Biofelsefe — Karbon
NFA 2020 / Aziz Yardımlı

 

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Biofelsefe — KARBON


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📹 Why is carbon the element of life? (VİDEO)

📹 Why is carbon the element of life? (LINK)

Carbon is the element of life. But, out of 92 naturally occurring elements, what makes carbon essential for making organic molecules that ultimately form living organisms? In this video, learn where carbon comes from, its atomic structure, and the importance of functional groups.

 



📹 Hybrid Orbitals explained — Valence Bond Theory (VİDEO)

📹 Hybrid Orbitals explained — Valence Bond Theory (LINK)

This video explains the hybridization of carbon's, nitrogen's, and oxygen's valence orbitals in a bond, including single, double, and triple bonds. Explained are orbital overlap, sigma and pi bonds, and hybrid orbitals in carbon using methane, ethene, and ethyne; in nitrogen using ammonia; and in oxygen using water.

 



  Carbon — CHEMICAL ELEMENT (B)

🛑 KARBON

KARBON —
  • Karbon hücrelerin kuru ağırlığının yarısında daha çoğunu oluşturur.

 

  • Organik kimya C ve C arasındaki ya da C ve başka bir atom arasındaki bağ çevresinde örgütlenir.
  • Karbon ve başka atomlar arasındaki bağlar çoğunlukla kovalenttir.
  • Biomoleküllerin kovalent bağlı karbon atomları doğrusal zincirler, dallı zincirler ve döngülü yapılar oluşturabilir.
  • Karbondan başka hiçbir kimyasal element büyüklükte, şekilde ve bileşimde böylesine geniş bir türlülük gösteren moleküller oluşturamaz.
  • Kimyasal tepkimeler kimyasal bağların yapılmasına ve bozulmasına yol açan elektron aktarımlarını içerir.
 
 
   
Karbon dört-bağ sınırını aşabilir
Hexamethybenzen iki elektron yitirince düzlem hexagonal halkadan burada gösterilen beş kenarlı piramid yapıya geçer. Piramid bağlarının tepesindeki karbon atomu olağan dört atom yerine altı başka karbon atomu ile bağlanır. (Science News)
 
 
 

Karbon bağlarının geometrisi
(a) Karbon atomunun dört tekil bağının karakteristik tetrahedral düzenlemesi.
(b) Karbon—Karbon tekil bağları dönüş yapabilir (burada etan (CH3—CH3) bileşimi için gösterildiği gibi)
(c) İkili bağlar daha kısadır ve dönüşe izin vermezler. Çifte bağlı iki karbon ve A, B, X ve Y ile belirtilen atomlar aynı katı düzlem üzerinde yatar. (Lehninger Principles of Biochemistry, 2017.)


Karbon atomunun oluşturduğu dört tekli bağ çekirdekten bir tetrahedronun dört ucuna doğru uzanır. Herhangi iki bağ arasındaki açı 109,5 derece ve ortalama bağ uzunluğu 0,154 nanometredir ( 0,134 nm uzunluğundaki ikili bağ daha kısadır ve ekseni üzerinde sınırlı dönüşe izin verir).

 
 

Karbon bağlarının değişkenliği
Karbon özellikle başka karbon atomları ile kovalent, tekli, ikili ve üçlü bağlar oluşturabilir (kırmızı ile gösterilen bağlar). Biomoleküller durumunda üçlü bağlar seyrektir. (Lehninger Principles of Biochemistry, 2017.)

 




📹 The versatility of carbon — The tremendous variety of organic compounds on earth (VİDEO)

📹 The versatility of carbon — The tremendous variety of organic compounds on earth (LINK)

Why do the vast majority of compounds on earth contain carbon? This video looks at carbon's unique properties that give it its versatility with bonding and thus its enormous potential to create novel compounds both biologically and synthetically, as well as the shorthand of skeletal structures to show basic structures of complex molecules.

 



Carbon — CHEMICAL ELEMENT (B)

Carbon — CHEMICAL ELEMENT (B)

Carbon (C), nonmetallic chemical element in Group 14 (IVa) of the periodic table. Although widely distributed in nature, carbon is not particularly plentiful—it makes up only about 0.025 percent of Earth’s crust—yet it forms more compounds than all the other elements combined. In 1961 the isotope carbon-12 was selected to replace oxygen as the standard relative to which the atomic weights of all the other elements are measured. Carbon-14, which is radioactive, is the isotope used in radiocarbon dating and radiolabeling.


Carbon and its properties.
 
   
atomic number 6
atomic weight 12.0096 to 12.0116
melting point 3,550 °C (6,420 °F)
boiling point 4,827 °C (8,721 °F)
density
diamond 3.52 g/cm3
graphite 2.25 g/cm3
amorphous 1.9 g/cm3
oxidation states +2, +3, +4
electron configuration 1s22s22p2
 

 

 

 
Properties And Uses

Properties And Uses

Properties And Uses (W)


On a weight basis, carbon is 19th in order of elemental abundance in Earth’s crust, and there are estimated to be 3.5 times as many carbon atoms as silicon atoms in the universe. Only hydrogen, helium, oxygen, neon, and nitrogen are atomically more abundant in the cosmos than carbon. Carbon is the cosmic product of the “burning” of helium, in which three helium nuclei, atomic weight 4, fuse to produce a carbon nucleus, atomic weight 12.

In the crust of Earth, elemental carbon is a minor component. However, carbon compounds (i.e., carbonates of magnesium and calcium) form common minerals (e.g., magnesite, dolomite, marble, or limestone). Coral and the shells of oysters and clams are primarily calcium carbonate. Carbon is widely distributed as coal and in the organic compounds that constitute petroleum, natural gas, and all plant and animal tissue. A natural sequence of chemical reactions called the carbon cycle—involving conversion of atmospheric carbon dioxide to carbohydrates by photosynthesis in plants, the consumption of these carbohydrates by animals and oxidation of them through metabolism to produce carbon dioxide and other products, and the return of carbon dioxide to the atmosphere—is one of the most important of all biological processes.

Carbon as an element was discovered by the first person to handle charcoal from fire. Thus, together with sulfur, iron, tin, lead, copper, mercury, silver, and gold, carbon was one of the small group of elements well known in the ancient world. Modern carbon chemistry dates from the development of coals, petroleum, and natural gas as fuels and from the elucidation of synthetic organic chemistry, both substantially developed since the 1800s.

 
   
Bituminous coal.  
   

Elemental carbon exists in several forms, each of which has its own physical characteristics. Two of its well-defined forms, diamond and graphite, are crystalline in structure, but they differ in physical properties because the arrangements of the atoms in their structures are dissimilar. A third form, called fullerene, consists of a variety of molecules composed entirely of carbon. Spheroidal, closed-cage fullerenes are called buckerminsterfullerenes, or “buckyballs,” and cylindrical fullerenes are called nanotubes. A fourth form, called Q-carbon, is crystalline and magnetic. Yet another form, called amorphous carbon, has no crystalline structure. Other forms—such as carbon black, charcoal, lampblack, coal, and coke—are sometimes called amorphous, but X-ray examination has revealed that these substances do possess a low degree of crystallinity. Diamond and graphite occur naturally on Earth, and they also can be produced synthetically; they are chemically inert but do combine with oxygen at high temperatures, just as amorphous carbon does. Fullerene was serendipitously discovered in 1985 as a synthetic product in the course of laboratory experiments to simulate the chemistry in the atmosphere of giant stars. It was later found to occur naturally in tiny amounts on Earth and in meteorites. Q-carbon is also synthetic, but scientists have speculated that it could form within the hot environments of some planetary cores.

 
   
Fullerene
Two fullerene structures: an elongated carbon nanotube and a spherical buckminsterfullerene, or “buckyball.”
 
   
The word carbon probably derives from the Latin carbo, meaning variously “coal,” “charcoal,” “ember.” The term diamond, a corruption of the Greek word adamas, “the invincible,” aptly describes the permanence of this crystallized form of carbon, just as graphite, the name for the other crystal form of carbon, derived from the Greek verb graphein, “to write,” reflects its property of leaving a dark mark when rubbed on a surface. Before the discovery in 1779 that graphite when burned in air forms carbon dioxide, graphite was confused with both the metal lead and a superficially similar substance, the mineral molybdenite.

Pure diamond is the hardest naturally occurring substance known and is a poor conductor of electricity. Graphite, on the other hand, is a soft slippery solid that is a good conductor of both heat and electricity. Carbon as diamond is the most expensive and brilliant of all the natural gemstones and the hardest of the naturally occurring abrasives. Graphite is used as a lubricant. In microcrystalline and nearly amorphous form, it is used as a black pigment, as an adsorbent, as a fuel, as a filler for rubber, and, mixed with clay, as the “lead” of pencils. Because it conducts electricity but does not melt, graphite is also used for electrodes in electric furnaces and dry cells as well as for making crucibles in which metals are melted. Molecules of fullerene show promise in a range of applications, including high-tensile-strength materials, unique electronic and energy-storage devices, and safe encapsulation of flammable gases, such as hydrogen. Q-carbon, which is created by rapidly cooling a sample of elemental carbon whose temperature has been raised to 4,000 K (3,727 °C [6,740 °F]), is harder than diamond, and it can be used to manufacture diamond structures (such as diamond films and microneedles) within its matrix. Elemental carbon is nontoxic.

Each of the “amorphous” forms of carbon has its own specific character, and, hence, each has its own particular applications. All are products of oxidation and other forms of decomposition of organic compounds. Coal and coke, for example, are used extensively as fuels. Charcoal is used as an absorptive and filtering agent and as a fuel and was once widely used as an ingredient in gunpowder. (Coals are elemental carbon mixed with varying amounts of carbon compounds. Coke and charcoal are nearly pure carbon.) In addition to its uses in making inks and paints, carbon black is added to the rubber used in tires to improve its wearing qualities. Bone black, or animal charcoal, can adsorb gases and colouring matter from many other materials.

Carbon, either elemental or combined, is usually determined quantitatively by conversion to carbon dioxide gas, which can then be absorbed by other chemicals to give either a weighable product or a solution with acidic properties that can be titrated.

 



 
Nuclear Properties

Nuclear Properties

Nuclear Properties (B)


Carbon has two stable isotopes, carbon-12 (which makes up 98.93 percent of natural carbon) and carbon-13 (1.07 percent); 14 radioactive isotopes are known, of which the longest-lived is carbon-14, which has a half-life of 5,730 ± 40 years.

The notation used for the nucleus of atoms places the atomic mass as a presuperscript to the symbol of the element and the atomic number as a presubscript; thus, the isotope carbon-12 is symbolized 126C. Of the stable nuclides, the isotope carbon-13 is of particular interest in that its nuclear spin imparts response in a device called a nuclear magnetic resonance spectrometer, which is useful when investigating the molecular structures of covalently bonded compounds containing carbon. This isotope is also useful as a label in compounds that are to be analyzed by mass spectrometry, another device that is used extensively to identify atoms and molecules. Of the unstable nuclides, only carbon-14 is of sufficiently long half-life to be important. It is formed by the interaction of neutrons, produced by cosmic radiation, with nitrogen (N) in the atmosphere in a reaction that may be written as follows (neutron is symbolized as 10n, the nitrogen atom as 147N, and a hydrogen nucleus, or proton, as 11H):

   
 
   

The carbon-14 atoms from this reaction are converted to carbon dioxide by reaction with atmospheric oxygen and mixed and uniformly distributed with the carbon dioxide containing stable carbon-12. Living organisms use atmospheric carbon dioxide, whether with stable or radioactive carbon, through processes of photosynthesis and respiration, and thus their systems contain the constant ratio of carbon-12 to carbon-14 that exists in the atmosphere.

Death of an organism terminates this equilibration process; no fresh carbon dioxide is added to the dead substance. The carbon-14 present in the dead substance decays in accordance with its 5,730-year (± 40 years) half-life, while the carbon-12 remains what it was at death. Measurement of the carbon-14 activity at a given time thus allows calculation of the time elapsed after the death of the organism. Measurement of the carbon-14 activity in a cypress beam in the tomb of the Egyptian Pharaoh Snefru, for example, established the date of the tomb as circa 2600 BCE. Many other items of archaeological significance have been dated similarly (see carbon-14 dating).

The nuclides carbon-12 and carbon-13 are of importance in the CNO cycle of energy creation in certain stars. The cycle can be summarized in terms of nuclear equations, the separate steps being:

   
 
   
Summation of the equations allows the fusion process to be written as a reaction among four atoms of hydrogen to yield one atom of helium (He), two positrons (0+1e), and energy:
   
 
   
this equation does not show that the process uses up and regenerates the carbon-12. In a sense, carbon acts as a catalyst for this mode of converting mass to energy.

 



 
Compounds

Compounds

Compounds (B)


More than one million carbon compounds have been described in chemical literature, and chemists synthesize many new ones each year. Much of the diversity and complexity of organic forms is due to the capacity of carbon atoms for bonding with one another in various chain and ring structures and three-dimensional conformations as well as for linking with other atoms. Indeed, carbon’s compounds are so numerous, complex, and important that their study constitutes a specialized field of chemistry called organic chemistry, which derives its name from the fact that in the 19th century most of the then-known carbon compounds were considered to have originated in living organisms.

 
   
Carbon cycle
The generalized carbon cycle.
 
   

All organic compounds, such as proteins, carbohydrates, and fats, contain carbon, and all plant and animal cells consist of carbon compounds and their polymers. (Polymers are macromolecules consisting of many simple molecules bonded together in specific ways.) With hydrogen, oxygen, nitrogen, and a few other elements, carbon forms compounds that make up about 18 percent of all the matter in living things. The processes by which organisms consume carbon and return it to their surroundings constitute the carbon cycle.

Carbon is present as carbon dioxide in Earth’s atmosphere at a concentration of about 0.04 percent by volume, an amount that is increasing. Carbon dioxide is a greenhouse gas, and it is dissolved in all natural waters. Carbon occurs in the crust of Earth in the form of carbonates in such rocks as marble, limestone, and chalk and in hydrocarbons—the principal constituents of coal, petroleum, and natural gas. Carbonate minerals are important sources of various metals, such as sodium, magnesium, calcium, copper, and lead.

 
   
Keeling Curve
The Keeling Curve, named after American climate scientist Charles David Keeling, tracks changes in the concentration of carbon dioxide (CO2) in Earth's atmosphere at a research station on Mauna Loa in Hawaii. Although these concentrations experience small seasonal fluctuations, the overall trend shows that CO2 is increasing in the atmosphere.
 
   

At ordinary temperatures, carbon is very unreactive — it is difficult to oxidize — and it does not react with acids or alkalies. At high temperatures it combines with sulfur vapour to form carbon disulfide, with silicon and certain metals to form carbides, and with oxygen to form oxides, of which the most important are carbon monoxide, CO, and carbon dioxide, CO2. Because at high temperatures carbon combines readily with oxygen that is present in compounds with metals, large quantities of coke (an inexpensive form of carbon) are used in metallurgical processes to reduce (remove oxygen from) metal oxide ores, such as those of iron and zinc.

A type of chemical reaction in which one substance (an oxidizing agent) accepts electrons from another substance (a reducing agent) and is thereby reduced (while the reducing agent is oxidized) is frequently observed with carbon and its compounds. Although carbon is usually a reducing agent, under acidic conditions elemental carbon is a moderately strong oxidizing agent. The large energy of the carbon–carbon bond makes activation energy requirements for the reaction so high that direct reduction of carbon — e.g., to methane (formula CH4) — is impractical. Reduction of carbon monoxide to elemental carbon and oxidation of carbon monoxide to carbon dioxide are both feasible but impractical in solution. Under alkaline conditions, only the oxidation of formate ion (HCO2) to carbonate ion (CO32−) is a reasonable process.

Carbon monoxide (CO) is both more readily absorbed and more firmly bound to the hemoglobin of the blood than is oxygen and is thus, even in small concentrations, a dangerous asphyxiant. Carbon dioxide (CO2), however, is an asphyxiant of significance only in relatively large concentrations; in small concentrations, it stimulates breathing. Hydrogen cyanide (HCN) and its derivatives (cyanogen compounds, cyanides) are all very toxic as protoplasmic poisons through the inhibition of tissue oxidation. Carbon tetrachloride (CCl4) and other chlorinated hydrocarbons damage the nervous system. Among organic compounds the most toxic are derivatives that contain the halogen elements (fluorine, chlorine, bromine, and iodine), sulfur, selenium, tellurium, nitrogen, phosphorus, arsenic, lead, and mercury. Most organometallic compounds are toxic, while oxygen-containing derivatives of the hydrocarbons are usually less toxic.

 



 







 
  Carbon (W)


Diagram of the carbon cycle. The black numbers indicate how much carbon is stored in various reservoirs, in billions tonnes ("GtC" stands for gigatonnes of carbon; figures are circa 2004). The purple numbers indicate how much carbon moves between reservoirs each year. The sediments, as defined in this diagram, do not include the ≈70 million GtC of carbonate rock and kerogen.

📹 Geometry of carbon bonds (VİDEO)

📹 Geometry of carbon bonds (LINK)

Tetrahedral and trigonal planar bond geometries, conformation and rotation.

 



Carbon (W)

Carbon (W)

📂 DATA

Carbon, 6C

 



Graphite (left) and diamond (right), two allotropes of carbon
Carbon
Allotropes graphite, diamond, others
Appearance
  • graphite: black
  • diamond: clear
Standard atomic weight Ar, std(C) [ 12.0096, 12.0116] conventional: 12.011
Carbon in the periodic table
 
Atomic number (Z) 6
Group group 14 (carbon group)
Period period 2
Block p-block
Element category Reactive nonmetal, sometimes considered a metalloid
Electron configuration [He] 2s2 2p2
Electrons per shell 2, 4
Physical properties
Phase at STP solid
Sublimation point 3915 K ​(3642 °C, ​6588 °F)
Density (near r.t.) amorphous: 1.8–2.1 g/cm3
graphite: 2.267 g/cm3
diamond: 3.515 g/cm3
Triple point 4600 K, ​10,800 kPa
Heat of fusion graphite: 117 kJ/mol
Molar heat capacity graphite: 8.517 J/(mol·K)
diamond: 6.155 J/(mol·K)
Atomic properties
Oxidation states −4, −3, −2, −1, 0, +1, +2, +3, +4 (a mildly acidic oxide)
Electronegativity Pauling scale: 2.55
Ionization energies
  • 1st: 1086.5 kJ/mol
  • 2nd: 2352.6 kJ/mol
  • 3rd: 4620.5 kJ/mol
  • (more)
Covalent radius sp3: 77 pm
sp2: 73 pm
sp: 69 pm
Van der Waals radius 170 pm
Spectral lines of carbon
Other properties
Natural occurrence primordial
Crystal structure graphite: ​simple hexagonal (black)
 
Crystal structure diamond: ​face-centered diamond-cubic (clear)
 
Speed of sound thin rod diamond: 18,350 m/s (at 20 °C)
Thermal expansion diamond: 0.8 µm/(m·K) (at 25 °C)
Thermal conductivity graphite: 119–165 W/(m·K)
diamond: 900–2300 W/(m·K)
Electrical resistivity graphite: 7.837 µΩ·m
Magnetic ordering diamagnetic
Magnetic susceptibility −5.9·10−6 (graph.) cm3/mo
Young's modulus diamond: 1050 GPa
Shear modulus diamond: 478 GPa
Bulk modulus diamond: 442 GPa
Poisson ratio diamond: 0.1
Mohs hardness graphite: 1–2
diamond: 10
CAS Number
  • graphite: 7782-42-5
  • diamond: 7782-40-3
History
Discovery Egyptians and Sumerians (3750 BCE)
Recognized as an element by Antoine Lavoisier (1789)
Main isotopes of carbon
Iso­tope Abun­dance Half-life (t1/2) Decay mode Pro­duct
11C syn 20 min β+ 11B
12C 98.9% stable
13C 1.1% stable
14C trace 5730 y β− 14N

 




Structural formula of methane, the simplest possible organic compound.
 
   

Carbon (from Latin: carbo "coal") is a chemical element with the symbol C and atomic number 6. It is nonmetallic and tetravalent — making four electrons available to form covalent chemical bonds. It belongs to group 14 of the periodic table. Three isotopes occur naturally, 12C and 13C being stable, while 14C is a radionuclide, decaying with a half-life of about 5,730 years. Carbon is one of the few elements known since antiquity.

Carbon is the 15th most abundant element in the Earth's crust, and the fourth most abundant element in the universe by mass after hydrogen, helium, and oxygen. Carbon's abundance, its unique diversity of organic compounds, and its unusual ability to form polymersat the temperatures commonly encountered on Earth enables this element to serve as a common element of all known life. It is the second most abundant element in the human body by mass (about 18.5%) after oxygen.

The atoms of carbon can bond together in diverse ways, resulting in various allotropes of carbon. The best known allotropes are graphite, diamond, and buckminsterfullerene. The physical properties of carbon vary widely with the allotropic form. For example, graphite is opaque and black while diamond is highly transparent. Graphite is soft enough to form a streak on paper (hence its name, from the Greek verb "γράφειν" which means "to write"), while diamond is the hardest naturally occurring material known. Graphite is a good electrical conductor while diamond has a low electrical conductivity. Under normal conditions, diamond, carbon nanotubes, and graphene have the highest thermal conductivitiesof all known materials. All carbon allotropes are solids under normal conditions, with graphite being the most thermodynamically stable form at standard temperature and pressure. They are chemically resistant and require high temperature to react even with oxygen.

The most common oxidation state of carbon in inorganic compounds is +4, while +2 is found in carbon monoxideand transition metal carbonyl complexes. The largest sources of inorganic carbon are limestones, dolomites and carbon dioxide, but significant quantities occur in organic deposits of coal, peat, oil, and methane clathrates. Carbon forms a vast number of compounds, more than any other element, with almost ten million compounds described to date, and yet that number is but a fraction of the number of theoretically possible compounds under standard conditions. For this reason, carbon has often been referred to as the “king of the elements.”

 
Characteristics

Characteristics

Characteristics (W)

The allotropes of carbon include graphite, one of the softest known substances, and diamond, the hardest naturally occurring substance. It bonds readily with other small atoms, including other carbon atoms, and is capable of forming multiple stable covalent bonds with suitable multivalent atoms. Carbon is known to form almost ten million compounds, a large majority of all chemical compounds. Carbon also has the highest sublimation point of all elements. At atmospheric pressure it has no melting point, as its triple point is at 10.8±0.2 MPa and 4,600 ± 300 K (4,330 ± 300 °C; 7,820 ± 540 °F), so it sublimes at about 3,900 K (3,630 °C; 6,560 °F). Graphite is much more reactive than diamond at standard conditions, despite being more thermodynamically stable, as its delocalised pi system is much more vulnerable to attack. For example, graphite can be oxidised by hot concentrated nitric acid at standard conditions to mellitic acid, C6(CO2H)6, which preserves the hexagonal units of graphite while breaking up the larger structure.

Carbon sublimes in a carbon arc, which has a temperature of about 5800 K (5,530 °C or 9,980 °F). Thus, irrespective of its allotropic form, carbon remains solid at higher temperatures than the highest-melting-point metals such as tungsten or rhenium. Although thermodynamically prone to oxidation, carbon resists oxidation more effectively than elements such as iron and copper, which are weaker reducing agents at room temperature.

Carbon is the sixth element, with a ground-state electron configuration of 1s22s22p2, of which the four outer electrons are valence electrons. Its first four ionisation energies, 1086.5, 2352.6, 4620.5 and 6222.7 kJ/mol, are much higher than those of the heavier group-14 elements. The electronegativity of carbon is 2.5, significantly higher than the heavier group-14 elements (1.8–1.9), but close to most of the nearby nonmetals, as well as some of the second- and third-row transition metals. Carbon's covalent radii are normally taken as 77.2 pm (C−C), 66.7 pm (C=C) and 60.3 pm (C≡C), although these may vary depending on coordination number and what the carbon is bonded to. In general, covalent radius decreases with lower coordination number and higher bond order.

Carbon compounds form the basis of all known life on Earth, and the carbon–nitrogen cycle provides some of the energy produced by the Sun and other stars. Although it forms an extraordinary variety of compounds, most forms of carbon are comparatively unreactive under normal conditions. At standard temperature and pressure, it resists all but the strongest oxidizers. It does not react with sulfuric acid, hydrochloric acid, chlorine or any alkalis. At elevated temperatures, carbon reacts with oxygen to form carbon oxides and will rob oxygen from metal oxides to leave the elemental metal. This exothermic reaction is used in the iron and steel industry to smelt iron and to control the carbon content of steel:

Fe3O4 + 4 C(s) → 3 Fe(s) + 4 CO(g)

 

Carbon monoxide can be recycled to smelt even more iron:

Fe3O4 + 4 CO(g) → 3 Fe(s) + 4 CO2(g)

 

with sulfur to form carbon disulfide and with steam in the coal-gas reaction:

C(s) + H2O(g) → CO(g) + H2(g).

 

Carbon combines with some metals at high temperatures to form metallic carbides, such as the iron carbide cementite in steel and tungsten carbide, widely used as an abrasive and for making hard tips for cutting tools.

 
Graphite is one of the softest materials known. Synthetic nanocrystalline diamond is the hardest material known.
Graphite is a very good lubricant, displaying superlubricity. Diamond is the ultimate abrasive.
Graphite is a conductor of electricity. Diamond is an excellent electrical insulator, and has the highest breakdown electric field of any known material.
Some forms of graphite are used for thermal insulation (i.e. firebreaks and heat shields), but some other forms are good thermal conductors. Diamond is the best known naturally occurring thermal conductor
Graphite is opaque. Diamond is highly transparent.
Graphite crystallizes in the hexagonal system. Diamond crystallizes in the cubic system.
Amorphous carbon is completely isotropic. Carbon nanotubes are among the most anisotropic materials known.

 



 
Compounds

Organic compounds

Organic compounds (W)

Carbon can form very long chains of interconnecting carbon–carbon bonds, a property that is called catenation. Carbon-carbon bonds are strong and stable. Through catenation, carbon forms a countless number of compounds. A tally of unique compounds shows that more contain carbon than do not. A similar claim can be made for hydrogen because most organic compounds contain hydrogen chemically bonded to carbon or another common element like oxygen or nitrogen.

The simplest form of an organic molecule is the hydrocarbon — a large family of organic molecules that are composed of hydrogen atoms bonded to a chain of carbon atoms. A hydrocarbon backbone can be substituted by other atoms, known as heteroatoms. Common heteroatoms that appear in organic compounds include oxygen, nitrogen, sulfur, phosphorus, and the nonradioactive halogens, as well as the metals lithium and magnesium. Organic compounds containing bonds to metal are known as organometallic compounds (see below). Certain groupings of atoms, often including heteroatoms, recur in large numbers of organic compounds. These collections, known as functional groups, confer common reactivity patterns and allow for the systematic study and categorization of organic compounds. Chain length, shape and functional groups all affect the properties of organic molecules.

In most stable compounds of carbon (and nearly all stable organic compounds), carbon obeys the octet rule and is tetravalent, meaning that a carbon atom forms a total of four covalent bonds (which may include double and triple bonds). Exceptions include a small number of stabilized carbocations (three bonds, positive charge), radicals (three bonds, neutral), carbanions (three bonds, negative charge) and carbenes (two bonds, neutral), although these species are much more likely to be encountered as unstable, reactive intermediates.

Carbon occurs in all known organic life and is the basis of organic chemistry. When united with hydrogen, it forms various hydrocarbons that are important to industry as refrigerants, lubricants, solvents, as chemical feedstock for the manufacture of plastics and petrochemicals, and as fossil fuels.

When combined with oxygen and hydrogen, carbon can form many groups of important biological compounds including sugars, lignans, chitins, alcohols, fats, and aromatic esters, carotenoids and terpenes. With nitrogen it forms alkaloids, and with the addition of sulfur also it forms antibiotics, amino acids, and rubber products. With the addition of phosphorus to these other elements, it forms DNA and RNA, the chemical-code carriers of life, and adenosine triphosphate (ATP), the most important energy-transfer molecule in all living cells.

 


Correlation between the carbon cycle and formation of organic compounds. In plants, carbon dioxide formed by carbon fixation can join with water in photosynthesis (green) to form organic compounds, which can be used and further converted by both plants and animals.

 







 
  Carbon bonding (B)

Carbon bonding

Carbon bonding (B)

The carbon atom is unique among elements in its tendency to form extensive networks of covalent bonds not only with other elements but also with itself. Because of its position midway in the second horizontal row of the periodic table, carbon is neither an electropositive nor an electronegative element; it therefore is more likely to share electrons than to gain or lose them. Moreover, of all the elements in the second row, carbon has the maximum number of outer shell electrons (four) capable of forming covalent bonds. (Other elements, such as phosphorus [P] and cobalt [Co], are able to form five and six covalent bonds, respectively, with other elements, but they lack carbon’s ability to bond indefinitely with itself.) When fully bonded to other atoms, the four bonds of the carbon atom are directed to the corners of a tetrahedron and make angles of about 109.5° with each other (see chemical bonding: Bonds between atoms). The result is that not only can carbon atoms combine with one another indefinitely to give compounds of extremely high molecular weight, but the molecules formed can exist in an infinite variety of three-dimensional structures. The possibilities for diversity are increased by the presence of atoms other than carbon in organic compounds, especially hydrogen (H), oxygen (O), nitrogen (N), halogens (fluorine [F], chlorine [Cl], bromine [Br], and iodine [I]), and sulfur (S). It is the enormous potential for variation in chemical properties that has made organic compounds essential to life on Earth.

The structures of organic compounds commonly are represented by simplified structural formulas, which show not only the kinds and numbers of atoms present in the molecule but also the way in which the atoms are linked by the covalent bonds—information that is not given by simple molecular formulas, which specify only the number and type of atoms contained in a molecule. (With most inorganic compounds, the use of structural formulas is not necessary, because only a few atoms are involved and only a single arrangement of the atoms is possible.) In the structural formulas of organic compounds, short lines are used to represent the covalent bonds. Atoms of the individual elements are represented by their chemical symbols, as in molecular formulas.

Structural formulas vary widely in the amount of three-dimensional information they convey, and the type of structural formula used for any one molecule depends on the nature of the information the formula is meant to display. The different levels of sophistication can be illustrated by considering some of the least complex organic compounds, the hydrocarbons. The gas ethane, for example, has the molecular formula C2H6. The simplest structural formula, drawn either in a condensed or in an expanded version, reveals that ethane consists of two carbon atoms bonded to one another, each carbon atom bearing three hydrogen atoms. Such a two-dimensional representation correctly shows the bonding arrangement in ethane, but it does not convey any information about its three-dimensional architecture. A more sophisticated structural formula can be drawn to better represent the three-dimensional structure of the molecule. Such a structural formula correctly shows the tetrahedral orientation of the four atoms (one carbon and three hydrogens) bonded to each carbon, and the specific architecture of the molecule.

Larger organic molecules are formed by the addition of more carbon atoms. Butane, for example, is a gaseous hydrocarbon with the molecular formula C4H10, and it exists as a chain of four carbon atoms with 10 attached hydrogen atoms. As carbon atoms are added to a molecular framework, the carbon chain can develop branches or form cyclic structures. A very common ring structure contains six carbon atoms in a ring, each bonded in a tetrahedral arrangement, as in the hydrocarbon cyclohexane, C6H12. Such ring structures are often very simply represented as regular polygons in which each apex represents a carbon atom, and the hydrogen atoms that complete the bonding requirements of the carbon atoms are not shown. The polygon convention for cyclic structures reveals concisely the bonding arrangement of the molecule but does not explicitly convey information about the actual three-dimensional architecture. It should be noted that the polygon is only a two-dimensional symbol for the three-dimensional molecule.

Under certain bonding conditions, adjacent atoms will form multiple bonds with each other. A double bond is formed when two atoms use two electron pairs to form two covalent bonds; a triple bond results when two atoms share three electron pairs to form three covalent bonds. Multiple bonds have special structural and electronic features that generate interesting chemical properties. The six atoms involved in a double bond (as in ethene, C2H4) lie in a single plane, with regions above and below the plane occupied by the electrons of the second covalent bond. Atoms in a triple bond (as in acetylene, or ethyne, C2H2) lie in a straight line, with four regions beside the bond axis occupied by electrons of the second and third covalent bonds.

 







 
  Carbon group element (B)

Carbon group element

Carbon group element (B)

Introduction (B)

Carbon group element, any of the six chemical elements that make up Group 14 (IVa) of the periodic table—namely, carbon (C), silicon (Si), germanium (Ge), tin (Sn), lead (Pb), and flerovium (Fl).

Except for germanium and the artificially produced flerovium, all of these elements are familiar in daily life either as the pure element or in the form of compounds, although, except for silicon, none is particularly plentiful in the Earth’s crust. Carbon forms an almost infinite variety of compounds, in both the plant and animal kingdoms. Silicon and silicate minerals are fundamental components of the Earth’s crust; silica (silicon dioxide) is sand. Tin and lead, with abundances in the crust lower than those of some so-called rare elements, are nevertheless common in everyday life. They occur in highly concentrated mineral deposits, can be obtained easily in the metallic state from those minerals, and are useful as metals and as alloys in many applications. Germanium, on the other hand, forms few characteristic minerals and is most commonly found only in small concentrations in association with the mineral zinc blende and in coals. Although germanium is indeed one of the rarer elements, it assumed importance upon recognition of its properties as a semiconductor (i.e., limited ability to conduct electricity).

 


Modern version of the periodic table of the elements.
 
Comparative Chemistry

Comparative Chemistry

Comparative Chemistry (B)


In the periodic table, the elements with eight electrons outermost form the group known as the noble gases (Group 18 [0]), the least reactive of the elements. The carbon group elements (Group 14), with four electrons, occupy a middle position. Elements to the left of Group 14 have fewer than four electrons in the valence shell and tend to lose them (with their negative charges) to become positively charged ions, represented by the symbol for the element with a superscript indicating the number and sign of the charges; such elements are called metals. The nonmetals (except boron) are in the groups to the right of Group 14; each has more than four electrons in its outermost shell and tends to acquire electrons to complete its octet, forming negatively charged ions.

Chemical reactions result from the exchange of electrons among atoms. In general, if a metal loses its few valence electrons to a nonmetal, the resulting oppositely charged ions are attracted to one another and form a bond, classified as ionic or electrovalent. Two nonmetals, neither of which can actually lose its valence electrons in chemical reaction, may nevertheless share them in pairs in such a way that what is called a covalent bond results. Metal atoms will bond to one another in a third type of bond, which releases their valence electrons in a way that allows them to conduct electricity.

All the carbon group atoms, having four valence electrons, form covalent bonds with nonmetal atoms; carbon and silicon cannot lose or gain electrons to form free ions, whereas germanium, tin, and lead do form metallic ions but only with two positive charges. Even lead, the most metallic of the carbon group atoms, cannot actually lose all four of its valence electrons, because, as each one is removed, the remainder are held more strongly by the increased positive charge. Because the distinction between covalent and ionic (electrovalent) bonds is often a matter of convenience for the chemist, and because the actual bond structure within a molecule may be quite complicated, it is often useful instead simply to count the total number of electrons an element gains or loses in bonding without regard to the nature of the bonds. This number is called the oxidation number, or oxidation state, of the element; many elements have more than one oxidation state possible, each oxidation state being found in different compounds. The oxidation state of an element is conventionally written as a Roman numeral following the name of the element in a compound—for example, lead(II) means lead in the +2 oxidation state. An alternative system of representation uses an Arabic number after the element name; thus, lead in the +2 state is written lead(+2). With the chemical symbol of the element, the oxidation state may be written as a superscript, as in Pb2+. When the compounds are ionic, the oxidation state is also the actual ionic charge. Covalent bonds generally are considered to be formed by interaction of the orbitals (in most cases, only the s, p, and d orbitals) in specific and varied ways. The most common are called sigma and pi bonds, written σ and π, respectively. The sigma bonds are symmetrical with respect to the axis of the bond, whereas the pi bonds are not. Examples of sigma and pi bonding as well as of ionic bonding can be found among the compounds of the elements of the carbon group.

 



 
 

General Properties Of The Group

General Properties Of The Group (B)


The properties of the carbon group elements and those of their compounds are intermediate between properties associated with the elements of the adjacent boron and nitrogen groups. In all groups the metallic properties, resulting from the tendency to hold valence electrons more loosely, increase with atomic number. Within the carbon group, more than in any other, the change from nonmetallic to metallic character with increasing atomic number is particularly apparent. Carbon is a true nonmetal in every sense. Lead is a true metal. Silicon is almost completely nonmetallic; tin is almost completely metallic. Germanium is metallic in appearance and in a number of its other physical properties (see Table), but the properties of many of its compounds are those of derivatives of nonmetals. These changes are consequences of increase in atomic size with substantial screening of the larger nuclear charge by intervening electronic shells, as evidenced by decrease in ionization energy (energy required to remove an electron) and electronegativity power to attract electrons with increasing atomic number.

 
Some properties of the carbon group elements


carbon
silicon germanium tin lead
atomic number 6 14 32 50 82
atomic weight 12.011 28.086 72.64 118.71 207.2
colour of element colourless (diamond), black (graphite) gray white metallic white metallic (beta), gray (alpha) bluish white metallic
melting point (°C) 3,700 1,414 938.25 231.93 327.5
boiling point (°C) 4,027 3,265 2,833 2,602 1,749
density (grams per cubic centimetre) 1.9–2.3 (graphite), 3.15–3.53 (diamond) 2.33 (25 °C) 5.32 (25 °C) 5.75 (alpha), 7.31 (beta) 11.35
oxidation states −4, (+2), +4 −4, (+2), +4 −4, +2, +4 (−4), +2, +4 (−4), +2, +4
mass number of most common isotopes (terrestrial abundance, percent) 12 (98.89), 13 (1.11) 28 (92.23), 29 (4.68), 30 (3.09) 70 (20.84), 72 (27.54), 73 (7.73), 74 (36.28), 76 (7.61) 112 (0.97), 114 (0.66), 115 (0.34), 116 (14.54), 117 (7.68), 118 (24.22), 119 (8.59), 120 (32.58), 122 (4.63), 124 (5.79) 204 (1.4), 206 (24.1), 207 (22.1), 208 (52.4)
radioactive isotopes (mass numbers) 8–11, 14–22 22–27, 31–44 60–69, 71, 75–89 100–111, 113, 121, 123, 125–137 181–205, 209–215
heat of fusion (calories per mole/kilojoules per mole) 25,100 (105) 12,000 (50.2) 7,600 (31.8) 1,700 (7) 1,140 (4.77)
heat of vaporization (kilojoules per mole) 715 359 334 290 178
heat of sublimation (kilocalories per gram atom) 170 85 78 47.5
heat capacity (joules per gram Kelvin) 0.709 0.712 0.32 0.227 0.13
critical temperature (°C) about 4,920
critical pressure (atmospheres) 1,450
electrical resistivity (microhm-centimetres) 1,375 10 4.6 × 107 11 20.648
hardness (Mohs scale) 0.5 6.5 6 1.5 1.5
crystal structure cubic (diamond), hexagonal (graphite) cubic cubic cubic, tetragonal close-packed, metallic
radius: covalent (angstroms) 0.76 1.11 1.2 1.39 1.46
radius: ionic (angstroms) 0.3 0.54 0.67 0.83 0.92
ionization energy (kilojoules per mole): first 1,086.50 786.5 762 708.6 715.6
ionization energy (kilojoules per mole): second 2,352.60 1,577.10 1,537.50 1,411.80 1,450.50
ionization energy (kilojoules per mole): third 4,620.50 3,231.60 3,302.10 2,943.00 3,081.50
ionization energy (kilojoules per mole): fourth 6,222.70 4,355.50 4,411 3,930.30 4,083
electronegativity (Sanderson) 2.75 2.14 2.62 1.49 2.29
electronegativity (Pauling) 2.55 1.9 2.01 1.96 2.33

 



Crystal structure

Crystal structure (B)

In the solid state, elemental carbon, silicon, germanium, and gray tin (defined as alpha [α] tin) exist as cubic crystals, based upon a three-dimensional arrangement of bonds. Each atom is covalently bonded to four neighbouring atoms in such a way that they form the corners of a tetrahedron (a solid consisting of four three-sided faces). A practical result is that no discrete small molecules of these elements, such as those formed by nitrogen, phosphorus, or arsenic, can be distinguished; instead, any solid particle or fragment of one of these elements, irrespective of size, is uniformly bonded throughout, and, therefore, the whole fragment can be considered as a giant molecule. Decreasing melting points, boiling points, and decreasing heat energies associated with fusion (melting), sublimation (change from solid to gas), and vaporization (change from liquid to gas) among these four elements, with increasing atomic number and atomic size, indicate a parallel weakening of the covalent bonds in this type of structure. The actual or probable arrangement of valence electrons is often impossible to determine, and, instead, relative energy states of the electrons, in the ground, or least energetic, state of the atom are considered. Thus, the same trend of nonmetallic toward metallic states is indicated by decreasing hardness and decreasing single-bond energy between atoms. Carbon crystallizes in two forms, as diamond and as graphite; diamond stands apart from all other elemental forms in the extreme stability of its crystal structure, whereas graphite has a layer structure. As may be expected, cleavage between layers of graphite is much easier to effect than rupture within a layer. The crystal structures of white beta (β) tin and elemental lead are clearly metallic structures. In a metal, the valence electrons are free to move from atom to atom, and they give the metal its electrical conductivity.

 



Electron configurations

Electron configurations (B)


The ground-state electronic configurations of atoms of these carbon group elements show that each has four electrons in its outermost shells. As has been explained, if n represents the outermost shell (n being two for carbon, three for silicon, etc.), then these four electrons are represented by the symbols ns2np2. Such a configuration suggests the importance of referring to the relatively stable noble-gas-atom configuration preceding each element in determining the properties of the element, in particular its chemical properties. The loss of four electrons by either a carbon atom or a silicon atom to give ions having a positive charge of four (or +4, written C4+ or Si4+) with the electron configurations of the preceding noble-gas atoms is precluded by the sizable ionization energies. Ions of +4 charge do not exist, nor is there any evidence that carbon or silicon ions of charge +2 can form by the loss of only two unpaired (np, or outermost) electrons. Electron loss by atoms of the heavier elements of the family is easier, but it cannot lead to ions with noble-gas-atom configurations because of the presence of underlying (i.e., d10) arrangements of electrons inside the outermost shell. It is again unlikely that the +4 ions of germanium, tin, and lead (in symbols Ge4+, Sn4+, and Pb4+) exist in known compounds, but it is true that the inertness of the ns2 pair of electrons (which are, in terms of energy states, closer to the nucleus than the np2 pair) increases substantially with increasing atomic number in the family and thus allows the np2 electrons to be removed separately, to form at least the ions, Sn2+ and Pb2+. Oxidation states of +2 and +4 can be assigned in covalent compounds of each of these elements with elements that are more electronegative (i.e., having greater affinity for electrons).

 



Catenation

Catenation (B)

Carbon is unique among the elements in the almost infinite capacity of its atoms to bond to each other in long chains, a process called catenation (Latin catena, chain). This characteristic reflects the strength of the bond between adjacent carbon atoms in the molecule, both in relationship to similar bonds involving other elements of the carbon family and in relationship to bonds between carbon atoms and atoms of many other elements. Only the carbon–hydrogen, carbon–fluorine, and carbon–oxygen single bonds (C―H, C―F, and C―O) are stronger than the carbon–carbon single bond (C―C), and each of these is weaker than the carbon–carbon multiple bonds (C=C or C≡C). On the other hand, the silicon–silicon single bond (Si―Si) is weaker than other single bonds involving an atom of other elements with the silicon atom. The same is undoubtedly true of the germanium–germanium and tin–tin single bonds (Ge―Ge, Sn―Sn) in relationship to single covalent bonds between atoms of these elements and atoms of other elements. Experimentally, there appears to be no practical upper limit to catenation involving carbon. This phenomenon in three dimensions produces the diamond and in two dimensions the layers in graphite. Catenation is also exhibited to a high degree by elemental silicon, germanium, and tin, but it is strictly limited in compounds of these elements; silicon may have up to 14 atoms in a chain; germanium, 9; and tin, 2 or 3 only, largely in hydrides (compounds containing hydrogen). Double and triple bonds in catenated arrangements are limited to carbon.

Catenation, via single or multiple bonds or both, combined with several other factors allows carbon to form more compounds than any other element. These factors are: (1) the stability of certain carbon bonds, in particular of the C―H bond; (2) the existence of carbon in both sp2 and sp3 hybridizations; (3) the ability of carbon to form both chain and cyclic compounds (in which the chain of atoms is joined end to end to form a ring) based upon either carbon atoms alone or carbon atoms in combination with those of other nonmetals (e.g., oxygen, sulfur, nitrogen) and either upon single- or multiple-bond arrangements; and (4) the capability of many carbon compounds to exist in isomeric forms (isomers are molecules with identical numbers of the same atoms bonded in different arrangements; such molecules have quite different properties). All but a very few carbon compounds are called organic compounds, and they are discussed in the article chemical compound.

 



Atomic size

Atomic size (W)

Reference has been made to some of the physical properties of the carbon group elements. Most of the variations in properties from carbon through lead parallel increase in atomic size and are comparable with those of elements in the boron, nitrogen, oxygen, and fluorine groups. The general trends are roughly those found for the adjacent boron group and nitrogen group elements. The significantly higher melting and boiling points of the carbon group elements reflect their tendency to exist as giant molecules, as opposed to the tendencies of elements in the adjacent families to exist as smaller, discrete molecules.

As is true of the lightest element in each group of elements, the physical properties of carbon differ substantially from those of the other members of its family. To a large degree, these differences reflect the substantially higher concentration of the positive charge on the carbon nucleus relative to the size of the carbon atom. That is, the nucleus of carbon holds only six electrons in two shells and, therefore, holds them close; the nucleus of lead, on the other hand, has 82 electrons distributed in six shells. The attraction between the nucleus of lead and its outermost electrons is less than in carbon, because intervening shells in lead shield the outer electrons. Structural differences between diamond and graphite produce profound differences between them in hardness, conductivity, density, heat capacity, and other properties. Inasmuch as graphite is a unique crystalline formation among the elements, its properties should not be compared directly with those of the other elements in the family.

 



Reactions

Reactions (B)

With a given reagent, diamond is generally less reactive than graphite and, thus, requires more rigorous conditions for reaction, such as a higher temperature; the ultimate products, however, are the same. Crystalline silicon is less reactive than finely divided and, possibly, amorphous silicon. Elemental germanium resembles silicon quite closely. Tin and lead behave in general as metals and thus yield at least some ionic products in reactions that are quite different from those of the other elements. Elemental carbon is of particular importance as a high-temperature reducing agent (a reagent that donates electrons) in metallurgical processing for metal oxides, a reaction that frees the metal. For example, tin can be obtained from its ore cassiterite by reduction with carbon in the form of charcoal. Thus to cite only a few of carbon’s more important applications, carbon is used directly in the production of elemental phosphorus, arsenic, bismuth, tin, lead, zinc, and cadmium, and indirectly, as carbon monoxide, in the production of iron. Elemental silicon, in the iron–silicon alloy ferrosilicon, is also a strong reducing agent and has been used as such to liberate magnesium from its oxide.

 



 







 
  Oxocarbon

📥 Oxocarbon (W)

 








 
  Carbon cycle


Movement of carbon between land, atmosphere, and ocean in billions of tons per year. Yellow numbers are natural fluxes, red are human contributions, white are stored carbon. The effects of volcanic and tectonic activity are not included.

📥 Carbon cycle (W)

 









 


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