead selenide and lead antimonide are finding applications in photovoltaic (solar energy) cells and infrared detectors.

Lead, in either pure form or alloyed with tin, or antimony is the traditional material for bullets and shot in firearms use.

Former applications

Lead pigments were used in lead paint for white as well as yellow, orange, and red. Most uses have been discontinued due of the dangers of lead poisoning. Beginning April 22, 2010, US federal law requires that contractors performing renovation, repair, and painting projects that disturb more than six square feet of paint in homes, child care facilities, and schools built before 1978 must be certified and trained to follow specific work practices to prevent lead contamination. Lead chromate is still in industrial use. Lead carbonate (white) is the traditional pigment for the priming medium for oil painting, but it has been largely displaced by the zinc and titanium oxide pigments. It was also quickly replaced in water-based painting mediums. Lead carbonate white was used by the Japanese geisha and in the West for face-whitening make-up, which was detrimental to health.

Lead is the hot metal that was used in hot metal typesetting. It was used for plumbing (hence the name) as well as a preservative for food and drink in Ancient Rome. Until the early 1970s, lead was used for joining cast iron water pipes and used as a material for small diameter water pipes.

Tetraethyllead was used in leaded fuels to reduce engine knocking, but this practice has been phased out across many countries of the world in efforts to reduce toxic pollution that affected humans and the environment.

Lead was used to make bullets for slings. Lead was used for shotgun pellets in the US until about 1992 when it was outlawed (for waterfowl hunting only) and replaced by non-toxic shot, primarily steel pellets. In the Netherlands, the use of lead shot for hunting and sport shooting was banned in 1993, which caused a large drop in lead emission, from 230 tonnes in 1990 to 47.5 tonnes in 1995, two years after the ban.

Lead was a component of the paint used on children’s toys – now restricted in the United States and across Europe (ROHS Directive). Lead was used in car body filler, which was used in many custom cars in the 1940s–60s. Hence the term Leadsled. Lead is a superconductor with a transition temperature of 7.2 K, and therefore IBM tried to make a Josephson effect computer out of a lead alloy.

Lead was also used in pesticides before the 1950s, when fruit orchards were treated especially against the codling moth. A lead cylinder attached to a long line was used by sailors for the vital navigational task of determining water depth by heaving the lead at regular intervals. A soft tallow insert at its base allowed the nature of the sea bed to be determined, further aiding position finding.

Health effects

Lead is a highly poisonous metal (regardless if inhaled or swallowed), affecting almost every organ and system in the body. The main target for lead toxicity is the nervous system, both in adults and children. Long-term exposure of adults can result in decreased performance in some tests that measure functions of the nervous system. Long-term exposure to lead or its salts (especially soluble salts or the strong oxidant PbO₂) can cause nephropathy, and colic-like abdominal pains. It may also cause weakness in fingers, wrists, or ankles. Lead exposure also causes small increases in blood pressure, particularly in middle-aged and older people and can cause anemia. Exposure to high lead levels can severely damage the brain and kidneys in adults or children and ultimately cause death. In pregnant women, high levels of exposure to lead may cause miscarriage. Chronic, high-level exposure have shown to reduce fertility in males. Lead also damages nervous connections (especially in young children) and cause blood and brain disorders. Lead poisoning typically results from ingestion of food or water contaminated with lead; but may also occur after accidental ingestion of contaminated soil, dust, or lead-based paint. It is rapidly absorbed into the bloodstream and is believed to have adverse effects on the central nervous system, the cardiovascular system, kidneys, and the immune system. The component limit of lead (1.0 μg/g) is a test benchmark for pharmaceuticals, representing the maximum daily intake an individual should have. However, even at this low level, a prolonged intake can be hazardous to human beings. The treatment for lead poisoning consists of dimercaprol and succimer.

The concern about lead’s role in cognitive deficits in children has brought about widespread reduction in its use (lead exposure has been linked to learning disabilities). Most cases of adult elevated blood lead levels are workplace-related. High blood levels are associated with delayed puberty in girls. Lead has been shown many times to permanently reduce the cognitive capacity of children at extremely low levels of exposure.

During the 20th century, the use of lead in paint pigments was sharply reduced because of the danger of lead poisoning, especially to children. By the mid-1980s, a significant shift in lead end-use patterns had taken place. Much of this shift was a result of the U.S. lead consumers’ compliance with environmental regulations that significantly reduced or eliminated the use of lead ion-battery products, including gasoline, paints, solders, and water systems. Lead use is being further curtailed by the European Union’s RoHS directive. Lead may still be found in harmful quantities in stoneware, vinyl (such as that used for tubing and the insulation of electrical cords), and Chinese brass. Older houses may still contain substantial amounts of lead paint. White lead paint has been withdrawn from sale in industrialized countries, but the yellow lead chromate is still in use. Old paint should not be stripped by sanding, as this produces inhalable dust.

Lead salts used in pottery glazes have on occasion caused poisoning, when acidic drinks, such as fruit juices, have leached lead ions out of the glaze. It has been suggested that what was known as “Devon colic” arose from the use of lead-lined presses to extract apple juice in the manufacture of cider. Lead is considered to be particularly harmful for women’s ability to reproduce. Lead(II) acetate (also known as sugar of lead) was used in the Roman Empire as a sweetener for wine, and some consider this a plausible explanation for the dementia of many Roman emperors, and, that chronic lead poisoning contributed to the empire’s gradual decline. (see Decline of the Roman Empire#Lead poisoning).

Biochemistry of poisoning

In the human body, lead inhibits porphobilinogen synthase and ferrochelatase, preventing both porphobilinogen formation and the incorporation of iron into protoporphyrin IX, the final step in heme synthesis. This causes ineffective heme synthesis and subsequent microcytic anemia. At lower levels, it acts as a calcium analog, interfering with ion channels during nerve conduction. This is one of the mechanisms by which it interferes with cognition. Acute lead poisoning is treated using disodium calcium edetate: the calcium chelate of the disodium salt of ethylene-diamine-tetracetic acid (EDTA). This chelating agent has a greater affinity for lead than for calcium and so the lead chelate is formed by exchange. This is then excreted in the urine leaving behind harmless calcium. According to the Agency for Toxic Substance and Disease Registry, a small amount of ingested lead (1%) will store itself in bones, and the rest will be excreted by an adult through urine and feces within a few weeks of exposure. However, only about 32% of lead will be excreted by a child.

Exposure to lead and lead chemicals can occur through inhalation, ingestion and dermal contact. Most exposure occurs through ingestion or inhalation; in the U.S. the skin exposure is unlikely as leaded gasoline additives are no longer used. Lead exposure is a global issue as lead mining and lead smelting are common in many countries. Most countries have stopped using lead-containing gasoline by 2007. Lead exposure mostly occurs through ingestion. Lead paint is the major source of lead exposure for children. As lead paint deteriorates, it peels, is pulverized into dust and then enters the body through hand-to-mouth contact or through contaminated food, water or alcohol. Ingesting certain home remedy medicines may also expose people to lead or lead compounds. Lead can be ingested through fruits and vegetables contaminated by high levels of lead in the soils they were grown in. Soil is contaminated through particulate accumulation from lead in pipes, lead paint and residual emissions from leaded gasoline that was used before the Environment Protection Agency issue the regulation around 1980. The use of lead for water pipes is problematic in areas with soft or (and) acidic water. Hard water forms insoluble layers in the pipes while soft and acidic water dissolves the lead pipes. Inhalation is the second major pathway of exposure, especially for workers in lead-related occupations. Almost all inhaled lead is absorbed into the body, the rate is 20–70% for ingested lead; children absorb more than adults. Dermal exposure may be significant for a narrow category of people working with organic lead compounds, but is of little concern for general population. The rate of skin absorption is also low for inorganic lead.

 

Practical skills

Qualitative tests on borate and carbonate ions

1. B(OH)₃ + 3C₂H₅OH ® B(C₂H₅O)₃­ + 3H₂O – born-ethyl ester burns green flame      colour.

2. СаСО₃ + 2НCl ® СaCl₂ + CO₂­ + H₂O

CO₂­ + Ca(OH)₂ = СаСО₃¯ + H₂O – white precipitation

СаСО₃¯ + CO₂­ + H₂O = Са(HСО₃)₂ – colorless

 

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 IVA group elements. It’s cool stuff! Check it out!

www.youtube.com/watch?v=MJ3oCT_HMoE

www.youtube.com/watch?v=FF4T0PhBZ7A

www.youtube.com/watch?v=9Ds5aypNeXg

 

Prepared by PhD Falfushynska H.