Properties of Carbon
Carbon has very diverse physical and chemical properties due to the nature of its bonding.
Discuss the properties of carbon.
- Carbon has several allotropes, or different forms in which it can exist. These allotropes include graphite and diamond, which have very different properties.
- Despite carbon’s ability to make 4 bonds and its presence in many compounds, it is highly unreactive under normal conditions.
- Carbon exists in 3 main isotopes: 12C, 13C, 14C. 14C is radioactive and used in dating carbon-containing samples ( radiometric dating ).
- allotropes: Different forms of a chemical element found in its natural state.
- radiometric dating: A technique used to date materials by comparing the natural abundance of radioactive atoms to their remaining decay products.
- half-life: In a radioactive decay process, the amount of time required to end up with half of the original (undecayed) material.
Carbon is the chemical element with the symbol C and atomic number 6 (contains 6 protons in its nucleus ). As a member of group 14 on the periodic table, it is nonmetallic and tetravalent—making four electrons available to form covalent chemical bonds. The most common isotope of carbon has 6 protons and 6 neutrons, and has an atomic mass of 12.0107 amu. Its ground state electron configuration is 1s22s22p2. Its oxidation state ranges from 4 to -4, and it has an electronegativity value of 2.55 on the Pauling scale. It is a solid, and sublimes at 3,642 °C (it has the highest sublimation point of all the elements).
Carbon has several allotropes, or different forms in which it exists. Interestingly, carbon allotropes span a wide range of physical properties: diamond is the hardest naturally occurring substance, and graphite is one of the softest known substances. Diamond is transparent, the ultimate abrasive, and can be an electrical insulator and thermal conductor. Conversely, graphite is opaque, a very good lubricant, a good conductor of electricity, and a thermal insulator. Allotropes of carbon are not limited to diamond and graphite, but also include buckyballs (fullerenes), amorphous carbon, glassy carbon, carbon nanofoam, nanotubes, and others.
Chemical Reactivity of Carbon
Carbon compounds form the basis of all known life on Earth, and the carbon-nitrogen cycle provides some energy produced by the sun and other stars. Carbon has an affinity for bonding with other small atoms, including other carbon atoms, via the formation of stable, covalent bonds. Despite the fact that it is present in a vast number of compounds, carbon is weakly reactive compared to other elements under normal conditions. At standard temperature and pressure, it resists oxidation; it does not react with sulfuric acid, hydrochloric acid, chlorine, or any alkali metals. At higher temperatures, carbon will react with oxygen to give carbon oxides, and metals to give metal carbides.
Carbon has the ability to form very long chains of strong and stable interconnecting C-C bonds. This property allows carbon to form an almost infinite number of compounds; in fact, there are more known carbon-containing compounds than all the compounds of the other chemical elements combined, except those of hydrogen (because almost all organic compounds contain hydrogen as well).
Carbon has two stable, naturally occurring isotopes: carbon-12 and carbon-13. Carbon-12 makes 98.93% and carbon-13 forms the remaining 1.07%. The concentration of 12C is further increased in biological materials because biochemical reactions discriminate against 13C. Identification of carbon in NMR experiments is done with the isotope 13C. 14C is a radioactive isotope of carbon with a half-life of 5730 years. It has a very low natural abundance (0.0000000001%), and decays to 14N through beta decay. It is used in radiometric dating to determine the age of carbonaceous samples (of physical or biological origin) up to about 60,000 years old.
In total, there are 15 known isotopes of carbon and the shortest-lived of these is 8C, which decays through proton emission and alpha decay, and has a half-life of 1.98739 x 10−21 seconds. The exotic 19C exhibits a nuclear halo, which means its radius is appreciably larger than would be expected if the nucleus were a sphere of constant density.
Carbides are a class of compounds composed of carbon and an electropositive atom.
Review the different types of carbides.
- Carbides are generally formed at high temperatures (> 1500 °C).
- Carbides are generally quite stable and exhibit high melting points.
- Carbides can be classified as salt-like, interstitial, and covalent.
- electronegativity: The tendency of an atom to attract electrons to itself.
Types of Carbides
Carbides are compounds composed of carbon and less electronegative elements and they are distinguished by their chemical bonding (ionic, covalent). They are generally prepared from metals or metal oxides at high temperatures (1500 °C or higher) by combining the metal with carbon. Carbides are used in key industrial applications.
Salt-like (saline) carbides are composed of the highly electropositive atoms, such as the alkali, alkali earth, and group -III metals, mixed with carbon. Aluminum forms carbides, but other elements from group XIII do not. These materials have isolated carbon centers, often described as “C4-” in the metanides, “C22-” in the acetylides, and “C34-” in the sesquicarbides.
- Methanides are carbides that decompose in water and generate water; aluminum carbide (Al4C3) and beryllium carbide (Be2C) are examples of this class of carbides.
- Acetylides are formed from alkali, alkali earth, and lanthanoid metals with the acetylide anion C22-. Lanthanoids also form carbides with the formula M2C3. Metals from group XI also form acetylides, such as copper(I) acetylide and silver acetylide. Carbides of the actinide elements, which have the structure MC2 and M2C3, are also described as salt-like derivatives of C22-.
- The polyatomic ion C34- is referred to as an allylenide or sesquicarbide and is found in Li4C3 and Mg2C3. The allylenide is linear and isoelectronic with CO2.
Covalent carbides are found in carbides of silicon and boron. The reason these two elements form “covalent” carbides is due to their similar electronegativity and size to carbon. Because of this, their association is completely covalent in character. Silicon carbide has two similar crystalline forms, which are both related to the diamond structure. Boron carbide (B4C), on the other hand, has an unusual structure that includes icosahedral boron units linked by carbon atoms. In this respect, boron carbide is similar to the boron-rich borides. Both silicon carbide (also known as carborundum) and boron carbide are very hard and refractory materials. Both materials have important industrial applications.
Interstitial carbides describe the carbides of the group-IV, -V, and VI transition metals. These carbides are metallic and refractory. They are formed so that the carbon atoms fit into octahedral interstices in a close-packed metal lattice when the metal atom’s radius is greater than ~135 pm. When the metal atoms are cubic-close-packed (ccp), then filling all of the octahedral interstices with carbon achieves 1:1 stoichiometry with the rock-salt structure. When the metal atoms are hexagonal-close-packed, (hcp), since the octahedral interstices lie directly opposite each other on either side of the layer of metal atoms, filling only one of these with carbon achieves 2:1 stoichiometry. As a result of the packing, they are quite stable and have very high melting points and low electrical resistance.
Intermediate Transition Metal Carbides
In intermediate transition metal carbides, the transition-metal ion is smaller than the critical 135 pm, and the structures are not interstitial but are more complex. Multiple stoichiometries are common. For example, iron forms a number of carbides: Fe3C, Fe7C3 and Fe2C. The best known is cementite, Fe3C, which is present in steels. These carbides are more reactive than the interstitial carbides; for example, the carbides of Cr, Mn, Fe, Co, and Ni are all hydrolyzed by dilute acids (and sometimes by water) to yield a mixture of hydrogen and hydrocarbons. These compounds share features with both the inert interstitials and the more reactive salt-like carbides.
Metal complexes containing Cn fragments are well known. These molecular carbides often have carbon-centered clusters. Some metals such as tin and lead are not believed to form carbides.
Carbon Oxides and Carbonates
Oxocarbons are compounds containing carbon and oxygen; they exist as carbon oxides and carbonates.
Discuss the chemical properties of oxocarbon compounds.
- Carbon monoxide is the simplest carbon oxide, consisting of one carbon atom bonded to an oxygen atom. It is highly toxic.
- Carbon dioxide is a linear compound composed of a carbon atom bonded to two oxygen atoms. It exists predominately as a gas and is a product of the human metabolism.
- Carbon dioxide is soluble in water, in which it readily and reversibly converts to carbonic acid. The conjugate bases of a carbonic acid are known as the bicarbonate and carbonate ions.
- Carbonates are the salts of carbonic acids. They form when a positively charged metal ion comes into contact with the oxygen atoms of the carbonate ion. These compounds are often insoluble in water and exhibit some level of basicity or acidity in aqueous solutions.
- oxocarbon: A compound containing only atoms of carbon and oxygen.
- oxide: A binary chemical compound of oxygen with another chemical element.
- centrosymmetric: Having a center of symmetry.
Properties of Carbon Oxides
Carbon oxides, or oxocarbons, are a class of organic compounds containing only carbon and oxygen. The most basic oxocarbons are carbon monoxide and carbon dioxide. Many other stable and metastable oxides of carbon are known but are rarely encountered.
The simplest oxocarbon is carbon monoxide (CO). Carbon monoxide is a colorless, tasteless gas that is slightly lighter than air. It is toxic to humans and animals when encountered in higher concentrations, despite the fact that it is produced in the metabolism and is thought to have some biological functions.
Carbon monoxide consists of one carbon and one oxygen atom connected by a triple bond. The distance between the carbon and oxygen atom is 112.8 pm, consistent with the presence of a triple bond. The bond dissociation energy of CO is 1072 kJ/mol and represents the strongest chemical bond known. CO has three resonance structures, but the structure with the triple bond is the best approximation of the real distribution of electron density in the molecule.
CO is naturally produced by the human body as a signaling molecule. Abnormalities in its metabolism have been linked to a variety of diseases, including hypertension and heart failure. CO is present in small amounts in the atmosphere, mostly as a result of the burning of fossil fuels and fires. Through natural processes in the atmosphere, it is eventually oxidized to carbon dioxide (CO2).
Carbon dioxide, or CO2, is a naturally occurring linear compound composed of two oxygen atoms covalently bonded to a carbon atom. The two C=O bonds are equivalent and short (116.3 pm), consistent with double bonding. The compound is centrosymmetric and so has no net dipole. CO2 is colorless; at high concentrations it has a sharp, acidic odor, but at lower concentrations it is odorless. At standard temperature and pressure, its density is 1.98 kg/m3, about 1.5 times that of air. It has no liquid state at pressures below 520 kPa; at 1 atm, the gas deposits directly to a solid at temperatures below -78.5 °C, and the solid sublimes directly to gas above this temperature. Solid CO2 is known as dry ice.
Carbon dioxide in Earth’s atmosphere currently occurs at an average concentration of about 390 parts per million by volume. Concentrations of the gas tend to fall during the northern spring and summer as plants consume the gas (during the process of photosynthesis), and rise during autumn and winter as plants go dormant, decay, or die. CO2 is an end product of the metabolism of organisms via the cellular respiration process, in which energy is obtained from the breaking down of sugars, fats, and amino acids. Despite the fact that the human body produces approximately 2.3 pounds of carbon dioxide per day, it is considered toxic, and concentrations up to 10 percent may cause suffocation.
Carbonic Acid and Its Conjugate Bases
Carbon dioxide is soluble in water; it reversibly converts to carbonic acid (H2CO3). The salt of carbonic acids are called carbonates and are characterized by the carbonate ion, CO32-. The carbonate ion is the simplest oxocarbon anion, consisting of one carbon atom surrounded by three oxygen atoms in a trigonal planar arrangement. The Lewis structure of the carbonate ion has two single bonds to negative oxygen atoms and one short double bond to a neutral oxygen. This structure is, however, incompatible with the ion’s observed symmetry, which implies that the three bonds and oxygen atoms are equivalent. The symmetry can best be represented by three resonance structures.
In aqueous solutions, carbonate, bicarbonate (HCO3–), carbon dioxide, and carbonic acid exist together in equilibrium. In strongly basic conditions, the carbonate ion predominates, while in weakly basic conditions, the bicarbonate ion predominates. In acidic conditions, aqueous CO2 (aq) is the main form and is in equilibrium with carbonic acid — the equilibrium lies strongly towards carbon dioxide.
Metal carbonates generally decompose upon heating, liberating carbon dioxide and leaving behind an oxide of the metal. Ionic compounds form when a positively charged metal ion, M+, attaches to the negatively charged oxygen atoms of the carbonate ion. Most salts are insoluble in water, with solubility constants (Ksp) less than 1 x 10-8, with the exception of lithium, sodium, potassium, and ammonium carbonates.
Sodium carbonate is basic when dissolved in water (meaning it results in a basic solution upon dissolution ), and sodium bicarbonate is weakly basic. These effects can be explained by considering that upon dissolution and subsequent dissociation of the salt into its ions, the carbonate or bicarbonate ions will react with H+ in the solution to form H2CO3 (which has a low Ka value – i.e., is a weak acid). On the other hand, carbon dioxide is weakly acidic (results in a slightly acidic solution) when dissolved in water. That’s because it reacts with water to produce H2CO3, a small amount of which will dissociate into H+ and a bicarbonate ion.
Although the carbonate salts of most metals are insoluble in water, this is not true of the bicarbonate salts. Under changing temperature or pressure, and in the presence of metal ions with insoluble carbonates, the equilibrium between carbonate, bicarbonate, carbon dioxide, and carbonic acid in water can result in the formation of insoluble compounds. This is responsible for the buildup of scale inside pipes caused by hard water.
Allotropes of Carbon
Various allotropes of carbon exhibit different properties and find applications in a variety of fields.
Describe the properties of the allotropes of carbon.
- Diamond is a well-known allotrope of carbon that exhibits hardness and high dispersion of light. It is the hardest known natural mineral and finds applications in cutting, drilling, and jewelry, and as a potential semiconductor material.
- Graphene is a single layer of carbon atoms arranged in one plane; layers of graphene make up graphite. Graphene is a material of interest due to its high electron mobility and its possible applications in electronics.
- Fullerenes are a class of carbon allotropes in which carbon takes the form of a hollow sphere, ellipsoid, or tube. This class of materials includes carbon nanotubes, buckyballs, and the newly discovered nanobuds.
- allotropes: Different forms of a chemical element.
Allotropy is the property of some chemical elements to exist in two or more different forms, or allotropes, when found in nature. There are several allotropes of carbon.
Diamond is probably the most well known carbon allotrope. The carbon atoms are arranged in a lattice, which is a variation of the face-centered cubic crystal structure. It has superlative physical qualities, most of which originate from the strong covalent bonding between its atoms. Each carbon atom in a diamond is covalently bonded to four other carbons in a tetrahedron. These tetrahedrons together form a three-dimensional network of six-membered carbon rings in the chair conformation, allowing for zero bond -angle strain. This stable network of covalent bonds and hexagonal rings is the reason that diamond is so incredibly strong as a substance.
As a result, diamond exhibits the highest hardness and thermal conductivity of any bulk material. In addition, its rigid lattice prevents contamination by many elements. The surface of diamond is lipophillic and hydrophobic, which means it cannot get wet by water but can be in oil. Diamonds do not generally react with any chemical reagents, including strong acids and bases. Uses of diamond include cutting, drilling, and grinding; jewelry; and in the semi- conductor industry.
Graphite is another allotrope of carbon; unlike diamond, it is an electrical conductor and a semi-metal. Graphite is the most stable form of carbon under standard conditions and is used in thermochemistry as the standard state for defining the heat of formation of carbon compounds. There are three types of natural graphite:
- Crystalline flake graphite: isolated, flat, plate-like particles with hexagonal edges
- Amorphous graphite: fine particles, the result of thermal metamorphism of coal; sometimes called meta-anthracite
- Lump or vein graphite: occurs in fissure veins or fractures, appears as growths of fibrous or acicular crystalline aggregates
Graphite has a layered, planar structure. In each layer, the carbon atoms are arranged in a hexagonal lattice with separation of 0.142 nm, and the distance between planes (layers) is 0.335 nm. The two known forms of graphite, alpha (hexagonal) and beta (rhombohedral), have very similar physical properties (except that the layers stack slightly differently). The hexagonal graphite may be either flat or buckled. The alpha form can be converted to the beta form through mechanical treatment, and the beta form reverts to the alpha form when it is heated above 1300 °C. Graphite can conduct electricity due to the vast electron delocalization within the carbon layers; as the electrons are free to move, electricity moves through the plane of the layers. Graphite also has self-lubricating and dry lubricating properties. Graphite has applications in prosthetic blood-containing materials and heat-resistant materials as it can resist temperatures up to 3000 °C.
A single layer of graphite is called graphene. This material displays extraordinary electrical, thermal, and physical properties. It is an allotrope of carbon whose structure is a single planar sheet of sp2 bonded carbon atoms that are densely packed in a honeycomb crystal lattice. The carbon-carbon bond length in graphene is ~0.142 nm, and these sheets stack to form graphite with an interplanar spacing of 0.335 nm. Graphene is the basic structural element of carbon allotropes such as graphite, charcoal, carbon nanotubes, and fullerenes. Graphene is a semi-metal or zero-gap semiconductor, allowing it to display high electron mobility at room temperature. Graphene is an exciting new class of material whose unique properties make it the subject of ongoing research in many laboratories.
Amorphous carbon refers to carbon that does not have a crystalline structure. Even though amorphous carbon can be manufactured, there still exist some microscopic crystals of graphite-like or diamond-like carbon. The properties of amorphous carbon depend on the ratio of sp2 to sp3 hybridized bonds present in the material. Graphite consists purely of sp2 hybridized bonds, whereas diamond consists purely of sp3 hybridized bonds. Materials that are high in sp3 hybridized bonds are referred to as tetrahedral amorphous carbon (owing to the tetrahedral shape formed by sp3 hybridized bonds), or diamond-like carbon (owing to the similarity of many of its physical properties to those of diamond).
Fullerenes and Nanotubes
Carbon nanomaterials make up another class of carbon allotropes. Fullerenes (also called buckyballs) are molecules of varying sizes composed entirely of carbon that take on the form of hollow spheres, ellipsoids, or tubes. Buckyballs and buckytubes have been the subject of intense research, both because of their unique chemistry and for their technological applications, especially in materials science, electronics, and nanotechnology. Carbon nanotubes are cylindrical carbon molecules that exhibit extraordinary strength and unique electrical properties and are efficient conductors of heat. Carbon nanobuds are newly discovered allotropes in which fullerene-like “buds” are covalently attached to the outer side walls of a carbon nanotube. Nanobuds therefore exhibit properties of both nanotubes and fullerenes.
Glassy or vitreous carbon is a class of carbon widely used as an electrode material in electrochemistry as well as in prosthetic devices and high-temperature crucibles. Its most important properties are high temperature resistance, hardness, low density, low electrical resistance, low friction, low thermal resistance, extreme resistance to chemical attack, and impermeability to gases and liquids.
Other allotropes of carbon include carbon nanofoam, which is a low-density cluster assembly of carbon atoms strung together in a loose three-dimensional web; pure atomic and diatomic carbon; and linear acetylenic carbon, which is a one-dimensional carbon polymer with the structure -(C:::C)n-.