GROUP 4A ELEMENTS [CARBON FAMILY]

Contents

• Introduction
• Group 14/4A elements
• Occurrence of Group 4A Elements
• Physical Properties of Group 4A Elements
• Chemical Properties of Group 4A Elements
• Carbon
• Silicon
• Germanium
• Tin
• Lead
• Flerovium
• Recommended Videos
• Test Questions
• Discuss And Explain

Learning Objectives

By the end of this section, you should be able to:

• Describe the physical and chemical properties of group 4A elements.
• Describe the allotropes of carbon.
• Describe glass, how it is made, and types of glass.

Introduction

I believe that this is not your first time to hear the word 'Periodic Table'. The Periodic Table contains all known elements; and these elements are arranged in columns and rows - technically known as groups and periods. 

Elements in the same group share a common characteristic - presence of the same number of electrons in their outermost or valence shell. On the other hand, elements in the same period have only the same number of electron shells. In our last discussion, we talked on elements in Group 13/3A, right now we're to discuss the elements in Group 14 or 4A of the periodic table.

Periodic table of elements

Group 14 / 4A Elements

Group 14 / 4A Elements are elements that have four electrons in their valence shell. A pair (of these valence electrons) can be found in the s orbital of the valence shell, while the remaining 2 electrons fill up the p orbitals; one is in p_x, and another, in p_y. Hence, group 14 / 4A elements usually have electronic configurations that end with ns² np², where n = period number.

As you can see, I used Group 14 / 4A; this is because this group follows right after group 13 that we discussed in the previous topic; I believe you know why group 13 is referred to as such. 

Group 14 of the periodic table is more often referred to as group 4A, because of the presence of four electrons in the valence shell of the elements present in the group.

The elements lose their 4 valence electrons to form compounds with +4 oxidation state, and moving down the group, the heavier elements start to rather form compounds with +2 oxidation state - I believe you know why? Yeah! Inert Pair effect. If you skipped the previous topic then this might be looking new to you.

Elements in Group 4A of the Periodic Table are: Carbon, Silicon, Germanium, Tin, Lead, and Flerovium.

Group 4A of the periodic table is sometimes referred to as The Carbon Family or Tetrels.

Occurrence of Group 4A Elements

Among all the elements of this group, carbon is the only one to occur in native state as diamond and graphite. Silicon is the second most abundant element in the earth’s crust. Silicates are present in rocks, clays are essentially alumino-silicates of Na and Ca, sand is an impure form of silica and Glass is also a mixture of silicates. Germanium is a rare element. It occurs in traces in coal, in rare mineral argyrodite, 4Ag₂S.GeS₂, in germanite, Cu₃(Ge.Fe)S₄, and as a mixture in zinc and tin ores. Tin occurs mainly as cassiterite or tin stone, SnO₂ and Lead occurs as galena, PbS. Flerovium on the other hand is radioactive.

Physical Properties of Group 4A Elements

• Group 4A elements are generally solids at room temperature.

• The metallic property of group 4A elements increases down the group.

• Ionic radii and atomic radii increases down the group (since there's addition of extra energy levels, or you can say new shells).

Chemical Properties of Group 4A Elements

• Group 4A elements tend to lose their 4 valence electrons. However, due to poor shielding of d and f electrons, heavier elements tend to lose only 2 electrons (which are available in the p subshell). The compounds formed (with +2 oxidation state) are however, unstable, until we reach Lead which has a stable, +2 oxidation state in its compounds, they are more stable than the +4 oxidation state's. This is due to inert Pair effect. This effect is commonly ascribed to the stability of an electron configuration with entirely filled subshells: in the inert pair effect, a metal loses all the p electrons in its outermost subshell, leaving a filled s² subshell; the pair of s electrons seems relatively “inert” and is less easily removed. The actual reasons for this effect are considerably more complex and will be understood better by CHM students in higher classes.

• Carbon undergoes many reactions, and can have an oxidation state of -4. By the way, carbon is regarded to as a non-metal, silicon and germanium are metalloids, while others beneath in the group are relatively metals.

• Carbon's valence electrons are held very tightly, close to the nucleus, leading to small size, thus, it is always hard for carbon to lose its valence electrons; carbon therefore, do undergo electron pairing to form covalent bonds and not ionic bonds. Silicon also forms covalent bonds, whereas other elements below in the group tend to form ionic bonds.

• The common coordination number of carbon is known to be four, however, there's always a misconception in many textbooks that this is carbon's maximum coordination number, whereas, many examples are known in which carbon has coordination numbers of 5, 6, or higher. 5-coordinate carbon is common, with methyl and other groups frequently forming bridges between two metal atoms, as in Al₂(CH₃)₆. There is also evidence for the 5-coordinate ion CH₅⁺. Other members of Group 4A can expand the octet due to their d orbitals. Thus, they can have higher coordination numbers and form complex ions e.g., SiF₆^2-, PbCl₄^2- etc.

• Catenation: Is a property by which an element bonds to itself (by single or multiple bond formation) to form stable chains, branches, and ring. E.g C–C bonds in hydrocarbons, Si–Si bonds in polysilicides.

The tendency to catenate decreases going down the group.

Carbon shows this tendency to a remarkable extent because of its small size and high bond energy. Si and Ge to a lesser extent and Sn and Pb to a very small extent.

• Ionization Energy: The first ionization energies of group 4A elements are higher than those of the corresponding group 3A elements, thus, it is harder for elements in group 4A to lose their valence electrons.

• Electronegativity: Group 4A elements are smaller in size when compared to group 3A elements, that’s why this group elements are slightly more electronegative (or hardly lose electron) than group 3A.

Carbon

Discovery

Carbon had been discovered since prehistory in several forms, but it was first identified as an element by Antoine Lavoisier during the 1700s. Carbon is reportedly said to be first isolated by H. Moissan, during late 1800s.

Occurrence

Carbon is found in nature as an element, in Earth’s atmosphere as carbon(IV) oxide gas, in Earth’s crust as mineral carbonates (e.g dolomite, magnesite, limestone), and in organic compounds produced in living cells. 

Carbon exists in several allotropic forms (allotrope = different physical structures of the same element).

Major Allotropes of carbon are:

Diamond

Diamond is a lustrous solid with a crystalline structure, composed of only carbon to carbon sigma bonds, which are arranged tetrahedrally within the diamond's structure. You can picture diamond's structure like a methane molecule, instead, all H (as in methane) are replaced by C, and are held tightly together. The carbon atoms in diamond are sp³ hybridized.

Internal structure of a diamond crystal

Since the structure of diamond possesses no free (mobile) electrons; it behaves as an insulator and it is inert to electricity.

The covalent bonds must also be broken before melting and boiling can occur, hence, diamond has an extremely high melting point (3550°C) and boiling point (4827°C).

Diamond is the hardest known naturally occurring substance, and this property makes it a good abrasive. Clear diamonds, occasionally in coloured form, are usually used as gems because of their high refractive index.

Diamond crystal | photo credit: iStock

Graphite

Graphite is a dull black, soft solid.

Graphite has a two dimensional layer structure. Its lattice consists of plane layers or sheets in which carbon atoms are arranged in regular planar hexagons.

Structure of graphite | photo credit: wikimedia commons

Graphite's Carbon atoms are sp² hybridized. Here, each carbon atom is linked to other three carbon atoms through 3 sigma bonds, and one leftover electron which has a pi bond overlap with others (remember what happens during sp² hybridization, recall from our discussion on hybridization). This unbonded electron (from each carbon atom) is delocalized (or free, mobile) over the whole planar layer.

These delocalized electrons make graphite a good conductor of electricity within the layer. 

Metallic lustre can also be attributed to these delocalized electrons.

Different layers are at a distance of 3.35Å from each other and are thus held by weak Van der Waals forces only. These layers can slide over one another easily, imparting softness and slippery touch to graphite. Hence it is used as a lubricant and in pencil.

A single isolated layer of graphite is called graphene, and its structure looks very much like a honey-comb.

Due to graphite's giant layer structure, it is very resistant to heat and hence used for crucibles.

It is chemically very inert, thus, it is employed as an electrode material in many electrolytic processes.

Graphite is known to be more stable than diamond at room temperature; Diamond is regarded to as metastable (=unstable, but without observable change). However, density of diamond is much more than of graphite. When graphite is subjected to high temperature and pressure, with the presence of metallic catalysts, it can be transformed into diamond.

Graphite | photo credit: istock

Fullerenes

Fullerenes are composed of carbon molecules, arranged in the form of a soccer ball shape. The first known fullerene, C₆₀, also called Bucky ball, or buckminsterfullerene, was discovered in 1985 by a team of 3 scientists: Sir Harold Kroto, Richard Smalley and Robert Curl by subjecting graphite to laser radiation at less than 10,000°C. The family of molecular fullerenes includes C₆₀, C₇₀, C₇₆, C₇₈, C₈₀ and C₈₄.

Bucky ball structure, C60 | photo credit: Wikimedia commons

The smallest known fullerene is C₂₀, synthesized by replacing the hydrogen atoms of dodecahedrane, C₂₀H₂₀, with bromines, followed by debromination.

Since 1985, fullerene chemistry has developed from its infancy into a broadly studied realm of science that is increasingly focused on practical applications. Fullerenes also form a variety of chemical compounds via reactions on their surfaces and can trap atoms and small molecules inside.

It was discovered in 2006 that a water-soluble C₆₀ derivative could be attached to a melanoma antibody; loading this derivative with anticancer drug molecules can deliver the drug directly into the melanoma. The use of fullerenes, nanotubes, quantum dots, and other nanoparticles for drug delivery has been reviewed.

Other chemistry of fullerenes are with organic chemistry, not to be discussed here.

Other allotropes of carbon are Lonsdaleite, C₅₄₀, single-walled carbon nanotube (also known as Bucky tube), glassy carbon, nanofoams, amorphous carbon (e.g soot, activated charcoal/carbon), and artificial carbon (e.g charcoal, coke, gas carbon, animal charcoal, carbon black etc)- artificial Carbon are referred to as amorphous but most of them show crystallinity.

Physical Properties

• Carbon exists in several allotropic forms (that we just discussed).

• Different allotropes of carbon have different densities. Amorphous carbon has a density of 1.8 – 2.1 g/cm³, graphite has a density of 2.27 g/cm³ and diamond has a density of 3.52 g/cm³.

• At atmospheric pressure, carbon does not have a melting point. It sublimes (=changes directly to vapour phase) at about 3650°C.

• Carbon burns with a bright orange flame.

Chemical Properties

• Carbon is known to form the most number of compounds in the universe, it is usually called the king of elements.

• Carbon has an atomic number of 6, and an atomic mass number of 12.

• Electronic configuration for Carbon in ground state is 1s² 2s² 2p² or [He] 2s² 2p².

• Carbon has the ability to catenate, form C–C bonds in its compounds, and this carbon property is one of the basis for which organic chemistry is founded.

• Carbon forms compounds covalently, by pairing its four valence electrons with other four electrons provided by other atoms or radicals. Carbon is known to be very obedient to the octet rule.

• Carbon is the only element in the group 4A to form stable derivatives with double and triple bonds, which show special characteristics.

• Carbon has a much higher electronegativity and higher ionisation energy than other elements in the group.

• Carbon forms stronger bonds not only with itself but also with elements like halogens, oxygen, nitrogen and sulphur than any other element in the group.

• Carbon not only reacts with non-metals, it behaves as a non-metal and reacts with metals.

Inorganic compounds of Carbon

• CO and CO₂, the most familiar oxides of carbon, are colourless and odourless gases. Carbon monoxide is a rarity of sorts, an unstable compound which is extremely toxic and hard to detect, forming a bright red complex with the iron in haemoglobin, which has a greater affinity for CO than for O₂, causing deaths like asphyxia. Carbon(IV) oxide is familiar as a component of Earth’s atmosphere—although only fifth in abundance, after nitrogen, oxygen, argon, and water vapor—and as the product of respiration, combustion, and other natural and industrial processes. It was the first gaseous component to be isolated from air, by Joseph Black in 1752. CO₂ has gained international attention because of its role in the greenhouse effect and the potential global warming and other climatic consequences of an increase in CO₂ abundance. Because of the energies of carbon(IV) oxide’s vibrational levels, it absorbs a significant amount of thermal energy and, hence, acts as a sort of atmospheric blanket. Since the beginning of the Industrial Revolution, the carbon(IV) oxide concentration in the atmosphere has increased substantially, an increase that will continue indefinitely unless major policy changes are made by the industrialized nations.

GreenHouse Effect: This is a term that describes the phenomenon whereby energy that is emitted by the sun (in form of light and radiation) reaches the earth's surface (in form of heat, light), which by method of radiation, warms up the (earth's) atmosphere, the energy is then re-emitted (from the earth's atmosphere) as heat (back to the earth's surface) by the greenhouse gas molecules (=methane, carbon(IV) oxide, dinitrogen oxide, water vapour).

Carbon monoxide is prepared by direct oxidation of Carbon in limited supply of oxygen.

$$2{\text{C}}_{(s)}+{\text{O}}_{2(g)}\to 2{\text{CO}}_{(g)}$$

It can also be prepared by dehydration of methanoic (or formic) acid. H₂SO₄ (which is a strong dehydrating agent) is commonly used:

$$\text{HCOOH}+{\text{H}}_2{\text{SO}}_4\to \text{CO}+{\text{H}}_2\text{O}$$

Commercially, it is prepared by the passage of steam over hot coke.

$${\text{H}}_2{\text{O}}_{(g)}+{\text{C}}_{(s)}\to {\text{CO}}_{(g)}+{\text{H}}_{2(g)}$$

Carbon(IV) oxide is prepared by complete combustion of carbon and carbon fuels in excess of air.

$${\text{C}}_{(s)}+{\text{O}}_{2(g)}\to {\text{CO}}_{2(g)}$$

In laboratory, it is prepared by the treatment of dilute HCl on CaCO₃.

$${\text{CaCO}}_{3(s)}+2{\text{HCl}}_{(aq)}\to {\text{CaCl}}_{2(aq)}+{\text{CO}}_{2(g)}+{\text{H}}_2{\text{O}}_{(l)}$$

Dry ice is the solid form of carbon(IV) oxide.

Dry ice pellet subliming | photo credit: Zephyris via wikimedia commons

It is a colourless, nonflammable solid that dissolves in water to form a weak acid.

We also have carbon suboxide, C₃O₂, dicarbon monoxide, C₂O, carbon trioxide, CO₃, etc

The oxides of carbon behave as weak acids.

• Carbon forms so many hydrides which are important in organic chemistry. We have the alkanes, alkenes, and alkynes.

• Carbon forms carbides with metals. There are basically 3 types of carbides:

Salt-like carbides, Interstitial carbides and Covalent carbides.

A Salt-like carbide is made by heating a more electropositive metal, its oxide or hydride with carbon, carbon monoxide or a hydrocarbon. Salt-like carbides are easily hydrolysed by water and they are classified according to the hydrocarbon they give:

Acetylides, e.g Ca₂, CrC₂, BaC₂, MgC₂ give acetylene when treated with water. CuC₂, Ag₂C₂, and AuC₂ are also formally known as acetylides, but are not hydrolysed by water.

$${\text{CaC}}_2+{\text{H}}_2\text{O}\to {\text{Ca(OH)}}_2+{\text{C}}_2{\text{H}}_2$$

Methanides, e.g Al₄C₃ and Be₂C yield methane on hydrolysis.

$${\text{Al}}_4{\text{C}}_3+12{\text{H}}_2\text{O}\to 4{\text{Al(OH)}}_3+3{\text{CH}}_4$$

Thorium and the lanthanides also form carbides. They have been reported as giving a mixture of acetylene, olefines and hydrogen on hydrolysis. Later experiments suggested that pure ThC₂ yields only acetylene, making it an acetylide.

Interstitial carbides are the carbides of the first 3 groups of the transition metals (with the exception of chromium). These carbides have metallic characters (e.g high melting, good electrical conduction, very hard) and are refractory. They are chemically inert except under oxidising conditions.

Silicon carbide and Boron carbide are majorly the ones referred to as covalent carbides, although virtually all compounds of carbon exhibit some covalent character. Covalent carbides are thermally stable, hard, chemically inert solids. Silicon carbide is widely used as an abrasive, while Boron carbide, which is harder, is used more as a radiation shield.

• Many of the transition metals of group 6B, 7B, and 8B form volatile, diamagnetic carbonyls in which the charge number of the metal is at zero. Examples are Ni(CO)₄, Fe(CO)₅, Cr(CO)₆, Mo(CO)₆

• Carbon forms oxohalides: COF₂ and COCl₂ are colourless gases made by the reaction of carbon monoxide and the halogen. The molecules appear to be trigonal planar, a form which suggests sp² hybridization. Carbonyl bromide COBr₂ is a colourless liquid best made by dropping concentrated H₂SO₄ onto CBr₄.

All these compounds are readily hydrolysed to form the hydrogen halides and carbon(IV) oxide.

$${\text{COX}}_2+{\text{H}}_2\text{O}\to 2\text{HX}+{\text{CO}}_2$$

• Carbon reacts with halogens to form carbon tetrahalides, CX₄, (where X= F, Cl, Br, I). The tetrahalides of carbon are relatively inert to hydrolysis, i.e, they don't decompose in water. 

• We have synthetic fluorocarbons (CF₄, C₂F₄, C₂F₆) which are resistant to attacks by acids, alkalis, oxidizing and reducing agents. Fluorocarbons are however attacked by hot metals like molten sodium.

• We also have chlorofluorocarbons, CFCs, (e.g CCl₂F₂, CFCl₃, CF₃Cl) which are volatile and chemically inert compounds.

• Carbon bonds with nonmetallic nitrogen to form C–N bonds. The most important species are cyanide, cyanate, and Thiocyanate ions. Hydrogen Cyanide (HCN) is a covalent, colourless and poisonous gas.

• Carbon also reacts with sulphur to yield carbonsulphide, or carbondisulphide.

• Carbon exists in the carbonate anion of compounds (HCO₃^-, CO₃^2-). E.g NaHCO₃, CaCO₃.

• There are millions of other carbon compounds in the world, especially organic compounds (discussed in organic chemistry).

Uses of Carbon

• Diamond is used as a gemstone in jewelleries, as a cutting and polishing tool for metals, and as an abrasive.

• Graphite is used as a lubricant, as an electrode for dry batteries, and together with clay, it is used to make the lead in pencils for writing and drawing.

• Coke is used as a reducing agent during extraction of metals.

• Coal is used as a fuel and source of energy.

• Carbon can form alloy with iron to form carbon steel.

• Carbon black is used as a pigment in printing ink, water colours, and carbon paper.

• Carbides of silicon, tungsten, boron and titanium, are among the hardest known materials, and are used as abrasives in cutting and grinding tools and some bulletproof materials.

• Activated carbon has enormous surface area and is used extensively in sugar industry as a decolourising agent. It is also used as an absorbent and adsorbent in filter materials.

Silicon

Discovery

Silicon's isolation was attempted by several scientists long time ago, however, credits for silicon's discovery is usually given to Jöns Jacob Berzelius.

Occurrence

Silicon is the second most abundant element in the earth's crust (after oxygen). It occurs in:

• Sand, SiO₂
• Olivine, Fe₂SiO₄
• Quartzite
• Pyrophyllite, AlSi₂O₅OH
• Andalusite, Al₂SiO₅
• And many others rocks.

Extraction

Silicon of high purity is produced by thermal reduction of silicon(IV) oxide with carbon.

Here, excess sand is reduced with pure coke in an electric arc furnace, in the presence of iron and small amounts of Phosphorus and sulphur.

Another method is thermal reduction of sand using Aluminium.

For semiconductor use, silicon is usually further purified by zone refining.

Physical Properties

• Silicon is a hard, brittle and lustrous crystalline solid. Pure silicon shines with a characteristic blue lustre.

Silicon | photo credit: wikimedia commons

• Silicon is grouped as a metalloid.

• Density of silicon is 2.3 g/cm³.

• Melting point of silicon is 1414°C and its boiling point is 3265°C.

• Silicon has a giant structure, similar to that of diamond.

Chemical Properties

• Silicon has an atomic number of 14 and a mass number of 28.

• The electronic configuration of silicon is 1s² 2s² 2p⁶ 3s² 3p² or [Ne] 3s² 3p²

• Silicon is more reactive than carbon, but it also forms covalent bonds.

• Silicon burns in oxygen at 400°C, the reaction being strongly exothermic. It combines directly with all the halogens at temperatures ranging from 300°C upwards, with sulphur vapour at 600°C, with nitrogen at 1300°C and with carbon at 2000°C.

• Silicon is acid resistant (except to HF), however, it is attacked by hot alkalis:

$$\text{Si}+\text{NaOH}+{\text{H}}_2\text{O}\to {\text{Na}}_2{\text{SiO}}_3+2{\text{H}}_2$$

and by steam at red heat:

$$\text{Si}+2{\text{H}}_2\text{O}\to {\text{SiO}}_2+2{\text{H}}_2$$

• Silicon exhibits catenation, however, Si–Si bonds are weaker than C–C bonds.

• Silicon forms mainly covalent compounds.

• Silicon makes use of its d orbitals to form compounds with coordination number of upto 6.

• Silicon forms stable tetravalent covalent compounds more often, the divalent variants are usually less stable and occur rarely.

• Silicon forms hydrides that are reactive and unstable. An example is Silane, SiH₄, which is always readily hydrolysed (reacts with water).

• Silicon forms a three dimensional covalent solid oxide, SiO₂, (as contrast to the simple gaseous oxide, CO₂, formed by carbon). SiO₂ behaves as an acid.

• Silicon dissolves readily in most metals (generally except bismuth, lead and thallium) to form silicides.

• Silicon forms tetrahalides with halogens. SiF₄ is a gas, SiCl₄ and SiBr₄ are liquids, and SiI₄ is a solid.

• Silicon tetrafluoride is made by treating a mixture of fluorite and silica with concentrated H₂SO₄:

$$2{\text{CaF}}_2+2{\text{H}}_2{\text{SO}}_4+{\text{SiO}}_2\to 2{\text{CaSO}}_4+{\text{SiF}}_4+2{\text{H}}_2\text{O}$$

The other silicon tetrahalides are made by direct combination with the halogen.

• Mixed tetrahalides such as SiF₃Cl and SiCl₂Br₂ have also been obtained.

• While carbon halides resist hydrolysis, silicon halides do hydrolyse:

$${\text{SiX}}_4+2{\text{H}}_2\text{O}\to {\text{SiO}}_2+4\text{HX}$$

However, when silicon fluoride undergo hydrolysis, some HF formed reacts with some of the tetrafluoride to form fluorosilicic acid.

2SiF₄ + 4H₂O $\to$ SiO₂ + 2H₃O⁺ + SiF₆^2- + 2HF

The octahedral SiF₆^2- ion is the only halogeno-complex of silicon; the bonding involves sp³d² hybrids and accordingly, carbon does not form such a compound. Fluorosilicic acid, H₂SiF₆, known only in solution, is a strong acid.

• Silicon, together with sulphur forms silicon disulphide, SiS₂.

• Silicon nitride, Si₃N₄, is a refractory material made by direct combination of Si and N above 1300°C. Another method is by reacting the hydrides or halides of silicon with ammonia and to heat the amino- and imino- silanes produced. These give the polymer [Si(NH)₂]_n which yields Si₃N₄.

• Zeolite: Zeolite is an alumino-silicate of metal. Metal cations participating in formation of Zeolite are usually Na⁺, K⁺, etc. Zeolites are usually used to remove permanent hardness of water.

• A silicone is a polymer composed of −R₂Si−O−SiR₂− where R = Organic group. They are usually colourless oils or rubber-like substances.

Uses of Silicon

• Silicon in its pure form is used as a semiconductor.

• Silica is used to make fire brick, a type of ceramic.

• Silicones are widely used as lubricants. Their inert nature makes them ideal for use in gas chromatography.

• Silicones are also used in breast implants, rubber accessories, contact lenses, space suits, explosives and pyrotechnics.

• Silicon is a main component in glass and cement.

Glass

Glass is a material with the internal structure of a liquid, but the hardness of a solid. In most solids, the particles are arranged in an orderly lattice, however, in glass, the molecules remain disordered, as in a liquid. The main ingredient in most glass is silica (SiO₂), which is one of the few substances that can cool without crystallizing.

Structure of a typical solid vs structure of glass

How is Glass made?

Glass is made by heating silica (or sand) to very high temperatures until it melts and turns into a liquid, the liquid is then cooled to form a frozen liquid (with a structure that is connected in a random manner), commonly referred to as an amorphous solid or glass.

However, because silica has a very high melting point, and its liquid form is difficult to control, a glass needs to have other components (which also help in diversity):

3 basic components make up a glass, these components are:

Formers: this is primarily the silica.

Fluxes: this lowers the melting point of silica. They form the various types of glasses.

Stabilisers: without a stabiliser, water and humidity would attack and dissolve glass. A stabiliser makes glass strong and water resistant.

Types of Glass

Silicate Glass or Quartz Glass

Silicate glass is formed by cooling melted, pure silica. It has low thermal expansion and excellent resistance to thermal shock. No flux is used here. Fused quartz is used for high-temperature applications such as furnace tubes, lighting tubes, crucibles, etc. However, its high melting temperature (1723°C) and viscosity make it difficult to work with.

Quartz glass | photo credit: wikimedia commons

Soda-Lime Glass

Sodium carbonate (Na₂CO₃) is used as a flux in the making of this type of glass. However, sodium silicate is soluble in water, which means this glass will destabilise on contact with water, hence, lime, or calcium oxide (CaO), generally obtained from limestone, is added to improve the stability of this glass. Soda-lime-silicate glass is transparent, easily formed, and most suitable for window, light bulb and table glasses. However, it expands when heated which makes it poorly resistant to heat.

Example of a soda lime glass

Borosilicate Glass

In this type of glass, Boron trioxide (B₂O₃) is used as a flux. Due to the hard nature and high heat resistance of Boron silicate, Borosilicate glasses are resistant to heat (more than soda-lime glass, but less than silicate glass). Thus, Borosilicate glass are less vulnerable to cracking from thermal shock. Popular examples of Borosilicate glass are Duran and Pyrex, and they are commonly used for laboratory equipments, kitchenwares and car head lamps.

Pyrex Borosilicate glass | photo credit: Wikimedia commons 

Lead Crystal Glass

Lead crystal glass contain lead(II) oxide (PbO) as a flux. Potassium oxide (K₂O), sodium oxide (Na₂O), alumina (Al₂O₃), and Zinc oxide (ZnO) are normally used as stabilisers. The melting point of lead glass are usually very low compared to the others discussed above, and this makes it easier to be formed into shape. Also, this glass has a high refractive index and high density, making it look brilliant and shiny.

Lead crystal glass

Alumino-silicate glass

Alumino-silicate glass contains Alumina (Al₂O₃) as a flux. The melting point of alumino-silicate mixture is higher than Borosilicate mixture, hence this type of glass is a bit harder to shape. Nevertheless, alumino-silicate glass is very resistant to heat and also durable. They are used in production of fibreglass, glass-reinforced plastics and halogen bulb glass.

Alumino-silicate glass

Coloured Glasses

Coloured glasses are made by the addition of coloured additives, transition metals with coloured ions are usually used.

Coloured glasses

Copper(I) oxide is a red compound, and it makes glass red coloured.

Iron oxide, FeO, makes glass green coloured.

Cobalt oxide, CoO, makes glass blue coloured.

Copper(II) oxide, CuO, makes blue and green coloured glass.

CaF₂ produces an opaque white (or milk coloured) glass.

Manganese compounds make glass black coloured.

Sulphur gives glass a blue colour.

Chromium compounds make glasses dark green coloured.

Titanium makes glass yellowish brown coloured.

Nickel makes glass green coloured.

Uranium makes glass yellow coloured.

Selenium makes glass bronze coloured.

Germanium

Discovery

Germanium was discovered in 1886 by Clemens Winkler, reportedly by analysing argyrodite.

Occurrence

• Argyrodite, Ag₈GeS₆
• Germanite,
• Impurity in other ores

Extraction

Germanium is mostly gotten as a by product during smelting of sphalerite, zinc ore. It is also found in silver, lead, and copper ores. The ore concentrates are usually mostly germanium sulphide, which is converted to germanium oxide by roasting:

$${\text{GeS}}_2+3{\text{O}}_2\to {\text{GeO}}_2+2{\text{SO}}_2$$

After leaching processes, GeO₂ is converted to germanium tetrachloride by action of chlorine gas or hydrochloric acid.

$${\text{GeO}}_2+4\text{HCl}\to {\text{GeCl}}_4+2{\text{H}}_2\text{O}$$

$${\text{GeO}}_2+2{\text{Cl}}_2\to {\text{GeCl}}_4+{\text{O}}_2$$

Germanium tetrachloride is either hydrolyzed to the oxide (GeO₂) or firstly purified by fractional distillation and then hydrolyzed. Purified GeO₂ is finally reduced to the elemental Ge by reduction with hydrogen.

$${\text{GeO}}_2+2{\text{H}}_2\to \text{Ge}+2{\text{H}}_2\text{O}$$

Carbon can also be used as a reducing agent:

$${\text{GeO}}_2+\text{C}\to \text{Ge}+{\text{CO}}_2$$

Physical Properties

• Germanium exists in form of two allotropes known as alpha germanium and beta germanium.

• α-germanium is the crystalline form of germanium which has a metallic luster and a diamond cubic crystal structure, the same as diamond. It is brittle, silvery-white coloured.

• At pressures above 120 Kbar, germanium becomes the allotrope β-germanium which has the same structure as β-tin.

• Germanium, like silicon and Gallium, expands when it changes from molten state to solid state.

• Germanium has a density of 5.32 g/cm³.

• The melting and boiling points of germanium are 940°C and 2833°C respectively.

• Germanium burns with a pale blue flame.

Chemical Properties

• Germanium has an atomic number of 32 and a mass number of 73.

• Electronic configuration of germanium is [Ar] 4s² 3d¹⁰ 4p²

• Germanium is more electropositive than silicon. It displays metallic character that are superior to those above it in the group.

• Germanium forms compounds in both 2-valent and 4-valent states. However, its tetravalent compounds are more stable than its divalent's.

• Germanium has no tendency to catenate.

• Germanium burns in oxygen at around 250°C to form GeO₂. 

• Germanium oxide behave as an amphoteric oxide.

• Germanium does not react with dilute acids or dilute alkalis. However, germanium dissolves slowly in hot concentrated H₂SO₄ and HNO₃.

• Germanium also reacts with molten alkalis to produce germanates [GeO₃]^2-.

• Germanium forms different hydrides with hydrogen, called germane (gas) which are very flammable and explosive when exposed to air.

• Germanium forms tetrahalides with halogens. While GeI₄ is a solid, GeF₄ is a gas and others are volatile liquids.All the tetrahalides are readily hydrolyzed to hydrated GeO₂.

• Germanium reacts with chalcogens (e.g sulphur, selenium, etc) to form germanium dichalcogenide e.g germanium disulphide, GeS₂, Germanium diselenide, GeSe₂, etc.

Uses of Germanium

Germanium found use mostly in the semiconductor industry.

Tin

Discovery

Tin was discovered during prehistoric times, there's no definite discoverer.

Occurrence

Tin occurs majorly as:

• Cassiterite, SnO₂
• Stannite, Cu₂FeSnS₄
• Cylindrite, PbSn₄FeSb₂S₁₄

Extraction

Tin is produced by thermal reduction of its oxide ore with carbon.

Physical Properties

• Tin exists in several physical states. We have alpha-tin, beta-tin, and even gamma and sigma tin.

• Beta tin is the most stable form of tin at room temperature, alpha tin is stable below room temperature, while gamma and sigma tin are stable well above room temperature.

• Beta tin appears to be a ductile, silvery-white solid, while alpha tin appears to be a grey powder.

• Pure Beta tin can transform to alpha tin under cold conditions.

Pure crystalline beta-tin

• Beta tin has a density of 7.27 g/cm³ while alpha tin has a density of 5.78 g/cm³.

• Melting and boiling points of metallic tin are 232°C and 2602°C respectively. When in powdered form, its melting point is lowered to 177°C.

• Beta tin has a body centered tetragonal crystal structure and is paramagnetic, while alpha tin has a face-centered diamond cubic crystal structure and is diamagnetic.

• Tin burns with a blue-white flame when heated in a hot Bunsen burner.

Chemical Properties

•Tin has an atomic number of 50 and a mass number of 119.

• The electronic configuration for tin is [Kr] 5s² 4d¹⁰ 5p²

• Tin exhibits both +2 and +4 oxidation states, while the +4 oxidation state is quite stable than the +2 oxidation state.

• Metallic tin does not readily oxidize in air or water due to formation of passivation layer, however, it is readily attacked by acids and alkalis. Tin oxide behave as an amphoteric oxide.

• Alloying elements such as copper, antimony, bismuth, cadmium, and silver increase the hardness and durability of beta tin, preventing transformation to alpha tin.

• Tin does not mix with most metals and elements. However, Tin mixes well with bismuth, gallium, lead, thallium and zinc.

• All four halides of Sn(IV) are known; SnF₄, SnCl₄, SnBr₄ and SnI₄; while SnI₄ is coloured, others are colourless, volatile compounds except SnF₄ which is polymeric.

• All four halides are also known for Sn(II), SnF₂, SnCl₂, SnBr₂ and SnI₂; they are colourless polymeric solids except the iodide which is also coloured.

• Tin(II) chloride is the most commercially important tin halide, and is gotten by reaction of hydrochloric acid and tin:

$$2\text{HCl}+\text{Sn}\to {\text{SnCl}}_2+{\text{H}}_2$$

Direct combination of Tin and chlorine yields more stable Tin(IV) chloride:

$$\text{Sn}+2{\text{Cl}}_2\to {\text{SnCl}}_4$$

However, Tin(IV) chloride can be converted to Tin(II) chloride by a process called comproportionation (= a chemical reaction in which two reactants, containing the same element in different oxidation state, forms a product in which the element balances its oxidation state).

$${\text{SnCl}}_4+\text{Sn}\to 2{\text{SnCl}}_2$$

• Tin oxide forms when powdered tin is heated in air, and like said earlier, SnO₂ is amphoteric, meaning it can dissolve and react in both acids and alkalis.

• Tin can have a coordination number of upto 8, because it can extend electrons down to its d-orbitals.

• Tin(II) sulphide (SnS) and Tin(IV) sulphide (SnS₂) are also known.

• Stannane (SnH₄) with tin in the +4 oxidation state is unstable. Other stable stannanes are known as organotin hydrides, involving tin, hydrogen and carbon bonds. These are colourless solids or liquids that are stable to air and water, and some can be toxic.

Uses of Tin

• Tin is widely used in making lead free solder wires used for joining pipes or electric circuits.

• Lead, zinc, steel, and sometimes copper, are plated with tin to prevent corrosion.

• Tin is most commonly alloyed with copper, to form pewter, bearing metal, bronze, bell metal, etc.

• Tin is used as a target to create laser-induced plasmas that act as the light source for extreme ultraviolet lithography.

• Tin is used as a negative electrode in advanced Li-ion batteries. Its application is limited due to the fact that some tin surfaces can catalyze decomposition of carbonate-based electrolytes used in Li-ion batteries.

• Tin(II) fluoride is added to some dental care products.

Lead

Discovery

Lead has been discovered since ancient times, there's no definite discoverer.

Occurrence

Lead occurs naturally as:

• Galena, PbS
• Pyromorphite, Pb₅(PO₄)₃Cl
• Boulangerite, Pb₅Sb₄S₁₁
• Anglesite, PbSO₄
• Cerrusite, PbCO₃

Extraction

Lead is extracted from its ore by reduction with coke. However, impure lead is gotten as a result, which can be purified by several pyrometallurgical processes or by electrometallurgy (–this is usually expensive and used only when mandatory).

PRO TIP: while some people may pronounce lead as /lid/, that is wrong, lead element is pronounced as /led/.

Physical Properties

• Lead is a silvery, very dense solid, which rapidly changes to dull grey when exposed to air.

Fresh shiny lead (right) and tarnished lead (left)

• Density of lead is 11.3 g/cm³.

• Lead is soft and malleable, and can even be cut with knife with sufficient force applied.

• Lead has a face centered cubic crystal structure and is diamagnetic.

• Lead has a melting point of about 330°C and a boiling point of 1749°C.

• Lead burns with a blue or blue-white flame when heated in a hot Bunsen burner.

Chemical Properties

• Lead has an atomic number of 82 and a mass number of 207.

• Lead has an electronic configuration of [Xe] 6s² 4f¹⁴ 5d¹⁰ 6p²

• Inert pair effect is exerted significantly in lead, this makes it that the +2 oxidation state of lead is more stable than the +4 oxidation state. This is due to the reluctance of s-electrons to unpair. Lead displays some unusual behaviour: the sum of its first and second ionization energies (=the total energy required to remove the two 6p electrons) is nearly same as that of tin (which, as expected, is supposed to reduce). This is due to the poor shielding effect of the f electrons, making the outer electrons to be held more tightly to the nucleus. Infact, the sum of the first four ionization energies of lead exceeds that of tin.

• The metallic character of lead is also drastically subdued, as lead has tendency of forming covalent bonds instead of ionic. Typically, lead is relatively unreactive and also called toxic.

• Lead can catenate and bond to itself to form chains and polyhedral structures.

• A chunk of lead metal, when exposed to air will form protective passivation layer and remain unreactive. However, when in powder form, lead tends to oxidize and burn.

• Chlorine reacts with lead on heating to form lead(II) chloride, PbCl₂, reaction of lead with fluorine can occur at room temperature, this forms lead(II) fluoride.

• Lead reacts very slowly with dilute mineral acids. Lead slowly reacts with sulphuric acid. It reacts with concentrated Nitric acid to form Pb(NO₃)₂, H₂O and oxides of nitrogen. Concentrated alkalis dissolve lead.

• Lead reacts with sulphur and selenium to form lead(II) sulphide and lead(II) selenide.

• Lead(II) oxide is amphoteric, and also insoluble in water.

• Lead can also form lead(II) cyanide, cyanate, and Thiocyanate.

• Lead does not form binary compounds with other metals. Lead hydrides, plumbane (PbH₄), are known to certain levels; very unstable gases.

Uses of Lead

• Lead is used in Car batteries.

• Lead is the main material in bullets, alloyed with other metals as hardeners.

• Lead is used as a protective sheath for underwater cable insulation.

• Lead is used in the production of lead crystal glass.

• Lead is used as a shielding material from radiation.

• Lead is present in alloys for making solder wires.

• Lead, together with other metals, is used in the production of organ pipes for music.

• Lead chromate is used for making yellow pigment for paint.

Flerovium

Flerovium is a synthetic element and has not been discovered naturally. It was successfully synthesized in December 1998 by a team of scientists in Russia.

Flerovium was made by bombarding Plutonium– 244 with Calcium– 48, in an accelerator.

One atom of Flerovium was reportedly produced during this process, and had a very short decay time (not upto a minute). Properties of Flerovium are predicted, it is believed that Flerovium's property would be the same as other elements above it in the group 4A, but there will be slight differences.

Recommended Videos

Group 4A Elements

Carbon

Silicon

Germanium

Tin

Lead

Flerovium

Glass

Test Questions

Discuss And Explain

• Explain Catenation. What element of group 14 exhibits maximum tendency for catenation?

• Give reason: C and Si are always tetravalent but Ge, Sn, Pb show divalency.

• Suggest a reason as to why CO is poisonous.

• What do you understand by:

1. Inert pair effect
2. Allotropy

• How is excessive content of CO₂ responsible for global warming?

• Describe the differences between structure of glass and a structure of a solid. Explain the formation of glass.

List and explain the different types of glass.